Chemistry 7: Atomic Structure

Section 1:

Atom Simplest form of matter

Element Made up of only one type of atom

Compound Made up of two or more different types of atoms chemically bonded.

Mixture Two or more atoms not chemically bonded.

Section 2a:

Radius of an atom Distance from the centre of the nucleus to the outer shell ~ 1 x10-10

Radius of a nucleus Distance from the middle of the nucleus to the outside ~1x10 -10

IsotopeAn atom of the same element that has the same number of protons but adifferent number of neutrons.

Relative atomic mass

average value that takes account of the abundance of the isotopes of that element.= sum (isotope abundance x isotope mass) / sum of abundance of all isoptopes

Section 2: Atomic Structure

Section 4: Development of the theory of the atom

Person Year Theory

Dalton Early 1800sAtoms were hard spheres that couldn’t be broken down into anything

more simple.

Thompson 1897Plum pudding model’ – ball of positive charge with electrons spread

throughout it.

Rutherford 1911Alpha scattering experiment this gave unexpected results and gave us the

nuclear model – positive nucleus with electrons around the outside.

Bohr 1913 Electrons are in fixed shells around the nucleus.

Chadwick 1932 Discovered neutrons.

Proton Neutron Electron

Mass 1 1 Negligible (very small)

Charge +1 0 -1

Location Nucleus Nucleus Shells

Section 3 Separation techniques

Separation technique What is it used to separate? How does it work?

Filtration An insoluble solid from a

liquid

The mixture is poured through filter paper. The liquid goes through the filter paper and the

solid remains in the filter paper.

Evaporation A soluble solid from a

solutionThe solution is slowly heated until all of the liquid has evaporated. You are left with a dry solid.

Crystallisation A soluble solid from a

solution

The solution is slowly heated until the point of crystallisation. The concentrated solution is left

to cool until crystals form. The crystals are filtered and dried.

Distillation

A liquid from a solution (can

be two liquids or a soluble

solid and a liquid).

The solution is heated until the part of the solution with the lowest boiling point evaporates.

The vapour is condensed and the liquid collected.

Fractional distillationA mixture of liquids with

different boiling points.

The solution is heated until the part of the solution with the lowest boiling point evaporates.

The vapour is condensed and the liquid collected. This is repeated at higher temperatures to

collect fractions with higher boiling points.

Chromatography

A mixture of dissolved solids

or liquids (such as dyes in

inks or food).

A mixture of the dissolved substance is put onto chromatography paper. A solvent is added and

this runs up the paper. The different substances in the mixture move at different speeds.

Section 4: Development of the theory of the periodic table

Stage People Idea

1Dalton, Newlands

attempted to classify the elements by arranging them in order of their atomic weightsThe early periodic tables were incomplete and some elements were placed in inappropriate

groups if the strict order of atomic weights was followed.

2 Mendeleev

Overcame some of the problems by leaving gaps for elements that he thought had not been discovered and in some places changed the order based on properties instead of

atomic mass.

3

Elements with properties predicted by Mendeleev were discovered and filled the gaps. Knowledge of isotopes made it possible to explain why the order based on atomic weights

was not always correct.

Section 5:Trends in groups

Group Properties Trends (going down the group) Reactions

0 Unreactive and do not easily form molecules because

their atoms have stable arrangements of electrons

(monatomic).

Colourless gases.

Non-flammable.

The noble gases have 8 electrons in their outer shell,

except for helium, which only has 2 electrons.

Boiling points increase.

Relative atomic mass increases.

Unreactive

1 Low density (the first three elements in the group are

less dense than water).

Soft.

1 electron in outer shell.

Very reactive.

Reactivity increases (electron is

more easily lost).

Melting and boiling points decrease.

Relative atomic mass increases.

React with water to produce a metal hydroxide

and hydrogen.

React with chlorine to produce a metal chloride

salt.

React with oxygen to produce a metal oxide.

7 Exist as pairs of atoms (e.g. Cl2).

Have 7 electrons in their outer shell.

Reactivity decreases (it is harder to

gain an electron).

Melting and boiling points increase.

Relative atomic mass increases.

Share electrons with other non-metals in

covalent bonding (e.g. HCl).

Form ionic bonds with metals (e.g. NaCl).

