AS level OCR Chemistry

Isotopes

Different atomic forms of the same element. Atoms with the same number of protons but a different number of neutrons

Relative atomic mass

The mean mass of an atom of an element, compared to 1/12th of the mass of an atom of carbon-12

Concentration=

Number of moles ÷ Volume

One mole

6.02×10^23

Number of moles=

mass ÷ molar mass

r.t.p

25 ºC
100 kPa (1 atm)

Molar gas volume

24 dm³ / mol
One mole of any gas always has the same volume at r.t.p

Number of moles=

Volume (dm^3) ÷ Molar gas volume (24 dm^3 / mol)

R- Gas constant

8.314 J / K / mol

Gas equation

pV = nRT
(Pa)(m³)(K)
Assumes forces between molecules are negligible and the molecules have a negligible size

Empirical formula

The smallest whole number ratio of atoms of each element in a compound

Molecular formula

The actual number of atoms of each element in a compound

Ions:
Nitrate
Carbonate
Sulfate
Hydroxide
Ammonium
Zinc ion
Silver ion

Formula:
NO3-
CO3²-
SO4²-
OH-
NH4+
Zn²+
Ag+

Acids

Proton donors - produce H+ ions in water

Alkalis

Proton acceptors - produce OH- ions in water

Acid + Base=
Metal oxide + Acid=
Metal hydroxide + Acid=

Salt + Water

Metal + Acid=

Metal salt + Hydrogen

Metal carbonate + Acid=

Metal salt + Carbon dioxide + Water

Ammonia + Acid=

Ammonium salt

Methyl orange

Yellow in alkali, red in acids

Phenolphthalein

Pink in alkali, colourless in acids

Oxidation number of oxygen

Nearly always -2, except in peroxides where it is -1 and 0 in molecular oxygen

Oxidation number of hydrogen

Nearly always +1, except in metal hydrides where it is -1 and 0 in molecular hydrogen

Sub-shells:
s
p
d
f

Orbitals:
1
3
5
7

Ionic bonding

The electrostatic forces of attraction between oppositely charged ions

Covalent bond

The electrostatic forces of attraction between a shared pair of electrons and the nuclei of the bonded atoms

Exceptions to covalent bonding

Boron trifluoride- 6 electrons in the outer shells
Sulfur hexafluoride- 12 electrons in the outer shells

Dative covalent bonding (coordinate bonding)

Both electrons from one atom

Shape of methane molecule

No lone pairs
Bond angle- 109.5º

Shape of ammonia molecule

1 lone pair
Bond angle- 107º

Shape of water molecule

2 lone pairs
Bond angle- 104.5º

Linear molecules

2 electron pairs around central atom
Bond angle- 180º

Trigonal planar

3 electron pairs around central atom
No lone pairs
Bond angle- 120º

Tetrahedral

4 electron pairs around central atom
No lone pairs
Bond angle- 109.5º

Pyramidal

4 electron pairs around central atom
1 lone pair included
Bond angle 107º

Nonlinear

4 electron pairs around central atom
2 lone pairs
Bond angle- 104.5º

Trigonal bipyramidal

5 electron pairs around central atom
No lone pairs
Bond angle- 120º, 90º

Octahedral

6 electron pairs around central atom
No lone pairs
Bond angle- All 90º

Electronegativity

An atom's ability to attract the electron pair in a covalent bond

Three types of intermolecular forces

Induced dipole-dipole
Permanent dipole-dipole interactions
Hydrogen bonding

Induced dipole-dipole

All atoms and molecules are attracted

Permanent dipole-dipole interactions

Weak electrostatic forces of attraction between polar molecules

Hydrogen bonding

Only possible when hydrogen is bonded to fluorine, nitrogen or oxygen

Periodic table blocks

First ionisation energy

The energy needed to remove one mole of electrons from one mole of gaseous atoms

Factors affecting ionisation energy

Nuclear charge
Atomic radius
Shielding

Ionisation energy drop between groups 2 and 3 is due to sub-shell structure

The outer electron in group 3 elements is in a p orbital rather than an s orbital.
A p orbital has a slightly higher energy than an s orbital in the same shell so the electron is further from the nucleus.
The p orbital has additional shielding provided by the s electrons which override the effect of increased nuclear charge.

Ionisation energy drop between groups 5 and 6 is due to p orbital repulsion

The repulsion between two electrons in a p orbital in group 6 elements means they are easier to remove than a singly-occupied p orbital in group 5 elements.

Carbon allotropes- high melting and boiling points, also insoluble

Diamond- Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral shape.
Silicon also forms a crystal lattice structure with similar properties as each silicon atom can form four covalent bond

Graphite- Each carbon atom forms three covalent bonds so there is one delocalised outer electron

Graphene- One layer of graphite in a hexagonal sheet, one atom thick. Transparent and incredibly light in one layer

Halogens (25ºC):
Fluorine- gas
Chlorine- gas
Bromine- liquid
Iodine- solid

Colour:
Pale yellow
Green
Red-brown
Grey

Test for halides

Silver nitrate solution
First add dilute nitric acid

Silver halides

Silver chloride- White precipitate, dissolves in ammonia
Silver bromide- Cream precipitate, dissolves in concentrated ammonia
Silver iodide- Yellow precipitate, insoluble in concentrated ammonia

Disporportionation

The same element is both oxidised and reduced

Halogen + alkali ---->
X2 + 2NaOH ---->
Disproportionation reaction

Metal halogen-ate + Metal salt + Water
NaXO + NaX + H2O

Chlorine and sodium hydroxide make bleach, Sodium chlorate(I)

2NaOH + Cl2 ----> NaClO + NaCl + H2O
0 +1 -1

Chlorine + Water ↔
Cl2 + H2O ↔

Hydrochloric acid + Chloric(I) acid
HCl + HClO

Chloric(I) acid + Water ↔
HClO + H2O ↔

Chlorate ion + Hydronium
ClO- + H3O+
Chlorate ions kill bacteria

Chlorine alternatives

Ozone (O3)- Strong oxidising agent, but is expensive and has a short half-life
Ultraviolet light- Damages DNA of microorganisms, but ineffective in cloudy water

Test for carbonates

First add dilute hydrochloric acid
If carbonate ions (CO3 2-) are present carbon dioxide will be released and will turn limewater cloudy

Test for sulfates

Barium chloride solution
First add dilute hydrochloric acid
A white precipitate (Barium sulfate) will form if sulfates are present

Test for ammonium compounds

Warm the mixture and add sodium hydroxide
If the damp red litmus paper turns blue, ammonia (alkali) is given off which means ammonium compounds are present.

