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AQA AS Chemistry Revision

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1 AQA AS Chemistry Revision
Chapter 10 Redox AS Unit 2 Revision

2 AS Unit 2 Revision

3 Oxidation and Reduction
Oxidation is loss of electrons. Reduction is gain of electrons. An oxidising agent accepts electrons and is itself reduced. A reducing agent donates electrons and is itself oxidised. AS Unit 2 Revision

4 Oxidation states The oxidation state of an atom in an element is always zero. The sum of the oxidation states of all the atoms and ions in a compound is always zero. The oxidation state of a simple ion is it’s charge. The sum of the oxidation states of all the atoms and polyatomic ion is it’s charge. The oxidation state of Flourine is always -1 as it is the most electronegative element. The oxidation state of oxygen is nearly always -2. (Except in peroxides where it is -1) The oxidation state of chlorine in a compound is usually -1 (unless it is combined with F or O). The oxidation state of hydrogen is +1 (except when it is bonded to a metal ion in which case it is -1) AS Unit 2 Revision

5 Oxidation and reduction of s and p block elements
S block - Group 1 and 2 Group 1 elements have one electron in their outer shell so always makes +1 ions. Oxidation state is always +1. They are reducing agents. Group 2 elements have one electron in their outer shell so always makes +2 ions. Oxidation state is always +2. They are reducing agents. AS Unit 2 Revision

6 Examples For the following reactions
Potassium reacts with oxygen to form K2O Calcium reacts with water to form Ca(OH)2 Write full balanced equations and the Write two half equations. Assign oxidation states to each element and determine which has been oxidised and which reduced. AS Unit 2 Revision

7 Oxidation and reduction of s and p block elements
P block - Group 5, 6 and 7 Group 5, E.g. Nitrogen can form compounds with metals or non-metals. They can have positive or negative oxidation states. The more electronegative element in a compound is always taken as the negative element when assigning oxidation states. AS Unit 2 Revision

8 Examples of Nitrogen Assign oxidation states and determine which is oxidised, which is reduced and name the oxidising agents and reducing agents. 3Mg(s) + N2(g) → Mg3N2(s) N2(g) + 3H2(g) ↔ 2NH3(g) N2(g) + O2(g) → 2NO AS Unit 2 Revision

9 Oxidation and reduction of s and p block elements
P block - Group 5, 6 and 7 Group 6, Sulfur can exist in several oxidation states. Group 7, Halogens always take -1 oxidation state by gaining 1 electron. When they react with metals that can take on variable oxidation states, the metal will tend to adopt its higher oxidation state. E.g. passing chlorine gas over hot iron forms FeCl3 instead of FeCl2 They can also react with non-metals to form covalent compounds but usually still retain the -1 oxidation state. AS Unit 2 Revision

10 Redox equations Two good power points for this
Balancing redox equations Redox equations Now try PPQs AS Unit 2 Revision

11 AQA AS Chemistry Revision
Chapter 11 Group 7, the halogens AS Unit 2 Revision

12 AS Unit 2 Revision

13 When warmed, iodine crystals sublime (turn directly to a gas), forming a purple vapour.
AS Unit 2 Revision

14 Atomic Radius The atomic radius increases down Group 7.
Going down the group, there are more filled energy levels between the nucleus and the outer electrons. This results in the outer electrons being shielded more from the attraction of the nucleus. AS Unit 2 Revision

15 The boiling point increases down Group 7.
The strength of the instantaneous dipole−induced dipole forces between the molecules increases as the size of the molecules increases. AS Unit 2 Revision

16 Electronegativity Electronegativity decreases down Group 7.
Fluorine is the most electronegative element in the periodic table. The atomic radius increases, the outer electrons are more shielded, so bonding electrons are less strongly attracted to the nucleus. AS Unit 2 Revision