More reactive halogens will displace less

reactive ones.

Chemistry 7: Atomic Structure: Chemistry ONLY

Transition metals: Comparing with Group 1

Group 1 Transition metals

Melting points

Low High

Reactivity High Low

StrengthSoft or liquid

Hard

Density Low High

CompoundsWhite or colourless

Coloured

Ions +1Can have more than one ion e.g Cu+ and Cu2+

UsesCatalysts e.g Iron for Haber process or Ni for hydrogenation of alkenes

Ionic compound

RatioFormula of the

compound

Sodium chloride

Na+ : Cl-

1 : 1NaCl

Calcium chloride

Ca2+ : Cl-

1 : 2CaCl2

Section 1: Types of bond

Type Between What happens

Ionic Metals and non metalsElectrons are transferred – electrostatic attraction between oppositely charged ions

Covalent Non metal and non metal Electrons are shared

Metallic Metal and metalAttraction between positive ion and delocalised electrons

Section 2: Ionic formula Section 2a: Common ions

Section 4: Bonding and properties

chemical

bondExists between Example Properties

Ionic bondingMetal and non-

metal

NaCl Ionic compounds

Giant lattice structure.

Strong electrostatic forces of attraction between oppositely charged ions.

High melting and boiling points.

Do not conduct as a solid – ions cannot move

Do conduct when melted/dissolved – ions can move.

Covalent

bondingNon-metals

Simple covalent

molecules

H2

Simple covalent molecules

Do not conduct electricity.

Strong covalent bonds between atoms but weak intermolecular forces between molecules – low melting and boiling

points.

Giant covalent

molecules

(macromolecules)

Giant covalent molecules (macromolecules)

Do not conduct electricity – no ions or delocalised electrons (exceptions include graphite).

Strong covalent bonds and all atoms bonded to each other – high melting and boiling points.

Metallic

bondingMetals

Sodium Delocalised electrons carry current and thermal energy well - conduct electricity and heat.

Strong electrostatic forces of attraction between positive ions and electrons - high melting/boiling points.

Layers can move around so metal is malleable/ductile.

Each metal puts in electrons, the more electrons in the cloud = stronger bonding.

Alloys (a metal mixed with another element) – stronger as the different sized atoms distort the layers so they cannot

slide over each other.

Section 3: States of matter

State Diagram Properties

Solid

Strong forces of attraction between particles.Particles are close together in lattice arrangement.The particles don’t move from their positions so solids keep a definite shape.The particles vibrate in their positions – the hotter the solid, the more they vibrate (solids expand when hot).

Liquid

Weak force of attraction between particles. They are randomly arranged and free to move past each other (but they stick closely together).Liquids have a definite volume but not a definite shape – they can flow.The particles are constantly moving with random motion. The hotter the liquid gets, the faster they move (liquids expand when hot).

Gas

Very weak force of attraction between particles. They are free to move and are far apart. The particles in gases travel in straight lines.Gases don’t keep a definite shape or volume.The particles move constantly with random motion. The hotter the gas gets, the faster they move. Gases either expand when heated, or their pressure increases.

Name Ion

Hydroxide OH-

Sulfate SO42-

Carbonate CO32-

Nitrate NO3-

Chemistry 2: Bonding, structure and properties of matter

Type of particles Size in nm Size in m

Nanoparticles 1-100 1 x 10-9 to 1 x 10-7

Fine particles (PM2.5) 100-2500 1 x 10- 7 to 2.5 x 10-6

Coarse particles (PM10) 2500-10000 1 x 10-5 to 2.5 x 10-6

Section 5:Structure and bonding of carbon

Allotrope of carbon

Structure Properties Use

DiamondEach carbon atom has 4 covalent bonds – diamond is very hardThese strong bonds take a lot of energy to break – diamond has a very high melting point.It doesn’t conduct electricity as there are no free electrons or ions.

In drill bits.

Graphite

Each carbon atom has 3 covalent bonds (each atom has 1 delocalised electron) – the delocalised electron allows graphite to conduct electricity and thermal energy.The covalent bonds take a lot of energy to break – graphite has a high melting point.There aren’t any bonds between layers – this makes graphite soft and slippery.

As lubricants.

