1. C1 – Science Thursday 16th May 2019
Key Revision Points
Combined
HT = On Higher Tier Paper ONLY

2. Elements, Compounds and Mixtures
All substances are made of atoms. An atom is the smallest part of an element that can exist
Elements are found in the periodic table
Compounds are made from wo or more elements chemically bonded  compounds can ONLY be separated by chemical reactions into elements
Mixtures consist of 2 or more elements OR compounds that are not chemically combined, therefore the chemical properties of each substance in a mixture remain unchanged

3. Separating Mixtures (How to?)
Filtration
Distillation

4. Development of the Atom
1. Before electron discovered, atoms thought to be tiny undividable spheres 2. Electron discovered  Plum Pudding ball of positive charge with electrons embedded in it 3. Nuclear Model  Alpha Particle Scattering XP  Atom was concentrated at the centre (nucleus – was charged). WHY?  the alpha particles bounced back from the gold foil at different angles, the positive alpha particles were being repelled, this was unexpected as they thought they would travel straight through. 4. Niels Bohr  suggested that electrons orbit the nucleus at specific distances 5. Later experiments subdivided the positive charge into smaller positive charges called protons. 6. James Chadwick  provided evidence of neutrons in nucleus

5. Charge, Size and Mass of subatomic particles
Number of electrons = number of protons therefore atoms have no overall charge (nucleus is positive)
Number of protons = atomic number (small number)
Mass Number = Number of protons + number of neutrons (Massive number)
Atoms are very small, having a radius of about 0.1 nm (1 x m).
The radius of a nucleus is less than 1/ of that of the atom(about 1 x m)

6. Isotopes and RAM
Isotopes have the same proton/atomic number therefore same number of protons
but DIFFERENT number of neutrons
Calculating RAM from Isotopes
E.g.  in any sample of Chlorine 25% will be 37Cl and 75% 35Cl. 

7. Electronic Configuration
The electrons in an atom occupy the lowest available energy levels (innermost available shells).
You need to be able to draw the first 20 elements up to Ca and the electronic configuration (i.e. 2.8.X)

8. Periodic Table Elements are in order of PROTON/ATOMIC number
Elements with similar properties are placed in groups
Elements in the same group have the same number of electrons in the outer shell (i.e. Grp. 2=2; Grp 7=7)
Moving across the PT (Periods) the electrons in the outer shell increase by 1; BUT moving down the PT (groups) the number of shells increase by 1, but the number of electrons in the outer shell remains the same

9. Development of Periodic Table
1. Before P, E, N  elements were arranged by their atomic weight 2. The first PT were incomplete and elements were in wrong groups, if atomic weight was followed 3. Mendeleev  overcame problem  by placing gaps of undiscovered elements, and changed the order based upon atomic weights 4. Elements predicted by Mendeleev were then discovered and filled gaps. Isotope knowledge explained why the atomic weight wasn’t always correct

10. Group 0 These are called NOBLE GASES
They are unreactive because they have a full outer shell (8 apart from He which has 2)
They do not form molecules easily as they have a stable electron arrangement
The B.P. of Group 0 increases as you go down the group

11. Group 1 These are called ALKALI METALS  1 electron in outer shell
The reactivity of the elements increase as you go down the group  Why? (1) atoms larger (2) outer electron is further away from nucleus (3) attraction between nucleus and outer electron is weaker as you go down the group
Metal + Oxygen  Metal Oxide
(e.g. Lithium + Oxygen  Lithium Oxide)
Metal + Chlorine  Metal Chloride (e.g. Lithium Chloride)

12. Group 7
These are called HALOGENS  7 electrons in outer shell  Non-metals  molecule pairs
The RAM, B.P. and M.P. increase as you go down the group.
BUT reactivity of the elements decrease as you go down the group  Why? Same 3 reasons as Grp. 1 BUT  therefore harder to attract electron
1. Reactions of Halogens + Metals
2. Reactions of Halogens + Non-Metals (Displacement reactions)

13. Chemical Bonding
3 types of chemical bonding: (i) Covalent – sharing pairs of electrons (ii) Ionic – transfer of electrons and oppositely charged particles (iii) Metallic – delocalised electrons
See laminate poster on ALL examples of Covalent bonding and ALL 4 possible combinations of Ionic bonding  make sure you read the rules and write those as answers (especially for Ionic)