17 Oxidation ability What is Oxidation?
Oxidation is the loss of electrons. What is an oxidizing agent? An oxidizing agent is an electron acceptor, the agent is reduced during the course of the reaction. This forms a redox reaction. 17 AS Unit 2 Revision 17

18 Oxidising power trend: Cl2 > Br2 > I2
When a halogen acts as an oxidising agent, it gains electrons (taken from the oxidised species). X e- → 2 X- Cl Br I Going down the group it becomes harder to gain an electron because: atoms are larger & there is more shielding (due to extra electron shell) AS Unit 2 Revision

19 Reducing agents 2 X– → X2 + 2 e– What is reduction?
Reduction is the gain of electrons What happens when a Halide is used as a reducing agent? Give the half equation for the reaction 2 X– → X e– When a halide ion reduces another substance, the halide is oxidised to a halogen. 19 AS Unit 2 Revision

20 Reducing power trend: Cl– < Br– < I–
When a halide ion acts as a reducing agent, it loses electrons (given to the reduced species). 2 X– → X e– Cl– Br– I– Down the group it becomes easier to lose an electron because: ions are larger & there is more shielding (due to extra electron shell) AS Unit 2 Revision

21 Reactions with silver nitrate
The samples to be tested are first acidified using dilute nitric acid, then silver nitrate solution is added. Nitric acid is needed to get rid of carbonates or hydroxides. These would form silver carbonates or silver hydroxides which are soluble. Chloride Bromide Iodide White precipitate of silver chloride Cream precipitate of silver bromide Yellow precipitate of silver iodide Ag+(aq) + Cl−(aq) → AgCl(s) Ag+(aq) + Br−(aq) → AgBr(s) Ag+(aq) + I−(aq) → AgI(s) AS Unit 2 Revision

22 Uses of Chlorine and Chlorate
Chlorine and water Cl2(g) + H2O(l) ↔ HClO(aq) + HCl(aq) This is disproportionation. Chlorine and alkali Cl2(g) + 2NaOH(l) → NaClO(aq) + NaCl(aq) + H20(l) Oxidising agent that kills bacteria. NaClO is an oxidising agent and the active ingredient in bleach AS Unit 2 Revision

23 AQA AS Chemistry Revision
Chapter 11 Group 2, the alkaline earth metals AS Unit 2 Revision

24 AS Unit 2 Revision

25 Physical Properties of Group 2
All have 2 electrons in outer s orbital. As you go down the group there is more shielding and outer electron is farther from the nucleus. Atomic radius increases down the group. Melting point decreases down the group. Ionisation energy decreases down the group. AS Unit 2 Revision

26 Reactions of Group 2 Elements with Water
With steam Mg Ca Sr Ba Burns vigorously Insignificant reaction  white MgO solid + colourless H2 gas Moderately fast reaction Burns very vigorously  alkaline Ca(OH)2 (some white ppt) + colourless H2 gas  white CaO solid + colourless H2 gas Fast reaction Burns explosively  alkaline colourless Sr(OH)2 solution + colourless H2 gas  white SrO solid + colourless H2 gas Very fast reaction Burns explosively  alkaline colourless Ba(OH)2 solution + colourless H2 gas  white BaO solid + colourless H2 gas AS Unit 2 Revision concentrate on these

27 Ba2+(aq) + SO42-(aq) → BaSO4(s)
Test for Sulphates Reacting the sulphate with barium chloride forms an insoluble salt Barium sulfate, this is seen as a precipitate. Ba2+(aq) + SO42-(aq) → BaSO4(s) AS Unit 2 Revision

28  MgSO4 Mg(OH)2   CaSO4 Ca(OH)2 SrSO4 Sr(OH)2 BaSO4 Ba(OH)2
more soluble more soluble Remember Opposite trends Ca(OH)2 & CaSO4 sparingly soluble BaSO4 insoluble AS Unit 2 Revision