Graphene

Graphene is on one atom think – it is a 2D compound.The network of covalent bonds makes it very strong and very light – it can be added to materials to improve their strength without adding much weight.It contains delocalised electrons so can conduct electricity through the whole structure.

In electronics.

Fullerenes

Fullerenes are molecules of carbon, shaped like closed tubes or hollow balls.They are arranged in hexagons (or pentagons or heptagons).They can trap molecules inside them.They have a huge surface area.The first fullerene to be discovered was the Buckminsterfullerene (C60).

For drug delivery.To make carbon nanotubes.Lubricants (they roll)

Fullerenes as nanotubes

Nanotubes are tiny carbon cylinders.The ratio between the length and the diameter of nanotubes is very high.They can conduct both electricity and thermal energy.They have high tensile strength.

In electronics.For tennis rackets.

Section 6: Bulk and surface properties of matter Section 6a:nUses of nanoparticles

Use Why? (property)

Catalysts Due to their huge surface area to volume ratio.

Medicine (to deliver drugs) They are small and more easily absorbed into the body.

Computer chips Some nanoparticles conduct electricity.

Surgical masks and wound dressings Silver nanoparticles have antibacterial properties.

Cosmetics To improve moisturisers without making them oily.

Section 7: Surface area to volume ratio

As the side of cube decreases by a factor of 10 the surface area to volume ratio increases by a factor of 10.

How to calculate Surface area to volume ration = surface area ÷volume

Section 6b: Disadvantages of nanoparticles

Effects on the body are unknown so there may be some risks. All products containing them should be labelled.

Section 1:Moles

Name Definition Example

Mass number (relative atomic mass)

Amount of protons and neutrons – the bigger number on the element symbol

Relative formula mass All of the relative atomic masses added up in a compound. CaCO3 = 40+12+(16x3) = 100

Percentage mass of a compound

𝐴𝑟 𝑥 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 𝑜𝑓 𝑡ℎ𝑎𝑡 𝑒𝑙𝑒𝑚𝑒𝑛𝑡

𝑀𝑟 𝑜𝑓 𝑡ℎ𝑒 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑x 100

% of Na in NaCl = 23 𝑥 1

23+35.5x 100 = 39.3%

Mole

An amount of a substance. One mole of any substance contains 6.02x1023 atoms.

Moles = 𝑚𝑎𝑠𝑠 (𝑔)

𝑀𝑟

Number of moles in 66g of CO2 = 66

12+16+16 = 1.5 mol

Balancing equations with moles

1.Calculate the moles of each substance.2. Dive the number of moles of each substance by the smallest number of moles in the reaction. 3. Multiply if you don’t have whole numbers.4. Write the balanced equation by putting these numbers at the front

Limiting reactantThe reactant that is used up first and causes the reaction to stop. The amount of product is directly proportional to the amount of this.

Reacting masses (calculatinghow much product you can make)

Write out the balanced symbol equation.Work out relative formula mass.Work out number of moles of reactants .Use molar ratios to work out moles of product.Multiply the number of moles by the RFM of the product you’re trying to find and this is your answer in g.

. Calculate the mass of aluminium oxide formed when 135 g of aluminium is burned in air.

4Al + 3O2 → 2Al2O3

1. Number of moles of Al = mass/Mr 135/27 = 5 moles.2. Ratio from the equation shows that 4 moles of Al produce 2 moles of

Al2O3 so 5 moles of Al produce 2.5 moles of Al2O3.3. Mass of 2.5 moles of Al2O3 = moles x Mr = 2.5 x 102 = 255 g.

Conservation of massNo atoms are destroyed or created during a chemical reaction – no mass is lost.

In a chemical reaction the mass may increase which is probably due to one of the reactants being a gas.

In a chemical reaction the mass may appear to change which is probably due to a gas being given off as a product in the reaction.

ConcentrationHow much solute is dissolved in a certain volume. Unit =g/dm3

Concentration= 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔)

𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 (𝑑𝑚3)

What is the concentration of a solution containing 30g of NaCl dissolved in 0.2dm3 of water?30/0.2 = 150g/dm3

Concentration (Chemistry only)

How many moles are dissolved in a certain volume. Unit =mol/dm3

Concentration= 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑚𝑜𝑙)

𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 (𝑑𝑚3)

Converting mol/dm3 to g/dm3 = multiply the moles by the Mr.