14. Ionic Compounds
Ionic compound = giant structure of ions  held together by strong ELECTROSTATIC forces  between oppositely charged ions  act in all directions of LATTICE
Q  Describe the limitations of using dot and cross, ball and stick, two and three-dimensional diagrams to represent a giant ionic structure

15. Properties of Ionic Compounds
Regular giant structures  giant ionic lattices
High M.P and B.P.  large amount of energy needed to break strong bonds
When melted/dissolved in water  conduct electricity  because ions are free to move so charge can flow
Properties of small molecules
Usually gases/liquids  low M.P and B.P.
Weak InterMolecular (IM) forces; BUT, bigger the size of molecule the bigger the IM forces therefore the higher the M.P. and B.P. (IM forces weaker than covalent bonds)
They do not conduct electricity  why?  molecules have no overall charge

16. Metallic Bonding Metals  giant structures of atoms  regular pattern
(1) electrons in outer shell delocalised (2) they are free to move through whole structure (3) sharing of delocalised electrons = strong metallic bonds  therefore good conductors of electricity and thermal energy
Metals  High MP and BP  arranged in layers  can be bent and shaped (pure metals too soft)  so mixed with other metals to make  alloys
Alloys  harder  because of distorted layers of atoms

17. 3 states of matter
Inter-molecular forces stronger in solid than liquid and gas.
(HT only) Limitations of the simple model above include that in the model there are no forces, that all particles are represented as spheres and that the spheres are solid.

18. Polymers
Very large molecules  linked to other atoms by strong covalent bonds  the IM forces between polymer molecules are strong  they are therefore solids at room temp.
You need to be able to recognise polymer structures  they look like this

19. Giant Covalent Structures
Giant covalent structures  solids  high MP  strong covalent bonds  hard to melt/boil
Diamond  4 covalent C-bonds  v.hard, high MP, does not conduct electricity
Graphite  3 covalent C-bonds  hexagonal rings no covalent bonds between layers; 1 electron is delocalised from each C atom  does conduct
Graphene  single layer of graphite  electronics use
Fullerenes  C-atoms with hollow shapes  hexagonal rings (6), but can be 5 and 7 rings too. (Buckminsterfullerene C-80!)
Nanotubes -- .cylindrical fullerenes  high length to diameter ratios  useful in nanotech., electronics and materials
Graphene
Graphite
Fullerene
Diamond

20. Conservation of Mass & Relative Formula Mass (RFM)
The law of conservation of mass states that no atoms are lost or made during a chemical reaction
Mass of reactants = mass of products
If the reactant does not = product  the missing mass might be a gas and it was lost

21. Moles (HT ONLY) The mass of a mol = to the RFM of that substance
1 mole of a substance contains the same number of particles as another mole of a substance
The number of atoms, molecules or ions in a mole of a given substance is the Avogadro constant. The value of the Avogadro constant is 6.02 x 1023 per mole.
Calculate the number of moles of carbon dioxide molecules in 22 g of CO2.
Calculate the mass of 2 mol of carbon dioxide (CO2).

22. Reacting Masses (HT ONLY)
Chemical equations can be interpreted in moles
In this equation
1 mole of Mg + 2 moles of HCl react to produce 1 Mole of Magnesium Chloride and 1 mole of H2
Calculate the mass of sodium sulfatemade when 20 g of sodium hydroxide reacts with excess sulfuric acid. (Ar of H = 1, Ar of O = 16, Ar of Na = 23, Ar of S = 32)

23. Concentration of solutions
Calculate the mass of solute in a given volume of solution of known concentration
Use equation: concentration = mass / volume
Units for concentration: grams per dm3 (g/dm3)
HT ONLY
Rearrange the above equation – you will need to remember it!

