The Perfect Answer
Revision Guide To…
Chemistry
Edexcel IGCSE
9-1
Triple Award
1st Edition
Copyright © 2019 Hazel Lindsey & Martin Bailey
Hazel Lindsey & Martin Bailey
For use by Abdul Basit back2basit@gmail.com ONLY. Not for redistribution.
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Contents
1. Principles of Chemistry
States of matter...................................................................................................................... 3
Elements, compounds and mixtures......................................................................................4
Atomic structure ..................................................................................................................... 5
The periodic table................................................................................................................... 6
Chemical formulae, equations and calculations .................................................................... 7
Ionic bonding..........................................................................................................................9
Covalent bonding ................................................................................................................... 9
Metalling bonding................................................................................................................. 11
Electrolysis ...........................................................................................................................11
2. Inorganic Chemistry
Group 1 (alkali metals) ..........................................................................................................13
Group 7 (halogens) ............................................................................................................... 14
Gases in the atmosphere .....................................................................................................15
Reactivity series ................................................................................................................... 16
Extraction and uses of metals.............................................................................................. 17
Acids, alkalis and titrations................................................................................................... 19
Acids, bases and salt preparations ......................................................................................20
Chemical tests......................................................................................................................22
3. Physical Chemistry
Energetics .............................................................................................................................25
Rates of reaction .................................................................................................................. 27
Reversible reactions and equilibria ......................................................................................27
4. Organic Chemistry
Introduction ..........................................................................................................................29
Crude oil ............................................................................................................................... 29
Alkanes................................................................................................................................. 31
Alkenes................................................................................................................................. 31
Alcohols................................................................................................................................ 32
Carboxylic Acids .................................................................................................................. 33
Esters ................................................................................................................................... 33
Synthetic polymers............................................................................................................... 34
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Note: Content in italics will not be examined on Paper 1
1. Principles of Chemistry
States of matter
Describe the structure of solids, liquids and gases
- solids - particles arranged regularly and packed closely together. Vibrate in fixed
positions. Strong forces between particles
- liquids - particles are mostly touching with some gaps. Particles move about at
random. Medium forces between particles
- gases - particles move at random and quickly. Particles are far apart. Weak forces
between particles.
Give the proper name for the following conversions:
- solid —> liquid (melting)
- liquid —> gas (boiling)
- gas —> liquid (condensing)
- liquid —> solid (freezing)
How does evaporation occur?
- particles have differing amounts of energy
- particles with the greatest amount of K.E (kinetic energy) break away from the surface
of the liquid
- the average K.E. of remaining particles is lowered
- in a closed container both evaporation and condensation occur simultaneously
Define diffusion
- net movement of particles
- from an area of high concentration to low concentration
e.g. ammonia and hydrogen chloride
- white ring forms closer to HCl end
- this tells you that NH3 diffuses faster (the reason being it has a lower Mr)
Define the following words:
- solvent - liquid in which a solute dissolves
- solute - a solid which dissolves in a solvent
- solution - mixture of a solute and a solvent
- saturated solution - a solution where no more solute can dissolve in the solvent
Define solubility
- the mass of solute which must dissolve in 100g of solvent at that temperature to form a
saturated solution
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Elements, compounds and mixtures
What is an element?
- a substance which contains one TYPE of atom only
- cannot be split into anything simpler by any chemical means
What is a compound?
- A substance made up of two or more elements chemically combined
What is a mixture?
- A substance made up of two or more elements NOT chemically bonded together
Examples of elements, compounds and mixtures:
Element
Compound
Mixture
iron
calcium carboate
honey
lead
ammonia
air
sulfur
carbon dioxide
sea water
nitrogen
water
blood
oxygen
iron sulfide
soup
What is a pure substance?
- contains one type of material only e.g. one type of element or molecule
Describe the melting and boiling points of pure substances
- fixed
Describe the melting and boiling points of mixtures
- may melt and boil over a range of temperatures
What method is used to separate an insoluble solute from a solvent?
- filtration
What method is used to separate a soluble solute from a solvent?
- evaporation or distillation
What method is used to separate liquids of different boiling points?
- fractional distillation
What method is used to separate pure water from sea water?
- simple distillation
What method is used to separate petrol and water and why is this used?
- separating funnel
- petrol and water are immiscible (don’t mix)
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What is paper chromatography used to separate?
- dyes/inks i.e. liquids with different solubilities
About chromatography:
- use a pencil line as reference line (pencil doesn’t move)
- add dots of ink
- dip filter paper in water
- furthest dot has the greatest solubility
How do you calculate the Rf value?
- Rf = distance travelled by component
distance travelled by solvent
Atomic structure
What is an atom?
- the smallest particle of a chemical element that can exist
What is a molecule?
- two or more atoms chemically bonded together
Protons, neutrons and electrons
Proton
Neutron
Electron
Relative charge
1
0
-1
Relative mass
1
1
1/1836
What is the atomic number?
- the number of protons
What is the mass number?
- the total number of protons and neutrons
What is the nucleon number
- same as mass number i.e. total number of protons and neutrons
What is an isotope?
- atoms of the same element with the same number of protons but different number of
neutrons
What is relative atomic mass?
- the ratio of the average mass of one atom of an element compared with one atom of
carbon-12
How do you calculate the relative abundance of a particular isotope?
