IGCSE Chemistry Revision Notes

Section 1: Principles of chemistry

a) States of matter

1.1 The arrangement, movement & energy of particles in each of three states of matter: solid, liquid, gas:

SOLID LIQUID GAS

Tightly packed

Vibrate about their fixed positions

In constant movement

Bounce off each other

Break and reform clusters

Move rapidly

Random direction

Very hight speeds

Strong forces of attraction between particles

Weaker forces of attraction between particles

Very weak forces of attraction between particles

M.P. + B.P GREATER than room temp.

M.P. LESS than room temp

B.P. GREATER than room temp

M.P.+ B.P. LESS than room temp

1.2 Interconversions of solids, liquids & gases:

SOLID LIQUID [melting]

LIQUID SOLID [freezing]

LIQUID GAS [evaporation/ boiling]

GAS LIQUID [condensation] GAS SOLID

SOLID GAS

1.3 Changes in arrangement, movement and energy of particles during interconversions

Melting: [melting point]

Particles vibrate faster

Gain enough energy to BREAK attractive forces

Freezing: [freezing point]

Particles move more slowly

Lose energy- come closer together

Forces of attraction take over

[sublimation]

Melting point and freezing point are exactly the SAME temps

Evaporation: [always happening at all temps]

Particles from the surface of liquid

Move faster than other particles

Have enough energy to break forces

Boiling: [specific temp]

All particles have enough energy to break forces

Escape from surface

Boiling point: shows strength of attractive forces between particles

Depends on surrounding pressure

Lower pressure lower B.P.

Sublimation: [fixed temp]

Solid directly to gas

b) Atoms

1.4 i) showing that particles are very small

Dissolve potassium maganate (VII) in water

Dilute solution until faint colour:

1 cm3 10 cm3 by adding water

Large number of particles in a very small mass

Must be very tiny

ii) Showing that particles move

Diffusion in gases:

Also do this with hydrogen and air (use lighted splint)

Diffusion in liquids:

iii) Showing that particles in different gases travel at different speeds:

Lid from each gas jar removed

Bromine gas particles bounce around until they reach the top

Brown colour evenly spread

Very slow

Crystal of potassium manganate (VII) into water

Crystal particles collide with water particles and SPREAD

NH4 molecules faster lighter [cloud closer to HCL)

LOWER THE RELATIVE MOLECULAR MASS FASTER DIFFUSION RATE

iv) Showing changes on a graph

Heating curve:

Cooling curve:

1.4 Atom smallest particle that can exist on its own + take part in a chem. Reaction

Molecule two/ more atoms chemically combined

1.5 Differences between elements, compounds & mixtures

Element Pure substance that cannot be split into anything simpler

Melting

Evaporation/ boiling

Condensation

Solidification/ crystallising

Compound 2/more

elements chemically joined [PURE]

Mixture 2/ more elements

mixed together [no chemical reaction]

Fixed composition No fixed composition

Own properties Same props are components

Fixed M.P. + B.P. Melts + boils over range of temps

Cannot be broken down by physical techniques

Separated by physical techniques

Undergoes chemical reaction No chemical reaction

1.6 Describe techniques for separation of mixtures

i) Simple distillation

Substance with the lower B.P. evaporates, rises up the flask and enters the condenser.

It condenses back into liquid and collects into the beaker

Substance with the highest B.P. is left behind in the distillation flask

A pure solvent from a solution

2 miscible liquids with a large difference

(more than 30°C) in their boiling points

ii) Filtration

iii) Evaporation

filtrate

residue

An insoluble solid from a liquid [solid

required]

2 solids, [one soluble and one insoluble in

particular solvent]

Uses: drinking water, penicillin

To obtain residue: rinse with water and dry

between filter paper or use desiccator

To obtain filtrate: heated to the point of

crystallisation

Evaporating basin

iv) Crystallisation

v) Fractional distillation

A solute (soluble solid) from an aqueous

solution

Point of crystallisation: until a saturated

solution is obtained.

Saturated solution: no more solute can

dissolve at a specific temperature in a specific

volume of water.

Remove a small quantity of the solution with

glass rod + drop on glass for cooling.

If crystals form quickly, the solution is saturated

faster the rate of crystallisation smaller +

impure crystals formed and vice-versa.

