T1: Sub-atomic particles• Atoms are made from smaller particles called

subatomic particles.• There are three types we need to know about,

summarised below.

T1: Mendeleev• Arranged elements by increasing atomic mass

but….• He broke this rule and left some gaps if an

element’s properties weren’t similar to the one above it.

• He thought the gaps were for elements that hadn’t been discovered yet and predicted their properties.

• When they were discovered, the properties matched the predictions

PERIODS….increasing atomic mass, differing properties

GRO

UPS…

…sim

ilar properties

Element Type

= non-metal = metal

Particle Relative charge

Relative mass

Found?

Proton 1 Positive, +1 In nucleusNeutron 1 Neutral, 0 In nucleusElectron Neglible () Negative, -1 In shells orbiting

nucleus

T1: Reading the Periodic Table

• Note: on some periodic tables, they are the wrong way up, just remember that the smaller number is the proton number.

Relative Atomic Mass (aka nucleon number):

The total number of protons and neutrons added together.

Atomic number (aka proton number):

The number of protons or electrons.

T1: What’s in my atom?

Protons = atomic numberElectrons = atomic numberNeutrons = relative atomic mass . – atomic number

Atomic number = 9Relative Atomic mass = 19Protons = 9Electrons = 9Neutrons = 19-9 = 10

T1: Atoms and Elements•Element = substance containing only one type of atom.•Protons and electrons: same for every atom of an element…it is the number of protons that decides the element.•Neutrons: can differ…atoms with the same number of protons but different numbers of neutrons are called isotopes

T1: Relative Atomic Mass•This is the mass of an element relative to 1/12th the mass of 12C.•Element: substance containing only one type of atom.•Protons and electrons: same for every atom of an element…it is the number of protons that decides the element.•Neutrons: can differ…atoms with the same number of protons but different numbers of neutrons are called isotopes.

T1: Isotopes (HT)• Versions of an element with same atomic number

but different atomic mass.• Number of protons is the same, but number of

neutrons is different.• Relative Atomic Mass is average of the masses of the

isotopes, weighted by their relative abundance

• For example, Neon has three isotopes

• Relative atomic mass of Neon =

• This is why some atoms have a relative atomic mass with a decimal point.

T1: Electron Configuration• Electrons orbit the nucleus in shells.• First shell holds two electrons• Second and third shell hold 8 electrons• Note: the third shell can actually hold more,

but we won’t worry about this until A-level.

Example: SiliconAtomic number is 14, so it has 14 electrons.You build up electrons from the first shell outwards, so in this case: - First shell has 2 - Second shell has 8 - Third shell has 4

This can be written as: 2.8.4; or drawn as:

Neon Isotope Mass

Relative Abundance (%)

20 90.5

21 0.3

22 9.2

Note: Si is in period three and group four of the periodic table; it also has three electron shells and four electrons in the outer shell – this is no coincidence!

T2: Forming IonsCations are positive (cat…pussitive!) ionsThey are formed when atoms lose electrons.Metals form cations by losing the electrons in their outer shellsIn the example, aluminium loses its three outer-shell electrons to become Al3+…each lost electrons cause 1 ‘+’ charge.

Anions are negative ionsThey are formed when atoms gain electrons.Non-metals form anions by filling their outer shells.Name ends with ‘-ide’ to show it is a negative ion,In the example, oxygen gains two outer-shell electrons to become O2-, giving it 8 electrons in its outer shell.

3+

T2: Making Ionic Compounds• An ionic bond is the attraction between a

positive and a negative ion.• The overall number of positive and

negative charges must cancel out.• Form between a metal and a non-metal• Ionic compounds do not form molecules

Example 1: Magnesium reacting with chlorine.• Anion: Cl forms Cl- ions• Cation: Mg forms Mg2+ ions• Formula = MgCl2 • Why: two Cl- gives a 2- charge to balance

2+ from Mg2+. • Name: magnesium chloride

Example 2: aluminium reacting with oxygen.• Anion: O forms O2- ions• Cation: Al forms Al3+ ions• Formula = Al2O3 • Why: Two Al3+ gives a 6+ charge, three

O2- gives a 6- charge. • Name: aluminium oxide

T2: Ionic Structures (HT)•A repeating 3D lattice of positive and negative ions.•Strong electrostatic bonds between ions.

