collecting gases / igcse /gcse chemistry notes

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

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

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

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

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

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

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

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

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

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

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

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

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

. _____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)

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

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)

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.