- (% of isotope 1 × mass of isotope 1) + (% of isotope 2 × mass of isotope 2) ÷ 100
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Example calculation: a naturally occurring sample of the element chlorine contains 75% of
Cl-35 and 25% of Cl-37. Calculate the relative abundance of chlorine.
= (75 x 35) + (25 x 37)
100
= 35.5
The periodic table
Give the electronic configurations of sodium, chlorine, oxygen, and magnesium.
- Na = 2.8.1
- Cl = 2.8.7
- O = 2.6
- Mg = 2.8.2
What is the period number of the periodic table?
- tells you the number of shells of electrons
- e.g. Ca 2.8.8.2 has four shells of electrons and is therefore in period 4
What is the group number of the periodic table?
- tells you the number of electrons in the outer shell
- e.g. F has 7 electrons in its outer shell and is therefore in group 7
Why do elements in the same group have similar chemical properties?
- same number of electrons in outer shell
Why are Noble gases (group 0) unreactive?
- they have a full outer shell of electrons
Describe the layout of the periodic table
- hydrogen by itself
- metals found left of stepped line
- non-metals found right of stepped line
Define malleable
- may be hammered into shape
Define ductile
- may be drawn into a wire
Describe the properties of metals
- good conductors of heat and electricity
- shiny
- malleable
- sonorous
- malleable
- ductile
- form positive ions in ionic compounds
- partake in ionic bonding
- form basic oxides
- solids at room temperature (except mercury)
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Describe the properties of non-metals
- poor conductors of heat and electricity
- brittle
- form negative ions in ionic compounds
- partake in ionic and covalent bonding
- form acidic oxides
Chemical formulae, equations and calculations
Define the mole
- amount of a substance that contains same number of units as the number of carbon
atoms in 12g of carbon-12
What is the Avogadro constant?
- the number of elements/molecules in a mole
- 6.02 x 1023
How do you calculate the number of moles in a given mass?
- number of moles = mass / mass of 1 mole (Mr)
Example questions using moles, mass and Mr
1. Find the Mr of MgCO3
= 24+12+(3x16)
= 84
2. Find the mass of 0.2 moles of CaCO3
mass = Mr x moles
mass = 0.2 x (40+12+(3x16)
=20g
3. Find the number of moles in 54g of H2O
number of moles = Mass/Mr
moles = 54/18
moles = 3
4. Find the empirical formula of a compound which contained 5.85g K, 2.10g N and 4.8g O
K
N
O
Mass
5.85
2.1
4.8
Mr
39
14
16
Moles
0.15
0.15
0.3
divide by smallest
number
0.15/0.15 = 1
0.15/0.15 = 1
0.3/0.15 = 2
ratio
1
1
2
Answer = KNO2
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How do you calculate the percentage by mass of element in a compound?
= total mass of element x100
mass of compound
5. Find the percentage by mass of magnesium in magnesium oxide
= 24
x 100
(24+16)
=60%
6. 4.5g of hydrochloric acid, HCl, reacted with calcium hydroxide, Ca(OH)2. Calculate the
mass of Calcium Chloride, CaCl2, formed.
2HCl + Ca(OH)2 —>
CaCl2
Mass
4.5
x
Mr
1+35.5 = 36.5
40+(35.5x2) = 111
Moles
0.123287…
0.06164….
Answer —>
x= 6.84
+
H2O
7. Calculate the amount, in moles, of 25cm3 of HCl with a concentration of 2 mol/dm3
- 25cm3 = 0.025dm3
- moles = concentration x volume
- moles = 2 x 0.025
- moles = 0.05
What is the volume that one mole of gas occupies?
- 24 dm3 (at room temperature and pressure)
- = 24000 cm3
8. Calculate the volume in cm3 of 3 mol of O2
- Volume = moles x 24000
- Volume = 3 x 240000
- Volume = 72000cm3
What is the equation for calculating percentage yield?
- percentage yield = amount of product produced
x100
maximum amount of product possible
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Ionic bonding
How is an ion formed?
- an atom either loses (to form a positive ion) or gains electrons (to form a negative ion)
Working out the charge on an ion, remember:
1. for groups 1-3, the charge on the ion is the same as the group number e.g. Mg is in
group 2 and therefore forms Mg2+
2. for groups 5,6 and 7, the charge on the ion is 8 - group number. e.g. N is in group 5
and therefore 8-5=3, therefore N3Learn the following ions off by heart:
- Ag+, Cu2+, Fe2+, Fe3+, Pb2+, Zn2+, H+, OH-, NH4+, CO32-, NO3-, SO42What is an ionic bond?
- the electrostatic forces of attraction between oppositely charged ions
Why do ionic structures have such high melting points?
- strong forces of attraction between oppositely charged ions
- requires lots of energy to break
Why don’t ionic substances conduct when solid?
- ions are held tightly in fixed positions
- not free to move
Why do ionic substances conducts when molten/dissolved?
- ions are free to move
Covalent bonding
What is a covalent bond?
- basic definition: a pair of electrons shared between two atoms
- detailed definition: strong electrostatic forces of attraction between nuclei (positively
charged) and shared pair of electrons (negatively charged)
Why do simple molecular substances have low melting points?
- weak intermolecular forces of attraction
- do not require a lot of energy to break
Why does the boiling point of simple molecular substances increase with increasing
relative molecular mass?
