chemistry notes / igcse-gcse- Olevel

collecting gases / igcse /gcse chemistry notes

2009-11-27T08:10:00.000-08:00

We can collect gases in
4 ways :
1- Downwards into a testube or a gas jar , this is used for gases which denser than air .
2- Upwards into a testube are a gas jar , this used to collect gases which less dense than air and are soluble in water .
3- Over water , - to collect gases which are less dense than air and are insoluble in water- or to collect gases which are having the same dense like air and are insoluble in water .
4- In a gas syringe , a gas syringe is used to collect the gas and to measure the volume of the gas collected .

2- Normal test – pure water boils at 100 degrees Celsius

TITRATION OF UNKNOWN ACID

2009-11-22T06:59:00.000-08:00

Titration of an Unknown (if done as per your instructor's directions)
1. Obtain an unknown solid and record its number.
2. Weigh out 0.1-0.2 g and place it in an Erlenmeyer flask.
3. Dissolve the acid in about 20 mL boiled distilled water.
4. Add two drops of the phenolphthalein indicator.
5. Titrate to the pale pink endpoint.6. Carry out at least two titrations; repeat until two consecutive values for equivalent weight of the acid differ by no more than 1%.

standardization ഓഫ് NaoH

2009-11-22T06:55:00.000-08:00

Standardization of NaOH1.
Clean a buret and rinse it with the NaOHyou just prepared.
2. Fill the buret with the NaOH.
3. Weigh a clean dry beaker on a balance.
4. Weigh out 0.2 - 0.4 g of KHC8H4O4 and place it in the beaker you just weighed. Make sure you record the weights!
5. Pour the acid into an Erlenmeyer flask; use a stream of water from the wash bottle to complete the transfer.
6. Dissolve this acid in about 20 mL of the boiled distilled water.
7. Add two drops of the phenolphthalein indicator to the flask.
8. Read the buret and record the initial volume making sure there is not a drop on the tip of the buret; if there is, use an extra beaker to get rid of this drop by touching the side of the beaker to the tip of the buret.
9. Place the Erlenmeyer flask under the buret, and place a piece of white paper under it.
10. Position the tip of the buret so it is just beneath the rim of the flask.
11. Add several mL (probably no more than ten) of base solution (NaOH) rapidly with constant swirling.
12. Close buret stopcock to only allow a rapid stream of drops and swirl; at the site where the solution in the buret drops in the acid solution, you will see a pink color, which will disappear

13. When the pink color lingers longer, close the stopcock on the buret more so that you add a drop and swirl until the color changes before adding another drop.
14. Occasionally wash down the sides of the flask with the wash bottle.
15. Near the end point you will want to use partial drops. To do this, open the stopcock so that a drop forms on the tip of the buret. Close the stopcock and use a stream of water from the water bottle to wash the droplet into the sol.
16. Keep adding base until a very light pink color persists for at least 20 seconds.
17. Record the final buret reading.18. Repeat this procedure at least two more times; repeat the titrations until two values for molarity differ by no more than 1% C.

Metallic bonding work sheet IGCSE

2009-11-18T04:39:00.000-08:00

1. DESCRIBE METTALIC BONDING
2. WHY METALS ARE MALLIABLE?
3. DEFINE ALLOYS? GIVE EX .
4. WHY ALLOYS ARE HARD IN NATURE?

IONIC BONDING / WORK SHEET / IGCSE / GCSE / CHEMISTRY

2009-11-18T04:33:00.000-08:00

1.Draw the dot and cross diagram dor the following
Calcium oxide
Magnesiumm chloride
Lithum chloride
2. Give the properties of ionic com pounds
3. sodium chloride conduct electricity in liquid state . why/
4. sodium chloride do not conduct electricity in solid state . why/

The Noble Gases- igcse / gcse/ o level

2009-11-18T04:20:00.000-08:00

1.What is the Group Number of the Noble Gases?
2.What does Monatomic mean?
3.Going down the Group, do the Noble Gases have a Higher Boiling Point?
4.Give one Use of Helium?
5.Give one Use of Neon?
6.Give one Use of Argon?

The Halogens- igcse/ gcse/ o level

2009-11-18T04:17:00.000-08:00

The Halogens- igcse/ gcse/ o level
1.What is the Group Number of the Halogens?
2.What does Diatomic mean? 37.What Colour is Chlorine?
3.What Colour is Bromine?
4.Is Iodine a Liquid?
5.Going down the Group, do the Halogens become More Reactive?
6.Going down the Group, do the Halogens have a Higher Boiling Point?
7.Write the Balanced Equation for the reaction between Aluminium and Chlorine.
8.Write the Balanced Equation for the reaction of Chlorine with Potassium Iodide.
9.Write the Ionic Equation for the reaction of Chlorine with Potassium Iodide.
10.Give one Use of Fluoride?
11.Give one Use of Chlorine? 47Give one Use of Bromide?
12.What does Hydrogen Chloride make when it is Dissolved in Water?
13.How is Hydrogen Chloride Safely Dissolved in Water?

The Transition Metals- igcse/ gcse/olevel

2009-11-18T04:16:00.000-08:00

1.Where do you find the Transition Metals in the Periodic Table?
2.What is their Group Number?
3.Do the Transition Metals Conduct Electricity?
4.Do the Transition Metals form Coloured Compounds?
5.Give one Use of Copper?
6.Give one Example of a Transition Metal Used as a Catalyst.

Electrolysis of Sodium Chloride in Water

2009-11-18T04:12:00.000-08:00

1.What is Brine?
2.Which Gas is given off at the Cathode?
3.Which Gas is given off at the Anode?
4.Give the Ionic Equation for the Gas given off at the Anode.
5Why don't you get Sodium Metal at the Cathode?
6.What Substance is left in Solution after Electrolysis?
7.Give one Use of this Substance?

The Alkali Metals- igcse/ gcse /olevel

2009-11-18T04:09:00.000-08:00

1.What is the Group Number of the Alkali Metals?

2.Why are the Alkali Metals stored under oil?

3.Give two Properties of the Alkali Metals?

4.Going down the Group, do the Alkali Metals become More Reactive?

5.Write the Word Equation for the reaction when Potassium burns in Air.

6.What Colour is the Flame from Potassium?

7.Write the Balanced Equation for the reaction when Potassium burns in Air.

8.Write the Balanced Equation for the reaction between Sodium and Water.

9.Write two things you would See in the reaction between Sodium and Water.

10.Write the Balanced Equation for the reaction between Lithium and Chlorine.

11.Give one Property of an Alkali Metal Compound?

12.Give one Use of Sodium Chloride?

13.Give one Use of Sodium Carbonate?

Preparation of NaOH solution

2009-11-18T03:58:00.000-08:00

InstructionsA.
Preparation of NaOH സോലുറേന്‍
1. Measure about 500 mL of distilled water and boil it for about five minutes.
2. Allow the beaker to cool enough so you can pick it up with a towel.
3. Using a graduated cylinder, measure 290 mL of this water into a large flask.
4. Add 10 mL of 3M NaOH to this water.CAUTION! NaOH IS CORROSIVE!!!!!!! If you spill any on yourself, wash it off immediately!!!!5. Stopper the flask, making sure not to let the stopper come in contact with the solution.6. Swirl the flask to mix the NaOH and water.7. Pour the rest of the water you boiled into a wash bottle. You can use this later to dissolve the acid sample and wash down the sides of the flask during the titration.B.

alcohols

2009-04-09T03:21:00.001-07:00

Alcohols make good fuels because they burn easily and they release a lot of heat energy, which can be harnessed to power machines etc...

The combustion of alcohol is an exothermic reaction and this means that heat is given out from the reaction, opposed to being taken in, which would occur in an endothermic reaction. The following equation shows the result of combustion of alcohol (in this case, ethanol).

Alcohol + Oxygen ----> Carbon Dioxide + Water

C2H5OH + 3O2 ----> 2CO2 + 3H2O

The general formula for alcohols is as follows:
CnH2n+1OH (where "n" is a constant) The Alcohols form a homologous series where they all have similar chemical structures and properties. They are produced from either anaerobic respiration or by reacting water with the alkenes (Hydration).

The first five members of this homologous series are:

CH3OH (Methanol)C2H5OH (Ethanol)C3H7OH (Propanol)C4H9OH (Butanol)C5H11OH (Pentanol)

contact process- IGCSE/GCSE NOTES

2009-04-09T03:18:00.000-07:00

The contact process is the name given to the process by which sulphuric acid is produced. Sulphuric acid has many uses such as the manufacturing of:
Paints / Pigments
Soaps / Detergents
Fibres
Plastics
Fertilisers The following reactions are involved:
S (s) + O2 (g) ----> SO2 (g) (sulphur dioxide)
2SO2 (g) + O2 (g) ----> 2SO3 (g) (sulphur trioxide)

work sheet - balancing equations

2009-04-09T03:09:00.001-07:00

. _____Zn + _____HCl --> _____ZnCl2 + _____H2
2. _____NH3 + _____HCl --> _____NH4Cl
3. _____Al + _____HCl --> _____AlCl3 + _____H2
4. _____Mg + _____H3PO4 --> _____Mg3(PO4)2 + _____H2
5. _____Cu + _____AgNO3 --> _____Cu(NO3)2 + _____Ag
6. _____Ca + _____Pb(NO3)2 --> _____Pb + _____Ca(NO3)2
7. _____Al + _____Pb(NO3)2 --> _____Pb + _____Al(NO3)3
8. _____Zn + _____Sn(NO3)4 --> _____Zn(NO3)2 + _____Sn
9. _____Cl2 + _____AlI3 --> _____AlCl3 + _____I2
10. _____Br2 + _____CuI --> _____CuBr + _____I2
11. _____NH4OH + _____FeCl3 --> _____NH4Cl + _____Fe(OH)3
12. _____KBr + _____Pb(NO3)2 --> _____KNO3 + _____PbBr2
13. _____AlCl3 + _____H2SO4 --> _____Al2(SO4)3 + _____HCl
14. _____Al2(SO4)3 + _____BaCl2 --> _____BaSO4 + _____AlCl3
15. _____Na2CO3 + _____CaCl2 --> _____CaCO3 + _____NaCl

AS level chemistry practical (GCE)

2009-01-09T11:13:00.000-08:00

Chemistry 101 Laboratory
Experiment 10: Acid/ Base Titrations
Pre-Laboratory
Pre-lab Questions
1. How would your calculated molarity of NaOH be affected by each of the following errors? Elaborate.

a.) The buret is not rinsed with solution before filling.
b.) Some solution splashes out of the flask during titration.
c.) You go past the end point in the titration.
2. How would your calculated value for the equivalent weight of the acid be affected by each of the following errors? Elaborate.
a.) Some acid is spilled out of the flask after it is weighed.
b.) Some solution splashes out of the flask during titration.

c.) You go past the end point of the titration.
3. What is the effect of allowing a solution of NaOH to remain in contact with the air in the laboratory? Be specific as to the two, main effects and their sources.
Laboratory
The reaction of an acid and a base in water involves the reaction of hydroxide and hydrogen ions to form water (neutralization) and the formation of the resulting salt from the conjugate base and acid. The latter may or may not be soluable.

In general, the reaction is given by the following reaction for simple acids and bases:
HA + HOX <-> H2O + A- + X+
In this example, just one hydrogen and one hydroxl ion were available for the reaction. The ionization of the acid therefore produces one "equivalent" (one mole per mole) of hydrogen ions and ionization of the base produces one equivalent of hydroxyl ions. ( It is important to remember that, in water, the hydrogen ion and hydroxyl ion are the strongest acid and base, respectively. For this reason, ionization of any acid or base in water leads to the production of soley these acidic and basic ions.)

It is possible that one mole of the acid or base will produce more than one mole (one equivalent mole) of hydrogen ions or hydroxyl ions. An example of this is the diprotic acid sulfuric H2SO4.

In this instance, a total of 2 moles of hydrogen ions are produced when the acid is completely ionized to the sulfate.
When a titration is carried out, an acid or base of unknown concentration is allowed to neutralize a base or acid known to produce a known number of moles of hydroxyl or hydrogen ions.

As the acid or base is added to the solution of a known number of moles, the acids and bases react to neutralize one another.

At the point where the number of moles of acid equals the number of moles of base, we have reached what is called equivalence.

This is also known as the equivalence point and is shown either by a indicating dye or pH meter.

At the equivalence point we know the number of moles of acid or base reacted (from the fact that we know the number of moles of base or acid originally present from the sample of known amount) and the number of liters of the unknown which were added to neutralize the solution.

From the definition of molarity we have the following relationship:
Equivalent molarity = (# moles originally )/ ( liters of unknown added)
Notice that this differs slightly from our usual definition of molarity.

The reason is that we have not taken into account the fact that our unknown acid or base may have more than one equivalent per mole.

In today's experiment, you will deal with a monoprotic acid and a simple hydroxyide (NaOH) so that the concentration you obtain will be the true molarity of the base.
One of the important concepts in any aspect of chemistry in which quantitative answers are required is that of the primary standard.

This is an agreed upon set of conditions, chemical compound, or anything else which is agreed to represent a measure to which other things are compared.

One such standard concerns hydrogen ion concentration and is the use of potassium hydrogen phthalate (KHC8H4O4, molecular weight 204.224) as a standard for the concentration of hydrogen ions in a solution.

One mole of this salt produces one mole of hydrogen ions in one liter of water.
A second thing which is required in the use of a primary standard is a means of knowing when all of the (in this case) hydrogen ions from the standard have been neutralized with the base.

One way of doing this is to use an indicator dye.

These are organic chemicals which are also weak acids and bases. On their ionization or protonation, they undergo a chemical change which also results in a change in color.

One such dye is phyenylphtalene, a close relative of the primary standard you are using. This dye is clear in acidic solutions, slightly pink in neutral pH solutions, and dark purple in basic solutions.

Addition of a small amount of the dye to the solution you are testing allows you to determine when an equivalent number of acids and bases were present.

AS level chemistry practical (GCE) Titration

2009-01-09T11:09:00.000-08:00

Titration

Titration that involve H2SO4 and NaOH is an examples of acid-base titration of a strong base acid and strong base, both the titrant and the analyte are completely ionized. The balanced reaction and its ionic reaction are:

H2SO4 + 2NaOH -> Na2SO4 + 2H2OH+ + SO42- + 2 Na+ + 2OH- -> 2Na+ + SO42- + 2H2OH+ + + 2OH- -> H2O

The H+ combined with OH- to form H2O, and other ions (Na+ and SO42-) remain unchanged.

