1	Introductory Chemistry: Concepts & Connections 4^th Edition by Charles H. Corwin Acids and Bases
2	The Chemistry of Acids and Bases Chapter 17
3	Fig. 11.15
4	Fig. 11.16
5	Properties of Acids An acid is any substance that releases hydrogen ions, H⁺, into water. Blue litmus paper turns red in the presence of hydrogen ions. Blue litmus is used to test for acids. Acids have a sour taste; lemons, limes, and vinegar are acidic.
6	Acid and Bases
7	Properties of Bases A base is a substance that releases hydroxide ions, OH^–, into water. Red litmus paper turns blue in the presence of hydroxide ions. Red litmus is used to test for bases. Bases have a slippery, soapy feel. Bases also have a bitter taste; milk of magnesia is a base.
8	Acid and Bases
9	Acid and Bases
10	Fig. 11.14
11	Acid/Base Neutralization Recall that an acid and a base react with each other in a neutralization reaction. When an acid and a base react, water and a salt are produced. For example, nitric acid reacts with sodium hydroxide to produce sodium nitrate and water: HNO₃(aq) + NaOH(aq) → NaNO₃(aq) + H₂O(l)
12	The pH Scale A pH value expresses the acidity or basicity of a solution. Most solutions have a pH between 0 and 14. Acidic solutions have a pH less than 7. As a solution becomes more acidic, the pH decreases. Basic solutions have a pH greater than 7. As a solution becomes more basic, the pH increases.
13	Acid/Base Classifications of Solutions A solution can be classified according to its pH:
14	Buffers A buffer is a solution that resists changes in pH when an acid or a base is added. A buffer is a solution of a weak acid and one of its salts: Citric acid and sodium citrate make a buffer solution When acid is added to the buffer, the citrate reacts with the acid to neutralize it. When base is added to the buffer, the citric acid reacts with the base to neutralize it.
15	Arrhenius Acids and Bases Svante Arrhenius proposed the following definitions for acids and bases in 1884: An Arrhenius acid is a substance that ionizes in water to produce hydrogen ions. An Arrhenius base is a substance that ionizes in water to release hydroxide ions. For example, HCl is an Arrhenius acid and NaOH is an Arrhenius base.
16	Strengths of Acids Acids have varying strengths. The strength of an Arrhenius acid is measured by the degree of ionization in solution. Ionization is the process where polar compounds separate into cations and anions in solution. The acid HCl ionizes into H⁺ and Cl^– ions in solution.
17	Strengths of Bases Bases also have varying strengths. The strength of an Arrhenius base is measured by the degree of dissociation in solution. Dissociation is the process where cations and anions in an ionic compound separate in solution. A formula unit of NaOH dissociates into Na⁺ and OH^– ions in solution.
18	Strong & Weak Arrhenius Acids Strong acids ionize extensively to release hydrogen ions into solution. HCl is a strong acid and ionizes nearly 100% Weak acids only ionize slightly in solution. HF is a weak acid and ionizes only about 1%
19	Arrhenius Acids in Solution All Arrhenius acids have a hydrogen atom bonded to the rest of the molecule by a polar bond. This bond is broken when the acid ionizes. Polar water molecules help ionize the acid by pulling the hydrogen atom away: HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl^–(aq) (~100%) HC₂H₃O₂(aq) + H₂O(l) → H₃O⁺(aq) + C₂H₃O₂^–(aq) (~1%) The hydronium ion, H₃O⁺, is formed when the aqueous hydrogen ion attaches to a water molecule.
20	Strong & Weak Arrhenius Bases Strong bases dissociate extensively to release hydroxide ions into solution. NaOH is a strong base and dissociates nearly 100% Weak bases only ionize slightly in solution. NH₄OH is a weak base and only partially dissociates
21	Arrhenius Bases in Solution When we dissolve Arrhenius bases in solution, they dissociate giving a cation and a hydroxide anion. Strong bases dissociate almost fully and weak bases dissociate very little: NaOH(aq) → Na⁺(aq) + OH^–(aq) (~100%) NH₄OH(aq) → NH₄⁺(aq) + OH^–(aq) (~1%)
22	Neutralization Reactions Recall, an acid neutralizes a base to produce a salt and water. HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) The reaction produces the aqueous salt NaCl. If we have an acid with two hydrogens (sulfuric acid, H₂SO₄), we need two hydroxide ions to neutralize it. H₂SO₄(aq) + 2 NaOH(aq) → Na₂SO₄(aq) + 2 H₂O(l)
23	Predicting Neutralization Reactions We can identify the Arrhenius acid and base that react in a neutralization reaction to produce a given salt such as calcium sulfate, CaSO₄. The calcium must be from calcium hydroxide, Ca(OH)₂, and the sulfate must be from sulfuric acid, H₂SO₄. H₂SO₄(aq) + Ca(OH)₂(aq) → CaSO₄(aq) + 2 H₂O(l)
24	Brønsted-Lowry Acids & Bases The Brønsted-Lowry definitions of acids and bases are broader than the Arrhenius definitions. A Brønsted-Lowry acid is a substance that donates a hydrogen ion to any other substance. It is a proton donor. A Brønsted-Lowry base is a substance that accepts a hydrogen ion. It is a proton acceptor.