What is the concentration of a solution containing 1 mol of NaCldissolved in 0.2dm3 of water?1/0.2 = 5 mol/dm3

Chemistry 3: Quantitative Chemistry

Section 1:Moles

Name Definition Example

Molar gas volume

At the same temperature and pressure, equal numbers of moles of any gas will occupy the same volume. At room temperature this is 24dm3

)

Volume of gas Volume of gas = 𝑀𝐴𝑠𝑠 𝑜𝑓 𝑔𝑎𝑠(𝑔)

𝑀𝑟 𝑜𝑓 𝑔𝑎𝑠x 24 What’s the volume of 319.5g of chlorine at r.t.p =

319.5

71x 24 = 108dm3

Atom economy

A Measure of how “green a process is” The higher the atom economy the more green the process.

Atom economy =𝑟𝑒𝑙𝑎𝑡𝑖𝑣𝑒 𝑓𝑜𝑟𝑚𝑢𝑙𝑎 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑑𝑒𝑠𝑖𝑟𝑒𝑑 𝑝𝑟𝑜𝑑𝑢𝑐𝑡 𝑓𝑟𝑜𝑚 𝑡ℎ𝑒 𝑒𝑞𝑢𝑎𝑡𝑖𝑜𝑛

𝑠𝑢𝑚 𝑜𝑓 𝑡ℎ𝑒 𝑟𝑒𝑙𝑎𝑡𝑖𝑣𝑒 𝑓𝑜𝑟𝑚𝑢𝑙𝑎 𝑚𝑎𝑠𝑠𝑒𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 𝑓𝑟𝑜𝑚 𝑡ℎ𝑒 𝑒𝑞𝑢𝑎𝑡𝑖𝑜𝑛x100

100% atom economy only one product.

Low atom economy – use up resources quickly, lots of waste, unsustainable, less profitable.

Calculate the atom economy for the production of CaO.

CaCO3 CaO + CO2

(40+16)

(40+16+12+16+16)x 100 = 56%

% Yield

Atom economy =𝑀𝑎𝑠𝑠 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡 𝑎𝑐𝑡𝑢𝑎𝑙𝑙𝑦 𝑚𝑎𝑑𝑒

𝑀𝑎𝑥𝑖𝑢𝑚𝑢𝑚 𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡x100

100% yield = All product you expected to get.0% yield = no product.

Zn + 2HCl ZnCl2 + H2. The theoretical yield of zinc chloride is 2.72g but after purification the sample weighed 2.31g. What was the % yield?

2.31

2.71x 100 = 84.9%

Reasons for yield not being 100%

Not all reactants react to make a product. Reversible reactions – the products can return back to reactants

There might be side reactions.Reactants may react with other things in the air or impurities in the reaction mixture.

May lose product whilst separating

When you filter a liquid you will always leave a bit of solid on the paper. Some of the liquid is absorbed onto the paper. Transferring liquids also causes a reduction in yield – you can never get all of the liquid out.

Chemistry 3: Quantitative Chemistry – Chemistry only

Section 2: Reactivity

Element Reaction Reactivity

12 PotassiumWhen potassium is added to water, the metal melts and floats. It moves around very quickly. The metal is also set on fire, with sparks and a lilac flame.

13 SodiumWhen sodium is added to water, it melts to form a ball that moves around on the surface. It fizzes rapidly.

14 LithiumWhen lithium is added to water, it floats. It fizzes steadily and becomes smaller.

15 Calcium Fizzes quickly with dilute acid.

16 Magnesium Fizzes quickly with dilute acid.

17 (Carbon)

18 Zinc Bubbles slowly with dilute acid.

19 Iron Very slow reaction with dilute acid.

20 (Hydrogen)

21 Copper No reaction with dilute acid.

Chemistry 4: Chemical Changes

Section 1: Key Terms

1 Metal oxide Metals react with oxides to produce metal oxides. This is an oxidation reaction.

2 Displacement reactionA more reactive metal can displace a less reactive metal from acompound.

3 OxidationTwo definitions:Chemicals are oxidised if they gain oxygen in a reaction.Chemicals are oxidised if they lose electrons in a reaction. (HT)

4 ReductionTwo definitions:Chemicals are oxidised if they lose oxygen in a reaction.Chemicals are oxidised if they gain electrons in a reaction. (HT)

