Acids and bases

1. Acids and bases Acids and bases key theory: Brønsted-Lowry theory of acids and bases: Acid-base reactions are ones in which protons are given and taken. Acids are proton donors Bases are proton acceptors

2. Conjugate pairs Every acid has what is called a conjugate base. This is particle left when the acid has given away its proton. In the same way, every base has a conjugate acid, which is the particle made when the proton has arrived. So the conjugate base of HCl is Cl- and the conjugate acid of ammonia is NH4+. Acids and bases Familiar conjugate pairs Acid Base H2O OH- H3O+ H2O H2SO4 HSO4- HSO4- SO42- CH3COOH CH3COO- HNO3 NO3- H2CO3 HCO3- HCO3- CO32- NH4+ NH3

3. Acids and bases To be a base, a substance must have a lone pair which can become a sigma pair as it accepts a proton e.g. :NH3 To be able to give protons, an acid must have a hydrogen atom which can be lost: e.g. HCl; H2SO4

4. Acids and bases Ampholytes The water in the first reaction above is acting as a base. One of its lone pairs accepts a proton. In the second reaction the water is being an acid, giving one of its protons to ammonia. Substances like water, with hydrogens that can be lost and lone pairs that can be protonated, are called ampholytes. They can be either acids or bases depending on what they react with. You can see some other common ampholytes in the list on the top right, because they appear in both the acid and base columns. Even ammonia can be an ampholyte in the right conditions, having NH4+ as its conjugate acid and NH2- as its conjugate base.

5. STRONG AND WEAK ACIDS AND BASES The equilibrium constants for the two reactions above are very different. The first one, 5.5×108, tells us that the reaction lies well over to the right. That means that there are practically no hydrogen chloride molecules left because they have almost all become ions. This solution would conduct strongly and so the hydrochloric acid is described as a strong acid. In the second reaction the constant is 3.12×10-8. This tells us that about one in a thousand ammonia molecules have reacted with water. There are hardly any ions in this solution and so it conducts weakly. Ammonia is called a weak base. So in water strong acids and bases exist mainly as ions, weak acids and bases exist mainly as molecules. Acids and bases

6. Acids and bases Common strong acids are: HCl; H2SO4; HNO3 Common weak acids are: ethanoic, CH3COOH; carbonic, H2CO3 Common strong bases are: NaOH; KOH Ammonia is a weak base.

7. Acids and basesWords Words and Expressions donor; acceptor; donate; accept conjugate: conjugate acid; conjugate base Ampholyte

8. Strength of acids The strength of an acid (how well it protonates) is measured by the equilibrium constant for the reaction: HA + H2O → H3O+ + A- Strength of acids [H2O] is nearly constant at 55.5 mol dm-3 in most bench solutions, so this expression can be simplified Ka is called the acidity or dissociation constant for the acid. If Ka is greater than 102 the acid is strong, while if it is less than 10-2 the acid is weak.

9. In the same way, it is possible to write a Kb expression for a base and a Kw expression for water. Compare the three: For an acid HA: Strength of acids For the conjugate base A-: For water Now look what happens when we multiply the first two together and cancel out terms on the top and bottom lines: Kw is called the dissociation constant for water and has the value 1×10-14 mol2 dm-6 at 25oC.

10. Complementary strengths of acids and their conjugate bases. Ka×Kb = Kw, which is a constant. So it follows that if Ka is large, Kb must be small, and if Ka is small Kb must be large. This means that if an acid is strong, its conjugate base must be weak, and vice versa. So strong acids have weak conjugate bases. Strength of acids

11. WHAT CONTROLS THE STRENGTH OF AN ACID? An energy cycle helps us focus on the two key ideas that help us answer this question: Strength of acids 1. How easy is it to break the bond joining the hydrogen to the rest of the acid particle? If the bond is weak, then the acid is likely to be strong. 2. How stable is the anion formed? If the anion is stable, the acid is likely to be strong. SO42- (sulfate ion) is a more stable anion than SO32- (sulfite ion). This is why H2SO4 (sulfuric acid) is a stronger acid than H2SO3 (sulfurous acid)

12. Strength of acidsWords Words and Expressions protonate: proton; protonation acid; acidity; base (alkali); basicity; alkaline bench: bench-scale; pilot-plant-scale; production

13. Buffer solution Buffer solutionsresist changes in pH on addition of acid or alkali Buffer solutions tend to resist changes in pH when small amounts of acid or base are added and their pH is not affected by dilution. Acidic buffers are made of a solution of a weak acid and its conjugate base, e.g. a solution of ethanoic acid and sodium ethanoate. Basic buffers are made of a solution of a weak base and its conjugate acid, e.g. a solution of ammonia and ammonium chloride.