- boiling breaks the intermolecular forces of attraction between molecules
- substances with greater Mr have greater intermolecular forces of attraction which need
breaking
- therefore more heat energy is needed to overcome these forces
Why don’t simple molecular substances conduct electricity?
- no overall electric charge
- no free electrons
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What is an allotrope?
- different forms of the same element
Give 3 allotropes of carbon
- diamond, graphite, C60 fullerene
Why does diamond have such a high melting point?
- giant covalent structure
- each carbon atom is bonded to 4 others
- many strong covalent bonds
- require lots of energy to break
Why don’t most covalent substances conduct electricity?
- no free electrons
- each electron in outer shell is bonded
Why doesn’t diamond conduct electricity?
- no free electrons
- each electron in outer shell is bonded
Why does graphite have such a high melting point?
- many strong covalent bonds
- require lots of energy to break
Why does graphite conduct electricity?
- each carbon atom is only bonded to 3 others
- 4th electron free to move
Why is graphite used as lubricant?
- Carbon atoms are arranged in layers
- layers are held together by weak intermolecular forces
- do not require a lot of energy to break
- layers slide off each other
Why does C60 fullerene have a lower melting and boiling point than graphite and diamond?
- simple molecular structure
- weak intermolecular forces
- require little energy to break
Why does C60 fullerene not conduct electricity?
- although each carbon atom is only bonded to 3 other, the 4th electron is not free to
move as it stays within each C60 molecule
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Metalling bonding
What is a metallic bond?
- electrostatic forces of attraction between positive metal ions and delocalised electrons
Why are metals good conductors of heat?
- delocalised electrons
- as electrons move around in the metal, heat energy is transferred throughout the
structure
Why do metals conduct electricity?
- delocalised electrons
- free to move
Why do metals have high melting and boiling points?
- strong electrostatic forces of attraction
- require a lot energy to break
Describe the structure of a metal
- postive ions
- delocalised electrons
- giant lattice
Why are metals malleable?
- layers of ions can slide over each other
Why are metals ductile?
- layers of ions can slide over each other
Electrolysis
Define electrolysis
- the breaking down of a substance using electricity
What sort of substances undergo electrolysis?
- giant ionic structures
Why does the electrolyte need to be molten?
- so the ions are free to move
What is an ion?
- charged particle - an atom which has either gained or lost electrons
What is an anion?
- negative ion
What is a cation?
- positive ion
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Use PANC to help you:
- Positive Anode
- Negative Cathode
Electrolysis rules:
- positive electrode attracts negative ions
- negative electrode attracts positive ions
- the least reactive element forms at the negative electrode i.e. jewellery metals or
hydrogen
- halogens form before all other elements at the positive electrode
What are the electrodes made out of?
- inert substances e.g. graphite, platinium
What does oxidation and reduction mean? (Use OIL RIG to help you)
- oxidation - loss of electrons
- reduction - gain of electrons
Predict the products at the cathode and anode of the following electrolysis reactions:
1. molten lead (II) bromide
Pb2+ + 2e- —> Pb (cathode - reduction)
2Br- - 2e- —> Br2 (anode - oxidation)
2. molten aqueous sodium chloride
2H+ + 2e- —> H2 (cathode - reduction)
2Cl- - 2e- —> Cl2 (anode - oxidation)
More information:
- hydrogen gas forms at the negative electrode because hydrogen is less reactive than
sodium. Hydrogen is used as a fuel.
- chlorine forms at the positive electrode because it is a halogen. Chlorine is used to
make bleach and to kill pathogens in swimming pools.
- sodium hydroxide is left over in the solution and is used in making paper and bleach
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2. Inorganic Chemistry
Group 1 (alkali metals)
What is the name given to group 1 elements?
- alkali metals
Describe the physical properties of group 1 elements
- soft (can be cut with a knife)
- low melting and boiling points
- low density (float on water)
- shiny (tarnish when exposed to air)
How should group 1 elements be stored and why?
- in oil
- very reactive
Describe the chemical properties of group 1 elements
- react with oxygen to form metal oxides e.g. Li2O
- form ionic compounds e.g. NaCl
- react with halogens e.g. KCl
- react with water to form metal hydroxides e.g. NaOH
Why do group 1 elements have similar chemical properties?
- they all have 1 electron in their outer shell
Describe the observations when group 1 metals are added to water
- float, move, fizz (releasing hydrogen), melt, turn UI blue (due to release of OH-)
- K - lilac flame, Na - orange flame
What is the word equation for when group 1 metals are added to cold water?
- metal + water —> metal hydroxide + hydrogen
What is the balanced symbol equation for when sodium (for example) is added to cold
water?
- 2Na + 2H2O —> 2NaOH + H2
Why do group 1 elements become more reactive down the group?
- elements are larger
- outer shell electron is further from the nucleus and more shielded
- electron more easily lost
What are the reactions of group 1 metals with air?