29 Uses of Group 2 compounds
Barium meal BaSO4 Milk of magnesia Mg(OH)2 Lime Ca(OH)2 AS Unit 2 Revision

30 AQA AS Chemistry Revision
Chapter 13 Extraction of metals AS Unit 2 Revision

31 AS Unit 2 Revision

32 Converting sulfide ores to oxides
Roasting in oxygen producing SO2(g) 2CuS + 3O2(g)  2CuO(s) + 2SO2(g) Problem of acid rain. SO2 + H2O  H2SO3 and SO3 + H2O  H2SO4 Reducing oxides Possible reductants Coke – impure form a carbon. Pros – cheap, cons – requires high temp and can form brittle metal carbides. Hydrogen – used to reduce W tungsten Electrolysis – Pros – gives pure product, cons – expensive, lots of energy required. More reactive metals as reductants. Con – expensive. AS Unit 2 Revision

33 The Blast Furnace 1. A mixture of limestone, coke and haematite is added at the top of the blast furnace. 3. The carbon reacts with the hot air to form carbon dioxide. 4. The coke reacts with the carbon dioxide to form carbon monoxide. 5. The carbon monoxide reduces the iron ore to form iron and carbon dioxide. 8. Waste gases are removed from the blast furnace. 2. Hot air is ‘blasted’ into the blast furnace. 6. The iron, pig iron, collects at the bottom of the blast furnace to be tapped off. 1900°C 7. Impurities collect at the bottom in a layer, ‘slag’, and can be tapped off. AS Unit 2 Revision 33

34 Extracting aluminium – redox equations
Boardworks GCSE Additional Science: Chemistry Electrochemistry At the negative electrode: Al3+ + 3e-  Al (reduction) At the positive electrode: 2O2-  O2 + 4e- (oxidation) aluminium oxide  aluminium + oxygen 2 Al2O3 (l)  4 Al (l) + 3 O2 (g) Cryolyte needed to form a solution so Al can melt at a lower temperature. Carbon electrodes burn away to CO2 and must be replaced. AS Unit 2 Revision 34

35 Extraction of Titanium
Titanium ore is mainly the oxide TiO2, which is converted into titanium tetrachloride TiCl4 by heating with carbon and chlorine. Titanium chloride is a gas and must be separated using fractional distillation under an inert atmosphere. TiO2(s) + 2C(s) + 2Cl2(g) → TiCl4(l) +2CO(g) AS Unit 2 Revision

36 Extraction of titanium from its ore
The titanium tetrachloride is then reacted with sodium or magnesium TiCl4 + 2Mg Ti + 2MgCl2 or TiCl4 + 4Na  Ti + 4NaCl This reaction is carried out in an atmosphere of inert argon gas so none of the metals involved becomes oxidised by atmospheric oxygen. Na is used in UK AS Unit 2 Revision

37 Cooled in between and Ti sponge removed. (Takes days to cool)
Batch process Cooled in between and Ti sponge removed. (Takes days to cool) Sponge has air spaces and contains 30% impurities, e.g. MgCl2. Wash with HCl to remove MgCl Removed by evaporation at high T. AS Unit 2 Revision

38 Disadvantages of this process
Cl, Na and Mg are expensive, must be produced by electrolysis. High temperatures are needed. TiCl4 is highly reactive with water and must be handled with care. An inert atmosphere is required to avoid reactions with oxygen. AS Unit 2 Revision

39 Tungsten Also known as wolfram, symbol W, from the ore wolframite.
High MP (3410°C) therefore used in light bulbs. Cannot use carbon as tungsten carbide would form. Must be reacted with Hydrogen at high temperatures to be reduced. Draw back is that H2 is so flammable. °C WO3(S) + 3H2(g)  W(s) + 3H2O(l) AS Unit 2 Revision

40 Scrap metals can be used instead of ores. Iron
Less land fill All ready extracted from ore, easily separated as it is magnetic. Produces less CO2 than extraction from ore. Aluminium Avoids pollution Uses about 5% energy needed to extract from ore. AS Unit 2 Revision

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