25	Brønsted-Lowry Acids & Bases Lets look at two acid/base reactions: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) HCl(aq) + NH₃(aq) → NH₄Cl(aq) HCl donates a proton in both reactions and is a Brønsted-Lowry acid. In the first reaction, the NaOH accepts a proton and is the Brønsted-Lowry base. In the second reaction, NH₃ accepts a proton and is the Brønsted-Lowry base.
26	Amphiprotic Compounds A substance that is capable of both donating and accepting a proton is an amphiprotic compound. NaHCO₃ is an example: HCl(aq) + NaHCO₃(aq) → NaCl(aq) + H₂CO₃(aq) NaOH(aq) + NaHCO₃(aq) → Na₂CO₃ (aq) + H₂O(l) NaHCO₃ accepts a proton from HCl in the first reaction and donates a proton to NaOH in the second reaction.
27	Acid-Base Indicators A solution that changes color as the pH changes is an acid-base indicator. Three common indicators are methyl red, bromothymol blue, and phenolphthalein. Each has a different color above and below a certain pH.
28	Acid-Base Indicators Shown below are the three indicators at different pH values
29	Acid-Base Titrations A titration is used to analyze an acid solution using a solution of a base. A measured volume of base is added to the acid solution. When all of the acid has been neutralized, the pH is 7. One extra drop of base solution after the endpoint increases the pH dramatically. When the pH increases above 7, phenolphthalein changes from colorless to pink indicating the endpoint of the titration.
30	Titration Problem Consider the titration of acetic acid with sodium hydroxide. A 10.0 mL sample of acetic acid requires 37.55 mL of 0.223 M NaOH. What is the concentration of the acetic acid? HC₂H₃O₂(aq) + NaOH(aq) → NaC₂H₃O₂(aq) + H₂O(l) We want concentration acetic acid, we have concentration sodium hydroxide. conc NaOH Þ mol NaOH Þ mol HC₂H₃O₂ Þ conc HC₂H₃O₂
31	Titration Problem Continued The molarity of NaOH can be written as the unit factor 0.233 mol NaOH / 1000 mL solution.
32	Another Titration Problem A 10.0 mL sample of 0.555 M H₂SO₄ is titrated with 0.233 M NaOH. What volume of NaOH is required for the titration? We want mL of NaOH, we have 10.0 mL of H₂SO₄. Use 0.555 mol H₂SO₄/1000 mL and 0.233 mol NaOH/1000 mL.
33	Problem Continued H₂SO₄(aq) + 2 NaOH(aq) → Na₂SO₄(aq) + H₂O(l)
34	Acid-Base Standardization A standard solution is a solution where the concentration is precisely known. Acid solutions are standardized by neutralizing a weighed quantity of a solid base. What is the molarity of a hydrochloric acid solution if 25.50 mL are required to neutralize 0.375 g Na₂CO₃? 2 HCl(aq) + Na₂CO₃(aq) → 2 NaCl(aq) + H₂O(l) + CO₂(g)
35	Standardization Continued
36	Ionization of Water Water undergoes an autoionization reaction. Two water molecules react to produce a hydronium ion and a hydroxide ion: H₂O(l) + H₂O(l) → H₃O⁺(aq) + OH^-(aq) or H₂O(l) → H⁺(aq) + OH^-(aq) Only about 1 in 5 million water molecules is present as ions so water is a weak conductor. The concentration of hydrogen ions, [H⁺], in pure water is about 1 × 10^-7 mol/L at 25°C.