- lithium burns with a red flame to form lithium oxide
- sodium burns with a yellow flame to form sodium oxide
- potassium burns with a lilac flame to form potassium oxide
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Predict the properties of Francium
- as with other group 1 elements, Francium is soft, has low melting and boiling point, low
density, shiny (tarnishes when exposed to air)
- it’s near the bottom of the periodic table which means its reactions will be more violent
Group 7 (halogens)
States at room temperature:
Physical state at room
temperature
Colour at room temperature
fluorine
gas
yellow
chlorine
gas
green
bromine
liquid
red-brown liquid
iodine
solid
grey solid, purple vapour
Describe the reactions of halogens with hydrogen
- hydrogen halides formed e.g. H2(g) + Br2(g) —> 2HBr(g)
- hydrogen halides are acidic and highly poisonous
- very soluble in water e.g. HCl(g) —> HCl(aq)
Halogen displacement reactions:
KCl
KBr
KI
Cl
x
reaction
reaction
Br
no reaction
x
reaction
I
no reaction
no reaction
x
Displacement summary equations (more reactive halogens displace less reactive
elements from their compounds)
E.g.
- Cl2 + 2KBr —> 2KCl + Br2
- Br2 + 2KI —> 2KBr + I2
Ionic equations of halogen displacement reactions:
E.g.
- Cl2 + 2KBr —> 2KCl + Br2
- Cl2 + 2K+ + 2Br- —> 2K+ + 2Cl- + Br2
- Cl2 + 2Br- —> Br2 + 2ClDescribe the properties of halogens
- low melting points and boiling points
- poor conductors of heat energy and electricity
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Why do group 7 elements become less reactive down the group?
- elements are larger
- outer shell electron is further from the nucleus and more shielded
- harder to gain an electron
Gases in the atmosphere
Name the gases present in the air and their approximate percentage by volume
- Nitrogen 78%
- Oxygen 21%
- Argon 0.9%
- Carbon dioxide 0.04%
Describe how copper can be used to show the percentage of oxygen in the air
- copper is put in a large tube that is attached to two gas syringes
- 1 syringe contains 100cm3 of air
- 1 syringe empty
- copper heated strongly
- air passed over copper using syringes
- oxygen reacts with copper
- volume of air decreases to 80cm3
Describe how iron can be used to show the percentage of oxygen in the air
- iron filings placed in a burette that is full of air
- end of burette placed in trough of water
- iron reacts with oxygen
- water moves into burette
Describe the observations when magnesium burns in oxygen
- bright, white light
- white solid forms
- 2Mg(s) + O2(g) —> 2MgO(s)
- magnesium oxide dissolves in water to form an alkaline solution
- MgO(s) + H2O(l)—> Mg(OH)2(aq)
Describe the observations when sulfur burns in oxygen
- blue flame
- poisonous, colourless sulfur dioxide gas forms
- S(s) + O2(g) —> SO2(g)
- sulfur oxide dissolves in water to form acidic solution of sulfurous acid
- SO2(g) + H2O(l) —> H2SO3(aq)
Describe the observations when hydrogen burns in oxygen
- pale blue flame
- water forms
- 2H2(g) + O2(g)—> 2H2O(l)
Define thermal decoposition
- breaking down of a substance using heat
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What type of reaction occurs when metal carbonates are heated?
- thermal decomposition
Describe the thermal decomposition of copper carbonate
- copper carbonate is a green solid
- decomposes to form copper oxide - a black solid
- CuCO3(s) —> CuO(s) + CO2(g)
Write word and symbol equations for the reactions that occur when copper(II) carbonate
and calcium carbonate are heated strongly
- Calcium carbonate —> calcium oxide + carbon dioxide
- CaCO3 —> CaO + CO2
- copper(II) carbonate —> copper (II) oxide + carbon dioxide
- CuCO3—> CuO + CO2
Explain the effect of carbon dioxide on the environment
- greenhouse gas —> contributes to global warming
- polar ice caps melt
- sea levels rise
- floods low lying land
- loss of biodiversity
Reactivity series
What is the reactivity series?
- a list of metals in order of their reactivity with the most reactive metals at the top and
the least reactive at the bottom
How can you determine the reactivity of a metal?
- place in cold water (most will not react) - those which react are the most reactive
metals
- if no reaction, test with steam
- if no reaction, test with acid (note: only metals above hydrogen in the reactivity series
will react with hydrogen)
Reactivity series order:
- potassium
- sodium
- lithium
- calcium
- magnesium
- aluminium
- carbon
- zinc
- iron
- hydrogen
- copper
- silver
- gold
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What conditions are needed for rusting?
- water
- oxygen
What methods are used to prevent iron from rusting?
- barrier methods e.g. paint, grease, oil
- galvanising (sacrificial protection)
What is sacrificial protection?
- a method used to stop iron from rusting
- iron is coated in a more reactive metal which undergoes oxidation in preference to iron
Describe the process of galvanising
- coating iron in a more reactive metal e.g. zinc
- zinc reacts forming Zn2+
- it’s impossible for Fe to form Fe3+ as the Zn donates electrons
Define oxidation
- gain of oxygen
- loss of electrons
Define reduction
- loss of oxygen
- gain of electrons
Define redox
- a reaction where both reduction and oxidation take place at the same time
Define oxidising agent
- something which causes another substance to be oxidised i.e. gain oxygen/lose
electrons
- an oxidising agent is itself reduced (it loses oxygen/gains electrons)
Define reducing agent
- something which causes another substance to be reduced i.e. lose oxygen/gain
electrons
- a reducing agent is itself oxidised (it gains oxygen/loses electrons)
Extraction and uses of metals
How are un-reactive metals obtained?
- found native i.e. exist naturally in Earth’s crust
Why is iron obtained by reduction with coke (carbon) in the Blast Furnace?