5 Acid A chemical that dissolves in water to produce H+ ions.

6 Base A chemical that reacts with acids and neutralise them. E.g. metal oxides, metal hydroxides, metal carbonate

7 Alkali A base that dissolves in water. It produces OH- ions in solution.

8 NeutralisationWhen a neutral solution is formed from reacting an acid and alkali. General equation: H+ + OH-

H2O

9 pHA scale to measure acidity/ alkalinity. A decrease of one pH unit causes a

10x increase in H+ ions. (HT)

10 Strong acid (HT)A strong acid is completely ionised in solution. E.g. hydrochloric, nitric and sulfuric acids.

11 Weak acid (HT)A weak acid is only partially ionised in solution. E.g. ethanoic, citric andcarbonic acids.

Section 4: Extracting Metals

22 Very unreactive metals Found naturally in the ground. Don’t need extracting.

23 Metals less reactive than carbon Extracted by reduction with carbon.

24 Metals more reactive than carbon Extracted by electrolysis.

Section 5: Reactions of Acids

25 With metal Acid + Metal Salt + Hydrogen

26 With alkaliAcid + Metal Hydroxide Salt + Water (Neutralisation reaction)

27 With metal oxide

Acid + Metal Oxide Salt + Water (Neutralisation reaction)

28 With carbonate

Acid + Metal Carbonate Salt + Water + Carbon Dioxide (Neutralisation reaction)

Section 6: Making a Soluble Salt

29 Add solid metal, metal carbonate, metal oxide or metal hydroxide to an acid.

30 Add solid until no more reacts.

31 Filter off excess solid.

32 Evaporate to remove some of the water.

33 Leave to crystallise.

34 Remove all water in a desiccator/ oven.

35 Acidic pH 0-6 36 Neutral pH 7 37 Neutral pH 8-14

Section 7 Electrolysis key terms

38 Electrolysis The process of splitting an ionic compound by passing electricity through it.

39 ElectrolyteAn ionic compound that is molten (melted) or dissolved in water. The ionsare free to move.

40 ElectrodeAn electrical conductor that is placed in the electrolyte and connected to the power supply.

41 Cathode The electrode attached to the negative terminal of the power supply.

42 Anode The electrode attached to the positive terminal of the power supply.

Section 8: What is discharged in electrolysis?

Electrolyte Cathode Anode

43 Molten Compound

Metal Non-metal

44 Dissolved compound (aqueous solution)

The metal if the metal is less reactive than hydrogen.Hydrogen is produced if themetal is more reactive than hydrogen.

Oxygen is produced unless the solution contains halide ions (chloride, bromide, iodide) when the halogen (chlorine, bromine, iodine) is produced.

Section 9: Aluminium Electrolysis

45 Cryolite

Aluminium oxide is dissolved in cryolite to lower its melting point. This saves money on energy costs.

46 Cathode

Positive Al3+ ions move to the cathode. Aluminium is produced. Al3+ + 3e- Al

47 Anode

Negative O2- ions move to the anode. Oxygen is made. 2O2- O2 + 4e-

Wears away as the carbon anode reacts with oxygen to form carbon dioxide.

42 Anode41 Cathode

39 Electrolyte

Chemistry 4: Chemical Changes Chemistry 5: Energy Changes

Section 7 Energy Changes Key Terms

1 Conservation of energy

Energy is not created or destroyed, only transferred from one store to another

2 ExothermicA reaction that transfers energy to the surroundings so the temperature of the surroundings increases, e.g. combustion and neutralisation reactions. Used in self-heating cans and hand warmers.

3 EndothermicA reaction that takes in energy from the surroundings so the temperature of the surroundings decreases, e.g. thermal decomposition. Used in sports injury packs.

4 Activation energy The energy needed for particles to successfully react.

5 Breaking bonds Energy is needed to break bonds.

6 Forming bonds Energy is released when bonds are formed.

9 Energy released from forming bonds is greater than the energy needed to break bonds. (HT)

10 Energy released from forming bonds is less than the energy needed to break bonds. (HT)

7 Exothermic Energy Profile 8 Endothermic Energy Profile

Chemistry 6: Rates of reaction

Section 3: Factors affecting the rate of reaction

5 ConcentrationIncreasing the concentration increases the number of particles in the same volume. This increases the collision frequency between reactant particles.