This is an example of neutralization reaction, the net result of this neutralization is conversion of H2SO4 into a neutral solution Na2SO4 in the equivalent point.

A titration curve is constructed by plotting the pH of the solution as a function of the volume of titrant added.

The volume changes during titration must be employed for determining the concentrations of the species in the solutions (H+, OH- ). This is the example of the titration curve of 0.1 M H2SO4 100 mL versus 0.1 M of NaOH.
We began with 0.1 mL H2SO4 100 mL in Erlenmeyer flask and then 0.1 mL of NaOH in the buret.
At the start of the titration there is only H2SO4 in the solution so the pH of this solution is:
[H2SO4] = 0.1 M[H+] = 2 x 0.1 = 0.2 MpH = - log [H+] = -log(0.2) = 0.699

as the titration begins, NaOH solution is added from the burette into the H2SO4 solution in the Erlenmeyer.

The pH of the solution when the 20 mL NaOH 0.1 M is added to the solution is:
moles of H2SO4 = 0.1 M x 100 mL = 10 mmolmoles of NaOH = 0.1 M x 20 mL = 2 mmol

AS level chemistry practical (GCE)

2009-01-09T10:58:00.000-08:00

A level chemistry practicals( GCE)
1. Titration -Acid base titration – Hcl AND NaOH
Purpose:

To determine the molarity of an unknown acid solution by titration using phenolphthalein as an indicator.

Materials:

Burets
Clamp
Stand
Beakers
1M HCl solution
1M NaOH solution
Unknown HCl solution
Phenolphthalein indicator solution

Procedure:

Clean and dry the burets and beaker, and clamp the two burets to the ring stand. Fill one of the two burets with 1M HCl solution, and the other with the NaOH solution.
Use the buret to measure out 20 mL of HCl into an empty beaker. Add 2-3 drops of the indicator solution.
Titrate slowly with the NaOH solution, with constant swirling, until one single drop of NaOH causes a permanent pink color that does not fade onswirling. Record the volume of NaOH used.
Use the formula M1V1=M2V2 to determine the concentration of the NaOH solution. This solution may now be used to titrate the unknown acid sample.
Replace the buret containing the 1M HCl with the buret containing the HCl solution of unknown concentration. Refill the NaOH buret, and wash out the beaker.
Repeat the titration from steps 2-4 using 20 mL of the unknown acid solution to determine the concentration of the HCl solution.

revision work sheet- Periodic Table

2008-12-12T04:44:00.000-08:00

1.Why did Dmitri Mendeleev leave gaps in his Periodic Table?

2.What is a Column called?

3.What is a Row called?

4.Where do you find Non-Metals in the Periodic Table?

5.What is important about the Group Number of an Element?

6.What is the Electron Structure of Potassium?

ATOMIC STRUCTURE igcse -chemistry / revision work ഷീറ്റ്

2008-12-12T04:27:00.000-08:00

1.What is an Atom?
2.Which Particles are found in the Nucleus?
3.What is the charge on an Electron?
4.What is the charge on a Neutron?
5.Does a Proton have more mass than a Neutron?
6.Does a Proton have more mass than an Electron?
7.What is the Atomic Number?
8.What is an Isotope?
9.What is the maximum number of Electrons in each Shell?

Making metals useful-gcse chemistry

2008-11-27T06:46:00.000-08:00

Making metals useful

The reactivity of aluminium and anodisingAluminium is high in the reactivity series but does not corrode in air or water.

The reason is that aluminium covers itself with a thin layer of aluminium oxide which protects it from further corrosion.

Anodising is the industrial process of coating aluminium objects with a thicker layer of aluminium oxide.

The aluminium object forms the anode (positive electrode) in a sulfuric acid electrolyte during electrolysis.

Oxygen atoms form at the anode and join with the aluminium. This is oxidation.

Draw a labelled diagram for apparatus suitable for anodising an aluminium rod.

Show the electrodes, the electrolyte, and the cell.

Explain duralumin use in aircraft instead of aluminium.

Explain magnalium use in window frames instead of aluminium.

Important uses of aluminium and its alloys
metal or alloy

aluminium -overhead power cables

good electrical conductor, low density

aluminium -drinks cans -Does not react with water

aluminium -cooking pots -good heat conductor

duralumin -aircraft and bicycle parts -high strength, low density and hard

magnalium -aircraft parts -high strength, low density and corrosionresistance

Chemical reactions in different parts of the blast furnace (high tier)Iron is made in the blast furnace.

Write equations for each of the reactions in the blast furnace.

The limited uses of pure iron and impure iron from the blast furnaceImpure iron from the blast furnace (only 93% pure) is called cast iron.

Pure or wrought iron is not now mass produced but is still available.

cast iron

brittle, high compression strength
car engine blocks, man hole covers, gas stoves.

wrought iron
soft, bends easily, easily worked, low corrosion
nails, bolts, chains, garden gates, decorative ironwork

The production of mild steelIron from the blast furnace contains impurities like carbon, sulfur, silicon and phosphorus.

These are removed in the basic oxygen process. In this process a water cooled lance is put into impure molten iron and pure oxygen is blown through it.

The impurities are changed to oxides which come out as gases such carbon dioxide, and sulfur dioxide.

Solid oxides formed react with added calcium oxide to form a slag which floats on top of the iron.

When the amount of carbon drops to about 1 or 2% the process is stopped and the result is called mild steel.

The uses of mild steelMild steel containing a small amount of carbon eg 0.5% is has the hardness and strength for making machines, rails, ship's plates and girders for bridges and buildings.

The uses of alloy steelsThe properties of steels can be controlled by carefully changing the amounts of carbon removed and amounts of other metals put into them.

Quantitative chemistry( chemical calculation)-IGCSE/GCSE

2008-11-27T06:34:00.000-08:00

Avogadro’s Law and use it to

calculate volumes of gases in reactionsAvogadro's law states that equal volumes of gases measured at the same temperature and pressure contain equal numbers of molecules.

The numbers of molecules shown in a chemical equation give the ratio of volumes of reacting gases. e.g. if steam is made the equation is:2H2(g) +O2(g) ---> 2H2O(g)

This means that 2 molecules of hydrogen react with 1 molecule of oxygen to form 2 molecules of steam.So 2dm3 of hydrogen reacts with 1 dm3 of oxygen and form 2dm3 of steam.

Or 4dm3 of hydrogen reacts with 2dm3 of oxygen and form 4dm3 of steam etc.volume of H2(g)/ volume of O2(g)= molecules of H2(g)/ molecules of O2(g)Write two other equations linking volumes and molecules in the above equation.

Calculate the volume of steam formed if 10cm3 of hydrogen is burned in oxygen.volume of steam/volume of hydrogen = molecules of steam/molecules of hydrogenvolume of steam= volume of hydrogen * molecules of steam/molecules of hydrogenvolume of steam = 10cm3 * 2/2 = 10cm3

Calculate the volume of steam formed if 10cm3 of oxygen is used to burn hydrogen.Calculate the volume of HCl gas formed if 2dm3 of hydrogen is burned in chlorine.

What volume of hydrogen and nitrogen is needed to make 30dm3 of ammonia by the equationN2 (g) + 3H2 (g)= 2NH3(g)

When ammonia is oxidised by oxygen what volume of NO and steam is formed by 25cm3 of ammonia? 4NH3(g) + 5O2(g) = 4NO(g) + 6H2O(g)

The chemical amount represents a number of particles.

One mole is a special very large number called the Avogadro number. It is 6*1023. We cannot count particles but can weigh substances. The molar mass is the mass of one mole of particles e.g. atoms, of a substance.
The molar mass of an element made of atoms is the relative atomic mass of the element in grams.
The molar mass of a compound of element made of molecules is the relative molecular mass or relative formula mass of that substance in grams.

These 3 quantities are connected:amount = mass/molar masse.g

What amount of methane molecule CH4 is there is 4g of methane?C=12 H=1 so molar mass of methane = 12 + (4*1) = 16g/molAmount = mass/molar mass = 4g/16g/mol = 0.25mol

What are the molar masses of the following; He, Na, Cl, Cl2, O2, N, FeS, MgO, KF, HCl, H2O, NH3, NaOH, HNO3, H2SO4, Ca(OH)2.

What amounts are the following masses; 2g of H, 2g of He, 7g of Cl2, 11.2g of FeS, 73g of HCl, 8g of NaOH, 25g of CaCO3.

What is the mass of: 1 mol of Li, 2 mol of C, 3 mol of S, 1 mol of O2, 1 mol of O3, 1 mol of NaCl, 0.1 mol of NH3, 0.5 mol of H2O, 0.2 mol of CaCO3.C7.22

Calculating the volume of a given mass of gas and vice versaThe volume of one mole of any gas is a constant known as the molar volume.At room temperature and pressure it is 24dm3/mol.amount of gas molecules = volume of gas/molar volumee.g.

What is the mass of 6dm3 of hydrogen H2 at room temperature and pressure if the molar volume under these conditions is 24dm3/mol?

H=1 so Molar mass of H2 = 1*2 = 2g/molamount of H2 = volume of H2/molar volume =6dm3/24dm3/mol =0.25molmass of H2 =amount of H2 * molar mass of H2 = 0.25mol*2g/mol = 0.5g

What is the mass of: 24dm3 of He, 12dm3 of N2, 6dm3 of CO2, 4dm3 of O2, 3dm3 of F2, 48dm3 of SO2, 2dm3 of H2, 1dm3 of H2S.What is the volume of 1g of He, 2.8g of N2, 22g of CO2, 64g of O2, 19g of F2, 64g SO2, 0.5g H2, 0.34g of H2S. (assume that the molar volume is 24dm3 at room temp)

Calculating reacting masses of substances or volumes of gasesA balanced chemical equation shows the amounts which react so masses or volumes of gases can be worked out from an equation. E.g. What mass of aluminium oxide can be made from 216g of aluminium and what volume of oxygen is needed?

Method 1

4Al(s) + 3O2(g) ---> 2Al2O3(s)Al = 27, so relative formula mass of 4Al = 4*27 =108O = 16 so relative formula mass of 2Al2O3 = 2(27*2 + 16*3) = 204so 108g of aluminium forms 204g of aluminium oxideso 1g of aluminium forms 204/108g of aluminium oxideso 216g of aluminium forms 216*204/108g of aluminium oxideso 216g of aluminium forms 408g of aluminium oxide

Method 2

4Al(s) + 3O2(g) ---> 2Al2O3(s)mass of Al = 216gamount of Al = mass/molar mass = 216g/27g/mol = 8molfrom equation: amount Al2O3/amount Al =2/4amount Al2O3 =amount of Al * 2/4= 8 mol *2/4 =4molmass Al2O3 = amount*molar mass =4mol*204gmol =408gFrom equation amount of O2/amount of Al =3/4amount of O2 =amount of Al*3/4 = 8mol *3/4mol = 6molmolar mass of O2 = 32g/molmass of O2 = amount of O2 * molar mass of O2mass of O2 = 6mol *32g/mol = 192gFor the reaction N2(g) +3H2(g) ---> 2NH3(g)

Calculate the masses and volumes of nitrogen and hydrogen needed to make 17g of ammonia NH3.

Converting mass-concentration into mol dm-3 and vice versaconcentration = mass/volume OR concentration = amount/volume

To convert them just change masses into amounts or vice versa.e.g. What is the concentration in mol/dm3 of a solution of sodium hydroxide NaOH of concentration 4g/dm3?

amount of NaOH in 1 dm3 = mass of NaOH/molar mass of NaOHamount of NaOH in 1 dm3 = 4g/40g/mol = 0.1mol so concentration of NaOH is 0.1mol/dm3 = 0.1MNB a concentration of 1M = 1mol/dm3.

Some dilute sulphuric acid, H2SO4, had a concentration of 4.90gdm-3. What is its concentration in mol dm-3?2. What is the concentration in gdm-3 of some potassium hydroxide, KOH, solution with a concentration of 0.200 mol dm-3?3.
What mass of sodium carbonate, Na2CO3, would be dissolved in 100cm3 of solution in order to get a concentration of 0.100 mol dm-3?

Simple calculations from the results of titrationse.g.

What is the concentration of a solution of sodium hydroxide NaOH if 10cm3 of the NaOH require 20cm3 of a 0.5M solution of sulphuric acid H2SO4 for neutralisation in a titration.

amount of H2SO4 = concentration of H2SO4 * volume of H2SO4amount of H2SO4 = 0.5mol/dm3 *20/1000dm3 = 0.01molH2SO4 + 2NaOH ---> Na2SO4(aq) + 2H2O(l)so amount of NaOH/amount of H2SO4=2/1so amount of NaOH = amount of H2SO4*2/1 = 0.01mol *2/1 = 0.02molconcentration of NaOH = amount of NaOH/volume of NaOH concentration of NaOH = 0.02mol/20/1000dm3 = 1mol/dm3NB 1cm3 = 1/1000dm3

What is the concentration of hydrochloric acid, 25.0cm3 of which neutralise 20.0cm3 of sodium hydroxide solution of concentration 0.15moldm-3. (ans = 0.15M)2.

What is the concentration of sulphuric acid, 20.0cm3 of which neutralise 30.0cm3 of a potassium hydroxide solution of concentration 0.1moldm-3 (ans=0.1M)3.

What is the concentration of sodium hydroxide, 10.0cm3 of which neutralise 15.0cm3 of hydrochloric acid solution of concentration 2.5moldm-3? (ans=2.5M)4.

What is the concentration of nitric acid, 10.0cm3 of which react with 25.0cm3 of a solution of sodium carbonate of concentration 0.2moldm-3? (ans = 0.2M)

Hard water- GCSE CHEMISTRY

2008-11-27T06:29:00.000-08:00

Hard water
definition of hard water
Soft water e.g. distilled water, easily forms a lather with soap.
Hard water e.g. London tap water, does not easily form a lather and forms a scum.
Soaps and soapless detergentsDetergents are chemicals with large molecules which help clean.
They have one end which dissolves in oil and one end which is ionic and dissolves in water.
oil soluble end C17H35COO-Na+ water soluble ionic end Hard water contains soluble calcium and/or magnesium salts. It contains calcium ions Ca2+(aq), which can be detected by a flame test or by testing with sodium hydroxide and/or magnesium ions Mg2+(aq).A soap has ions which react with ions in hard water to form a precipitate (a scum).
e.g. sodium stearate C17H35COO-Na+ is a soap.
stearate ion + calcium ion ---> calcium stearate
Calcium stearate is insoluble so forms a precipitate (a scum) in hard water.

