1. 1 Introduction to Acids and Bases The earliest definition was given by Arrhenius: An acid contains a hydrogen atom and dissolves in water to form a hydrogen ion, H +. A base contains hydroxide and dissolves in water to form OH −. HCl(g) acid H + (aq) + Cl − (aq) NaOH(s) base Na + (aq) + OH − (aq)

2. 2 Introduction to Acids and Bases The Arrhenius definition correctly predicts the behavior of many acids and bases. However, this definition is limited and sometimes inaccurate. For example, H + does not exist in water. Instead, it reacts with water to form the hydronium ion, H 3 O +. H 3 O + (aq)H + (aq) + H 2 O(l) hydrogen ion: does not really exist in solution hydronium ion: actually present in aqueous solution

3. 3 Introduction to Acids and Bases The Brønsted–Lowry definition is more widely used: A Brønsted–Lowry acid is a proton (H + ) donor. A Brønsted–Lowry base is a proton (H + ) acceptor. H 3 O + (aq) + Cl − (aq)HCl(g) + H 2 O(l) This proton is donated. HCl is a Brønsted–Lowry acid because it donates a proton to the solvent water. H 2 O is a Brønsted–Lowry base because it accepts a proton from HCl.

4. 4 Introduction to Acids and Bases Brønsted–Lowry Acids A Brønsted–Lowry acid must contain a hydrogen atom. Common Brønsted–Lowry acids (HA): HCl hydrochloric acid HBr hydrobromic acid H 2 SO 4 sulfuric acid HNO 3 nitric acid

5. 5 Introduction to Acids and Bases Brønsted–Lowry Bases A Brønsted–Lowry base is a proton acceptor, so it must be able to form a bond to a proton. A base must contain a lone pair of electrons that can be used to form a new bond to the proton. N H H H + H 2 O(l) N H H H H + + OH − (aq) Brønsted–Lowry base This e − pair forms a new bond to a H from H 2 O.

6. NH 3 ammonia 6 Introduction to Acids and Bases Brønsted–Lowry Bases Common Brønsted–Lowry Bases(B ) NaOH sodium hydroxide KOH potassium hydroxide Mg(OH) 2 magnesium hydroxide Ca(OH) 2 calcium hydroxide H 2 O water Lone pairs make these neutral compounds bases. The OH − is the base in each metal salt.

7. 7 Proton Transfer The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base HA+ B A − HB+B+ + gain of H + acidbase loss of H + This e − pair stays on A. This e − pair forms a new bond to H +.

8. 8 Proton Transfer The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base The product formed by loss of a proton from an acid is called its conjugate base. HA+ B A − HB+B+ + gain of H + acidbaseconjugate base conjugate acid loss of H + The product formed by gain of a proton by a base is called its conjugate acid.

9. 9 Proton Transfer The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base HBr+ + gain of H + acidbaseconjugate base conjugate acid loss of H + H2OH2O Br − H3O+H3O+ HBr and Br − are a conjugate acid–base pair. H 2 O and H 3 O + are a conjugate acid–base pair. The net charge must be the same on both sides of the equation.

10. 10 Proton Transfer The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base Amphoteric compound: A compound that contains both a hydrogen atom and a lone pair of e − ; it can be either an acid or a base. HOH H 2 O as a base add H + HOH H + conjugate acid HOH H 2 O as an acid remove H + HO − conjugate base

11. 11 Acid and Base Strength Relating Acid and Base Strength When an acid dissolves in water, the proton transfer that forms H 3 O + is called dissociation. When a strong acid dissolves in water, 100% of the acid dissociates into ions. Common strong acids are HI, HBr, HCl, H 2 SO 4, and HNO 3. A single reaction arrow is used, because the product is greatly favored at equilibrium. H 3 O + (aq) + Cl − (aq)HCl(g) + H 2 O(l)

12. 12 Acid and Base Strength Relating Acid and Base Strength When a weak acid dissolves in water, only a small fraction of the acid dissociates into ions. Unequal reaction arrows are used, because the reactants are usually favored at equilibrium. H 3 O + (aq) + CH 3 COO − (aq)CH 3 COOH(l) + H 2 O(l)

13. 13 Acid and Base Strength Relating Acid and Base Strength When a strong base dissolves in water, 100% of the base dissociates into ions. Na + (aq) + OH − (aq)NaOH(s) + H 2 O(l) Common strong bases are NaOH and KOH. When a weak base dissolves in water, only a small fraction of the base dissociates into ions. NH 4 + (aq) + OH − (aq)NH 3 (g) + H 2 O(l)