6 Surface areaA powder has a greater surface area than a solid. Increasing the SA will make the particles more exposed which increases the collision frequency between reactant particles.

7 TemperatureIncreasing the temperature of the reaction increases the energy of the reactant particles meaning more collisions will have the required energy to be successful.

8 PressureThis is the same as concentration but for gases. Reactant particles occupy a smaller volume.

9 Catalyst Reduces the activation energy by providing an alternative reaction pathway

Section 1: Organisation

1 Reactant The starting materials in a reaction that undergo a chemical change.

2 Product What you end up with in a reaction.

3 Collision Theory Particles must collide with enough energy and the correct orientation to react.

Section 7 Reversible reactions

The reaction is reversible

Closed system None of the products or reactants can escape and nothing can get in.

Dynamic equilibrium

In a closed system, the rate of the forward and backward reaction are the same.

Section 4: Measuring rates of reaction

Method How Negatives

10 Precipitation and colour changeMeasure the time taken for the solution to go cloudy

Colour change is subjective

11 Change in mass (usually gas) Reaction takes place on a balanceMost accurate but gas escapes into the room

12 Measuring the volume of gas given off

Use a gas syringe to collect the gas

If the reaction is too vigorous the plunger may blow off.

Section 2

Section 8: Le Chatelier’s principleChange Effect

TemperatureHeating the reaction moves the equilibrium to the endothermic side. Reducing the temperature moves equilibrium to the exothermic side.

PressureIncreasing the pressure shifts the equilibrium to the side with fewer molecules Decreasing the pressure shifts equilibrium to the side with more molecules.

ConcentrationIncreasing the concentration of reactants shifts equilibrium to the side of products. Decreasing product concentration shifts the equilibrium to the side of reactants.

Section: Required practicals

𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡 𝑓𝑜𝑟𝑚𝑒𝑑

𝑡𝑖𝑚𝑒

𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡 𝑢𝑠𝑒𝑑

𝑡𝑖𝑚𝑒or

Mean rate of reaction =

1. Add 40cm3 hydrochloric acid.2. Add a metal3. Place cotton wool in the top.4. Record the mass at set time intervals.5. Vary the concentration, repeat

1. Add sodium thiosulfate to conical flask on the cross2. Add sodium hydroxide.3. Time how long it takes for the cross to disappear.4. Vary the concentration of thiosulfate, repeat.

Section 6

Chemistry 7: Organic Chemistry

Section 1:

1 Hydrocarbon A molecule consisting of hydrogen and carbon only

2 Completecombustion

Molecule burns completely in oxygen to produce carbon dioxide and water

3 Incomplete combustion

Molecules don’t get enough oxygen so produces carbon monoxide, carbon and water after burning with a smoky yellow flame and less energy than complete .

Section 7 : Properties of alcohols

Flammable Undergo complete combustion to form carbon dioxide and water.

Solubility First four alcohols are soluble in water with neutral pH.

React with Na Produces sodium hydroxide and hydrogen

Can be oxidised React with oxygen to produce carboxylic acids. E.g. ethanol ethanoic acid

Section 7a Uses of alcohol

Solvents Can dissolve things water can’t like fats and oils

Fuels Burnt in spirit burners fairly cleanly and non-smelly

Section 7b Making alcohol

Equation Catalyst Conditions

FermentationSugar ethanol + carbon dioxide

C6H12O6 2C2H5OH + 2CO2

Yeast

37°C, slightly acidic and with no oxygen (anaerobic

respiration). Too hot enzyme denatures. Too cold reaction

too slow.

Section 3: Cracking

CrackingBreaking larger alkanes to produce smaller more useful alkanes + an alkene. Through thermal decomposition.

Homologous seriesA series of compounds that have similar chemical properties and the same general formula

Catalytic cracking Vaporise the hydrocarbon, pass over an aluminium oxide catalyst

Steam cracking Vaporise the hydrocarbon, mix with steam and heat to high temperatures

Section 2: Fractional distillation

Crude oil Fossil fuel made from plant and animal remains, drilled up.

Fractionaldistillation

Process of separating crude oil by boiling point.