2C17H35COO-(aq) + Ca2+(aq) ---> (C17H35COO)2Ca(s)
A soapless detergent has ions which do not react with the ions in hard water.e.g. sodium 3-dodecylbenzene sulfonate
C18H29SO3-Na+ is a soapless detergent.
C18H29SO3-Na+ and (C18H29SO3)2Ca are both soluble so no scum forms in hard water.
Draw labelled diagrams to show beakers with the ions and molecules that they contain for
(a) soft water, hard water,
soft water + soap,
soft water + soapless detergent, hard water + soap,
hard water + soapless detergent.
Limestone, chalk and gypsum and hard water
Hard water forms when calcium or magnesium salts in rock dissolve in rain water as it flows through the rock.
Gypsum is a rock containing calcium sulfate.
Calcium sulfate is insoluble but a little does dissolve to leave some calcium ions in the water where is has past over gypsum.Limestone and chalk are rocks containing calcium carbonate.
Calcium carbonate is very insoluble and none of it dissolves as water passes over limestone or chalk.
Calcium ions from limestone and chalk do dissolve to make water hard in a reaction with carbonic acid.
water + carbon dioxide ---> carbonic acid
H2O(l) + CO2(g) ---> H2CO3(aq)
calcium carbonate + carbonic acid ---> calcium hydrogencarbonate
CaCO3(s) + H2CO3(aq) ---> Ca(HCO3)2(aq)does not exist!

Industrial and domestic problems caused by scaleHard water normally contains dissolved calcium hydrogencarbonate. This can slowly change back into insoluble calcium carbonate but the change is speeded up by heating.
The solid calcium carbonate formed is called scale.
calcium hydrogencarbonate ---> calcium carbonate + carbon dioxide +water

Ca(HCO3)2(aq)) ---> CaCO3(s) + H2O(l) + CO2(g)
scale
Scale forms on a kettle's heating element insulating it and wasting energy when it boils.Scale forms inside hot water pipes which can even block a pipe.
Stalactites and stalagmites are made of calcium carbonate and form in caves in hard water areas. Explain how this might happen.
Benefits of hardness in water
The formation of thin layers of scale on the inside of pipes can be useful.Corrosion is reduced because the water in the pipes is not in contact with the metal.Poisonous metal salts of lead or copper from the metal surface of pipes cannot enter drinking water if the metal surface is covered in scale.Calcium is needed in the diet for healthy bones and teeth.
The treatment of water to make it suitable for domestic use

Collection of gases-IGCSE /GCSE CHEMISTRY

2008-11-27T06:25:00.000-08:00

Collection and identification of gases

Methods for collecting gasesdownward deliverysuitable for heavy dense gas solubility not important e.g. chlorine, upward deliverysuitable for light low density gas, solubility not important e.g hydrogen

collection over waterSuitable for any gas with a low solubility e.g. nitrogencollection using a gas syringeSuitable for any gas especially if volume is to be measured.Learning activity - use the data in the table and the rules above to select a collection method and give a reason.

Hazards with common gasesHydrogen - flammable, explosive when mixed with air. Use in small amounts.Hydrogen chloride and sulphur dioxide- poisonous as they are very acidic. Use small amounts and only in the fume cupboard.

Titration-IGCSE /GCSE-CHEMISTRY

2008-11-27T06:16:00.000-08:00

Titration
Reasons for using titrationWhen an acid and an alkali react it is not possible to see the end of the reaction unless an indicator is used. If an indicator is used the salt made will be impure. Titration is used to find out exactly what volume of acid and alkali should be mixed.

Methods for carrying out titration

1.Fill a 25cm3 pipette up to the line with the alkali solution of known concentration.
2. Transfer exactly 25.0cm3 of alkali to a conical flask and add an indicator.
3. Rinse a burette with acid then fill it with the acid using a funnel.Burette
4. Carefully run acid from the burette int the conical flask until the indicator changes colour.

method for salt preparation-IGCSE CHEMISTRY

2008-11-27T06:13:00.000-08:00

Choosing a method for salt

preparationSalts can be made by the following methods:

direct combination not normally used in school labs

e.g. iron + sulphur ----> iron sulphide

adding a carbonate to an

acidcarbonate + acid ---> salt + water + carbon dioxide

calcium carbonate + hydrochloric acid ---> callcium chloride + water + carbon dioxide

normally used to make soluble salts from insoluble carbonate, the reaction is complete (the acid is neutralised) when the effervescence (fizzing) finishes and some undissolved solid remains, filter, crystallise, filter again, wash and dry to obtain saltadding a metal to an

acidmetal + acid ---> salt + hydrogeniron + sulfuric acid ---> iron II sulfate + hydrogen

normally used to make soluble salts, the reaction is complete (the acid is neutralised) when the effervescence (fizzing) finishes and some undissolved solid remains, filter and crystallise, filter again, wash and dry to obtain saltadding an base to an acidacid + base ---> salt + water

normally used to make soluble salts from an insoluble base (most oxides and hydroxides),nitric acid + copper oxide ---> copper nitrate + waterthe reaction is complete (the acid is neutralised) when some undissolved solid remains, filter and crystallise, filter again, wash and dry to obtain saltIf a solution of a soluble base (an alkali e.g. soluble hydroxide) is used titration is carried out first to find out how alkali to add, crystallise, filter, wash and dry to obtain saltlearning activity -
Forming precipitatesprecipitationsoluble salt1 + soluble salt2 ----> insoluble salt + soluble salt3
sodium chloride + lead nitrate ----> lead chloride + sodium nitrateused to make insoluble salts from two solutions of soluble saltslearning activity - choose suitable pairs of solutions to make the following insoluble salts:calcium carbonate, barium sulfate, silver chloride, zinc carbonate, and lead chloride.

Purification of insoluble saltsWhen made an insoluble salt made is filtered, washed with distilled water and dried.

The Earth and its atmosphere-GCSE CHEMISTRY

2008-11-27T06:11:00.000-08:00

The Earth and its atmosphere
The composition of the atmosphereThe atmosphere is made up of the following gases, Nitrogen 78%, Oxygen 21%, Argon 1%, Carbon dioxide 0.03%, water - variable
Earth's early atmosphere and volcanoesThe primary atmosphere of the Earth was hydrogen and helium. These light gases were slowly lost. They were replaced by a secondary atmosphere produced by the action of volcanoes.
Composition of the Earth's early atmosphereThe Earth's secondary atmosphere was made up of some left over hydrogen, carbon dioxide, water vapour, nitrogen, carbon monoxide, sulphur dioxide ammonia and methane.Task C6.10 Match these formulae to secondary atmosphere gases: CH4, SO2, CO2, H2, NH3, CO, N2, H2O.
Origin of the oceansAs the Earth cooled to below 100oC oceans were formed when water vapour condensed and formed liquid water. Oceans are reservoirs for carbon dioxide because they can store the gas when it dissolves in them. The new oceans dissolved a great deal of the carbon dioxide in the atmosphere.

The oceans still play a part in keeping carbon dioxide levels constant. If there is a lot of carbon dioxide in the air then more can dissolve. If there is less carbon dioxide in the air then some comes out of solution back into the air.

The release of oxygen into the atmosphereAs the temperature of the Earth cooled simple green plants evolved in the oceans to use the carbon dioxide in the environment. These green plants steadily removed carbon dioxide and produced oxygen by photosynthesis. Oxygen levels in the atmosphere slowly increased.

The carbon cycleThe carbon cycle helps to keep the atmospheric composition constant by adding carbon dioxide to the atmosphere and also taking carbon dioxide away from the atmosphere. Carbon dioxide is taken away from the atmosphere or out of the cycle by photosynthesis, dissolving in water and by chemical reactions, for example with rock. It is brought into the atmosphere or into the cycle by respiration, combustion, volcanic activity and decay.

Formation of igneous rocksIgneous rocks are formed when magma pushes up into the crust and cools. It is made up of crystals It does not contain any fossils. Any living thing falling into the molten rock, from which it is made would be burnt and leave no trace. Igneous rock forms as magma cools slowly under the surface e.g. Granite. Magma reaching the surface through a volcano cools quickly e.g. basalt.

Crystal size and igneous rockIgneous rocks which cool slowly have large crystals e.g. granite but rock forming quickly has smaller crystals e.g. balsalt.

The formation of sedimentary rockThis rock is formed in shallow seas. After long periods of time sediment layers pile up and the lower ones come under great pressure.
This pressure pushes the water out of the layers or sediments and salt crystallizes and sticks the particles together to form sedimentary rocks. This process is called lithification. Living things falling into the sediments leave an impression as the rock forms (a fossil). Fossils show that a rock was made from sediments.

Manufacture of ammonia and fertilisers-GCSE /IGCSE CHEMISTRY

2008-11-27T06:07:00.000-08:00

Manufacture of ammonia and fertilisers

The chemical reaction used to make ammonia

Conditions used in the Haber processThe optimum (best) conditions for the Haber process that turns nitrogen and hydrogen into ammonia are:

350 atmospheres; high pressure increases yield

about 450ºC ; high temperature cuts yield but increases rate

and the use of a catalyst, which is usually iron; increases rate

ammonia is:nitrogen + hydrogen --> ammoniaN2 + 3H2 ---> 2NH3Ammonia can also decompose to form nitrogen and hydrogen:2NH3 ---> N2 +3H2

This reaction is reversible. The reaction is never complete but does reach a state when no more change can be seen.
This state is called equilibrium.Although no change is seen at equilibrium the reaction still carries on with some ammonia molecules being made and some decomposed. This is called dynamic equilibrium.

Temperature, pressure and the position of dynamic equilibriumTemperatureThe reaction below is exothermic, energy is given out when ammonia forms, but energy is taken in if ammonia
breaks up.
N2(g) + 3H2(g) ---> 2NH3(g) ; DH = - 92kJ/mol

Reactions resist changes.

If the temperature goes up the reaction tries to prevent this by taking in energy.
The reaction can take in energy by breaking up ammonia.
The position of the equilibrium moves to the left.So if the temperature goes up then ammonia breaks up which is not helpful.

If the temperature goes down the reaction tries to prevent this by giving out energy.The reaction can give out energy by forming ammonia. The position of the equilibrium moves to the right.So if the temperature goes down then ammonia form which is helpful.PressureOn the left hand side of the equation there are 4 moleculesOn the right hand side there are only 2 molecules which take up less space than 4 moleculesPressure is reduced if the position of the equilibrium moves to the right, so an increase in pressure causes a shift to the right so more ammonia is formed.

As high pressure favours a big yield of ammonia so 200 atmospheres pressure is used.

State and explain the effect of

(a) increasing the temperature and

(b) decreasing the pressure on the following reactions:2NO(g) + O2(g) = 2NO2(g) ; DH = +57kJ/mol2H2(g) + O2(g) = 2H2O(g) ; DH = - 280kJ/molCH4(g) + H2O(g) = CO(g) + 3H2(g) ; DH = + 40kJ/mol

Reaction rates and equilibriumUsing a high pressure gives a small increase in rate because the gases are more concentrated. Using a low temperature gives a low rate so it takes a long time to reach equilibrium.
A catalyst like iron in the Haber process is needed to speed up the reaction.

Neutralising ammonia with nitric acidWhen nitric acid reacts with ammonia (an alkali) the acid is neutralised and a salt is formed.acid + alkali ---> salt + waternitric acid + ammonium hydroxide ---> ammonium nitrate + waterHNO3(aq) + NH4OH(aq) ---> NH4NO3(aq) + H2O(l)

Ammonium nitrate contains fixed (chemically combined) nitrogen and so is a good fertiliser.

Write equations for the neutralisation of ammonium hydroxide, potassium hydroxide and ammonia by nitric acid HNO3, sulphuric acid H2SO4, phosphoric acid H3PO4.

Nitrogenous fertilisers and plant growthPlants grow well when they can obtain fixed nitrogen from the soil.
Only a few plants like peas and beans can make use of nitrogen in the air. Most plants need fixed nitrogen in compounds like nitrates. Plants need nitrogen to make proteins which gives them strong stems and healthy leaves. Some nitrates find their way into the soil naturally but intensive farming removes a lot when crops are harvested. Fertilisers containing nitrogen are used to replace nitrogen lost from the soil during farming.

The leaching of artificial fertilisers
1. Fertilisers are very soluble.
2. Fertilisers dissolve in rain water.
3. Fertilisers are leached from the soil and washed into rivers.
4. Water plants grow very well in fertilised river water.
5. The over growth of plants like algae at the surface cuts out the light to plants below.
6. Plants without light stop growing and die.
7. Dead plants rot due to bacteria that use a lot of oxygen
8 The amount of oxygen in the water drops.
9 Fish and other animals start to die because of a lack of oxygen.
10 The process is called eutrophication.

energetics- IGCSE / GCSE

2008-11-27T06:06:00.000-08:00

Energy changes accompanying reactions

Temperature changes in reactionsMany reactions give off heat such as the burning of wood which causes a temperature increase. Other reactions take in energy and cause a temperature fall.

Exothermic reactionsAn Exothermic reaction is one which gives out energy to the surroundings, usually in the form of heat and usually shown by a rise in temperature. An example of an exothermic process is the burning of fuels.

Endothermic reactionsAn Endothermic reaction is one which takes in energy from the surroundings, usually in the form of heat and usually shown by a fall in temperature. An example of an endothermic is photosynthesis. This is because it takes in energy from the sun in the form of light.

Bond breaking and making in reactionsDuring a chemical reaction, old bonds are broken and new ones formed. Energy must be supplied to break existing bonds and so this is endothermic. Energy is released when new bonds are formed and so the formation of bonds is exothermic.Hydrogen reacts with oxygen to form water.

O=O + H-H + H-H --->O +O + H +H + H +H --->H-O-H + H-O-H Old bonds break single atoms with new bonds form in oxygen and no bonds in water molecules hydrogen molecules exothermicendothermic

Rates of reaction-igcse notes

2008-11-27T05:48:00.000-08:00

Rates of reaction

Fast and slow reactionsSome chemical reactions are slow like the rusting of iron.Some chemical reactions are faster like the burning of wood.Some chemical reactions are very fast like the explosion of gunpowder.