- iron is less reactive than carbon
Why can’t aluminium be extracted using the Blast Furnace?
- aluminium is more reactive than carbon
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Why is electrolysis not used to extract iron?
- electrolysis is very expensive
- high electricity demands
What is an alloy?
- a mixture of metal with either other metals or carbon
Why do alloys tend to be harder than individual metals?
- ions of different sizes
- more difficult for layers to slide over each other
Why is the manufacture of aluminium such an expensive process?
- high cost of electricity
- replacement of carbon electrodes which burn away due to reaction between carbon
and oxygen producing carbon dioxide
Summary of equations of the electrolysis of Aluminium:
- at the negative electrode (reduction): Al3+ + 3e- —> Al
- at the positive electrode (oxidation): 2O2- - 4e- —> O2
Give some uses of aluminium and link them with an appropriate property of aluminium
- aeroplanes (low density)
- electricity cables (good conductor of electricity, ductile)
- saucepans (good conductors of heat, malleable)
Describe the uses and properties of low-carbon steel
- contains 0.25% carbon
- hard, strong, malleable, ductile
- uses: nails, car bodies, bridges, ship building
What are the disadvantages of low-carbon steel?
- high density means it’s heavy
- rusts easily if exposes to oxygen and water
Describe the uses and properties of high-carbon steel
- contains 0.6 - 1.2% carbon
- harder and more resistant to wear the low-carbon
- brittle
- uses: cutting tools e.g. knives
Describe the uses and properties of stainless steel
- contains iron, chromium and nickel
- oxide layer prevents corrosion
- uses: sinks, saucepans, cutlery, gardening tools
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Acids, alkalis and titrations
Types of indicator:
indicator
colour in acid
colour in alkaline
methyl orange
red
yellow
phenolphthalein
colourless
pink
litmus
red
blue
What is universal indicator?
- indicator used to determine the pH of a solute
- green = neutral
- red = strongly acidic
- purple = strongly alkaline
Describe the pH scale
- ranges from 0 -14
- 0-3 - strongly acidic (red)
- 4-6 - weakly acidic (orange)
- 7- neutral
- 8-10 weakly alkaline (blue)
- 11-14 - strongly alkaline (purple)
What ion is responsible for making something acidic?
- H+
What ion is responsible for making something alkaline?
- OHHow do you carry out a titration?
- use pipette to add alkali to conical flask
- add indicator to flask e.g. methyl orange, phenolphthalein
- place on white tile
- use burette to add acid to the conical flask
- add acid to conical flask drop-wise towards the end
- swill contents of conical flask to mix
- record volume of acid that caused colour change
Define acid
- H+ donor
Define base
- H+ acceptor
- OH- donor
- e.g. metal carbonate, hydroxide, oxides, ammonia
Define alkali
- soluble base
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Acids, bases and salt preparations
What is a salt?
- an ionic compound formed by the neutralisation of an acid by a base
Salts background:
- all acids contain hydrogen e.g. HCl, H2SO4 and HNO3
- when hydrogen in an acid is replaced with a metal or ammonium a salt is formed
- e.g. Magnesium sulfate, zinc chloride, ammonium chloride and potassium nitrate
Examples of acid and the salts they form:
Acid
Formula
Example of salt
Name of salts
Hydrochloric acid
HCl
KCl
Chlorides
Nitric acid
HNO3
NaNO3
Nitrates
Sulfuric acid
H2SO4
MgSO4
Sulfates
Ethanoic acid
CH3COOH3
CH3COONa
Ethanoates
Phosphoric acid
H3PO4
Li3PO4
Phosphates
Describe the reactivity of metals with acids
- metals below hydrogen in the reactivity series don’t react with acids
- Metals above hydrogen in the reactivity series react with acids producing hydrogen
- Note: very reactive metals react very explosively with acids e.g. Potassium
Metal + acid —> salt + hydrogen
- E.g. magnesium + hydrochloric acid —> magnesium chloride + hydrogen
Metal hydroxide + acid --> salt + water
- E.g. Sodium hydroxide + hydrochloric acid --> sodium chloride + water
Metal oxide + acid --> salt + water
- E.g. magnesium oxide + hydrochloric acid —> magnesium chloride + water
Metal carbonate + acid —> salt + water + carbon dioxide
- E.g. potassium carbonate + nitric acid —> potassium nitrate + water + carbon dioxide
Ionic equation summarising this reaction H+ + OH- --> H2O
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Remembering which salts are soluble/insoluble:
- all nitrates are soluble
- all sulfates are soluble except lead (ii) sulfate, barium sulfate, and calcium sulfate
- all chlorides are soluble except lead (ii) chloride and silver chloride
- all carbonates are insoluble except ammonium, potassium and sodium salts
- all hydroxides are insoluble except ammonium, potassium and sodium salts
How do you make soluble salts (except ammonium, potassium and sodium salts)?
For the reactants, you can use:
- acid + metal oxide/hydroxide/carbonate
- acid + metal (not too reactive a metal though!)
Use the crystallisation method
- REACT
- FILTER
- EVAPORATE: heat to evaporate some water
- COOL: collect crystals that form
- DRY: allow the crystals to dry in a warm place or on filter paper
How do you make soluble salts - that do contain sodium, potassium or ammonium?
For the reactants, you can use:
- acid + metal hydroxide/carbonate
- acid + ammonia solution
Why can’t you use the crystallisation method above?