Avoiding false positives

Test for carbonates ----> Test for sulfates ----> Test for halides

Enthalpy change (ΔH)

The heat energy transferred in a reaction at constant pressure
Unit- kJ / mol

Standard enthalpy change of reaction

The enthalpy change when the reaction occurs in the molar quantities shown in the chemical equation, under standard conditions

Standard enthalpy change of formation

The enthalpy change when one mole of a compound is formed from its elements in their standard states, under standard conditions

Standard enthalpy change of combustion

The enthalpy change when one mole of a substance is completely burned in oxygen, under standard conditions

Standard enthalpy change of neutralisation

The enthalpy change when an acid and alkali react together to form one mole of water, under standard conditions

Average bond enthalpy

The energy needed to break one mole of bonds in the gas phase, averaged over many different compounds

Enthalpy change, q (Joules)=

m×c×ΔT

Specific heat capacity

The amount of energy needed to raise the temperature of one gram of a substance by one kelvin

Specific heat capacity of water

4.18 J / g / K

Hess's law

The total enthalpy change is (always the same) independent of the route taken

Enthalpy change of reaction=

Total energy absorbed (bond breaking) - Total energy released (bond making)

Catalyst

A substance that increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy. The catalyst remains chemically unchanged

Production of ethanol

Ethene and steam are reacted
60-70 atmospheres
300ºC
phosphoric (V) acid catalyst

Equilibrium constant, Kc

When you have a homogenous reaction that's reached dynamic equilibrium, the larger the value of Kc the further the equilibrium lies to the right and vice versa

Structural formula

The arrangement of atoms carbon by carbon with attached hydrogen and functional groups

Skeletal formula

The bonds of the carbon skeleton only, with any functional groups

Homologous series

Same general formulas and functional groups

Carbon skeleteon

Aromatic or aliphatic
Aromatic compounds contain a benzene ring
Aliphatic compounds contain carbon and hydrogen joined in straight chains, branched chains or non-aromatic rings (alicyclic)

Isomers

Same molecular formula but different structural formula, the atoms are arranged differently
Types- Structural isomers and stereoisomers

Structural isomers

Chain isomers- The carbon skeleton can be arranged differently (straight or branched)
Similar chemical properties but different physical properties

Positional isomers- The functional group could be attached to a different carbon atom
different physical and maybe chemical properties

Functional group isomers- The same atoms can be arranged into different functional groups
Very different physical and chemical properties

Alkane molecule shape

Alkane molecules are tetrahedral around each carbon atom
Each carbon atom has four pairs of bonding electrons around it
Bond angle- 109.5º

Halogen + Alkane=

Haloalkane (photochemical reaction)
Free-radical substitution reaction
1) Initiation- Free radicals are produced
2) Propagation- Free radicals are used up and created
3) Termination- Free radicals are mopped up

Alkene double bond

Sigma bond- Two s orbitals overlap, C-C or C-H in alkanes (high bond enthalpy)
Pi bond- Sideways overlap of two adjacent p orbitals (low bond enthalpy)

E/Z isomerism

Stereoisomerism because of the lack of rotation around the carbon-carbon double bond

E-isomer

The same groups are across the double bond
Trans isomer

Z-isomer

The same groups are on the same side of the double bond
Cis isomer

Adding hydrogen to C=C bonds produces alkanes

Ethene will react with hydrogen in an electrophilic addition reaction to produce ethane
Nickel catalyst
150ºC

Halogens react with alkenes to form dihaloalkanes

Electrophilic addition
Orange bromine water decolourises when mixed with an alkene and forms a dibromoalkane

Alkenes undergo electrophilic addition with hydrogen halides

Two haloalkanes are formed
Markownikoff's rule- The major product is the one where hydrogen adds to the carbon with the most hydrogens already attached

Addition polymers

Alkenes (monomers) join up and the double bond is removed

Alcohols

Primary
Secondary
Tertiary

Halogen + Alcohol ---->

Haloalkane
The -OH is substituted by the halide
Acid catalyst required such as sulfuric acid

Dehydration

Alcohols can be dehydrated to form alkenes
Concentrated sulfuric acid or phosphoric acid
Heated

Oxidising alcohols

Oxidising agent- acidified potassium dichromate (VI)
Primary alcohols ----> Aldehydes (distill) and then carboxylic acids (reflux)
Secondary alcohols ----> Ketones
Tertiary alcohols will only oxidise by being burnt

Hydrolysis

Haloalkanes can be hydrolysed to make alcohols in a nucleophilic substitution reaction
Warm aqueous alkali solution (Sodium hydroxide)

CFCs (Haloalkanes)

Stable, volatile, non-flammable and non-toxic

Refluxing

Vertical Liebig condenser
Prevents loss of volatile substances
Continuously boils,evaporates and condenses the vapours

Seperation

Separating funnel with water added to mixture
Anhydrous salt (Magnesium sulfate) can be added to remove water after separation​