State the rate for following reactions; TNT exploding, petrol burning, rock reacting with water, copper roof turning green, dynamite being used in a quarry, bread baking.

Experiments to investigate ratesConcentration-

Carry out reactions between sodium thiosulphate and hydrochloric acid in a conical flask using identical volumes and identical temperatures but at different concentrations. Record the concentrations and the times needed for a cross under the flask to disappear.

Temperature- Carry out the reaction between magnesium and hydrochloric acid in a conical flask for different acid temperatures but for identical volumes of acid, at identical concentrations and for identical masses of magnesium. Record the time for the magnesium to completely react.

Particle size-

Carry out the reaction between calcium carbonate and hydrochloric acid in a conical flask fitted with a stopper and a delivery tube to a measuring cylinder inverted in water. Use the same masses of powder, small chips and large lumps of calcium carbonate with the same volumes, concentrations and temperatures of acid. Record the volume of gas formed every 30 seconds for 10 minutes.

Interpreting experimental resultsTemperature From this graph between sodium thiosulphate and hydrochloric acid you can see that the rate of reaction is fastest at highest temperatures

ConcentrationThe rate increases as the concentration increases.

Particle sizeThe rate gets bigger as the pieces get smaller.The rate gets bigger as the surface area gets bigger.

Explaining the effect of changing temperature, concentration and surface area

(i)Concentration

The concentration of a substance, normally a solution, is the amount in a given volume. concentration = amount {units = mol/dm3 or M} volume In a higher concentration solution there are more particles to react therefore there are more collisions.

As a reaction depends on collisions happening, more collisions lead to a faster reaction rate.

If we were doing a reaction with acid and we double the number of acid particles, we double the number of collisions and therefore are likely to double the reaction rate.

Draw diagrams to show particles in a low concentration HCl solution and particles in a high concentration HCl solution. Show water molecules and HCl and magnesium particles in the diagrams. Show different numbers of collisions in each diagram.

(ii)Temperature

Increasing the temperature increases the speed of the reacting particles and faster particles collide more often than slow ones.

The increase in the number of collisions leads to an increase in the rate of reaction. Increasing the temperature also gives the particles more energy so that they collide with more violence. Energetic particles have a better chance of their collisions leading to a reaction.

Draw diagrams to show water molecules and HCl and magnesium particles in two diagrams at two different temperatures. Show different numbers of collisions and different energies for the collisions.

(iii) Particle size

Particle size is all to do with surface area. Powder has a higher surface area than lumps and therefore powder makes more collisions possible than lumps. This simple diagram explains the idea of surface area clearly:

Collision theoryReactions can only occur when particles collide and if you increase the frequency and/or energy of the collisions, you increase the rate of reaction.
If you increase the temperature you give the particles more energy so that they move quicker and have more collisions
Increasing the concentration means more frequent collisions as there are more particles in a certain volume

If you cut a lump into powder you give each individual particle a larger surface area, this gives an increased area in which the acid, for example, particles can react with.

few collisions in 1 second (small surface area, low concentration, low temperature) = slow reactionmany collisions in 1 second (large surface area, high concentration, high temperature) = fast reactionlow energy collisions (low temperature) = slow reactionhigh energy collisions (high temperature) = fast reaction

Describe what would happen on a motorway if cars were like particles in a chemical reaction and if

(a) there were more cars on the road,

(b) the cars travelled faster. Link the description to collision theory by using the words concentration and temperature in your answer.

The effects of catalysts

A catalyst is a substance that speeds up the rate of a reaction without being used up. Catalysts are usually transition metals or transition metal compounds. An example of a catalyst is iron which catalyses the reaction of nitrogen and hydrogen to produce ammonia. A catalyst usually works either by providing a surface for the reaction to take place or by forming intermediate compounds

Match the (catalysts) with the reactions that they affect.(iron, manganese dioxide, nickel, platinum, vanadium pentoxide)hydrogen peroxide ---> water + oxygenCarbon monoxide + oxygen ----> carbon dioxidesulfur dioxide + oxygen ---> sulfur trioxidenitrogen + hydrogen ----> ammoniavegetable oil + hydrogen ---> margarine

EnzymesEnzymes are biological catalysts. They speed up reactions. They work best around body heat (37oC). They break up (denature) and stop working if the temperature rises much above 40oC. Enzymes are found in biological washing powder. Enzymes in yeast are used to help fermentation in making bread and beer.

Useful products from crude oil- igcse

2008-11-27T05:39:00.000-08:00

Useful products from crude oil

HydrocarbonsHydrocarbons are compounds which contain only the elements carbon and hydrogen. Crude oil is a mixture of different sized hydrocarbon molecules. These hydrocarbons are basically fuels such as petrol.Task C4.11 Pick out the hydrocarbons in the following list CH4, CH3OH, C6H6, C2H6, H2O, C8H18, C4H8, C2H3N, C5H12, C6H14.

Fractional distillation of crude oilThe fractional distillation of crude oil is the process which gives us the different substances (fractions) made from crude oil. The crude oil is pumped in to the fractional column from the bottom. The heat is applied at the bottom of the fractionating column. The different fractions are obtained in different positions in the column.

Top 70oC, small molecules, light colour, runny, easy to light

Gases
Petrol
Naptha
Kerosene
Diesel
Lubricating oil
Fuel oil
BitumenBottom 360oC,

large molecules, dark colour, viscous, hard to light

Draw diagrams to show the relative sizes of molecules of gas, diesel and bitumen.

The size of molecules and boiling pointAs the size of a hydrocarbon molecule increases the boiling point increases. If it has a low boiling point it is very volatile (forms a vapour easily). If it has a high boiling point it is not volatile.

size of hydrocarbon molecule (carbon atoms)

Boiling point/oC

4-12 carbon

40

11-15carbon

180

15-19 carbon

260

Over 50 carbon

Over 340

Uses of fractions from crude oil

Fraction

Use

gases

Bottled gas for gas cookers, boilers, camping gas

Petrol

Cars, electricity generators

Kerosine

Jet fuel

Diesel oil

Trucks, and some cars

oil

lubrication

Fuel oil

Boilers in ships or buildings

Bitumen

Covering road surfaces

Butane gas for camping

Complete and incomplete combustionComplete combustionHappens with plenty of oxygen. All of the carbon and hydrogen in a hydrocarbon turns to carbon dioxide and water.

Hydrocarbon + oxygen carbon dioxide + water

e.g. CH4 + 2O2 --> CO2 + 2H2O

carbon dioxide

Not poisonous but build up in atmosphere. carbon dioxide molecules trap energy from the sun. This leads to global warming.

water
harmless product
Sulphur Dioxide

Produced by sulphur impurities in burning fossil fuels. Can be dangerous if inhaled. Also it can dissolve into clouds to form acid rain. When this precipitates it is harmful to the environment by killing fish in lakes, damaging forests and can corrode buildings and metals structures.
Nitrogen Dioxide

This is formed in car engines when oxygen and nitrogen combine. This is an acid gas and turns into nitric acid when it dissolves. Acid rain results.

Complete the following equations for complete combustion of hydrocarbonsethene + oxygen ---> C2H4 + 3O2 --->ethyne + oxygen --->2C2H2 + 5O2 ---> propane +oxygen --->C3H8 +5O2 --->Incomplete combustionHappens if there is not enough oxygen.

The hydrocarbon turns into soot (carbon) and poisonous carbon monoxide as well as carbon dioxide.

Hydrocarbon + oxygen ---> carbon monoxide + water 2CH4 + 3O2 ---> 2CO + 4H2O
carbon

Formed by incomplete combustion as soot. It is bad for the lungs and disfigures buildings.
Carbon Monoxide

This is produced by traffic and some gas fires. When inhaled by people it replaces oxygen in the haemoglobin, however, it does not give it up, this eventually leads to suffocation.

Complete the following equations for incomplete combustion of hydrocarbonsethyne +oxygen --->

Match these combustion products to their formulae: carbon, water, C, sulfur dioxide, CO2, nitrogen dioxide, NO2, carbon dioxide, SO2, carbon monoxide, H2O, CO.

Chemical tests for carbon dioxide and waterIf carbon dioxide is bubbled through limewater it turns milky. We can test for the presence of water using anhydrous copper sulphate which is white. The copper sulphate turns blue when water is added to it. Cobalt chloride can also be used. This turns from blue to pink when water is added.Task C4.16 Show the above information in a suitable table.

CrackingCracking is the splitting up of long chain hydrocarbons in to smaller chains. Cracking is a form of thermal decomposition. A lot of longer molecules produced from fractional distillation are cracked into smaller ones because there's more demand for products like petrol and Kerosine than for diesel oil. The products include compounds with double bonds such as ethene. For example kerosine could be broken down to octane and ethene.

C-C-C-C-C-C-C-C-C-C- ----cracking-----> C=C + C=C-C + C-C=C + C-Cbig molecules small moleculealkanes mostly alkenessingle bonds double bondssaturated unsaturatedonly good for fuels can be made into polymers

Industrial conditions for crackingIndustrial conditions for cracking vaporised hydrocarbons are to use a powdered catalyst at about 400ºC to 700ºC. The catalyst could be Aluminium oxide.

The products of crackingCracking turns big molecules into small molecules which are mostly alkenes. Some alkanes are also made.

Saturated and unsaturated hydrocarbonsAlkanes are saturated hydrocarbons. They have single C-C bonds only.Alkenes are unsaturated hydrocarbons. They have double C=C bonds.Saturated hydrocarbons are so called because their C atoms have no spare bonds left to join with any more hydrogen atoms. They contain C-C single bonds and each C atom is joined to 4 other atoms. Unsaturated hydrocarbons are so called because they have some spare bonds which could be used to add on some more hydrogen atoms.

Natural gasNatural gas is mostly made up of an alkane called methane CH4.

The formulae and structures of alkanesMethane CH4 H H-C-H H
Ethane C2H6 H H H-C-C-H H H
Propane C3H8 H H H H-C-C-C-H H H H
Butane C4H10 H H H H H-C-C-C-C-H H H H H

Formulae and Structures of alkenes
Ethene C2H4 H H C=C H H
Propene C3H6 H H H H-C-C=C-H H C4.24 Testing alkanes and alkenes with bromine water

Making polymers from small moleculesPolymers like polythene, polypropene and polystyrene are large molecules, which can be formed by combinations of many smaller molecules called monomers.

Addition polymerslots of monomers add together ---> 1 addition polymerEach has C=C bonds only single bondsunsaturated saturatedlots of ethene molecule ---> 1 polyethene moleculeH H H H H H H H H H H H H H C=C + C=C + C=C + .... ---> -(C- C- C- C-C- C- C- C)-n H H H H H H H H H H H H H H

Draw structures and equations as above for the formation of polypropene and polychloroethene.

Uses and properties of polymers
Material
Use
Properties
Poly(ethene)
Plastic bags, bottles, buckets and bowls
Softens when warm so easily formed and moulded
Poly(propene)
Crates, rope, carpets, car bumpers, fishing nets
tough and strong
poly(chloroethene)polyvinyl chloridePVC
gutter, drain pipes, window frames, covering for electrical wiring.
electrical and heat insulator, tough, not easliy decomposed by sunlight

Extraction and uses of metals-igcse

2008-11-27T05:38:00.000-08:00

Extraction and uses of metals

Oxidation and reduction
Oxidation is a reaction where oxygen is gained. eg iron + oxygen ---> iron oxide
Reduction is a reaction where oxygen is lost. Eg iron oxide is reduced to form ir

Classify the following reactions as oxidation or reduction: burning magnesium, making aluminium from aluminium oxide, burning carbon in air,
copper oxide + carbon --> copper + carbon dioxide
carbon monoxide forming carbon dioxide.

Electron loss and gain

Oxidation is also a loss of electrons eg Fe --> Fe2+ + 2e-
Reduction also the gain of electrons eg (when Fe ions becomes Fe atoms)

Identify all of the elements below as being oxidised or reduced in the following reactions:
Cu2+ + 2e- ---> Cu
lead ions forming lead atoms
Na ---> Na+ + e-
Ca + Cl2 --> CaCl2
Fe2+ ---> Fe3+ + e-

Reduction of metal ores

An ore is a material found in the ground which contains a metal. An ore is often a metal oxide mixed with rock . When a metal is extracted its ore is reduced. Metal loses oxygen from its oxide.

Extraction and uses of metals

Oxidation and reduction

Oxidation is a reaction where oxygen is gained. eg iron + oxygen ---> iron oxideReduction is a reaction where oxygen is lost.

Eg iron oxide is reduced to form iron.

Classify the following reactions as oxidation or reduction: burning magnesium, making aluminium from aluminium oxide, burning carbon in air,

copper oxide + carbon --> copper + carbon dioxide

carbon monoxide forming carbon dioxide.

Electron loss and gainOxidation is also a loss of electrons eg Fe --> Fe2+ + 2e- Reduction also the gain of electrons eg (when Fe ions becomes Fe atoms)

Identify all of the elements below as being oxidised or reduced in the following reactions:

Cu2+ + 2e- ---> Cu lead ions forming lead atoms

Na ---> Na+ + e-Ca + Cl2 --> CaCl2Fe2+ ---> Fe3+ + e-

Reduction of metal oresAn ore is a material found in the ground which contains a metal. An ore is often a metal oxide mixed with rock . When a metal is extracted its ore is reduced. Metal loses oxygen from its oxide.

Examples include:
haematite which is mostly iron III oxide Fe2O3

bauxite which is mostly aluminium oxide Al2O3

malachite which is mostly copper carbonate CuCO3

Extraction and position in reactivity seriesThe method of extraction of a metal depends on the reactivity of the metal. Unreactive metals like silver can be found uncombined. More reactive metals like iron can be reduced with carbon because carbon is more reactive than iron. Very reactive metals like aluminium can only be extracted by electrolysis. Carbon is not used as it is less reactive than aluminium.

The cost of electricity for extracting metalsElectrolysis is an extraction method that uses a lot of electrical energy. Electrical energy is more expensive than energy from burning carbon.

The extraction of aluminiumAluminium is found in the ground in an ore called bauxite. Bauxite is aluminium oxide (Al2O3) with iron oxide impurities. After purification aluminium oxide is mixed with cryolite to lower the melting point from 2000º to 1000º, which saves money. This mixture is heated and the molten liquid used as the electrolyte.

