What’s in my atom?

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

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

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

summarised below.

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

Property Eka-aluminium - Ea(the prediction)

Gallium - Ga(the one discovered)

Atomic mass about 68 70

Density (g/cm3) 6.0 5.9

Melting point (OC) Low 29.8

Formula of oxide Ea2O3 Al2O3

Density of oxide 5.5 5.9

Reacts with acids and alkalis?

Yes Yes

PERIODS….increasing atomic mass, differing properties

GRO

UPS…

…sim

ilar properties

Element Type = non-metal = metal

C2 Topic 1Atomic Structure and

the Periodic Table

Particle Relative charge

Relative mass

Found?

Proton 1 Positive, +1

In nucleus

Neutron 1 Neutral, 0

In nucleus

Electron Neglible ()

Negative, -1

In shells orbiting nucleus

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.

Atomic number = 9Relative Atomic mass = 19

Protons = 9Electrons = 9Neutrons = 19-9 = 10

Atomic number = 16Relative Atomic mass = 32

Protons = 16Electrons = 16Neutrons = 32-16 = 16

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.

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.

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.

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!

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

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.

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+

C2 Topic 2Ionic

Compounds and Analysis

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-

Solubility• Soluble: a compound dissolves in a given

liquid.• Insoluble: a compound does not dissolve.

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

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.

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:

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

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.

3+

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

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

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.

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.

C2 Topic 3Covalent Compounds and

Separation Techniques

Chromatography• Separates compounds based on differences in their solubility in a solvent, producing a

chromatogram.• Small spots of sample are placed near the bottom of a sheet of chromatography paper, which

is then suspended in a solvent.• As the solvent soaks up the paper, the different components of the sample move up at

different speeds, causing the sample to separate.• Retention Factor (Rf) is a measure of how far the spots move.• Rf is fixed for a given compound and solvent.

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

Separating Immiscible Liquids

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.

H HO

O OC

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

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.

• 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.

Uses:• Identifying colourings used in

foods• Analysing DNA collected from

crime scenes• Analysing the make-up of

paints and dyes.

Transition Metals• High melting points• Form brightly coloured compounds

C2 Topic 4Groups in the Periodic Table

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 flameExplaining 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.

Metallic Bonding

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.

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.

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 waterConduct electricity?

Only when molten or dissolved

No No (except graphite)

• 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

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

C2 Topic 5Chemical Reactions

Endothermic and Exothermic

Exothermic 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)

Catalysts• Catalysts: increase the rate of a reaction

without getting used up.• Catalysts are often used in industry to speed

up chemical processes.• For example:

Chem

ical

Ene

rgy Reactants

Products

Energy released so gets hotter

Reactants

Products

Energy absorbed

so gets colder

EXOTHERMIC ENDOTHERMIC

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.

Rates of Reaction (Intro)• 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.

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

Catalyst Process

Iron Production of ammonia from hydrogen and nitrogen

Nickel Production of margarine from vegetable oil.

Aluminium oxide

Cracking alkanes in an oil refinery

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.

C2 Topic 6Quantitative Chemistry

Relative Masses

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)

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.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:

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?

Waste and Profit

Waste Disposal• Many chemical products produce useful waste

products that can be sold, these are called by-products.

• Other products have no use and are called waste products.

• There are various problems with waste products:

• May cause social problems – due to unpleasant smells, or eyesores.

• May be damaging to the environment.

• May be expensive – due to treatment to make them safe, or transport to landfill sites.

Maximising Profit (HT only)• Chemical factories have to make money, so

chemists work hard to make sure the most product is produced for the least cost.

• They try to:• Increase the % yield to reduce

waste.• Increase the speed of the reaction

so more product can be made in a given time.

• Make sure all products have uses so can be sold.

Element Relative MassHydrogen, H 1

Carbon, C 12Oxygen, O 16

Sodium, Na 23Chlorine, Cl 35.5

Empirical FormulaeRelative 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 H2OEthane, C2H6 CH3

Glucose, C6H12O6 CH2O

Mg BrMass 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

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

Relative atomic mass of C 12Percentage by mass of C = 52.1%

Where:• 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.