- because sodium, ammonium and potassium compounds are soluble in water
- this means that when added to acid they would react both with the acid and water and
as such the salt would constantly dissolve
- there would be no visible excess to filter off
Instead, use the titration method:
- REACT: an acid (from a burette) with an alkali (in a conical flask)
- INDICATOR: requires an indicator to show when the alkali has been neutralised (all alkali
has been reacted)
- REPEAT: once amounts required have been worked out, add required volumes of acid
to alkali without indicator
- EVAPORATE: heat to evaporate some water, this concentrates the solution
- COOL: collect crystals that form
- DRY: allow the crystals to dry in a warm place or on filter paper
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How do you make insoluble salts?
For the reactants, you can use:
- 2 soluble salts mixed together to form an insoluble salt and a soluble one
- e.g. to make insoluble silver chloride, mix together silver nitrate and sodium chloride
Use the precipitation method
- REACT
- FILTER
- WASH
- DRY: allow the precipitate to dry in a warm place or on filter paper
Give an explanation of what happens in precipitation reactions using the reaction of sliver
nitrate and sodium chloride as an example.
- in silver nitrate solution the Ag+ and NO3- are attracted weakly
- in the sodium chloride the Na+ and Cl- are attracted weakly
- when you mix the 2 solutions, the 4 ions mix together
- Ag+ and Cl- attract strongly forming AgCl
- Na+ and NO3- attract weakly and remain in solution
What is the Brønsted-Lowry theory?
- states that a base is a H+ acceptor
- states that an acid is a H+ donor
Provide an example of the Brønsted-Lowry theory in context
- dissolving hydrogen chloride in water
- H2O(l) + HCl(aq) —> H3O+(aq) + Cl-(aq)
- HCl donates H+ (acts as an acid)
- H2O accepts H+ (acts as a base)
What is a hydroxonium ion?
- H3O+
Chemical tests
What is the test for hydrogen?
- lighted splint pops
What is the test for oxygen?
- glowing splint relights
What is the test for carbon dioxide?
- limewater turns cloudy
What is the test for chlorine?
- bleaches damp blue litmus paper
What is the test for ammonia?
- damp red litmus paper turns blue
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How do you carry out a flame test?
- dip nichrome wire in hydrochloric acid to clean
- dip in sample
- hold in roaring blue Bunsen flame
Results:
- lithium (Li+) - red
- potassium (K+) - lilac
- sodium (Na+) - yellow
- calcium (Ca2+) - orange-red
- copper (Cu2+) - blue green
How do you test for copper (Cu2+), iron (II) (Fe2+) and iron (III) (Fe3+)?
- use precipitation reactions - add sodium hydroxide
Cu2+ Copper hydroxide - light blue precipitate
Fe2+ Iron(II) hydroxide - green precipitate
Fe3+ Iron(III) hydroxide - brown precipitate
Detecting ammonium ions (NH4+)
- add sodium hydroxide solution, no precipitate is formed, but smell of ammonia is given
off
- test with damp red litmus paper, it should turn blue
Detecting halides:
- add dilute nitric acid (to remove carbonate ions)
- add silver nitrate
Results of adding silver nitrate to halides:
- silver chlorides - white precipitate
- silver bromides - cream precipitate
- silver iodides - yellow precipitate
- e.g. Ag+(aq) + Br-(aq) —> AgBr(s)
Detecting sulfates:
- add dilute hydrochloric acid (removes carbonate ions)
- add barium chloride
- result: barium sulfate is a white precipitate
- Ba2+(aq) + SO42-(aq) —> BaSO4(s)
Detecting Carbonates:
- add dilute nitric acid
- fizzing indicates carbon dioxide
- test for carbon dioxide using lime water (turns milky/cloudy)
- 2H+(aq) + CO32-(aq) —> CO2(g) + H2O(l)
What is the chemical test for water?
- white anhydrous copper sulfate turns blue
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How do you do a physical test for water?
- check boiling point
- water boils at 100C
How do you show that water is pure?
- check boiling point
- pure water has single boiling point at 100C
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3. Physical Chemistry
Energetics
Define exothermic
- heat energy is released
- more energy is needed to make the bonds than break the bonds
- ΔH is negative
Define endothermic
- heat energy is taken in
- less energy is needed to make the bonds than break the bonds
- ΔH is positive
Define activation energy
- the minimum amount of energy required for a reaction to occur
Define bond energy
- the energy needed to break the bond between two atoms
Describe a simple experiment to investigate temperature change during a combustion
reaction
- measure cold water into a copper calorimeter
- record starting temperature of water
- heat water using flame from burning fuel
- record final temperature of water
Describe a simple experiment to investigate temperature change during a displacement
reaction
- measure volume of copper sulfate solution and place in insulated calorimeter (e.g.
polystyrene cup)
- record initial temperature of copper sulfate solution
- add measured mass of zinc powder to copper sulfate solution
- measure temperature change in copper sulfate solution
Describe a simple experiment to investigate temperature change during a dissolving
reaction
- measure volume of water and place in insulated calorimeter (e.g. polystyrene cup)
- record initial temperature of water
- add measured mass of solid solute to water
- when the solid has dissolved, measure temperature change in water
Describe a simple experiment to investigate temperature change during a neutralisation
reaction
- measure volume of acid and place in insulated calorimeter (e.g. polystyrene cup)
- place polystyrene cup into beaker
- measure temperature of the acid
- add measured volume of alkali to the acid
- record temperature rise to calculate energy change.