Both electrodes are made of graphite (carbon). The anode (+ve) is graphite and the cathode (-ve) is a graphite lining to a steel case.

The carbon anodes react with oxygen so have to be replaced.C + O2 --> CO2

At cathode -

positive aluminium ions attracted, gain electrons and become atoms.

Al3+ + 3e- ---> Al

At anode - negative oxide ions attracted, lose electrons and become atoms.

2O2- ----> O2 + 4e-

Uses and properties of aluminium

overhead power cables
good electrical conductor, low density
drinks cans
Does not react with water
aircraft parts
high strength and low density

Carbon and carbon monoxide for reducing oxidesCarbon and carbon monoxide can both remove oxygen from other compounds so are good for reduction. They are used to reduce the ores of metals below carbon in the reactivity series. E.g. zinc, iron, tin and lead.

Iron extraction using the blast furnace

A blast furnace is used in the process of extracting iron.

The raw materials iron ore, coke and limestone are put in at the top. Hot air is blasted into this furnace at the bottom making the coke (carbon) burn faster and the temperature rises to about 1500º.

When the coke burns, carbon dioxide is produced.

C + O2 ---> CO2

CO2 reacts with the unburnt coke to form carbon monoxide

CO CO2 + C ---> 2CO

Iron oxide Fe2O3 in the ore is reduced to iron by the reaction with the carbon monoxide.

3CO + Fe2O3 ---> 3CO2 + 2Fe

Molten iron is a dense liquid, so runs to the bottom of the furnace and is tapped off.

Limestone CaCO3 helps remove impurities during the extraction by forming calcium oxide CaO.

CaCO3 ---> CaO + CO2 The rock impurities silicon dioxide SiO2 are then removed by the following reaction.

CaO + SiO2 ---> CaSiO3 CaSiO3 is known as slag and can be used in making cement and road building.

The purification of copper

Very pure copper is needed for copper wires. Electrolysis is needed to purify copper. The anode is a mass of impure copper and the cathode is pure copper. The electrolyte is sulphuric acid. The impurities drop at the anode as sludge during electrolysis.

At anode Cu ---> Cu2+ + 2e-

At cathode Cu2+ + 2e- ---> Cu

Chemical from calcium carbonate- gcse only

2008-11-27T05:36:00.000-08:00

Chemical from calcium carbonate

Thermal decomposition of calcium carbonatecalcium carbonate ---heat---> calcium oxide + carbon dioxideCaCO3(s) ---heat---> CaO(s) +CO2(g)(limestone) (quicklime)

Water and calcium oxidecalcium oxide + water ---> calcium hydroxideCaO(s) + H2O(l) ---> Ca(OH)2(s) (slaked lime)

looks flaky and gets hotWith more water the calcium hydroxide dissolves to make limewater Ca(OH)2(aq)

Neutralising soil acidityCalcium oxide and calcium hydroxide are bases. (like all metal oxides and hydroxides)acid + base ---> salt + waterso acid soil is neutralised by calcium oxide or calcium hydroxide

e.g. nitric acid + calcium oxide --> calcium nitrate + water 2HNO3(aq) + CaO(s) ---> Ca(NO3)2(aq) +H2O(l)

Farmers can add calcium oxide or calcium hydroxide to their acid soil to neutralise it.The neutral soil is better for growing plants.Task C3.25 Write word equations and balanced chemical equations for neutralising the following acids using calcium oxide and calcium hydroxide: nitric acid, sulphuric acid, hydrochloric acid, phosphoric acid (H3PO4)

Uses of calcium carbonateCement is made from calcium carbonate (limestone).

limestone --heat--> calcium oxide + carbon dioxideCalcium oxide + clay ----> cementGlass is made using calcium carbonate.

Calcium carbonate + sand + sodium carbonate ---heat--> glassIron is manufactured using calcium carbonate in the blast furnace.calcium carbonate --heat--> calcium oxide + carbon dioxideCalcium oxide + silicon dioxide ----> calcium silicate (mostly sand) (slag) (floats on top of molten iron)

Chemicals from salt- gcse/igcse

2008-11-27T05:32:00.000-08:00

The electrolysis of aqueous sodium chloride solution

When the electrolysis of aqueous sodium chloride takes place, hydrogen and chlorine are given off as gases and sodium hydroxide is left. Aqueous sodium chloride contains hydrogen ions H+ and hydroxide ions OH- (from the water) and sodium ions Na+ and chloride ions Cl-.

The positive sodium and hydrogen ions go to the cathode and the negative chloride and hydroxide ions go to the anode. Hydrogen is formed at the cathode and chlorine is formed at the anode.

substance test result

hydrogen

lighted splint squeaky pop

chlorine

damp blue litmus turns red then white

sodium hydroxide

damp red litmus turns blue/purple

Write a word and a balanced chemical equation for the reaction in which hydrogen burns.
Complete: Chlorine is (acidic/alkaline) because it turns litmus red.
Chlorine is a (dye/bleach) because it turns litmus white.
This type of reaction is (oxidation/reduction).

Uses of sodium chloride, hydrogen, chlorine and sodium hydroxide

sodium chloride preserve food, flavouring, stops ice on roads because salt lowers freezing point of water to -5oC,

manufacture of sodium, chlorine and sodium hydroxide

chlorine kills bacteria in swimming pools and drinking water, bleach, making hydrochloric acid and solvents.

hydrogen rocket fuel, making margarine, making ammonia
sodium hydroxide detergents, bleach, paper, fibres, purifying bauxite

Transition metals-igcse

2008-11-27T02:00:00.000-08:00

Transition metals

The colours of transition metal compoundsAll compounds of transition metals are coloured. Some examples are:copper sulphate - blue, iron III oxide - brown, cobalt chloride pink.

Transition metals as catalysts
Transition metal or compound can be used as Catalyst

manganese (IV) oxide

decomposing hydrogen peroxide

platinum

catalytic converter in car exhaust

nickel

making margarine from vegetable oil and hydrogen

vanadium (V) oxide

making sulphuric acid

platinum with rhodium

making nitric acid

iron

making ammonia

Uses and properties of titanium, iron and copper

iron

Bridge Construction/High Strength

iron

permanent magnets /Magnetic

copper

water pipes/does not react with water

copper

Electrical wires/Good conductor of electricity

titanium

propeller shafts on ships, hip joints/does not react with water

titanium

aircraft engines/high strength, low density, high mp

Halogens -igcse

2008-11-27T01:56:00.000-08:00

Physical properties of halogens

The colour of the Halogens changes from a lighter colour to a dark colour as we go down the group and the melting and boiling points increase as you go down the group.
Fluorine at the top of the group is a yellow gas

This changes to a yellow-green gas for chlorine, a red liquid for bromine and finally black solid for iodine at the bottom of the group.

Reactions of halogens with metals

Metal + Halogen --> Metal Halide Iron reacts slowly with iodine to form iron II iodide. Iron reacts faster with bromine to form iron III bromide.

Iron reacts with chlorine reacts even faster to form iron III chloride.

Other metals like sodium also react.sodium + chlorine ---> sodium chloride

The reaction of halogens with hydrogenHydrogen reacts with halogens to form hydrogen

halides. E.g.Hydrogen + chlorine ---> hydrogen chlorideH2(g) +Cl2(g) --> 2HCl

(g) (a hydrogen halide)

Hydrogen halides are very soluble in water.

The gas hydrogen chloride forms hydrochloric acid when it dissolves in water.

Other hydrogen halides like hydrogen bromide, HBr and hydrogen iodide, HI also dissolve to form acid solutions. An acid solution has a pH less than 7 e.g. 1. An acid turns universal indicator red.

Displacement reactionsHalogens like chlorine are very reactive and displace less reactive halogens like bromine from halides like bromides.

Potassium bromide + chlorine ---> potassium chloride + bromine

2KBr(aq) + Cl2(aq) --> 2KCl(aq) + Br2(aq)

2Br-(aq) + Cl2(aq) --> 2Cl-(aq) + Br2(aq)

Chlorine and bromine also react with potassium iodide showing that the order of reactivity is: Chlorine > bromine > iodine

Uses of halogens and halidesFluorine compounds (fluorides) are put into toothpaste and some drinking water supplies.

Fluorides join with tooth enamel and make teeth resist attack by acid which prevents tooth decay.

Chlorine is used to in swimming pools and drinking water to kill bacteria.

Iodine is used as an antiseptic because it will kill the germs on the skin without damaging it.

Noble gases uses and properties-igcse

2008-11-27T01:53:00.000-08:00

Argon /Light bulbs/Doesn’t react with the metal filament
Helium/Used with O2 for deep sea dives/Low solubility of helium in the blood.
Helium/To inflate the tyres of large aircraft/Non-flammable

Helium/To fill airships and weather balloons/Low density, does not burn
Neon/In advertising signs because it glows red when electricity passes/Conductor of electricity/at high voltage

Krypton/Xenon/In lamps used in photographic flash units, in stroboscopic lamps used in lighthouses/Gives out a lot of light when electricity passes through

Representing reactions- igcse notes

2008-11-27T01:35:00.000-08:00

If a reaction occurs between magnesium and oxygen, magnesium oxide is produced, here is the word equation for this reaction: -

magnesium + oxygen --> magnesium oxide

Some other examples are: hydrochloric acid + calcium carbonate --> calcium chloride+ carbon dioxide+ water

sodium + water --> sodium hydroxide + hydrogen

hydrochloric acid + sodium hydroxide --> sodium chloride + water

Write word equations for the reactions in which the following compounds form from a halogen and another suitable element: hydrogen fluoride, hydrogen chloride, iron III chloride, iron III bromide, sodium chloride, copper chloride.

Formulae

The formula of an element or compound is simply the symbol of each element present and numbers to show how many atoms are present.
Carbon dioxide has the formula CO2.
This means that it has one carbon atom and two oxygen atoms in each molecule.

substance-methane-formula -CH4

substance -bromine -formula -Br2

ethane -C2H6

hydrogen -H2

propane -C3H8

ethanol -C2H5OH

butane -C4H10

State the name and the number of atoms of each element in the formulae above. Formulae can be worked out from valency.

Atom or ion with that valency 1

hydrogen,
group 1 e.g. sodium and potassium, group 7
e.g. fluorine and chlorine, ammonium NH4+, hydroxide OH-, nitrate NO3-

Atom or ion with that valency2

group 2

e.g. magnesium and calcium, group 6, sulphate SO42-, carbonate CO32-, copper , II Cu2+ iron II Fe2+

Atom or ion with that valency3

group 3

e.g. aluminium, group 5, phosphate PO43-, iron III Fe3+

Atom or ion with that valency4

group 4

e.g. carbon

Use valency to work out the formulae of the following compounds: sodium chloride, potassium bromide, magnesium oxide, calcium sulphide, aluminium nitride, calcium iodide, lithium oxide, aluminium chloride, aluminium sulphide, magnesium nitride.

Calculating relative formula massAdd up the relative atomic mass (found in periodic table) of each atom in the compound.e.g. Al203 relative atomic masses of Al = 27, O = 16 (found in periodic table).

The formula shows 2 atoms of aluminium and 3 atoms of oxygen so:
formula mass of = (2*27) + (3*16) =54 + 48 = 102

Work out the relative formula masses of the following: MgO, FeS, O2, H2O, CaBr2, Na2S, CaCO3, NaOH, HCl, (NH4)2SO4.

Relative atomic masses Mg=24, O=16, Fe=56, S=32, Ca=40, Br=80, C=12, Na=23, H=1, Cl=35.5.

Simple balanced equations

It is possible to write balanced equations for reactions.

For example substances such as hydrogen and magnesium combine with oxygen.

One method to write them is:Write a word equation first.Magnesium + oxygen --> magnesium oxideWrite in the formulae of the substances used.Mg + O2 --> MgO

Balance the equation so that each element has the same number of atoms on each side.
2Mg + O2 --> 2MgO

sodium + oxygen --> sodium oxide

4Na{s} + O2{g} --> 2Na2O{s}

(word equation) hydrogen + oxygen ---> water

(formulae) H2 + O2 -----> H2O

(balance) 2H2 + O2 -----> 2H2O

State symbols

The state symbols are put in a balanced equation to show whether something is a solid, liquid, gas or dissolved in water (aqueous solution).

Magnesium + oxygen --> magnesium oxide

2Mg{s} + O2{g} --> 2MgO{s}

hydrochloric acid + calcium carbonate --> calcium chloride + carbon dioxide + water

2HCl (aq) + CaCO3 ---> CaCl2(aq) + CO2(aq) + H20(l)

Balanced equations and ionic equations

Ionic equations only show ions which change in a reaction and ignore those which do not change.

E.g.word equationhydrochloric acid + sodium hydroxide --> sodium chloride + water

balanced chemical equationHCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)ionic equationH+(aq) + OH-(aq) ---> H2O(l)

E.g. in the electrolysis of sodium chlorideCl-(aq) --> Cl(g) + e-2Cl(g) --> Cl2(g)

Working out formulae from reacting masses

elements reacting magnesium chlorine

symbols of elements Mg Cl

masses reacting (from experiment) 2.4g 7.1g

amounts (amount = mass/molar mass) 2.4g/24g/mol 7.1g/35.5g/mol = 0.1mol 0.2mol

ratio of atoms (divide by smallest) 1 : 2

formula MgCl2

Work out formulae of compounds formed when the following react:56g of iron and 32g of sulphur (Fe =56, S =32)2g of hydrogen and 16g of oxygen (H=1, O=16)14g of lithium and 16g of oxygen (Li=7)32g of copper and 8g of oxygen (Cu=64)6.4g of copper and 0.8g of oxygen

Calculating reacting masses using equations

You can work out ratio of the masses of products and reactants by simply multiplying the number of moles shown in the equation by the formula mass of each substance.

Example 1: What mass of magnesium oxide can be made from 12g of magnesium? Relative atomic masses are Mg =24, O = 16.

equation 2Mg(s) + O2(g) --> 2MgO(s) formula 2*24 1(16*2) 2(24+16)
masses =48 =32 =80

reacting 48g of Mg forms 80g of MgOmasses 1g of Mg forms 80/48 g of MgO 12g of Mg forms 12*80/48 g of MgO = 20g

Example 2: What mass of magnesium oxide can be made from 12g of magnesium?