24. Extraction of metals & OILRIG
Unreactive metals = gold found by itself BUT most metals are found as compounds and need extraction RULE 1 Metals less reactive than Carbon  can be extracted by reduction (LOSS of OXYGEN) Metals that react with Oxygen  is OXIDATION (GAIN of OXYGEN) HOWEVER (HT ONLY) – RULE 2 OILRIG (Oxidation is Loss and Reduction is Gain of ELECTRONS too)

25. Acids + Metal (REDOX) (HT ONLY)
ACID + METAL  SALT + HYDROGEN (All – rest below HT)
Metals: Zinc, Magnesium and Iron [ARE OXIDISED]
Acids: Hydrochloric and Sulphuric
Zinc/Mg/Iron + Hydrochloric Acid  Zinc/Mg/Fe Chloride + Hydrogen
Zinc/Mg/Iron + Sulphuric Acid  Zinc/Mg/Fe Sulphate + Hydrogen
All the metals are Oxidised (loss of electrons) (OIL) – see ionic equation

26. Neutralisation of Acids (Acid + Alkali)
ACID + ALKALI  SALT + WATER
Alkali could be  Metal hydroxide or metal oxide
ACID + METAL CARBONATES  SALT + WATER + CO2
Acids: Hydrochloric; Sulphuric; Nitric
Salts produced: Chloride; Sulphate; Nitrate

27. Soluble Salts (Making crystals)
Soluble salts = acids + insoluble substances (metals, metal oxides, hydroxides and carbonates) – see prev.
REQ. Practical – Crystallisation Method
1. Excess solid (metal oxide/hydroxide/carbonate) added to acid (sulphuric, hydrochloric/nitric). (Volume and mass required)
2. Filtered  produce solution (filter paper/funnel)
3. Solution is placed in
evaporating dish and heated
(to evaporate water)
4. Left (certain amount of time)
to produce crystals
Most common example is
Copper Oxide + Sulphuric Acid

28. pH and Neutralisation Acids = H+ ions Alkalis = OH- ions HT ONLY
Neutralisation reactions between an Acid + Alkali produces Water: Ionic Equation below
HT ONLY
Strong acids (HCl, Suphuric & Nitric)– completely ionise in (aq) solution.
Weak acids (ethanoic, citric & carbonic) partially ionise
As the pH decrease by 1 unit, the H+ ion concentration increases by a factor of 10 (see below from pH3 to pH6.

29. Electrolysis (Req. Prac)
Ionic compounds needs to be melted or dissolved in water  so ions are free to move about  they can then conduct electricity  they are called electrolytes
Positive ions  move to negative Cathode & Negative ions  move to positive Anode [PANIC]
Uses
(i) ionic compounds
(ii) extract metals – if metal is too reactive to be extracted by carbon (reduction) e.g. Al  from aluminium oxide + cryolite (using Carbon as anode)
(iii) electrolysis of aqueous solutions (Rules) – 1- At the negative electrode (cathode), hydrogen is produced if the metal is more reactive than hydrogen. – 2- At the positive electrode (anode), oxygen is produced unless the solution contains halide ions when the halogen is produced.

30. Electrolysis (HT ONLY) (Half - Equations)
At the cathode (negative electrode), positively charged ions gain electrons and so the reactions are reductions.
At the anode (positive electrode), negatively charged ions lose electrons and so the reactions are oxidations.
Sodium
Chloride
Lead bromide

31. Exo and Endothermic reactions
Energy is conserved in a reaction (starts = end)
EXOTHERMIC reactions – transfer energy to surroundings (so the temp. of the surroundings increase – gets hotter
ENDOTHERMIC reactions – takes in energy from the surrounds (so the temp. of the surroundings decrease) – gets cooler
Exothermic examples: combustion, neutralisation and oxidation
Endothermic examples: thermal decomposition, electrolysis, sport injury packs & citric acid + sodium hydrogen carbonate
REQ Practical: investigate the variables that affect temperature changes in reacting solutions e.g. monitor the temperature rise as small volumes of sodium hydroxide solution are added to dilute hydrochloric acid in an insulated cup

32. Exo & Endo Reaction Diagrams
You will need to label Activation energy
Make sure you know whether products are lower or higher than reactants
Every reaction has an activation energy 'barrier' (the black line 'humps' of height Ea) that must be overcome before a particle collision can lead to a chemical change. The bigger the 'hump' the bigger the energy needed.

33. Bond Energy Calculations (HT ONLY)
Energy is needed to break bonds in reactants
Energy is released when bonds in the products are formed
EXOTHERMIC  energy released is GREATER than energy needed to break existing bonds
ENDOTHERMIC  energy needed to break bonds is GREATER than energy realised from making new bonds
EXO
ENDO