Miscible liquids with a small difference in boiling points

Liquids evaporate at the same time and mixture of vapours enters the fractionating column

Column heated up to temp of lower B.P. liquid in

mixture

The vapour of lower B.P. liquid can no longer

condense in the fractionating column and continues to

rise until it enters the condenser collected flask

On further heating the higher boiling point liquid will

distil over + be collected

vi) Chromatography

Separating Funnel

Only separates the mixture into its components

Filter/ chromatography paper

Rubber bung to ensure that the air in the container is saturated with the vapour

Most soluble solute in solvent travels through paper faster

Two immiscible liquids

Less dense liquid on top

Tap slowly opened + denser liquid collected in a beaker

Tap is then closed and liquid that remains in the separating funnel is

collected into another beaker

c) Atomic Structure

1.8 Atomic Structure

1.9 relative mass + relative charge of proton, neutron + electron

PROTON NEUTRON ELECTRON 1 amu 1 amu 1/2000 amu

+ No charge -

1.10

Atomic number: number of protons

Mass number: number of protons + number of neutrons

PROTONS

NEUTRONS

Cloud of ELECTRONS

Isotopes: atoms of the same element with the same no of protons but different no of neutrons

Same atomic number but different mass number

Relative atomic mass: the no of times heavier an atom of an element is compared to 1/12th the

mass of a carbon-12 atom

1.11 calculating the Ar from relative abundances of isotopes

(Abundance (%) x mass number) + (abundance (%) + mass number)

100

1.12 definition of periodic table

Periodic table: arrangement of elements in order of increasing atomic number

d) Relative Formula masses and molar volumes of gases

1.15 calculate Mr from Ar

By adding all Ar values of atoms in molecular formula

Unit: amu

1.16 understand the mole

Mole: measure of the amount of substance

1.17 mole as the Avogadro number of particles in a substance

1 mole = 6 x 1023 particles = Ar/Mr in grams (g)

Avogadro constant: the no of atoms present in 12g of carbon-12 isotope

1.18 examples of mole calculations

1. Calculate the mass given the number of moles:

1 mol Ca 40g

2.25 mol Ca x g

x= 2.25 x 40 = 90g

2. Calculate the number of moles given the mass (g):

5.5 g of lithium sulphate

Li2SO4 = 14+32+64= 110 g

1 mol 110 g

x mol 5.5 g

x= 5.5/110 = 0.05 mol

1.19 molar volume of a gas in calculations

Under the same conditions of temp + pressure, 1 mol of any gas will occupy the same volume

Room temp + pressure (rtp): 1 mol= 24000 cm3= 24 dm3 = 24 L

Examples:

1. Calculate no of moles in given volume

1 mol 24 dm3

x mol 1.5 dm3

x= 1.5/24 = 1/16 mol

2. Calculating mass in given volume

Avogadro’s Law: equal volumes of gas at same temp + pressure contain equal no of

molecules, hence equal no of moles.

Examples:

1. Calculate volume of H required to react completely 30cm3 of N + volume of ammonia

produced

3H2 + N2 NH3

3 : 1 : 2

3x30: 30: 2x30

H= 90cm3

NH3= 60 cm3

Example: sodium sulphite reacts with dilute HCl; the products are sodium chloride, sulphur dioxide and water. If 0.125 mol sodium chloride was produced during the reaction, calculate

a) The mass of sodium sulphite that reacted

Na2SO3 +2HCl 2NaCl + SO2 + H2O

2 mol NaCl from 1 mol Na2SO3

0.125 mol NaCl 0.0625 mol Na2SO3

Mr of Na2SO3 = 126

1 mol 126 g

0.0625 mol x g x= 126 x 0.0625 = 7.875 g

b) The volume of sulphur dioxide evolved at r.t.p

1 mol Na2SO3 1 mol SO2

0.0625 mol Na2SO3 0.0625 mol SO2

1 mol SO2 24 000 cm3

Calculations involving chem. Equations:

Balanced chem. Equation mole ratio of reactants + products

0.0625 mol x cm3 x= 0.0625 x 24000 = 15000 cm3

Concentration: the amount of solute dissolved in a given volume of solution [mol dm-3]

aka the number of moles of solute in 1000 cm3/ 1 dm3

Examples:

1. Calculate the number of moles given the volume and concentration of solution

3 dm3 of 2.0 mol dm-3 solution

2 mol 1 dm3

x mol 3 dm3

x = 3 x 2= 6 mol

2. Calculate concentration given the mass of solution and the volume

5.52 g K2CO3 to give 250 cm3 of solution

K2CO3 138 g

138 g 1 mol

5.52 g x mol x = 5.52/138 = 0.04 mol

0.04 mol 0.25 dm-3

x mol 1 dm-3 x= 0.16 mol dm-3

3. Calculate the mass required to make a given volume of specific concentration solution

Nitric acid to make 250cm3 of 0.100 mol dm-3 solution

0.100 mol dm-3 HNO3 1000 cm3

x mol HNO3 250 cm3 x= 250 x 0.1 / 1000 = 1/40 mol

HNO3= 63 g

1 mol HNO3 63 g

1/40 mol HNO3 x g x= 1.575 g

e) Chemical formulae and chemical equations

1.20 Write word equations and balanced chem. Equations

MUST KNOW:

Name of ion Symbol of ion Charge

Zinc

Zn

2+

Silver

Ag

1+

Ammonium

NH4

1+

Nitrite

NO2

1-

Nitrate

NO3

1-

Hydroxide

OH

1-

Sulphite

SO3

2-

Sulphate

SO4

2-

Hydrogen sulphate

HSO4

1-

Ethanoate

CH3COO

1-

Phosphate

PO4

3-

Carbonate

CO3

2-

Hydrogen carbonate

HCO3

1-

1.21 Uses of state symbols [s, l, g, aq]

Ionic equations:

-Break apart soluble molecules into the two ions that are formed (one positive and one negative)

-Insoluble molecules NOT broken apart

-Cancel out all common ions on both sides

Soluble Insoluble

ALL Na, K + NH4 salts

ALL nitrates

ALL hydrogen

carbonates

Chlorides except Ag + Pb chlorides

Sulphates except Ca, Ba + Pb sulphates

ALL carbonates

(Except Na, K +NH4)

Most Pb + Ag salts

ALL sulfides

Group 1, Ca + Ba except Most oxides + hydroxides Oxides + hydroxides

1.22 how formulae of simple compounds can be obtained experimentally including metal oxides, water + salts containing water of crystallization

Experiment to find the formula of an oxide of copper

Mass of oxygen present in oxide= mass of CuO at start – mass of Cu at end

1. How many mol in specific mass of oxygen 2. Determine mole ratio 3. Apply mole ratio 4. Find empirical formula

Experiment to determine the formula of water

- H gas passed over heated CuO reduced to Cu + H2O - H2O vapour through U-tube containing anhydrous calcium chloride absorbs water - Weighing mass of u tube before + after passing the H2O to calculate mass of H2O - Mass of O = Mass of CuO – Mass of Cu - Mass of H = mass of H2O – mass of O - Determine mole ratio of H + O - Divide by smallest number of two to attain whole numbers

Mass of copper oxide weighed + heated

Oxide: black red-brown (since only Cu left)

Stream of H2 passed continuously to prevent hot Cu reacting with air to form CuO

[ excess H2 gas ignited to prevent explosion ]

heating, cooling + weighing repeated until constant mass recorded

Combustion tube tilted (prevent condensed steam running back to hot tube)

1.23 calculate empirical + molecular formulae from experimental data

Empirical formula: the simplest whole number ratio of different atoms present in a molecule of a substance

- Mass of each element (from percentage composition) changed to MOLES - Simplest whole number ratio found by dividing by smallest number of moles found - If decimals obtained multiply by suitable number

Molecular Formula: the exact number of atoms of each element present in 1 molecule of the compound [multiple of empirical formula]

- Find empirical formula - Calculate relative molecular mass (Mr) - (Mass of empirical formula) n = Mr

Example:

Liquid Y molar mass 88 g mol-1 contains 54.5% C, 36.4% O + 9.1% H

C 54.5 % in 100g= 54.5 g = 4.5 mol

O 36.4% in 100 g= 36.4 g =2.275 mol

H 9.1%--> in 100g = 9.1 g = 9 mol

C : O : H

4.5/2.275 2.275/2.275 9/2.275

2 : 1 : 4

Empirical formula: C2H4O

Mr = 44 amu

Empirical: molecular

1 : 2

Molecular formula: C4H8O2

1.25 percentage yield

Percentage yield = Actual yield

Theoretical yield

X 100

f) Ionic compounds

1.27 formation of ions by gain/ loss of electrons

1.28 oxidation and reduction

Oxidation: gain of oxygen OR loss of electrons

Reduction: loss of oxygen OR gain of electrons

Oxidising agent: substance that causes another to be oxidized by getting reduced

Reducing agent: substance that causes another to be reduced by getting oxidised

The positive ion: CATION The negative ion: ANION

+ -

Always happen together

REDOX reactions

E.g. burning, respiration, rusting, spoiling of food

Half equations

Example of redox reaction:

Mg(s) + CuO (s) MgO(s) + Cu(s)