Equation 2Mg(s) + O2(g) --> 2MgO(s)

amounts 2 moles 1 mole 2 moles masses 2*24 1{16*2} 2{24+16} =48g =32g =80g so 48g Mg forms 80g MgO 1g Mg forms 80/48 g MgO 12g Mg forms 12*80/48 g MgO = 20g

Also note that the ratio of amounts of reactants and products in the equation above can be written as:

Amount of Mg/amount of O2 =2/1 Or Amount of O2/amount of MgO = 1/2

chemicalbonding-igcse notes

2008-11-20T02:42:00.000-08:00

Elements forming compounds with chemical bonds

Electron transfer and ionic bondsAtoms have no charge.
A charged particle is called an ion.
If an atom loses an electron, it becomes a positively charged (+) ion.
An ion that is positively charged is known as a cation. If an atom gains an electron, it becomes a negatively charged (-) ion.
An ion that is negatively is known as an anion. The negative and positive ions attract each other to form an ionic bond.

Complete the gaps in the text below:_____ have no charge. A charged particle is called an ___. If an atom loses an ________, it becomes a positively charged (+) ion. An ion that is positively charged is known as a ______. If an atom gains an electron, it becomes a negatively charged (-) ion. An ___ that is negatively is known as an anion. The negative and positive ions attract each other to form an _____ bond.

The formation of sodium and chloride ion

Draw atoms and ions for lithium, potassium, fluorine, magnesium, oxygen, sulfur and aluminiuM

Draw diagrams of ionic bonding in LiF, KF, LiCl, NaF, MgCl2, AlF3, MgO, MgS, Na2O and Al2O3.

Physical properties of giant ionic structuresIonic bonds form when metal and non-metal atoms join. A substance with ionic bonding has an ionic structure. Each ion is firmly held in place by strong ionic bonds so they have high melting and boiling points. If melted, charged ions become free to carry an electric current. The ions also become free if dissolved in water so solutions are also electrolytes. The solids are insulators because the ions are not free to move and cannot carry a current. Sodium chloride NaCl, and magnesium oxide MgO are good examples.

Pick out the substances which are (a) ionic (b) have covalent bonds (c) have high melting points (d) conduct electricity when molten: sodium chloride, sulfur dioxide, magnesium oxide, iron fluoride, carbon dioxide, NaBr, H2O, NH3, Al2O3, KCl.

Covalent bonds and electron sharingNon-metal atoms join using covalent bonds. When a covalent bond is formed, atoms share their electrons. The atoms then have full shells. One covalent bond needs one shared electron from each atom. Each atom involved has to make enough covalent bonds to fill up its outer shell. Sharing electrons is called covalent bonding. Below is a diagram to show hydrogen gas (H2).

Dot and cross diagrams

Draw atoms of F, H, O, N and C.Draw a dot and cross diagram for fluorine F2, hydrogen fluoride HF, water H2O, ammonia NH3, methane CH4, oxygen O2, nitrogen N2, ethene C2H4.

Physical properties of simple molecular substancesSimple Molecular Substances have low melting and boiling points and most are gases or liquids at room temperature. This is because of weak forces between the molecules. Molecular substances do not conduct electricity, because there are no ions. E.g. water.Draw diagrams to show how the molecular structures for the following might look: fluorine F2, hydrogen fluoride HF, water H2O, ammonia NH3, methane CH4, oxygen O2, nitrogen N2

Giant structures with covalent bondsIn giant covalent structures all the atoms are bonded to each other by strong covalent bonds so they have very high melting and boiling points. They do not usually conduct electricity even if in the liquid state. Diamond and graphite are two examples, which are made from carbon atoms. These two different types of the same element are called allotropes.

Diamond: Each carbon atom forms four covalent bonds in a very rigid giant covalent structure.
Diamond:

Physical properties of giant covalent structuresGiant molecular structures have very high melting points because all atoms are held firmly in place by strong covalent bonds. In graphite each carbon atom is held in place by three strong covalent bonds which gives graphite a high melting point. In diamond 4 strong covalent bonds holds each atom in place. This also gives diamond a very high melting point. The four bonds make diamond very hard. Graphite has weak bonds between layers so the layers slip over each other making graphite soft.
They do not usually conduct electricity even when molten because there are no charged particles to carry the current. There are free electrons between layers in graphite so it conducts electricity.

Explaining differences between propertiesSimple molecular substances like water have weak bonds between molecules so melt at low temperatures because little energy is needed to separate the molecules. Giant covalent structures like diamond have strong covalent bonds holding each atom in place. They melt at high temperatures because a lot of energy is needed to break these strong bonds.

The Extraction of Aluminium-igcse

2008-11-06T09:48:00.000-08:00

Aluminium is obtained from mining the mineral bauxite.
The purified bauxite ore of aluminium oxide is continuously fed in. Cryolite is added to lower the melting point and dissolve the ore.

Ions must be free to move to the electrode connections called the cathode (-, negative), attracting positive ions e.g. Al3+, and the anode (+, positive) which attracts negative ions e.g. O2-.

When the d.c. current is passed through aluminium forms at the negative cathode (metal*) and sinks to the bottom of the tank.

At the positive anode, oxygen gas is formed (non-metal*). This is quite a problem. At the high temperature of the electrolysis cell it burns and oxidises away the carbon electrodes to form toxic carbon monoxide or carbon dioxide. So the electrode is regularly replaced and the waste gases dealt with!
It is a costly process (6x more than Fe!) due to the large quantities of expensive electrical energy needed for the process.

* Two general rules: for electrolysis

Metals and hydrogen (from positive ions), form at the negative cathode electrode.
Non-metals (from negative ions), form at the positive anode electrode.

Raw materials for the electrolysis process:

Bauxite ore of impure aluminium oxide [Al2O3 made up of Al3+ and O2- ions]

Carbon (graphite) for the electrodes.

Cryolite reduces the melting point of the ore and saves energy, because the ions must be free to move to carry the current

Electrolysis means using d.c. electrical energy to bring about chemical changes e.g. decomposition of a compound to form metal deposits or release gases. The electrical energy splits the compound!

At the electrolyte connections called the anode electrode (+, attracts - ions) and the cathode electrode (-, attracts + ions). An electrolyte is a conducting melt or solution of freely moving ions which carry the charge of the electric current.

The redox details of the electrode processes:
At the negative (-) cathode, reduction occurs (electron gain) when the positive aluminium ions are attracted to it. They gain three electrons to change to neutral Al atoms.

Al3+ + 3e- ==> Al
At the positive (+) anode, oxidation takes place (electron loss) when the negative oxide ions are attracted to it. They lose two electrons forming neutral oxygen molecules.

2O2- ==> O2 + 4e-
or 2O2- - 4e- ==> O2

Note: Reduction and Oxidation always go together!
The overall electrolytic decomposition is ...
aluminium oxide => aluminium + oxygen
2Al2O3 ==> 4Al + 3O2

and is a very endothermic process, lots of electrical energy input!

anode graphite react with oxigen to form carbodioxide. so the it should be replaced time to time

chemical calculation - igcse practice

2008-11-05T06:20:00.000-08:00

Introducing the connection between moles, mass and formula mass

The mole is most simply expressed as the 'formula mass in g' of the defined chemical 'species', and that is how it is used in most chemical calculations.

Every mole of any substance contains the same number of the defined species.

The actual particle number is known and is called the Avogadro Constant and is equal to 6.023 x 1023 'defined species' per mole. This means there are that many atoms in 12g of carbon (C = 12) or that many molecules of water in 18g* (H2O = 1+1+16 = 18, H = 1; O = 16) * this is about 18cm3, so picture this number of molecules in a nearly full 20cm3 measuring cylinder or a 100ml beaker less than 1/5th full!

However, the real importance of the mole is that it allows you to compare ratios of the relative amounts of reactants and products, or the element composition of a compound, at the atomic and molecular level. If you have a mole ratio for A:B of 1:3

it means 1 particle of A to 3 particles of B irrespective of the atomic or formula masses of A and B. (see also section 6. for reacting masses not using moles)

Important Note. Relative is just a number based on the carbon-12 relative atomic mass scale. Molar mass is a term used to describe the mass of one mole i.e. the relative atomic/formula/molecular mass in grams (g).

Examples:

Example 7.1.1: 1 mole of ammonia, NH3,

consists of 1 mole of nitrogen atoms combined with 3 moles of hydrogen atoms.

Or you could say 2 moles of ammonia is formed from 1 mole of nitrogen molecules (N2) and 3 moles of hydrogen molecules (H2).

Example 7.1.2: 1 mole of aluminium oxide,

Al2O3, consists of 2 moles of aluminium atoms combined with 3 moles of oxygen atoms

(or 1.5 moles of O2 molecules).

For calculation purposes learn the following formula for 'Z' and use a triangle if necessary.

(1) mole of Z = g of Z / atomic or formula mass of Z,

(2) or g of Z = mole of Z x atomic or formula mass of Z

(3) or atomic or formula mass of Z = g of Z / mole of Z

where Z represents atoms, molecules or formula of the particular element or compound defined in the question.

Example 7.2.1: How many moles of potassium ions and bromide ions in 0.25 moles of potassium bromide?

1 mole of KBr contains 1 mole of potassium ions (K+) and 1 mole of bromide ions (Br-).

So there will be 0.25 moles of each ion.

Example 7.2.2: How many moles of calcium ions and chloride ions in 2.5 moles of calcium chloride?

1 mole of CaCl2 consists of 1 mole of calcium ions (Ca2+) and 2 moles of chloride ion (Cl-).

So there will be 2.5 x 1 = 2.5 moles of calcium ions and 2.5 x 2 = 5 moles chloride ions.

Example 7.2.3: How many moles of lead and oxygen atoms are needed to make 5 moles of lead dioxide?

1 mole of PbO2 contains 1 mole of lead combined with 2 moles of oxygen atoms (or 1 mole of oxygen molecules O2).

So 1 x 5 = 5 mol of lead atoms and 2 x 5 = 10 mol of oxygen atoms (or 5 mol oxygen molecules) are needed.

Example 7.2.4: How many moles of aluminium ions and sulphate ions in 2 moles of aluminium sulphate?

1 mole of Al2(SO4)3 contains 2 moles of aluminium ions (Al3+) and 3 moles of sulphate ion (SO42-).

So there will be 2 x 2 = 4 mol aluminium ions and 2 x 3 = 6 mol of sulphate ion.

Example 7.2.5: How many moles of chlorine gas in 6.5g? Ar(Cl) = 35.5)

chlorine consists of Cl2 molecules, so Mr = 2 x 35.5 = 71

moles chlorine = mass / Mr = 6.5 / 71 = 0.0944 mol

Example 7.2.6: How many moles of iron in 20g? (Fe = 56)

iron consists of Fe atoms, so moles iron = mass/Ar = 20/56 = 0.357 mol Fe

Example 7.2.7: How many grams of propane C3H8 are there in 0.21 moles of it? (C = 12, H = 1)

Mr of propane = (3 x 12) + (1 x 8) = 44, so g propane = moles x Mr = 0.21 x 44 = 9.24g

Example 7.2.8: 0.25 moles of molecule X was found to have a mass of 28g. Calculate its molecular mass.

Mr = mass X / moles of X = 28 / 0.25 = 112

Example 7.2.9: What mass and moles of magnesium chloride is formed when 5g of magnesium oxide is dissolved in excess hydrochloric acid?

reaction equation: MgO + 2HCl ==> MgCl2 + H2O

means 1 mole magnesium oxide forms 1 mole of magnesium chloride (1 : 1 molar ratio)

formula mass MgCl2 = 24+(2x35.5) = 95,

MgO = 24+16 = 40, 1 mole MgO = 40g, so 5g MgO = 5/40 = 0.125 mol

which means 0.125 mol MgO forms 0.125 mol MgCl2,

Mass = moles x formula mass = 0.125 x 95 = 11.9g MgCl2

Example 7.2.10: What mass and moles of sodium chloride is formed when 21.2g of sodium carbonate is reacted with excess dilute hydrochloric acid?

reaction equation: Na2CO3 + 2HCl ==> 2NaCl + H2O + CO2

means 1 mole sodium carbonate gives 2 moles of sodium chloride (1:2 ratio in equation)

Formula mass of Na2CO3 = (2x23) + 12 + (3x16) = 106

Formula mass of NaCl = 23 + 35.5 = 58.5

moles Na2CO3 = 21.2/106 = 0.2 mole

therefore 2 x 0.2 = 0.4 mol of NaCl formed.

mass of NaCl formed = moles x formula mass = 0.4 x 58.5 = 23.4g NaCl

Using the Avogadro Constant, you can actually calculate the number of particles in known quantity of material.

Example 7.3.1: How many water molecules are there in 1g of water, H2O ?

formula mass of water = (2 x 1) + 16 = 18

every mole of a substance contains 6 x 1023 particles of 'it' (the Avogadro Constant).

moles water = 1 / 18 = 0.0556

molecules of water = 0.0556 x 6 x 1023 = 3.34 x 1022

Since water has a density of 1g/cm3, it means in every cm3 or ml there are

33 400 000 000 000 000 000 000 individual H2O molecules or particles.

Example 7.3.2: How many atoms of iron (Fe = 56) are there in an iron filing of mass 0.001g ?

0.001g of iron = 0.001 / 56 = 0.00001786 mol

atoms of iron in the nail = 0.00001786 x 6 x 1023 = 1.07 x 1019 actual Fe atoms

(10.7 million million million atoms!)

Example 7.3.2: (a) How many particles of 'Al2O3' in 51g of aluminium oxide?

Atomic masses: Al =27, O = 16, f. mass Al2O3 = (2x27) + (3x16) = 102

moles 'Al2O3' = 51/102 = 0.5 mol

Number of 'Al2O3' particles = 0.5 x 6 x 1023 = 3 x 1023

(b) Aluminium oxide is an ionic compound. Calculate the number of individual aluminium ions (Al3+) and oxide ions (O2-) in the same 51g of the substance.

For every Al2O3 there are two Al3+ and three O2- ions.

So in 51g of Al2O3 there are ...

0.5 x 2 x 6 x 1023 = 6 x 1023 Al3+ ions, and

0.5 x 3 x 6 x 1023 = 9 x 1023 O2- ions.

More advanced use of the mole and Avogadro Number concepts (for advanced level students only)

You can have a mole of whatever you want in terms of chemical species e.g.

