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
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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
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