Chapter 20 Acids and Bases

Items from Chapter n Reversible Reactions - p. 539 –In a reversible reaction, the reactions occur simultaneously in both directions –double arrows used to indicate this 2SO 2(g) + O 2(g)  2SO 3(g) –In principle, almost all reactions are reversible to some extent

Items from Chapter n Le Chatelier’s Principle - p.541 –If a stress is applied to a system in dynamic equilibrium, the system changes to relieve the stress. –Stresses that upset the equilibrium in a chemical system include: changes in concentration, changes in temperature, and changes in pressure

Items from Chapter n Equilibrium Constants (K eq ) - p. 545 –Chemists generally express the position of equilibrium in terms of numerical values –These values relate to the amounts of reactants and products at equilibrium

Items from Chapter n Equilibrium Constants - p. 545 –consider this reaction: aA + bB  cC + dD –The equilibrium constant (K eq ) is the ratio of product concentration to the reactant concentration at equilibrium, with each concentration raised to a power (= the coefficient)

Items from Chapter n Equilibrium Constants - p. 545 –consider this reaction: aA + bB  cC + dD –Thus, the “equilibrium constant expression” has the general form: [C] c x [D] d [A] a x [B] b ( [ ] = molarity ) K eq =

Items from Chapter n Equilibrium Constants - p. 545 –the equilibrium constants provide valuable information, such as whether products or reactants are favored: K eq > 1, products favored at equilibrium K eq < 1, reactants favored at equilibrium n Sample Problem 19-2, p. 545

Section 20.1 Describing Acids and Bases n OBJECTIVES: –List the properties of acids and bases.

Section 20.1 Describing Acids and Bases n OBJECTIVES: –Name an acid or base, when given the formula.

Properties of acids n Taste sour (don’t try this at home). n Conduct electricity. –Some are strong, others are weak electrolytes. n React with metals to form hydrogen gas. n Change indicators (blue litmus to red). n React with hydroxides to form water and a salt.

Properties of bases n React with acids to form water and a salt. n Taste bitter. n Feel slippery (don’t try this either). n Can be strong or weak electrolytes. n Change indicators (red litmus turns blue).

Names and Formulas of Acids n An acid is a chemical that produces hydrogen ions (H 1+ ) when dissolved in water n Thus, general formula = HX, where X is a monatomic or polyatomic anion n HCl (g) named hydrogen chloride n HCl (aq) is named as an acid n Name focuses on the anion present

Names and Formulas of Acids 1. When anion ends with -ide, the acid starts with hydro-, and the stem of the anion has the suffix -ic followed by the word acid 2. When anion ends with -ite, the anion has the suffix -ous, then acid 3. When anion ends with -ate, the anion suffix is -ic and then acid n Table 20.1, page 578 for examples

Names and Formulas of Bases n A base produces hydroxide ions (OH 1- ) when dissolved in water. n Named the same way as any other ionic compound –name the cation, followed by anion n To write the formula: write symbols; write charges; then cross (if needed) n Sample Problem 20-1, p. 579

Section 20.2 Hydrogen Ions and Acidity n OBJECTIVES: –Given the hydrogen-ion or hydroxide-ion concentration, classify a solution as neutral, acidic, or basic.

Section 20.2 Hydrogen Ions and Acidity n OBJECTIVES: –Convert hydrogen-ion concentrations into values of pH, and hydroxide-ion concentrations into values of pOH.

Hydrogen Ions from Water n Water ionizes, or falls apart into ions: H 2 O  H 1+ + OH 1- n Called the “self ionization” of water n Occurs to a very small extent: [H 1+ ] = [OH 1- ] = 1 x M n Since they are equal, a neutral solution results from water n K w = [H 1+ ] x [OH 1- ] = 1 x M 2 n K w is called the “ion product constant”

Ion Product Constant H 2 O H + + OH - H 2 O H + + OH - n K w is constant in every aqueous solution: [H + ] x [OH - ] = 1 x M 2 n If [H + ] > then [OH - ] then [OH - ] < n If [H + ] n If we know one, other can be determined n If [H + ] > 10 -7, it is acidic and [OH - ] 10 -7, it is acidic and [OH - ] < n If [H + ] n Basic solutions also called “alkaline” n Sample problem 20-2, p. 582