Ionic equation: Mg(s) + Cu2+(s) Mg2+

(s) + Cu(s)

Mg atoms Mg ions

Mg(s) Mg2+(s) + 2e-

Cu ions Cu atoms

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

1.29 recall charges of common ions (see 1.20)

1.30 deduce charge of ion from electronic configuration of atom from which it’s formed

Example: Sodium

Oxidation

Loses 1 electron

Charge: 1+

Reduction

1.31 explain using dot & cross diagrams the formation of ionic compounds by electron transfer (combinations of elements from group 1,2,3,5,6,7)

Example 1: NaCl

Example 2: CaCl2

1.32 Ionic bonding definition [metal + non metal]

Ionic bonding: strong electrostatic force of attraction between oppositely charged ions

1.33 Properties of ionic compounds + explanations

High M.P. + B.P.

Very strong electrostatic forces of attraction (ionic bonds) require large amounts of heat energy to be broken

Hard

Very strong electrostatic forces of attraction between ions require large amounts of energy to be broken

Non-Conductors in SOLID state

Ions held in their fixed positions by strong ionic bonds not free to move around

Conductor in LIQUID (molten/aqueous) state

Ions free to move (strong ionic bonds broken) + act as mobile charge carriers

Soluble in water

Insoluble in organic solvents

High densities

Ions close together by strong ionic bonds

1.34 Relationship between ionic charge with M.P. + B.P. of ionic compound

The higher the charge of the ions in the lattice the stronger the ionic bonding the higher the M.P. + B.P.

1.35 the ionic crystal

Giant 3-D lattice structure held together by attraction between oppositely charged ions

1.36 Simple diagram representing positions of ions in crystal of sodium chloride

DRAW BOTH DIAGRAMS

g) Covalent substances

1.37 covalent bonding

Covalent bonding: 2/ more NON-METALLIC atoms share their unpaired outer shell electrons

Strong electrostatic force of attraction between the shared pair of electrons + the nuclei of the atoms

1.39 dot + cross digrams to show formation of covalent compounds by electron sharing

i. hydrogen

ii. chlorine

iii. Hydrogen chloride iv. water

v) methane vi) ammonia

vii) oxygen viii) nitrogen

ix) carbon dioxide x) ethane

xi) ethane

1.40 properties of covalent compounds

Low M.P. + B.P.

Held together by weak intermolecular forces of attraction can be broken easily with small amounts of heat energy

Non-Conductors of electricity in all states

Made up of neutral molecules no mobile charged particles to act are charge carriers

Soft

Weak intermolecular forces of attraction require small amounts of energy t be broken

Soluble in organic solvents

Insoluble in water

Relatively low densities

Weak intermolecular forces of attraction do not attract molecules close together

1.42 substances with giant covalent structures have high melting points

Giant molecular structure made up of millions of atoms covalently bonded in an ordered way in a 3-D arrangement

Large numbers of covalent bonds need to be broken large amount of heat energy

Allotropes: different crystalline forms of the same element which can exist in the same physical state

1.43 Diagrams of diamond & graphite atoms

Diamond

Very High M.P.

Large numbers of strong covalent bonds need to be broken

Non-conductor of electricity

4 outer shell electrons of each C atom used in bonding

no free electrons to act as charge carriers

Relatively high density

Strong covalent bonds pull atoms closer together

Very hard

used in cutting

Very strong covalent bonds hold carbon atoms together firmly

Graphite:

Very High M.P.

Large numbers of strong covalent bonds need to be broken

Good conductor of electricity

3 out of 4 outer shell electrons used in bonding

No free electrons to act as charge carriers

Relatively high density

Strong covalent bonds pull atoms in structure closer together

Lower than diamond contains van der Waals’ forces

Soft [lubricant]

Layers held by weak van der Waals’ forces which are easily broken layers can slide over

h) Metallic crystals

1.45 metal: giant structure of positive ions surrounded by a sea of delocalized electrons

1.46 properties of metals

High M.P. + B.P. (many exceptions Hg, Na, Pb)

Very strong metallic bonds large amounts of heat energy to be broken

Good conductors of electricity

Delocalized electrons free to move around + act as charge carriers transfer electric charge from one part to another

Malleable [hammered into thin sheets] + Ductile [drawn out into wires]

Layers of cations can slide over each other

i) Electrolysis

1.47 electric current: flow of electrons

1.48 + 1.49 See 1.33 + 1.40

1.50 distinguishing between electrolytes + non-electrolytes

Electrolyte: compound that is decomposed by electricity

Non-electrolyte: compound which does not conduct electric current since no ions present