14. Redox reactions Reduction and oxidation reactions involve the transfer or loss and gain of electrons. Redox reactions ACID-BASE AND REDOX Acid-base reactions involve proton transfer. Because protons are written as H it is usually easy to see which particle has lost a proton and which has gained one. H2O + HCl → H3O+ + Cl－ Redox reactions involve electron transfer. Because electrons do not appear in most written equations, it is not always obvious which has lost or gained electrons. But it is not hard to work this out. SO2+2H2S → 3S+2H2O gains a proton oxidized or reduced? lost or gained electrons? loses a proton

15. Redox reactions Oxidation Oxidation is defined as the loss of electrons Reduction Reduction is defined as the gain of electrons Remember ‘OIL RIG’! Oxidation is loss: reduction is gain

16. ELECTRON CONTROL The nucleus of every atom is surrounded by electrons. The electron density around the nucleus of an atom in a pure element is taken as the standard or normal situation. When an element reacts or combines with other elements, the electrons around the nucleus of each atom are rearranged so that the atom either: (i) loses some control over the electrons in its outer shell compared to the standard situation in the pure element Na → Na+ + e－ (ii) or gains some control over the electrons in its outer shell compared to the standard situation in the pure element. Cl + e- → Cl－

17. Redox reactionsWords Words and Expressions Redox: reduction-oxidation

18. Oxidation numbers The electron density around an atom in a substance is shown by oxidation numbers. These numbers indicate the electron density around an atom compared to its situation in the pure element. Oxidation numbers

19. Oxidation numbers AN OXIDATION NUMBER IS MADE OF TWO PARTS: (i) the sign ●if the sign is positive, the atom has lost control of electrons ● if the sign is negative, the atom has gained control of electrons (ii) the number is always written as a Roman number ● this gives the number of electrons over which electron control has changed compared to the situation in the pure element

20. RULES FOR WORKING OUT OXIDATION NUMBERS 1. The oxidation number of atoms in a pure element is zero. 2. In hydrogen compounds, the oxidation number of hydrogen is always +I, except in metal hydrides where it is –I. 3. In oxygen compounds the oxidation number of oxygen is always –II, except in peroxides where it is –I and in fluorine compounds where it is +II. 4. The oxidation number of an ion made from a single atom equals the value of the charge on the ion. 5. The sum of all the oxidation numbers of the atoms in the formula of a compound is zero. 6. The sum of all the oxidation numbers of the atoms in the formula of a molecular ion equals the charge on the ion. Oxidation numbers

21. RECOGNIZING REDOX EQUATIONS Redox equations are recognized by: (i) working out all the oxidation numbers of the atoms in the equation; (ii) seeing if the oxidation number of any atom has changed. If it has the reaction is a redox one. +II -II +I -I +II -I +I -II PbO + 2HCl → PbCl2 + H2O No change → not redox  +IV -II +I -I +II –I +I -II 0 PbO2 + 4HCl → PbCl2 + 2H2O + Cl2 Oxidation numbers → this is redox do change Oxidation numbers

22. Oxidation numbersWords Words and Expressions oxidation numbers hydride: metal hydride oxide; peroxide; superoxide (Na2O; Na2O2; NaO2)

23. Recognizing reduction and oxidation A change in the oxidation number of an atom during a reaction means that the atom has either lost or gained electron control. Recognizing reduction and oxidation REDUCTION If the oxidation number of an atom has become more negative, control over electrons has been gained. The substance containing that atom has been reduced. OXIDATION If the oxidation number of an atom has become more positive, control over electrons has been lost. The substance containing that atom has been oxidized.

24. Recognizing reduction and oxidation REDOX IS A TWO-WAY PROCESS Electron control lost by one atom is gained by another. For this reason, reduction and oxidation always happen at the same time. This is why the term redox is usually used. If the words oxidation or reduction are used it is because we are only focusing on one reactant.