T2: Precipitates and Precipitation•When an insoluble salt is formed from the reaction of two soluble salts.•Goes cloudy as small particles of solid are made.•Predicting precipitates: simply choose a combination of soluble salts where you tell that if the ions swapped over you would get an insoluble salt: use the solubility table for help.•Example:

T2: Common Ions• You should try to memorise the ions formed

by various species:

• There are also some ‘compound’ ions made of more than one atom with an overall charge:

• Hydroxide: OH-

• Nitrate: NO3-

• Sulphate, SO42-

• Carbonate, CO32-

• Ammonium, NH4+

Group Electrons in outer shell

Ion formed

Examples

1 1 + Li+, Na+, K+

2 2 2+ Be2+, Mg2+, Ca2+

3 6 2- O2-, S2-

4 7 - F-, Cl-, Br-, I-

T2: Solubility• Soluble: a compound dissolves in a given liquid.• Insoluble: a compound does not dissolve.

Soluble in water In soluble in waterAll sodium, potassium, ammonium saltsAll nitratesMost chlorides Except: silver and lead

chloridesMost sulfates Except: lead, barium

and calcium sulfates. Except: sodium, potassium and ammonium carbonates

Most carbonates

Except: sodium, potassium and ammonium hydroxides

Most hydroxides

T2: Properties of Ionic Compounds• Melting point: High due to strong bonds between ions.• Boiling point: Higher, due to strong bond between ions.• Solid: do not conduct electricity• Molten (liquid): do conduct electricity• Dissolved (aqueous): do conduct electricity

Why? (HT)Electrical Conductivity• Electricity is conducted when there are charged particles

that are free to move.• Solid: there are charged particles (the ions), but they are

not free to move, so they do not conduct.• Liquid/Aqueous: the ions are now free to move, so they

do conduct

High Melting/Boiling Points• Ionic bonds (attraction between positive and negative

ions) are very strong.• Melting and boiling require these bonds to be broken.• This takes lots of (heat) energy.

T2: Making Insoluble Salts

1. React solutions of (the right) two soluble salts together.

2. Filter the mixture to collect the precipitate.

3. Rinse the filter residue with distilled water to remove impurities.

4. Allow the residue to dry.

T2: Barium Meals

• A patient is given a drink containing barium sulfate.• This can show up on

a x-ray, helping doctors to investigate the digestive system.

T2: Flame tests1. Clean a metal loop in acid2. Did loop in a metal salt.3. Heat in roaring Bunsen flame.

• Sodium, Na+ Yellow• Potassium, K+ Lilac• Calcium, Ca2+ Red• Copper, Cu2+ Green-blue

Precipitation TestsChloride: add acidified silver nitrate to get a white precipitate if chloride is present.Sulfate: add acidified barium chloride to get a white precipitate if sulfate is present.

Carbonate Test1. Add acid to the sample2. Pass any gas produced through

limewater: will go cloudy if the sample contained carbonate

T3: Diamond vs Graphite (HT)Diamond:• Very hard, as all carbon atoms joined

with strong covalent bonds.• Used to make cutting tools• Insulator as all electrons locked-tight

in bonds, so can’t move.

Graphite:• Layers of hexagonal carbon mesh that

rub away from each other, as there are only weak forces between the layers.

• Used as a lubricant.• Conductor as the electrons between

the layers are free to move. This is very rare for a giant covalent structure.

T3: Separating Immiscible Liquids• Immiscible = when liquids do not dissolve in

each other….like oil and water, one floats on top of the other.• Can be separated with a separating funnel;

the denser layer is tapped-off at the bottom.