14. 14 Acid and Base Strength Relating Acid and Base Strength A strong acid readily donates a proton, forming a weak conjugate base. HCl strong acid Cl − weak conjugate base A strong base readily accepts a proton, forming a weak conjugate acid. OH − strong base H 2 O weak conjugate acid

15. 15 Dissociation of Water H OH base HOH H + conjugate acid HOH conjugate base HO − ++ loss of H + gain of H + Water can behave as both a Brønsted–Lowry acid and a Brønsted–Lowry base. Thus, two water molecules can react together in an acid–base reaction:

16. 16 Dissociation of Water Pure water contains an exceedingly low concentration of ions, H 3 O + and – OH. Since one H 3 O + ion and one – OH ion are formed in each reaction, their concentrations are equal in pure water. Multiplying these concentrations together gives the ion-product constant for water, Kw. K w = [H 3 O + ][OH − ] ion-product constant [H 3 O + ] = [OH − ] = 1.0 × 10 –7 M at 25 °C.

17. 17 Dissociation of Water Substituting the concentrations for H 3 O + and – OH into the expression for K w gives the following result. K w = [H 3 O + ] [OH − ] K w = (1.0 x 10 −7 ) x (1.0 x 10 −7 ) K w = 1.0 x 10 −14 K w is a constant, 1.0 x 10 −14, for all aqueous solutions at 25 o C.

18. 18 Dissociation of Water To calculate [ − OH] when [H 3 O + ] is known: To calculate [H 3 O + ] when [ − OH] is known: K w = [H 3 O + ][OH − ] [OH − ] = 1 x 10 −14 [H 3 O + ] [OH − ] = KwKw [H 3 O + ] [OH − ] = 1 x 10 −14 [H 3 O + ] [OH − ] = KwKw [H 3 O + ]

19. 19 Dissociation of Water If the [H 3 O + ] in a cup of coffee is 1.0 x 10 −5 M, then the [ − OH] can be calculated as follows: [OH − ] = KwKw [H 3 O + ] = 1 x 10 −14 1 x 10 −5 =1.0 x 10 −9 M In this cup of coffee, therefore, [H 3 O + ] > [OH − ], and the solution is acidic overall.

20. 20 Dissociation of Water

21. 21 The pH Scale Calculating pH pH = −log [H 3 O + ] The lower the pH, the higher the concentration of H 3 O +. Acidic solution: pH 1 x 10 −7 Basic solution: pH > 7  [H 3 O + ] < 1 x 10 −7 Neutral solution: pH = 7  [H 3 O + ] = 1 x 10 −7

22. 22 The pH Scale Calculating pH from [H 3 O + ] If [H 3 O + ] = 1.0 x 10 –5 M for a urine sample, what is its pH? pH = –log [H 3 O + ] = –log (1.0 x 10 –5 ) = –(–5.00) = 5.00 The urine sample is acidic because the pH < 7.

23. 23 The pH Scale Calculating [H 3 O + ] from pH If the pH of seawater is 8.50, what is the [H 3 O + ]? pH = −log [H 3 O + ] 8.50 = −log [H 3 O + ] −8.50 = log [H 3 O + ] antilog (−8.50 ) = [H 3 O + ] [H 3 O + ] = 3.2 x 10 −9 M The seawater is basic because [H 3 O + ] > 1 x 10 –7 M.

24. 24 The pH Scale Calculating pH A logarithm has the same number of digits to the right of the decimal point as are contained in the coefficient of the original number. [H 3 O + ] = 3.2 x 10 −9 M two significant figures pH = 8.50 two digits after decimal point pH = 8.50

25. 25 Focus on the Human Body The pH of Body Fluids

26. 26 Common Acid–Base Reactions Reaction of Acids with Hydroxide Bases Neutralization reaction: An acid-base reaction that produces a salt and water as products. HA(aq) + MOH(aq) acid base OH(l) + MA(aq) H watersalt The acid HA donates a proton (H + ) to the OH − base to form H 2 O. The anion A − from the acid combines with the cation M + from the base to form the salt MA.

27. 27 Common Acid–Base Reactions HOW TO Draw a Balanced Equation for a Neutralization Reaction Between HA and MOH Example Write a balanced equation for the reaction of Mg(OH) 2 with HCl. Step [1] Identify the acid and base in the reactants and draw H 2 O as one product. HCl(aq) + Mg(OH) 2 (aq) acid base H 2 O(l) + water salt

28. 28 Common Acid–Base Reactions HOW TO Draw a Balanced Equation for a Neutralization Reaction between HA and MOH Step [2] Determine the structure of the salt. The salt is formed from the parts of the acid and base that are not used to form H 2 O. HCl H + reacts to form H 2 O Cl − used to form salt Mg(OH) 2 Mg 2+ used to form salt 2 OH − react to form water Mg 2+ and Cl − combine to form MgCl 2.