Section 6: Functional groups

Saturatation Functional group General formula suffix

Alkane SaturatedSaturated

hydrocarbonCnH2n+2 ane

Alkene Unsaturated C=C CnH2n ene

Alcohol X OH CnH2n+1OH ol

Carboxylicacid

X CnH2n+1COOH Oic acid

Section 4: Naming compounds

Number of carbons

Prefix example

1 meth methane

2 eth ethane

3 prop propane

4 but butane

Section 5: Naming alkanes

Section 12a Types of polymerisation

Addition Polymerisation Condensation polymerisation

Example

Number of types of monomers

Only 1 monomer containing C=C

Two monomer types each containing two of the same functional group.

OrOne monomer with two different

functional groups.

Number of products 1 2 – the polymer and water

Functional groups involved in

polymerisation

C=C Two reactive groups on each monomer

Section 8: Reactions of alkenes

Name Conditions Example

Hydrogenation

React with hydrogen and a catalyst.

Making alcoholReact with steam and a catalyst. The

product is then condensed and purified by fractional distillation.

HalogenationReact with halogens to form

dihalocompounds. E.g dibromoethane

Section 10a: carboxylic acids

Acid Strength Weak acids because they only partially ionise in water.

Section 10b Reactions

Reaction with basesEthanoic acid + sodium carbonate Sodium ethanoate + water + carbon dioxide

Formation of estersAlcohol + Carboxylic acid ester + water

Eg. ethanol+ ethanoic acid Ethyl ethanoate + water

Section 9: Test for saturation Alkane Alkene

Add bromine water Solution stays orange-brownSolution turns from orange-

brown to colourless

Section 11: Naming organic molecules

Section 12: Polymers

Monomer Small molecules that react together to form a polymer

PolymerA large molecule formed from lots of different monomers bonding together. Usually

needs high pressure and a catalyst.

Repeating unit

The part of the polymer that repeats itself.

+2nH20

Section 13: Naturally occuring polymers

Amino Acid A monomer for proteins – contains two different functional groups.

Protein Polymers of amino acids. Structure determined by the order of amino acids.

BaseA,T,C,G. Pair together on a different polymer change to make the double helix

structure

Nucleotides A polymer chain. Two of these combine to make the double helix of DNA

Sugars Large carbohydrate polymers e.g cellulose

Chemistry 8: Chemical analysis

Section 1: Purity and formulations

1 Pure substance Only contains one compound or element

2 Melting and boiling point

Pure substances melt and boil at a specific temperature. Impurities alter this and give a wider range.

3 Formulations Mixtures with exact amounts of components each with a specific purpose.

Section 5: Instrumental methods

Advantages Disadvantages

Very sensitive – can detect tiny amounts More expensive

Very fast and tests can be automated

Very accurate

Section 5a Flame emission spectroscopy

Flame emission spectroscopy

Used to analyse metal ions in solutions to determine their concentration. The output is a line spectrum. Match the lines up in a mixture to identify the ions

Section 3: Tests for gases and anions

Chlorine gas Bleaches damp litmus paper turning it white

Oxygen gas Relights a glowing splint

Carbon dioxide gas Turns limewater cloudy

Hydrogen gas Gives off a squeaky pop if a lit splint is held next to a test tube of hydrogen

Carbonate (CO32-)

Add dilute acid to cause fizzing, connect to a test tube of limewater, turns cloudy as CO2 produced

Sulfates (SO42- )

Add hydrochloric acid followed by barium chloride White precipitate of Barium sulfate

Cl-Ag++Cl- AgCl(s) white precipitate

Br- Ag++Br- AgBr(s) cream precipitate

I- Ag++I- AgI(s) yellow precipitate

Section 2: Chromatography

Chromatography Used to separate the substances in a mixture.

Mobile phase Where the molecules can move. This is always liquid or gas.

Stationary phase Where the molecules can’t move. This can be a solid or a thick liquid.

SolubilitySubstances with high solubility will spend more time in mobile phase and travel further up the paper.

Purity Pure substances will only ever produce one spot on a chromatogram.

Rf Value Distance travelled by substance divided by distance travelled by solvent.