37	Autoionization of Water Since [H⁺] is 1 × 10^-7 mol/L at 25°C, the hydroxide ion concentration, [OH^-], must also be 1 × 10^-7 mol/L at 25°C: H₂O(l) → H⁺(aq) + OH^-(aq) At 25°C [H⁺][OH^-] = (1 × 10^-7)(1 × 10^-7) = 1.0 × 10^-14 This is the ionization constant of water, K_w.
38	[H⁺] and [OH^-] Relationship At 25°C, [H⁺][OH^-] = 1.0 × 10^-14. So, if we know the [H⁺], we can calculate [OH^-]. What is the [OH^-] if [H⁺] = 0.1 M? [H⁺][OH^-] = 1.0 × 10^-14 (0.1)[OH^-] = 1.0 × 10^-14 [OH^-] = 1.0 × 10^-13
39	The pH Concept Recall that pH is a measure of the acidity of a solution. A neutral solution has a pH of 7, an acidic solution has a pH less than 7, and a basic solution has a pH greater than 7. The pH scale uses powers of ten to express the hydrogen ion concentration. Mathematically: pH = –log[H⁺] [H⁺] is the molar hydrogen ion concentration
40	Calculating pH What is the pH if the hydrogen ion concentration in a vinegar solution is 0.001 M? pH = –log[H⁺] pH = –log(0.001) pH = –(–3) = 3 The pH of the vinegar is 3, so the vinegar is acidic.
41	Calculating [H⁺] from pH If we rearrange the pH equation for [H⁺], we get: [H⁺] = 10^–pH Milk has a pH of 6. What is the concentration of hydrogen ion in milk? [H⁺] = 10^–pH = 10^–6 = 0.000001 M [H⁺] = 1 × 10^–6 M.
42	Advanced pH Calculations What is the pH of blood with [H⁺] = 4.8 × 10^–8 M? pH = –log[H⁺] = –log(4.8 × 10^–8) = – (–7.32) pH = 7.32 What is the [H⁺] in orange juice with a pH of 2.75? [H⁺] = 10^–pH = 10^–2.75 = 0. 0018 M [H⁺] = 2.75 × 10^–3 M
43	Strong & Weak Electrolytes An aqueous solution that is a good conductor of electricity is a strong electrolyte. An aqueous solution that is a poor conductor of electricity is a weak electrolyte. The greater the degree of ionization or dissociation, the greater the conductivity of the solution.
44	Electrolyte Strength Weak acids and bases and insoluble ionic compounds are weak electrolytes.
45	Total Ionic Equation The concept of ionization allows us to portray ionic solutions more accurately by showing strong electrolytes in their ionized form. HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) Write strong acids and bases and soluble ionic compounds as ions: H⁺(aq) + Cl^-(aq) + Na⁺(aq) + OH^-(aq) → Na⁺(aq) + Cl^-(aq) + H₂O(l) This is the total ionic equation. Each species is written as it predominantly exists in solution.
46	Net Ionic Equation H⁺(aq) + Cl^-(aq) + Na⁺(aq) + OH^-(aq) → Na⁺(aq) + Cl^-(aq) + H₂O(l) Notice that Na⁺ and Cl^- appear on both sides of the equation. They are spectator ions. Spectator ions are in the solution, but do not participate in the overall reaction. We can cancel out the spectator ions to give the net ionic equation. The net ionic equation is: H⁺(aq) + OH^-(aq) → H₂O(l)
47	Writing Net Ionic Equations Complete and balance the non-ionized chemical equation. Convert the non-ionized equation into the total ionic equation Write strong electrolytes in the ionized form Write weak electrolytes, water, and gases in the non-ionized form Cancel all the spectator ions to obtain the net ionic equation. If all species are eliminated, there is no reaction.
48	Conclusions pH is a measure of the acidity of a solution. The typical range for pH is 0 to 14. Neutral solutions have a pH of 7. Below are some properties of acids and bases.
49	Conclusions Continued An Arrhenius acid is a substance that ionizes in water to produce hydrogen ions. An Arrhenius base is a substance that ionizes in water to release hydroxide ions. A Brønsted-Lowry acid is a substance that donates a hydrogen ion to any other substance. It is a proton donor. A Brønsted-Lowry base is a substance that accepts a hydrogen ion. It is a proton acceptor
50	Conclusions Continued In an aqueous solution, [H⁺][OH^-] = 1.0 × 10^-14. This is the ionization constant of water, K_w. pH = –log[H⁺] [H⁺] = 10^–pH Strong electrolytes are mostly dissociated in solution. Weak electrolytes are slightly dissociated in solution.