1. ACIDS AND BASES Topic 8

2. 8.1 Reactions of acids and bases Acids with metals Produces a salt and hydrogen gas Mg + 2HCl  MgCl 2 + H 2 Acids with carbonates and hydrogencarbonates Produces salt + carbon dioxide + water Na 2 CO 3 + H 2 SO 4  Na 2 SO 4 + H 2 O + CO 2

3. Con’t Acids with bases and alkalis Bases are metal oxides Produce a salt and water CuO + H 2 SO 4  CuSO 4 + H 2 O Alkalis are bases that dissolve in water Produce a salt and water NaOH + HNO 3  NaNO 3 + H 2 O

4. 8.2 Definitions of acids and bases BrØnsted-Lowry definitions Acid  proton (H + ) donor Base/ alkal  proton (H + ) acceptor Conjugate base  the base formed when an acid reacts and donates a proton to become a base Conjugate acid  the acid formed when a base reacts and accepts a proton to become an acid Referred to as conjugate acid- base pair

5. Conjugate acid-base pairs CH 3 COOH + H 2 O  CH 3 COO- + H 3 O + Which is the beginning acid? Base? Which is the acid’s conjugate? The base’s? Water is called amphoteric. What is that? Can act as an acid or base

6. Another way to phrase it… In the forward reaction the CH 3 COOH acts as the acid and the H 2 O acts as the base In the reverse reaction the CH 3 COO - acts as the base and the H 3 O + acts as the acid

7. Lewis theory of acids and bases Acid  electron pair acceptor Base  electron pair donor Must understand the Lewis structure of the compound to know which substance will accept the electrons Ex. NH 3 + H +  NH 4 + Which substance gained electrons? Which donated?

8. Con’t A dative covalent bond is always formed in a Lewis acid-base reaction What is a dative covalent bond? Both electrons come from the same atom For a substance to act as a Lewis base, it must have a lone pair of electrons For a substance to act as a Lewis acid, it must have space to accept a pair of electrons

9. 8.3 Strong and weak acids and bases When acid reacts with water it dissociates or ionizes Can use the Bronsted-Lowry theory to understand this Strong acids  completely dissociate in aqueous solution Which direction does the equilibrium dominantly lie? To the right (products) HA  H + + A - Uses a non-reversible arrow

10. Strong acids HCl is considered a monoprotic acid  it dissociates to form one proton per molecule H 2 SO 4 is considered diprotic  dissociates to form two protons per molecule H 2 SO 4 + H 2 O  HSO 4 - + H 3 O + HSO 4 - + H 2 O  SO 4 2- + H 3 O + Sulfuric acid is only considered a strong acid in the first dissociation

11. Weak acids Only partially dissociate in aqueous solution The equilibrium arrow is used for these equations HA  H + + A - Ex. Carbonated water is acidic due to dissolved CO 2, which acts as a weak acid

12. Bases Strong bases ionize completely in aqueous solution Ex. NaOH  Na + + OH - The group 1 hydroxides are strong bases; along with Ba(OH) 2 Weak bases  ionize partially in aqueous solution Equilibrium arrows are used in these equations Ex: NH 3 + H 2 O  NH 4 + + OH -

13. Distinguishing experimentally between strong and weak acids and bases Solutions of strong acids conduct electricity better than weak acids Why? There is a large concentration of ions to carry the electrical current Can also be called strong electrolytes or weak electrolytes The same concept is true for strong and weak bases

14. Con’t Strong acids have a lower pH than weak acids What does pH measure? The concentration of H + ions in solution Lower pH = more H + ions Would the pH for strong bases be higher or lower? Higher Why? There are very few H+ ions in the solution of strong bases

15. Con’t Strong acids react more violently with metals or carbonates The higher concentration of free H + ions cause a more rapid reaction with metal to form H 2 (g) There is a similar effect when a carbonate is added

16. Con’t strength vs. concentration Concentration  refers to the number of moles of acid in a certain volume (i.e. mol dm -3 ) Strength refers to what? How much the acid dissociates Ex. Ethanoic acid is considered a weak acid. No matter how concentrated the acid solution is, it will still not fully dissociate. Similar for bases

17. 8.4 pH Definition: pH is the negative logarithm to base 10 for the hydrogen ion concentration in aqueous solution pH= -log 10 [H + (aq)] The pH scale is used to indicate how acidic or alkaline a solution is The scale is from 1 to 14 One being the most acidic Fourteen is the most alkaline Seven is neutral

18. pH Since pH is on a log scale, a 1 unit change in pH means there is a tenfold change in H + ion concentration Calculating [H + ] from pH [H + ]= 10 -pH This is the inverse of the previous equation

19. Calculating pH of a strong acid It can be assumed that a strong acid fully dissociates and the [H + ] is equal to [acid] Ex: calculate the pH of a 0.00150 mol dm -3 solution of HCl. pH=-log 10 [H + ]= -log[0.00150]= 2.82 Just plug in the [acid] for hydronium ions

20. pH is not a measure of acid strength This is the measure of what? [H + ] ions It is possible for a dilute solution of a strong acid to have a higher pH than a concentrated solution of a weak acid pH can be used to compare the strength of acids,ONLY IF THE CONCENTRATIONS OF THE ACIDS ARE EQUAL