Tests for cations

Section 4: Flame tests

Metal Colour

Lithium crimson

Sodium Yellow

Potassium Lilac

Calcium Orange-red

Copper Green

Section 4:

Metal ionsColour of

precipitateWith NaOH

Ionic equation

Ca2+ White Ca2+(aq) + 2OH-

(aq) Ca(OH)2(s)

Cu2+ BlueCu2+

(aq) + 2OH-(aq) Cu(OH)2(s)

Fe2+ GreenFe2+

(aq) + 2OH-(aq) Fe(OH)2(s)

Fe3+ BrownFe3+

(aq) + 3OH-(aq) Fe(OH)3(s)

Al3+White but

colourless with excess NaOH

Al3+(aq) + 3OH-

(aq) Al(OH)3(s)

Mg2+ WhiteMg2+

(aq) + 2OH-(aq) Mg(OH)2(s)

Section 1: The Atmosphere

Chemistry 9: Chemistry of the Atmosphere

Section 4: Common Pollutants

Pollutant Formula Cause Effect

26 Carbon monoxide

COIncomplete combustion of a hydrocarbon fuel.

Toxic gas. Colourless and odourless so hard to detect.

27 Sulfur dioxide SO2Burning coal or petrol. Both contain sulfur which reacts with oxygen in the air.

Cause respiratory problems (e.g. for those with asthma). Combine with water vapour to cause acid rain.

28 Nitrogen oxides NOxIn car engines. N2 and O2 from air react at high temperatures.

29 Particulates CO2Incomplete combustion of a hydrocarbon fuel.

Global dimming (reduction in sunlight reaching Earth).

3 78% Nitrogen

1 21% Oxygen 2 1% Argon

4 Air also contains 0.04% carbon dioxide and variable amounts of water vapour

Section 2: Formation of the Atmosphere

5. Early Atmosphere

Atmosphere is mainly carbon dioxide with no

oxygen.

6. 4.6 – 3.6 Billion Years Ago

Volcanoes erupt releasing nitrogen and water vapour. Water vapour

condenses and forms the oceans. Some carbon dioxide dissolves in the oceans.

Carbon dioxide is also locked in fossil fuels and sedimentary rocks.

7. 2.7-1.7 Billion Years Ago

Plants evolve and release oxygen through photosynthesis. They

take in more carbon dioxide.

Section 3: Greenhouse Effect and Global Warming

11 Greenhouse effect

The process by which the temperature on Earth is kept high enough to support life by greenhouse gases absorbing radiation radiated by the Earth.

12 Greenhouse gasGreenhouse gases keep temperatures on Earth high enough to support life. Water vapour, methane and carbon dioxide are greenhouse gases.

13 Short wavelength radiation

The radiation from the Sun. Is able to pass through the Earth’s atmosphere and warm the surface of the Earth without being absorbed by greenhouse gases.

14 Long wavelength radiation

The radiation from the Earth’s surface. Some is absorbed by greenhouse gases and doesn’t escape the atmosphere.

15 Carbon footprintThe total amount of carbon dioxide and other greenhouse gases emitted over the full life cycle of a product or event.

16 Global warming The increase of the average temperature of the Earth.

18 How humans increase carbon dioxide in the atmosphere

19 How humans increase methane in the atmosphere

Combustion of fossil fuels Increased animal farming

Deforestation Decomposition of rubbish in landfill

20 How humans can decrease carbon dioxide concentration

21 How humans can decrease methane concentration

Use alternative forms of energy e.g. wind turbines

Alternative foods – non-animal based

Energy efficiency e.g. more efficient cars Increased recycling

Carbon capture – capturing CO2 from power stations and trapping it

Carbon off-setting – planting more trees

Effects of global warming

22 Some regions will not be able to produce enough food because of drought.

23 Changes to distribution of species and migration patterns.

24 Increase in sea levels because of melting of polar ice caps.

25 Reduction of water supplies in some regions.

Section 2a: Reduction of CO2 by formation of deposits

8 CoalPlants absorbed CO2. They died and decayed. This layer of decaying plants was compressed to form coal.

9 Oil and natural gas

Plankton absorbed CO2. Plankton died and were deposited in muds on the sea floor. They were covered over and compressed over millions of years.

10 LimestoneShelled animals absorbed CO2 to make their calcium carbonate shells. The remains of these animals were compressed to form limestone.

17 GreenhouseEffect