T3: Covalent Bonds• Form when non-metals share electrons

between them.• Attraction between each atom and the

shared electron pair.• Atoms share electrons to complete their

outer shells• One bond is formed for each ‘gap’ in the

outer shell• Bonding represented with dot-and-cross

diagrams showing only the outer-shell electrons.

Example 1: WaterEach hydrogen needs one more electron to complete it’s outer shell and the oxygen needs two more. Oxygen forms two single bonds: one to each hydrogen.

Example 2: Carbon dioxide (HT only)Carbon needs two more electrons to complete it’s outer shell and each oxygen needs two more. Carbon forms two double bonds: one to each oxygen.

H HO

O OC

T3: Covalent StructuresSimple Covalent Molecules• Molecule = A particle made of a

small group of atoms, covalently bonded together.

• Low melting and boiling point, due to weak attractive forces between molecules..

• Electrical insulator as no electrons free to move.

• Examples: water, ammonia, oxygen

Giant Covalent• Repeating pattern of

many millions of atoms covalently bonded.• High melting/boiling

point because much heat energy needed to break strong covalent bonds.• Electrical insulator as

no electrons free to move.• Examples: silicon

dioxide, diamond, graphite

Lead nitrate + potassium iodide lead iodide + potassium nitratePb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq)

T3: Separating Miscible Liquids•Miscible = when liquids dissolve in each other…like alcohol and water.•Separate with fractional distillation using a fractionating column.•The components of the mixture have different boiling points, so if you heat it, each component will boil at a different time, allowing you to collect and condense the pure vapour.•We can do this to separate the gases in air by first cooling the air to turn the gases to liquid.

.

T4: Halogens and Their Reactions• Group 7: Fluorine (F) – pale yellow gas, Chlorine (Cl) – pale

green gas, Bromine (Br) – orangey-brown liquid, Iodine (I) – grey solid.• Most reactive at top of group, and get less reactive as you go

down.• Form halide ions with a charge of ‘-1’

Reaction with metals• React with metals to form metal halides• General equation: metal + halogen metal halide• For example: magnesium + iodine magnesium iodide Mg(s) + I2(s) MgI2(s) Note: Mg forms a 2+ ion, so two I- ions are needed.

Reaction with hydrogen• React with hydrogen to form hydrogen halides.• Hydrogen halides dissolve in water to form acids.• General equation: metal + halogen hydrogen halide• For example: hydrogen + fluorine hydrogen fluoride H2(g) + F2(g) 2HF(g) Note: hydrogen fluoride dissolves to make hydrofluoric acid.

Displacement Reactions• More reactive halogens can react with the ions of less

reactive halogens and displace them from compounds.• For example: 2KI(aq) + Br2(aq) 2KBr(aq) + I2(aq)

• This reaction works because bromine is more reactive than iodine.

• The orange colour of bromine would change to the brown colour of aqueous iodine.

• The reverse reaction would not work.

Reactivity Series of Halogens• Displacement reactions can be used to determine the

order of reactivity of the halogens.• Try reacting each halogen with solutions of each halide

salt, the halogen that does most reactions is most reactive.

T4: Metallic Bonding• Electrons are delocalised, moving

freely between all the atoms creating a ‘sea of electrons’• All atoms have a positive charge

as their outer-shell electrons have left them.• The bond is the attraction

between the positive ions and the sea of electrons.• Conduct electricity as electrons

are free to move.• Malleable (change shape but

don’t shatter when hit) because rows of atoms slide past each other when hit

T4: Transition Metals

• High melting points• Form brightly coloured

compounds

Halide SaltPotassium

fluoridePotassium chloride

Potassium bromide

Potassium iodide

Halogen

Fluorine x Reaction Reaction ReactionChlorine No

reactionx Reaction Reaction

Bromine No reaction

No reaction x Reaction

Iodine No reaction

No reaction No reaction

x

T4: Alkali Metals

• Group 1: Lithium (Li), Sodium (Na), Potassium (K)…• Properties: low melting point, soft (can be cut with

a knife).• React with water as follows:

General equation: metal + water metal hydroxide + hydrogenFor example: 2K(s) + 2H2O(l) 2KOH(aq) + H2(g)

Reactivity• Reactivity increases down the group:• Lithium just fizzes before disappearing• Sodium fizzes and gets hot enough to melt

into a ball, occasionally catching fire• Potassium fizzes very vigorously, getting hot

enough to burn with a lilac flame

Explaining Reactivity (HT only)• All reactions require you to remove the outer-shell

electron/• Atoms get bigger going down the group outer-

shell electrons further from nucleus easier to remove the outer shell electron.

T4: Noble Gases• Group 0 in the periodic table.• Helium ((He, Neon (Ne), Argon (Ar), Krypton (Kr)

Xenon (Xe), Radon (Rn)• Full outer shells so extremely unreactive: inert.

Discovery:• Lord Rayleigh noticed the density of nitrogen made in

reactions was less than nitrogen made from air.• Sir William Ramsey hypothesised that the nitrogen in

the air must also contain a denser gas that had not yet been discovered.• Through careful experiments, Rayleigh and Ramsey

discovered a gas that they named ‘argon’.• They also discovered helium, and then later Ne, Kr and

Xe.

Uses:• He and Ar were used to stop in filament in old bulbs

burning.• Ar and He used in welding to stop hot metal oxidising.• Ar used in fire extinguishing systems in server rooms.• He used in airships/blimps due to low density.• Neon lights due to red colour of light produce by neon.

T5: Endothermic and ExothermicExothermic Reactions• Chemical energy is converted to heat energy.• The surroundings get hotter.• For example: combustion reactions: Methane + oxygen carbon dioxide + water CH4 + 2O2 CO2 + 2H2O• Explosions are just very fast exothermic

reactions.

Endothermic Reactions• Heat energy is converted to chemical energy.• The surroundings get colder.• Examples: ammonium nitrate dissolving in

water, photosynthesis

Making and Breaking Chemical Bonds• In reactions, old chemical bonds are broken,

and then new ones are made.• Breaking bonds takes in energy; making bonds

gives out energy.• Stronger bonds take more energy to break,

and give out more when made.• In exothermic reactions, weaker bonds are

broken and stronger bonds are made.• In endothermic reactions, stronger bonds are

broken and weaker bonds are made.

Energy Diagrams (HT only)

Chem

ical

Ene

rgy

Reactants

Products

Energy released so gets hotter

Reactants

Products

Energy absorbed

so gets colder

EXOTHERMIC ENDOTHERMIC

T5: Catalytic Converters

• Part of exhaust pipe that helps make car exhaust less environmentally damaging.

• Toxic carbon monoxide and unburned hydrocarbons (from petrol) are converted into carbon monoxide and water.• The catalytic converter has a fine honeycomb

structure coated with the catalyst.• The catalyst contains a mixture of platinum,

rhodium and palladium.• The metals are expensive, so only a very thin

coating is used.• The catalysts work best at high temperatures, so

car exhaust is more damaging when the car has only just started and hasn’t warmed up.

T5: Collision Theory (HT)• To react: particles must collide with enough

energy.• To increase rate: increase the amount of

collisions or the energy of the collisions.

Effect of Concentration:• Increasing concentration increases the number of

reacting particles.• This increases the number of collisions.

Effect of Surface Area:• Increasing the surface area increases the

proportion of (solid) particles available to react.• This increases the number of collisions.

Effect of Temperature:• Increasing the temperature increases the speed

that particles are moving• This means there are more collisions, and those

collisions have more energy.

T5: Rates of Reaction (Intro)

Note: you increase the surface area by breaking a large piece into many smaller pieces, with powder being the best.