29. 29 Common Acid–Base Reactions HOW TO Draw a Balanced Equation for a Neutralization Reaction between HA and MOH Step [3] Balance the equation. HCl(aq) + Mg(OH) 2 (aq) acid base H 2 O(l) + water salt MgCl 2 2 Place a 2 to balance Cl. Place a 2 to balance O and H. 2

30. 30 Common Acid–Base Reactions Reaction of Acids with Hydroxide Bases A net ionic equation contains only the species involved in a reaction. HCl(aq) + NaOH(aq) H—OH(l) + NaCl(aq) Written as individual ions: H + (aq) + Cl − (aq) + Na + (aq) + OH − (aq) H—OH(l) + Na + (aq) + Cl − (aq) Omit the spectator ions, Na + and Cl –. H + (aq) + − OH(aq)H—OH(l) What remains is the net ionic equation:

31. 31 Common Acid–Base Reactions Reaction of Acids with Bicarbonate Bases A bicarbonate base, HCO 3 −, reacts with one H + to form carbonic acid, H 2 CO 3. H + (aq) + HCO 3 − (aq) Carbonic acid then decomposes into H 2 O and CO 2. H 2 O(l) + CO 2 (g) H 2 CO 3 (aq) HCl(aq) + NaHCO 3 (aq) H 2 O(l) + CO 2 (g) NaCl(aq) + H 2 CO 3 (aq) For example:

32. 32 Common Acid–Base Reactions Reaction of Acids with Bicarbonate Bases A carbonate base, CO 3 2–, reacts with two H + to form carbonic acid, H 2 CO 3. 2 H + (aq) + CO 3 2– (aq) H 2 O(l) + CO 2 (g) H 2 CO 3 (aq) 2 HCl(aq) + CaCO 3 (aq) H 2 O(l) + CO 2 (g) 2 CaCl 2 (aq) + H 2 CO 3 (aq) For example:

33. 33 Buffers A buffer is a solution whose pH changes very little when acid or base is added. Most buffers are solutions composed of roughly equal amounts of a weak acid. the salt of its conjugate base. The buffer resists change in pH because added base, − OH, reacts with the weak acid. added acid, H 3 O +, reacts with the conjugate base.

34. 34 Buffers General Characteristics of a Buffer CH 3 COOH(aq) + H 2 O(l)H 3 O + (aq) + CH 3 COO − (aq) weak acid conjugate base If an acid is added to the following buffer equilibrium, Adding more product… …drives the reaction to the left. then the excess acid reacts with the conjugate base, so the overall pH does not change much.

35. 35 Buffers General Characteristics of a Buffer CH 3 COOH(aq) + − OH(aq) H 2 O(l) + CH 3 COO − (aq) weak acid conjugate base If a base is added to the following buffer equilibrium, Adding more reactant… …drives the reaction to the right. then the excess base reacts with the weak acid, so the overall pH does not change much.

36. 36 Focus on the Human Body Buffers in the Blood Normal blood pH is between 7.35 and 7.45. The principle buffer in the blood is carbonic acid/ bicarbonate (H 2 CO 3 /HCO 3 − ). CO 2 (g) + H 2 O(l) H 2 CO 3 (aq) H2OH2O H 3 O + (aq) + HCO 3 − (aq) CO 2 is constantly produced by metabolic processes in the body. The amount of CO 2 is related to the pH of the blood.

37. 37 Focus on the Human Body Buffers in the Blood Respiratory acidosis results when the body fails to eliminate enough CO 2, due to lung disease or failure. CO 2 (g) + 2 H 2 O(g) H 3 O + (aq) + HCO 3 − (aq) A lower respiratory rate increases [CO 2 ]. This drives the reaction to the right, increasing [H 3 O + ]. Blood then has a higher [H 3 O + ] and a lower pH.

38. 38 Focus on the Human Body Buffers in the Blood Respiratory alkalosis is caused by hyperventilating; very little CO 2 is produced by the body. CO 2 (g) + 2 H 2 O(g) H 3 O + (aq) + HCO 3 − (aq) A higher respiratory rate decreases [CO 2 ]. This drives the reaction to the left, decreasing [H 3 O + ]. Blood then has a lower [H 3 O + ] and a higher pH.