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What is specific heat capacity?
- the quantity of energy, (in Joules), needed to change the temperature of one gram of a
substance by one degree Celsius
How do you calculate heat energy change from a measured temperature?
- heat energy change = mass x specific heat capacity x change in temperature
- Q = mcΔT
Example: Calculate the specific heat capacity of an alloy if a 13.7 g sample absorbs 382 J
when it is heated from 0.0C to 38.1C
- m = 13.7g
- Q = 382J
- ΔT = 38.1 - 0.0 = 38.1
- 382 = 13.7 x c x 38.1
- c = 382/(13.7x38.1)
- 0.731 J/gC
How do you calculate molar enthalpy change (ΔH) from the heat energy change (Q)?
- molar enthalpy change = heat energy change / number of moles
- ΔH = Q/n
Example: 0.674g of C2H6 reacts with O2 to from CO2 + H20. The reaction heated 100cm3
of water from 24C to 33C. How much energy would be obtained from one mole of C2H6?
- C2H6(l) + 3.5O2(g) —> 2CO2(g) + 3H2O(l)
- Q = mcΔT
- q = 100 x 4.18 x 9
- = 3762J
- Moles = mass/Mr
- = 0.674/30
- =0.0224 mol
- ΔH = Q/n
- = 3762/0.0224
- = -167448 J/mol (-167.4 kJ/mol)
- negative because it is an exothermic reaction
Describe the type of process when bonds are broken
- endothermic
Describe the type of process when bonds are formed
- exothermic
How do you calculate enthalpy change during a chemical reaction?
- enthalpy change = ΔH bonds broken - ΔH bonds made
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Rates of reaction
What effect does high temperature have on the rate of reaction?
- collisions occur more frequently
- particles have greater kinetic energy
- particles collisions are harder and more successful
What effect does increasing the concentration have on the rate of reaction?
- more particles in the same volume
- collisions occur more frequently
What effect does increasing the surface area have on the rate of reaction?
- collisions occur more frequently
Define catalyst
- a substance that increases the rate of reaction
- chemically unchanged at the end of the reaction
How does a catalyst work?
- speeds up the rate of reaction without being used up
- offers an alternative reaction pathway with lower activation energy
Reversible reactions and equilibria
What does this
arrow represent?
- reaction is reversible
Describe the dehydration of copper (II) sulfate crystals
- heat blue (hydrated) copper (II) sulfate crystals
- blue crystals turn to white powder because water is lost i.e. water of crystallisation is lost
- anhydrous copper (II) sulfate is formed
Describe the heating of ammonium chloride
- when you heat ammonium chloride it splits into ammonia and hydrogen chloride
- the white crystals disappear from the bottom of the tube and appear higher up
- reaction reverses when conditions change from hot to cold
What is a dynamic equilibrium?
- forward and reverse reactions occur at the same rate
- the concentration of reactants and products remains the same
What effect does the addition of a catalyst have on the position of equilibrium?
- catalyst has no effect on the position of equilibrium
- because the rates of the forward and reverse reactions increase equally
Describe and explain the effect of increasing the temperature on a reversible reaction
- increasing the temperature favours the endothermic reaction
- position of equilibrium shifts in favour of the endothermic reaction
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Describe and explain the effect of decreasing the temperature on a reversible reaction
- decreasing the temperature favours the exothermic reaction
- position of equilibrium shifts in favour of the exothermic reaction
Describe and explain the effect of increasing the pressure on a reversible reaction
- increasing the pressure favours the side with fewer moles of gas
- position of equilibrium shifts to the side with fewer moles of gas
Describe and explain the effect of decreasing the pressure on a reversible reaction
- decreasing the pressure favours the side with more moles of gas
- position of equilibrium shifts to the side with more moles of gas
Example:
2NO2(g)
brown
N2O4(g)
delta H = -57KJ/mol
colourless
Describe and explain the effect of increasing the temperature on the colour of the reaction
vessel
- increasing the temperature favours the endothermic (reverse reaction)
- position of equilibrium shifts to the left
- mixture turns brown
Describe and explain the effect of increasing the pressure on the colour of the reaction
vessel
- increasing the pressure favours the side with fewer moles of gas (the right hand side)
- position of equilibrium shifts to the right
- mixture turns colourless
Why are the conditions used in the Haber Process described as ‘compromised’?
N2 + 3H2
2NH3
-
forward reaction is exothermic and therefore favoured by low temperatures
but rates of reaction are too slow at low temperature so 450C temperature is used
the forward reaction results in fewer moles of gas so is favoured by high pressures
but high pressures are dangerous and expensive so 200 atmospheres is used
What effect does the iron catalyst have on the yield of ammonia?
- no effect on yield but increases the rate of the forward and reverse reactions equally
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4. Organic Chemistry
Introduction
What is a hydrocarbon?
- a compound which contains hydrogen and carbon atoms ONLY
What does molecular formulae mean?
- the exact number of atoms of each element present in a compound
What does empirical formulae mean?
- the simplest ratio of atoms of each element present in a compound
What does displayed/structural formulae mean?
- a drawing of the bonds within a compound
What is a homologous series?