T6: Reacting Quantities (HT)• Combining relative masses with balanced equations lets us

work out the masses of chemicals involved in reactions.• We can use this mathematical relationship:

Example:• What mass of carbon dioxide can be produced by burning

15g ethene (C2H4) in excess oxygen (O2)?C2H4 + 3O2 2CO2 + 2H2O

• Substance 2 will be ethene, substance 1 will be carbon dioxide.• Calculate relative masses:

• Mr(ethene) = 2 x 12 + 4 x 1 = 28• Mr(carbon dioxide) = 12 + 2 x 16 = 44

• Then:

• m = mass of substance present• Mr = relative formula mass of

substance• n = number of substance in balanced

equation• 1 refers to the first substance• 2 refers to the second substance

Write out the equation.

Sub in the numbers

Rearrange to make m1 the subject.

T6: Empirical Formulae

Relative Atomic Mass, Ar• The lowest whole number ratio of atoms in a molecule.• For example:

• The empirical formula can be calculated from the masses of substances that react with each other as below.

• For example: 10.0g of magnesium reacts with 133.3 g of bromine.

Molecular Formula Empirical Formula

Water, H2O H2O

Ethane, C2H6 CH3

Glucose, C6H12O6 CH2O

Mg Br

Mass in g 10.0 133.3

Relative atomic mass 12 80

Divide by relative atomic mass

10 / 12 = 0.83 133.3 / 80 = 1.67

Divide both sides by smallest answer

0.83 / 0.83 = 1 1.67/0.83 = 2

Empirical formula MgBr2

T6: Yield• Theoretical yield: the amount of

product you would expect according to the calculation in the ‘Reacting Quantities’ box.• Actual yield: the amount of product

you actually get in practice.• Percentage yield: the proportion of

the theoretical yield that you actually achieve.

% yield is always less than 100 because:• The reaction may be incomplete• Some product may be lost during the

steps to prepare it.• Some reactants may also produce

products other than the desired one.

𝑚1

𝑀𝑟 1𝑛1=

𝑚2

𝑀 𝑟2𝑛2

T6: Percentage by Mass• This is the percentage of the mass of a compound due to a

particular element.

For example: what is the carbon in ethanol, C2H6O?

Calculate Mr of C2H6O Mr = (2 x 12) + (6 x 1) + 16 = 46

Number of C in C2H6O 2

Relative atomic mass of C

12

Percentage by mass of C = 52.1%

T6: Relative MassesElement Relative Mass

Hydrogen, H 1

Carbon, C 12

Oxygen, O 16

Sodium, Na 23

Chlorine, Cl 35.5

Relative Atomic Mass, Ar• The mass of atom relative to

the mass of 12C (carbon-12).• For example…

Relative Formula Mass, Mr• This is the sum of all the relative masses in a formula.• Relative formula mass of carbon dioxide, CO2:

Mr = Ar(C) + 2 x Ar(O) = 12 + (2 x 16) = 44

• Relative formula mass of sodium chlorate, NaClO3

Mr = Ar(Na) + Ar(Cl) + 3 x Ar(O) = 23 + 35 + (3 x 16)

• The rate of a reaction is its speed, how quickly products are made.• Reactions happen when particles collide with each

other.• Concentration: increasing concentration (the amount

of solute (dissolved stuff) in a given volume) will increase the rate.• Temperature: increasing temperature will increase

the rate.• Surface area: increasing surface area will increase the

rate.

Type of Bonding

Ionic Simple molecular Giant Molecular

How the bonds form

Swapping electrons to form ions

Sharing electrons Sharing electrons

Examples Sodium chloride, magnesium oxide

Water, methane, nitrogen

Quartz (silicon dioxide)

Bond strength

Strong Strong bonds, weak intermolecular forces

Strong bonds

Melting and boiling point

High Low High

Solubility Most in water Some in water Insoluble in water

Conduct electricity?

Only when molten or dissolved

No No (except graphite)