In terms of electric charge, 1 Faraday = 96500 C (coulombs) = 6 x 1023 electrons

If you have 2.5 moles of the ionic aluminium oxide (Al2O3) you have ...

2 x 2.5 = 5 moles of aluminium ions (Al3+) and 3 x 2.5 = 7.5 mol of oxide ions (O2-)

When you write ANY balanced chemical equation, the balancing numbers, including the un-written 1, are the reacting molar ratio of reactants and products.

Extra Advanced Questions - more suitable for Advanced AS-A2 students which can be completely tackled after ALSO studying section 9 on the molar volume of gases and ANSWERS to QA7.1

QA7.1 This question involves using the mole concept and the Avogadro Constant in a variety of situations.

The Avogadro Constant = 6.02 x 1023 mol-1. The molar volume for gases is 24dm3 at 298K/101.3kPa.

Atomic masses: Al = 27, O = 16, H = 1, Cl = 35.5, Ne = 20, Na = 23, Mg = 24.3, C = 12

Where appropriate assume the temperature is 298K and the pressure 101.3kPa.

Calculate ....

(a) how many oxide ions in 2g of aluminium oxide?

(b) how many molecules in 3g of hydrogen?

(c) how many molecules in 1.2 cm3 of oxygen?

(d) how many molecules of chlorine in 3g?

(e) how many individual particles in 10g of neon?

(f) the volume of hydrogen formed when 0.2g of sodium reacts with water.

(g) the volume of hydrogen formed when 2g of magnesium reacts with excess acid.

(h) the volume of carbon dioxide formed when the following react with excess acid

(1) 0.76g of sodium carbonate and (2) 0.76g sodium hydrogencarbonate

(i) the volume of hydrogen formed when excess zinc is added to 50 cm3 of hydrochloric acid, concentration 0.2 mol dm-3.

(j) the volume of carbon dioxide formed when excess calcium carbonate is added to 75 cm3 of 0.05 mol dm-3 hydrochloric acid.

reactivity series - IGCSE /GCSE

2008-11-05T06:13:00.000-08:00

It is possible to organise a group of similar chemicals that undergo either oxidation or reduction according to their relative reactivity.

Oxidation (and reduction) is a competition for electrons.

The oxidising species (agents) remove electrons from other species and can force them to become reducing agents (releasers of electrons)

A good example of this competition for electrons is the behaviour of metals.

Metals always react by losing electrons (oxidation) they are then reducing agents.

However if a metals is in competition with metal ions the more reactive metal can oblige the less reactive metal (in the form of ions) to accept electrons.

This is called a displacement reaction.

Example

Zinc reacts with a solution containing copper ions.

The zinc metal is more reactive than copper metal and so it can force the copper metal ions to accept electrons and become metal atoms.

Zn(s) Zn2+(aq) + 2e

Cu2+(aq) + 2e Cu(s)

The zinc metal passes its electrons to the copper ions.

We observe that the zinc develops a pink layer of coper on its surface and the blue copper ion solution fades in colour.

We say that the zinc displaces the copper ions from solution.

Experimental observations

If we observe that there is a reaction between a metal and another metal ion in solution this tells us that the solid metal is more reactive than the metal of the dissolved metal ions.

Iron displaces copper from a solution of copper II sulphate

Copper displaces silver from a solution of silver nitrate

Given this information we can deduce that the most reactive of the three metals is iron, followed by copper, followed by silver. This allows us to arrange the metals into a reactivity series based on these specific reactions

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Reduction of metal oxides by metals

metal A + metal B oxide metal A oxide + metal B

When a metal A is heated with a metal B oxide there will be a reaction if the free metal A is more reactive than the metal B of the metal B oxide. This is because the metal B in the metal B oxide is in the form of a metal ion - it has already lost electrons.

There is a competition between the metal ion (in the oxide) and the free metal for the electrons. The more reactive of the two metals will win the competition.

Consequently if there is a reaction between a metal and a metal oxide then this tells us that the free metal is more reactive than the metal in the metal oxide.

Experimental observations

Magnesium reacts with zinc oxide:- Mg + CuO MgO + Cu

Sodium reacts with magnesium oxide: 2Na + MgO Na2O + Mg

Zinc reacts with copper oxide:- Zn + CuO ZnO + Cu

We can use this information to arrange the metals in order of reactivity

Sodium most reactive

Magnesium

Zinc

Copper least reactive

Sodium has the greatest electron releasing power (and conversely the copper ions - Cu2+ - would have the greatest electron attracting power)

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Predictions from reactivity series

Once a reactivity series is produced it can be used to predict reactions of pairs of reactant.

For example in the table above it should be appreciated that magnesium will react with copper oxide reducing it to copper metal.

Any metal that is more reactive will react with compounds of less reactive metals.

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Reactivity series involving non-metals

Metals react by losing electrons - they are reducing agents. Non-metals react by gaining electrons - they are oxidising agents.

In the same way that metals can be ordered in terms of reducing strength, the non-metals can be ordered in terms of their oxidising strength.

The halogens are a typical example of a non-metal reactivity series.

Reactivity of the halogens

Fluorine most reactive

Chlorine

Bromine

Iodine least reactive

Fluorine is so reactive that we cannot isolate it in the laboratory very easily, as it reacts with both water and glass. As a result we don't usually deal with fluorine at pre-university level

but compare only the other three (astatine is very rare and radioactive)

Do not confuse this order of reactivity with that of the metals - these are non-metals, their reactivity is in terms of oxidising power - i.e. chlorine is the best oxidising agent out of chlorine, bromine and iodine.

1. Chlorine will displace bromine from solutions containing bromide ions

Cl2 + 2Br- Br2 + 2Cl-

In this reaction the chlorine is oxidising the bromide ions by removing an electron from them. Bromine is liberated from the solution and may be detected by its orange/red colour

2. Bromine will displace iodine from solutions containing iodide ions

Br2 + 2I- I2 + 2Br-

In this reaction the bromine is oxidising the iodide ions by removing an electron from them. Iodine is liberated from the solution and may be detected by its orange/brown colour which turns blue/black in the presence of starch indicator.

It is predictable, then, that chlorine will also displace iodine from a solution containing iodide ions

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10.2.2: Deduce the feasibility of a redox reaction from a given reactivity series.

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Prediction of feasibility

Once a reactivity series is constructed depending on the reduction or oxidation ability of each species, we can use it to predict the feasibility (probability) of a reaction occurring between any two pairs of reactants.

If one of the substances is a reducing agent - i.e. it reacts by losing electrons then this must react with an oxidising agent - i.e. a species that gains electrons.

Example

Potassium

K-------- K+ + 1e
best reducing agents (left hand side species)
Magnesium

Mg --------Mg2+ + 2e

Zinc

Zn------- Zn2+ + 2e

Iron

Fe---------- Fe2+ + 2e

Copper

Cu ----Cu2+ + 2e

Hydrogen

H2--------- 2H+ + 2e

Iodine

2I- -----I2 + 2e

Bromine

2Br- -----------Br2 + 2e

Chlorine

2Cl- ----------Cl2 + 2e
best oxidising agents (right hand side species)

Any species from the right hand side of one of the redox equilibria (the oxidising agent) can be predicted to react with any species above it on the left hand side of the redox equilibria (the reducing agent).

The species on the right hand side of the equilibria will gain electrons to go to the right hand side. They can only gain these electrons form species that are abopve them on the left hand side of the series.

We can therefore predict that chlorine (right hand side) will react with copper (left hand side) to form copper ions nad chloride ions according to the equation:

Cl2 + Cu Cu2+ + 2Cl-

Similarly, we can predict that iodide ions (left hand side) will NOT react with zinc ions (left hand side) as the zinc ions are poor oxidising agents and the iodide ions poor reducing agents.

Note: although a reaction may be predicted as feasible it does not mean that it will happen spontaneously. If the activation energy is high then it may need an extra "push" to get it going. - for example the reaction between chlorine and hydrogen needs a spark or ultraviolet light and then it is explosively fast.

chemical calculation-igcse

2008-11-05T06:00:00.000-08:00

Mole concept & Avogadro's constant

Mole of something is equivalent to 6.023 x 1023 (<-- Avogadro's constant) units of it...ie lots of atoms, molecules etc...the periodic table gives molar masses...ie the number of grams of a substance required for 1 mol of atoms.

This can be extrapolated to molecules of known molecular formula. Number of mols = mass / mass per mol (Usually found on periodic table)...

The coefficients in chemical equations give the molar ratios of reactants and products..

.ie 2A + 3B -> C.

There is 2/3 as much A as B, and 3 times more B than C involved in the reaction...

Assuming the reaction goes to completion, there must be 3/2times as much B as A for neither to remain...

If this ratio is not followed, one will be a limiting reactant, and so the reaction will have some of the other reactant left over when it completes.

Moles vs Mass...Moles is a number of something...every mol being 6.023x1023 individual elements.

Mass is the property which results in 'weight' in the presence of Gravity. Given a molar mass, M a mass m and a number of mols N then N=m/M.

An 'Empirical formula' is the formula describing the different atoms present in a molecules, and their ratios, but not the actual number present.

ie AxByZc could be an empirical formula if x, y, and z are in lowest common terms. The molar mass can then be used to calculate the actual numbers of each atom present per molecule. The empirical formula can be determined by percentage composition, or anything else which gives the ratios of atoms present.

A Molecular Formula is similar to an empirical formula except that it includes the the number of atoms present in each molecule, rather than the ratio.

It will be an integer multiple of the empirical formula ie KAxByZc and can be calculated from the empirical formula and the molar mass of the substance.

Chemical Equations
The mole ratio of two species in a chemical equation is the ratio of their coefficients...ie aX + bY -> cZ : The ratio of X/Yis a/b, Y/Z= b/cetc...

Balancing equations...change only the coefficients, not the subscripts to make sure all atoms, and charge is conserved (half equations can be balanced by addition of electrons to either side...2 half equations can be added by making the number of electrons equal in each, then vertically adding.)

State symbols -- (s)-Solid , (l)-liquid, (g)-gas, (aq)-aqueous solution...ie something dissolved in water. Should be included in all chemical reactions (but won't be penalized).

Mass relationships in chemical reactions

Mass is conserved throughout reactions. This fact allows masses to be calculated based on other masses in the reaction eg burning Mg in air to produce MgO and so to find the mass or Mg present in the original sample (ie purity)...can be extended to concentrations...ie titration.

When a reaction contains several reactants, some may be in excess...is more is present that can be used in the reaction. The first reactant to run out is the limiting reagent (or reactant). Knowing the number of mols of the limiting reagent allows all other species to be calculated, and so the yield, and remaining quantities of other reactants.

Solutions

Solvent - the stuff you're dissolving in...ie water. Solute - the stuff you're dissolving in it...an ionic compound or something...Solution - the two of them when mixed together. Concentration - the amount of solute per amount of solvent...in mols per dm3...ie mols per liter or grams per liter.

Apply the equation concentration = number/volume...rather obvious from the units of concentration, but remember to covert everything into the same units.

Use chemical equations to relate the amount of one species to the amounts of others.

states of matter

2008-11-05T05:54:00.000-08:00

States of Matter

Kinetic Theory — All matter is made up of particles which are in constant motion.

Solids

· Fixed shape and volume

· Particles are held together by relatively strong forces, and are incompressible

· Particles do not have free movement, but can vibrate about fixed positions

· An increase in temperature will give the particles more energy, the amplitude of vibration increases, therefore the solid expands

Liquids

· No fixed shape, but a fixed volume

· Particles are further apart (due to weaker forces), so there is slight compressibility

· Forces are present between particles, but weaker than in solid so particles can move throughout the bulk of the liquid

Gases

· No fixed shape or volume

· Particles are very far apart, so are easy to compress

· Intermolecular forces are negligible, but do exist. Particles move in a random fashion.

Diffusion

Diffusion is the process by which a gas fills all the space available to it. Gases will diffuse (mix with each other) because the particle are moving randomly and quickly in all directions. Lighter particles diffuse quicker than heavier particles.

sublimation- Change of state from soild to gas
example- dry ice

Changes from solid to liquid to gas are endothermic processes

Changes from gas to liquid to solid are exothermic processes
If energy is supplied to a solid, its particles will vibrate more. If they vibrate enough, they may separate from each other and become free to move. This is called melting. At the melting point, energy goes into breaking forces between the particles, and so there is no temperature increase. The temperature at which the solid melts is called the melting point.

Heating a liquid makes the particles move faster. At evaporation, some molecules move faster than others and have more energy, so they overcome the forces of attraction. If the temperature is increased further, the kinetic energy of the particles increases until the boiling point, when the forces between the particles are almost completely broken. At the boiling point energy goes into breaking the forces between the particle, and so there is no temperature increase.

For an impure sample the melting point is lowered and the boiling point is raised.

POLYMERS- NOTES

2008-10-14T23:58:00.000-07:00

<<strong>A polymer is a large molecules composed of repeating structural units (monomers) connected by covalent chemical bonds

It’s the breaking down of double bond carbon atoms , and the the electrons in it are used to join to neighboring molecules .

In its case polymerization take place ,where lots of little molecules(monomers) join up to make one big molecule (polymer)

In the case of ethene lots of ethene molecules join together to make poly(ethene)

When propene polymerized you get polypropene

Polymerising chloroethene gives you poly(chloroethene).

2- condensation polymer

Here the polimerisation method is totally different . if one monomer contains an – OH group and the other an –H , then they add by forming a molecule of water . when they polymerise , the water molecule is split out between them .
Also the same way in splitting out of HCL molecule.

*by removing water

HOOC- ___ ___ COOH + H2N- ___ ___ NH2 --
(carboxylic acid) (amino group)

HOOC ___ ___ CONH ___ ___NH2 +H2O
( amide linkage )

{another way }

COOH ___ ___COOH + NH2 ___ ___NH2---
(dicarboxylic acid) (diamine) -H2O

(___ ___CONH___ ___ )n

(amide linkage)

*by removing HCL

COCL ___ ___ COCL +NH2 ___ ___ NH2 -----
-HCL

(____ ____CONH ____ ____)n
(amide linkage)

*by removing water

HOOC ___ ___ COOH +HO- ___ ___OH ----
(carboxylic acid ) (alcohol)

HOOC- ___ ___ COO___ ___OH
( ester linkage )

{another way}
OH___ ___OH + COOH ___ ___COOH à
(diol) (dicarboxylic acid) -H2O

(___ ___COO___ ___)n
(ester linkage)

……………………………………………………………………….