- a group of compounds with the same chemical properties because they have the same
functional group e.g. alcohols -OH
What do members of the same homologous series have in common?
- same chemical properties
- trend in physical properties
- same functional group
- same general formula
What is a functional group?
- an atom or group of atoms which determine the chemical properties of a compound
What does isomerism mean?
- compounds with same molecular formula but different displayed/structural formula
Crude oil
What is crude oil made up of?
- mixture of hydrocarbons
What is a fuel?
- a substance which releases energy when burnt
What is a fraction?
- group of substances with similar boiling points
How is crude oil separated into its various fractions?
- fractional distillation
- crude is heated
- crude oil boils and vaporises
- vapour passed into bottom of fractionating column
- column hottest at the bottom - longest chain fractions condense here e.g. bitumen
- column coolest at the top - shortest chain fractions condense here e.g. refinery gases
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What are the main uses of the following fractions:
- refinery gases (mixture of methane ethane, propane) - bottled gas
- gasoline - fuel for cars
- kerosene - fuel for planes
- diesel - fuel for buses, lorries
- fuel oil - fuel for ships
- bitumen - road surfacing
What does viscosity mean?
- how readily a liquid flows
- honey - very viscous
- water - not very viscous
Compare the colour, viscosity and boiling point of bitumen and refinery gases
- bitumen darker in colour, refinery gases are lighter in colour
- bitumen high boiling point, refinery gases low boiling point
- bitumen very viscous, refinery gases not viscous at all
What is complete combustion?
- plentiful oxygen
- produces carbon dioxide and water
What is incomplete combustion?
- insufficient oxygen
- produces carbon monoxide and water
Why is incomplete combustion a problem?
- carbon monoxide is made - toxic
- combines irreversibly with haemoglobin
- less oxygen transported in blood
How is acid rain formed?
- high temperatures found in car engines cause nitrogen to react with oxygen forming
nitrogen oxides —> dissolve in rain water —> nitric acid
- sulfur impurities in crude oil —> sulfur dioxide —> dissolve in rain water —> sulfuric acid
What is cracking?
- the breaking down of long alkane chains into smaller, more useful chains of alkanes and
alkenes
What reaction conditions are needed for cracking
- 600-700C
- alumina or silica catalyst
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Alkanes
Alkanes: Number of Carbon
Atoms
Name
1
methane
2
ethane
3
propane
4
butane
5
pentane
Displayed Formula
What is the general formulae of an alkane?
- CnH2n+2
Describe alkanes reactions with bromine water
- substitution reaction
- UV radiation required
Alkenes
Alkenes: Number of Carbon
Atoms
Name
1
n/a
2
ethene
3
propene
4
butene
5
pentene
Displayed Formula
What is the general formulae of an alkene?
- CnH2n
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What does unsaturated mean?
- contains C=C
What does saturated mean?
- all carbon bonds are single
What is the test for an unsaturated compound/alkene?
- add to bromine water
- orange colour turns colourless
- addition reaction
- e.g. C2H4 + Br2 —> C2H4Br2
- e.g. CH4 + Br2 —> CH3Br + HBr
Alcohols
What is the functional group of the alcohols?
- OH
Alcohols. Number of Carbon
Atoms
Name
1
methanol
2
ethanol
3
propanol
Displayed
Formula
How can alcohols be oxidised?
- burning in oxygen (complete combustion)
- reaction with oxygen in air to form ethanoic acid (microbial oxidation)
- heating with potassium dichromate (VI) in dilute sulfuric acid to form ethanoic acid
Give 2 uses of alcohol
- fuels, perfumes
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Manufacture of ethanol:
Comparison
Hydration of Ethene
Fermentation of sugar
Rate of reaction
High
Low
Type of Process
Continuous
Batch
Renewable
Non-renewable
Renewable
Temperature/Pressure
High temperature (300C) and
pressure
Low temperature (30C) and
pressure
Purity of Alcohol
Pure
Impure
Equation
CH2CH2 + H2O —> CH3CH2OH
C6H12O6 —> 2C2H5OH + 2CO2
(this is anaerobic respiration of yeast)
Carboxylic Acids
What is the functional group of the carboxylic acids?
- -COOH
Carboxylic acids. Number
of Carbon Atoms
Name
1
methanoic acid
2
ethanoic acid
3
propanoic acid
Displayed
Formula
What acid is found in vinegar?
- ethanoic acid
Esters
What is the functional group of the esters?
- -COOWhat reaction conditions are needed in the formation of esters?
- strong acid catalyst (sulfuric acid)
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What is the general equation for the formation of an ester?
- carboxylic acid + alcohol
ester + water
e.g.
- ethanoic acid +methanol
methyl ethanoate + water
What does volatile mean?
- evaporates readily
Give uses of esters and how they are linked to their properties
- food flavourings and perfumes because they are highly volatile
Synthetic polymers
Define monomer
- a small molecule that joins together to form a polymer
Define polymer
- a large molecule formed from many small molecules
What is an addition polymer?
- the joining up of many small molecules called monomers
What does biodegradable mean?
- can be broken down using microorganisms
What are the difficulties with the disposal of addition polymers?
- inert (unreactive)
- non biodegradable
- when burned give off toxic gases
Note: biopolyesters are biodegradable
What is condensation polymerisation?
- a dicarboxylic acid reacts with a diol producing a polyester and water
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