25. OXIDIZING AGENTS OR OXIDANTS Oxidizing agents are substances that cause oxidation in the substance they react with. The substance being oxidized loses electrons to the oxidizing agent. As the oxidizing agent gains these electrons it is being reduced. So oxidizing agents can be defined as electron acceptors Common oxidizing agents are: reactive non-metals: i.e. oxygen and the halogens oxo-compounds e.g. potassium dichromate and especially in acid conditions: potassium manganate (VII) Recognizing reduction and oxidation

26. REDUCING AGENTS ON REDUCTANTS Reducing agents are substances that cause reduction in the substance they react with. The substance being reduced gains electrons from the reducing agent. As the reducing agent loses these electrons it is being oxidized. So reducing agents can be defined as electron donors Common reducing agents are: reactive metals: i.e. group 1 and 2 metals: zinc; iron; tin hydrogen, carbon, and carbon monoxide hydrogen sulphide and sulphur dioxide Recognizing reduction and oxidation

27. Recognizing reduction and oxidation TESTING FOR OXIDIZING AND REDUCING AGENTS Acids and bases are tested for using indicators. Indicators are weak acids whose conjugate base is a different color. Oxidizing and reducing agents are tested for in a similar way by using a reducing or oxidizing agent whose conjugate partner is a different color.

28. Recognizing reduction and oxidation Tests for reducing agents Add the substance to Look for the color change Oxidizing agent Conjugate reducing agent acidified potassium manganate (VII) solution purple → colorless MnO4-(aq) Mn2+(aq) acidified potassium dichromate solution orange → green/blue Cr2O72-(aq) Cr3+(aq) Tests for oxidizing agents Add the substance to Look for the color change Reducing agent Conjugate oxidizing agent iron (II) sulphate solution green → yellow Fe 2+ (aq) Fe3+(aq) potassium iodide solution colorless → yellow/brown I-(aq) I2(aq)

29. Recognizing reduction and oxidationWords Words and Expressions two-way: a two-way process oxidant: oxidizing agent; reductant: reducing agent zinc Tin acidified: acidify; basified; basify potassium dichromate potassium manganate

30. Conjugate redox pairs OXIDANT REDUCTANT Each oxidizing agent has a reduced form called its conjugate reducing agent. Cl2 Cl－ Each reducing agent has an oxidized form called its conjugate oxidizing agent. H2O H2 Strong oxidizing agents have weak conjugate reducing agents. MnO4－Mn2+ Cr2O72－Cr3+ While strong reducing agents have weak conjugate oxidizing agents. Mg2+ Mg CO2 CO

31. Conjugate redox pairs RELATIONSHIP BETWEEN CONJUGATE REDOX PAIRS AND A COMPARISON WITH ACID—BASE PAIRS OXIDANT: Causes oxidation; contains atoms in relatively high oxidation state 1. Can be easily reduced 2. Oxidation number decreasing 3. Particles good at gaining electrons Weak conjugate reductant Weak conjugate oxidant  1. Can be easily oxidized 2. Oxidation number increasing 3. Particles good at losing electrons REDUCTANT: Causes reduction; contains atoms in relatively low oxidation state

32. Conjugate redox pairs ACID: Causes protonation; contains Weak conjugate base slightly positive hydrogen atoms 1. Can be easily deprotonated 2. Good at losing protons Weak conjugate acid BASE: Causes deprotonation ; contains particles with prominent lone pairs 1. Can be easily protonated 2. Good at gaining protons

33. Conjugate redox pairsWords Words and Expressions conjugate redox pair comparison: compare

34. STANDARD ELECTRODE POTENTIALS The change in electron density around atoms during oxidation and reduction means that there is a difference in electrical potential between the oxidized and reduced forms of the element. Standard electrode potentials in calcium metal the ions are surrounded by delocalized electrons Oxidation in solution the ions are surrounded by the lone pairs of electrons from water molecules is the potential difference between the two

35. Standard electrode potentials Measuring this potential difference under standard conditions gives a comparative measure of the oxidizing or reducing power of a substance just as Ka gives a comparative measure of acid strength. Standard conditions are: 298 K 1 atmosphere pressure 1 mol dm-3 concentration The measured potential is called the standard electrode potential. It has the symbol and is measured in volts.