*by removing HCL

COCL ___ ___COCL + OH___ ___OH ---
(acid chloride) (diol) -H2O

(___ ___COO___ ___)n
(dicarboxylic acid)

USES

polyethene

Bags, plastic milk bottles
Not very strong , flexible

poly propene

To make ropes and crates
strong

polyvinaly chloride

Drainpipes, replacement windows
Quite strong and rigid

nylon

Make toothbrush and ropes
strong
To make clothing
Very strong

periodic table - groups -notes

2008-10-14T23:54:00.000-07:00

PERIODIC TABLE

The elements in the periodic table are arranged according ti the increasing atomic number. In the peridoic table there are eight main groups and seven periods.
The elements in the same groups have similar properties because they have the same number of electrons in their outer shell.
The group number of the elements is the same as the number of the outer electrons. 3/4 of the periodic table is filled with metals.

There are other elements called the semi metals whos hare both the properties of metals and non metals.They conduct like metals and are brittle like non-metals.

Reactions and electron arrangementsGroup 1 elements all have atoms with 1 outer electron. e.g. sodium 2,8,1All of these atoms try to lose 1 electron so all react in the same way.

They all form a positive ion and form ionic compounds. E.g. sodium Na form sodium ion Na+.Group 1 elements all react with oxygen (that is burn) to form oxides. The products like sodium oxide, are all ionic.

sodium burn

Group 0 elements all have atoms with complete outer shells e.g. helium 2, neon 2,8. None of these elements react with anything.

Name other elements in group 1 and say how they react with oxygen giving a reason. State how argon reacts with oxygen and give a reason.

Trends in group propertiesGroup 1 elements get more reactive with water as you go down the group. e.g.Lithium (gentle fizz) is less reactive than potassium (violent lilac flame).Group 7 elements get more reactive with iron as you go up the group.Iodine darkens iron when it is heated but chlorine makes iron burst into flames.

GROUP 1

They are called alkali metals because they react with water to form strong alkali. The increase in number of shells and atomic size causes increase ub the reactivity down the group because bigger the atom, the less its attractive force towards the nucleus.

Group 1 metals have one electron in the outermost shell which is why they are called group 1.

They are less dense than group 2 metals because they have only one electron in the outermost shell and they're stored under oil to prevent them from reacting iwth oxygen and water vapour form at the atnoshpere as it is dangerous.

Reaction with water

Sodium is a strong alkali, when it reacts with water :

1) It floats and melts.
2) Hydrogen gas is produced.
3) It forms sodium hydroxide.

Sodium + water ---------------> Sodium hydroxide + hydrogen
Na + H2O ---------------------> NaOH + H2
Na + HOH -------------> NaOH + H2

Potassium when reacting with water :

1) It floats.
2) It catches fire and burns with a lilac flame.
3) It forms potassium hydroxide and hydrogen.

Potassium + water --------------------> Potassium hydroxide + hydrogen
K + H2O --------------------------------> KOH + H2
K + HOH -------------------------------> KOH + H2

Lithium

Lithium forms Lithium hydroxide and hydrogen gas when reacting with water.

Li + H2O ----------------> LiOH + H2

Rubedium (Rb)

Rb + H2O -----------> Rb (OH) + H2

When rubedium reacts with water :

1) Heat is produced.
2) Explosion takes place.
3) It forms Rubedium hydroxide and hydrogen gas.

analysis - NOTES/ flame test

2008-10-14T23:35:00.000-07:00

Testing ions by using flame test :
A flame test is used to show the presence of certain metal ions in a compound. A wire of platinum or nichrome is first dipped into a concentrated Hydrochloric acid and then held in a hot bunsen flame ,,this process should be done before testing any ion to clean the wire.
Test for some ions : (by using a flame test)
flame colour
ion
Red
Lithium / Li
Yellow-orange (strong yellow)
Sodium / Na
Orange-red (brick red)
Calcium / Ca
Pale green (apple green)
Barium / Ba
Lilac colour
Potassium / K
Blue-green
Copper / Cu

PURIFICATION AND USES OF COPPER

2008-10-07T07:57:00.001-07:00

keep impure copper at anode and purecopper wire at cathode

At anode

Cu --------------Cu²+ + 2e-

At cathode

Cu²+ + 2e- ------------- Cu

* Salt solution of copper (CuSo4 ) is used as the electrolyte.

* The copper from the anode goes to the solution as C

* Cu²+ ion goes to the cathode and gets deposited.

Uses Of Copper:

1- Electrical wiring: good conductor of electricity.
2- Domestic plumbing: does not react with water.
3- Making alloys like brass (zinc & copper): harder than copper & zinc.
4- Making coins: long lasting without any corrosion.

electrolysis of sodium chloride

2008-10-07T07:54:00.001-07:00

Sodium hydroxide, chlorine and hydrogen are produced from natural deposits of sodium chloride (rocksalt). This is the basis of the chloralkali industry. Electrolysis of brine (salt solution) produces 1 tonne of chlorine at the same time as 1.13 tonnes of sodium hydroxide and 0.028 tonnes of hydrogen. There are many different processes used in the electrolysis of brine. An aqueous solution contains the following ions: Na+, Cl-, H+ and OH-.

at anode

2Cl- Cl2 + +2e-

At cathode:
2H+ + 2e- H2

* The remaining solution contains NaOH

electrolysis of lead bromide-notes

2008-10-07T07:43:00.000-07:00

molten lead bromide is taken in the cell. Electricity is passed.

cations (Pb²+) are moving towards the cathode and get discharged as:

Pb Pb²+ + 2e-

The anions (Br-) are moving towards the anode and get discharged as:

2Cl- 2Cl2 + 2e-

oxidation and reduction- notes

2008-10-07T07:41:00.000-07:00

Oxidation of ions or neutral molecules can take place at the anode, and the reduction of ions or neutral molecules at the cathode.

Oxidation and Reduction reactions takes place in the cell, so this is known as reduce reaction.

It can also be memorized as (OIL RIG) where Oxidation is losing of electrons and Reduction is gaining of electrons.

Substances that have the ability to oxidize other substances are said to be oxidative and are known as oxidizing agents, oxidants or oxidizers. Put another way, the oxidant removes electrons from another substance, and is thus reduced itself. And because it "accepts" electrons it is also called an electron acceptor.

Substances that have the ability to reduce other substances are said to be reductive and are known as reducing agents, reductants, or reducers. Put in another way, the reductant transfers electrons to another substance, and is thus oxidized itself. And because it "donates" electrons it is also called an electron donor.

electrolysis- notes

2008-10-07T07:40:00.000-07:00

* Electrolysis is the process of splitting substances by passing electricity.

* Electrolysis involves the passage of an electric current through an ionic substance that is either molten or dissolved in a suitable solvent, resulting in chemical reactions at the electrodes.

* Electrodes are metals used for electrical contact. The positive electrode is called the anode and the negative electrode is called the cathode. To be useful for electrolysis, the electrodes need to be able to conduct electricity, and metal electrodes are generally used. Graphite electrodes and semiconductor electrodes are also used.

* An ionic compound, or compound that reacts with solvent to produce ions (such as an acid) is dissolved in an appropriate solvent, or an ionic compound is melted by heat.

* Each electrode attracts ions that are of the opposite charge. Therefore, positively charged ions (called cations) move towards the electron-emitting (negative) cathode, whereas negatively-charged ions (called anions) move towards the positive anode.

* Electrolytes are substances which can be split by passing electricity. They are usually ionic compounds.
E.g.: NaCl, PbBr, etc …

* The energy required to separate the ions, and cause them to gather at the respective electrodes, is provided by an electrical power supply. At the electrodes, electrons are absorbed or released by the ions, forming a collection of the desired element or compound.

summary non metals- notes

2008-10-07T00:31:00.001-07:00

Non-metals

1- Non-metals have different physical and chemical properties to metals.
2- You can prepare oxygen, carbon dioxide, hydrogen and ammonia gases in the laboratory.
3- Ammonia can be oxidized to nitric acid.
4- Ammonium nitrate is both a fertilizer and an explosive.
5- Cooling air gives liquid air which can be separated.
6- Sulphur forms allotropes.

Experiment to show the % of oxygen in air- notes

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procedure

1- Syringe A contains 100cm of air from the atmosphere.

2- The silica tube contains cooper tunings.

3.The syringe B is closed

4- Light the burner D and press the syringe A.

5- Copper react with oxygen from the syringe and forms black color copper oxide.

6- The remaining gas will be collected in syringe B which is 79 cm3.

7- So the oxygen present in air = 100-79 = 21cm3.

Extraction of oxygen and nitrogen from liquid air

Composition of gases in air - notes

2008-10-07T00:28:00.001-07:00

GAS % IN ATMOSPHERE

Nitrogen 78%

Oxygen 21%

Argon 0.9%

Carbon dioxide 0.03%

preparation of hydrogen- notes

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equation

Zn+ HCl ------ ZnCl2 + H2

sulphur- notes

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S + O2--------SO2
S+ Fe ----------FeS (IRON SULPHIDE)

Sulphur : Three forms of sulphur
1- Rhombic sulphur
2- Monoclinic sulphur
3- Plastic sulphur

ammomium salts- notes

2008-10-07T00:25:00.002-07:00

Ammonium salt:

Ammonia + Hydrochloric acid = Ammonium ChlorideIt is a reversible reaction.
Heating the NH4CL gives back NH3 and Hcl gases.
NH3 + HCL= NH4CL
NH3 + H2SO4 = (NH4)2 SO4
(NH4)2 is an important nitrogen containing fertilizer.
NH3 + HNO3 = NH4NO3 (Ammonium nitrate- fertilizer)

ammonia-notes/IGCSE /GCSE -CHEMISTRY

2008-10-07T00:25:00.001-07:00

test for carbondioxide- notes

2008-10-07T00:23:00.000-07:00

Turns lime water milky.
Calcium hydroxide + carbon dioxide ---------calcium carbonate + water.
Ca (OH) 2(aq) + CO2 (g) CaCO3(s) + H2O (l)

properties and uses of carbondioxide - notes/IGCSE/GCSE NOTES

2008-10-07T00:22:00.000-07:00

1- Colorless gas.
2- More dense than air.
3- Solid below -78 degree.
4- Soluble in water.
H2o + Co2 --> H2CO3 (carbonic acid)
Carbonic acid is present in acid rain and carbonated drinks.
Uses of carbon dioxide
1- Fizzy drinks.
2- Fire extinguisher.
3- Dry ice.

4. PHOTO SYNTHEIS

preparation of carbondioxide- notes/IGCSE /GCSE/CHEMISTRY NOTES

2008-10-07T00:21:00.000-07:00

Equation
CaCO3 + 2HCl = CaCl2 + H2O + CO2

formation of oxides- Notes

2008-10-07T00:20:00.001-07:00

Mg + O2 --> Mgo (basic- All the metals forms basic oxides)

2- S + O2 --> So2 (Acidic- All the non-metals forms acidic oxides)

3- CH4 + O2 --> Co2 + H2o

Very important: Any hydrocarbon reacts with oxygen gives carbon dioxide and water.

reactions of copper- notes/ IGCSE / GCSE

2008-10-07T00:17:00.000-07:00

When copper is heated in a flame, it turns black with a coating of copper (II) oxide.
Copper+ oxygen --> copper(II) oxide
2- This is an example of thermal decomposition. Copper (II) oxide is a base- it reacts with acids to form salts.
Base+ acid--> salt + water
Copper(II) sulphate + nitric acid--> copper nitrate + water
3- The formation of the deep blue complex ion with ammonia is used as a test for the presence of copper in solution. Other metals give green-blue solutions for example nickel or cobalt, but only copper forms this colored complex.
Here, ammonia (NH3) is added to copper(II) sulpate (CuSO4). The result is a pale blue precipitate, which is Cu(OH)2.
NH3 Add more the product is
(Cu(H20)2 (NH3)4 ) 2+ )
Compound Color
Iron(II) chloride Green
Iron(III) chloride Brown
Copper sulphate Blue
Copper carbonate Blue green
Copper oxide Black
Copper nitrate Blue
Copper chloride Green
Other reactions:
Fe + Hcl--> Fecl2
Fe+Cl2-->Fecl3
Cuo+Hno3-->Cu(No3)2+H2o
Cuo + Hcl --> Cucl2 + H2o
Cu + Hno3--> Cu(No3)2 + H2o + No2

reactions of iron-igcse /gcse notes

2008-10-07T00:13:00.000-07:00

Iron reacts with solutions of dilute acids to form iron (II) salts.
Iron+ sulphuric acid --> iron (II) sulphate + hydrogen
Iron reacts slowly with water to form rust. This reaction is used to produce Hydrogen from rusted scrap iron.

atomic structure - work sheet

2008-10-06T08:10:00.000-07:00

Answer the following
1. The smallest repeating unit of an element is called-----------------
2. Central of the atom is called---------------------
3. Name the three particles in side the atom
4. Name the negative particle in side the atom ---------------------
5. Name the positive particle inside the atom --------------------------
6. Name the path of the electron around the nucleus------------------------
7. The number of --------------------------and------------------------particles are same in
inside the atom
Try to answer the following
1.Why the atom is neutral?
2.Define mass number.
3. Define atomic number
4. Define isotopes
5.why isotopes have similar properties
6.which one of the following are isotopes
35 37 37 35
Cl , Cl , Cl , Cl
17 16 17 16
7.Draw the electron arrangement in
Magnesium, calcium , sodium, lithium, sulphur
8.Relative atomic mass of an elementUse relative mass of isotopes and their relative abundance. E.g. Chlorine has two isotopes with mass numbers 35 and 37. 35 37 75% is Cl and 25% is Cl 17 17 Let there be 100 atoms Total mass of 100 atoms = (75 * 35) + (25 * 37) = 3550 Average mass of an atom (relative atomic mass of chlorine) = Total mass /Number of atoms =3550/100 so relative atomic mass of chlorine = 35.5
Task
1. Bromine exists as two isotopes with mass numbers 81 and 79. If there is 50% of each isotope in a sample what is the relative atomic mass of bromine?
2. Calculate the relative atomic mass of the following:
(a) gallium 60% Ga 69 and 40% Ga 71
(b) neon 90% Ne 20 and 10% Ne 22
(c) silver 50% Ag 107 and 50% Ag 109
(d) boron 20% B 10 and 80% B 11