36. MEASURING STANDARD ELECTRODE POTENTIALS Standard electrode potentials are relative values. They are always compared to the hydrogen electrode potential. is defined as the potential difference between a standard half cell and the standard hydrogen half cell. Standard electrode potentials A STANDARD CELL BEING MEASURED

37. Standard electrode potentials FACTORS AFFECTING ELECTRODE POTENTIALS A change in conditions produces a change in electrode potential. They will be changed by: 1. Pressure of gases 2. Temperature 3. Concentration of the ions 4. pH of solutions 5. Ligands present e.g. Fe(H2O)63+ + e－Fe(H2O)62+ +0.77 volts Fe(CN)63- + e－Fe(CN)64- +0.34 volts

38. STANDARD HALF CELLS Standard electrode potentials

39. Standard electrode potentials Words and Expressions electrode; anode; cathode electrode potential; standard electrode potential comparative; compare; comparison salt bridge cell; half cell black platinum

40. Predicting redox reactions The value of the standard electrode potential, E$, for a half cell gives a measure of the relative stability of the conjugate oxidizing and reducing agents. Predicting redox reactions A POSITIVE E$value tells us that the equilibrium position for the half-cell reaction lies to the right compared to the hydrogen half equation. It is positive because electrons are being used up .

41. Predicting redox reactions A NEGATIVE E$ value tells us that the equilibrium position for the half-cell reaction lies to the left compared to the hydrogen half equation. It is negative because electrons are being produced.

42. HOW TO PREDICT WHITHER A REDOX REACTION WILL HAPPEN Acid-base reactions can be predicted using Ka values. In the same way E$ values can be used to predict redox reactions. Predict whether a redox reaction will go like this: Predicting redox reactions 1. Look up the two redox half equations in the data book (remember that the oxidizing agents are always written on the left). 2. Write the half equations down with the one with the more positive E$ value on top. This gives the order of oxidizing and reducing strength. 3. Reaction is likely to occur between the stronger oxidizing agent and the stronger reducing agent. The products will be the weaker reducing and oxidizing agents. Fe3+ + e- → Fe2+ E = +0.77 V Cl2 + 2e- → 2Cl- E = +1.36 V stronger Cl2 + 2e- → 2Cl- stronger oxidant Fe3++ e- → Fe2+ reductant Reactants: Cl2 + Fe2+ products: Cl- + Fe3+ The bigger the difference in E$ values the more likely is the reaction to go.

43. Predicting redox reactions ‘WRONG’ PREDICTIONS Sometimes, reactions predicted by this method do not happen. The reason for this is either ● the reaction has a very high activation energy(E$ values only tell us about the relative energies of the reactants and products, not about the reaction pathway) or ● the actual conditions of the reaction are not standard and so the actual E values are different from the standard E$ values.

44. Balancing and using redox equations BALANCING REDOX EQUATIONS Once you have predicted that a reaction will happen: 1. Copy down the two half equations from the data book with ● the more positive one on top ● the more negative one underneath and reversedso that all the reactants are on the left. 2. Scale the two equations so that the number of electrons involved in each is the same. 3. Add the two half equations together. 4. Cancel out any ions or molecules which appear on both sides of the equation.

45. Balancing and using redox equations e.g. For the reaction between manganate (Vll) and hydrogen peroxide MnO4-/Mn2+ E = +1.58; O2/H2O2 E = +0.68 MnO4- + 8H+ + 5e- Mn2+ 4H2O +1.58 H2O2 O2 + 2H+ + 2e－+0.68 2MnO4- + 16H+ +10e- 2Mn2+ + 8H2O 5H2O2 5O2 + 10H+ + 10e－ 2MnO4- + 16H+ + 5H2O2 2Mn2+ + 8H2O + 5O2 + 10H+ 2MnO4- + 6H+ + 5H2O2 2Mn2+ + 8H2O + 5O2

46. PREDICTING INFORMATION ABOUT CELLS FROM E$ VALUES E$ values can be used to predict changes which would happen in a cell made up of two half cells. For examples, take a cell made up of a copper half cell and an iron(III)/iron(II) half cell. Look up the E$ values and write them on a number line like this Balancing and using redox equations

47. Balancing and using redox equations CHANGING CONDITIONS IN THE CELLS The reaction in each half cell is an equilibrium whose position can be changed by changing the concentrations of the ions. So increasing [Fe3+] will increase the cell potential while increasing [Cu2+] will decrease the cell potential.

48. Balancing and using redox equationsWords Words and Expressions underneath (below); on the top reverse scale cell potential