Logarithms and the pH concept n Logarithms are powers of ten. –Review from earlier lessons, and p. 585 n definition: pH = -log[H + ] n in neutral pH = -log(1 x ) = 7 n in acidic solution [H + ] > n pH < -log(10 -7 ) n pH < 7 (from 0 to 7 is the acid range) n in base, pH > 7 (7 to 14 is base range)

pH and pOH n pOH = -log [OH - ] n [H + ] x [OH - ] = 1 x M 2 n pH + pOH = 14 n Thus, a solution with a pOH less than 7 is basic; with a pOH greater than 7 is an acid

Basic Basic AcidicNeutral [OH - ] pH [H + ] pOH

Examples: n Sample 20-3, p.586 n Sample 20-4, p.586 n Sample 20-5, p.587 n Sample 20-6, p.588

Measuring pH n Why measure pH? –Everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc. n Sometimes we can use indicators, other times we might need a pH meter

Acid-Base Indicators n An indicator is an acid or base that undergoes dissociation in a known pH range, and has different colors in solution (more later in chapter) n Examples: litmus, phenolphthalein, bromthymol blue: Fig 20.8, p.590

Acid-Base Indicators n Although useful, there are limitations to indicators: –usually given for a certain temperature (25 o C), thus may change at different temperatures –what if the solution already has color? –ability of human eye to distinguish colors

Acid-Base Indicators n A pH meter may give more definitive results –some are large, others portable –works by measuring the voltage between two electrodes –needs to be calibrated –Fig , p.591

Section 20.3 Acid-Base Theories n OBJECTIVES: –Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis

Section 20.3 Acid-Base Theories n OBJECTIVES: –Identify conjugate acid-base pairs in acid-base reactions.

Svante Arrhenius n Swedish chemist ( ) - Nobel prize winner in chemistry (1903) n one of the first chemists to explain the chemical theory of the behavior of acids and bases n Dr. Hubert Alyea-last graduate student of Arrhenius. (link below)

Hubert N. Alyea ( )

1. Arrhenius Definition n Acids produce hydrogen ions (H 1+ ) in aqueous solution. n Bases produce hydroxide ions (OH 1- ) when dissolved in water. n Limited to aqueous solutions. n Only one kind of base (hydroxides) n NH 3 (ammonia) could not be an Arrhenius base.

Svante Arrhenius ( )

Polyprotic Acids n Some compounds have more than 1 ionizable hydrogen. n HNO 3 nitric acid - monoprotic n H 2 SO 4 sulfuric acid - diprotic - 2 H + n H 3 PO 4 phosphoric acid - triprotic - 3 H + n Having more than one ionizable hydrogen does not mean stronger!

Polyprotic Acids n However, not all compounds that have hydrogen are acids n Also, not all the hydrogen in an acid may be released as ions –only those that have very polar bonds are ionizable - this is when the hydrogen is joined to a very electronegative element

Arrhenius examples... n Consider HCl n What about CH 4 (methane)? n CH 3 COOH (ethanoic acid, or acetic acid) - it has 4 hydrogens like methane does…? n Table 20.4, p. 595 for bases

2. Brønsted-Lowry Definitions n Broader definition than Arrhenius n Acid is hydrogen-ion donor (H + or proton); base is hydrogen-ion acceptor. n Acids and bases always come in pairs. n HCl is an acid. –When it dissolves in water, it gives it’s proton to water. n HCl(g) + H 2 O(l) H 3 O + + Cl - n Water is a base; makes hydronium ion.

Johannes Bronsted / Thomas Lowry ( ) ( )

Acids and bases come in pairs... n A conjugate base is the remainder of the original acid, after it donates it’s hydrogen ion n A conjugate acid is the particle formed when the original base gains a hydrogen ion n Indicators are weak acids or bases that have a different color from their original acid and base

Acids and bases come in pairs... n General equation is: n HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) n Acid + Base Conjugate acid + Conjugate base n NH 3 + H 2 O NH OH 1- base acid c.a. c.b. base acid c.a. c.b. n HCl + H 2 O H 3 O 1+ + Cl 1- n acid base c.a. c.b. n Amphoteric - acts as acid or base

3. Lewis Acids and Bases n Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction n Lewis Acid - electron pair acceptor n Lewis Base - electron pair donor n Most general of all 3 definitions; acids don’t even need hydrogen! n Sample Problem 20-7, p.599

Gilbert Lewis ( )

Section 20.4 Strengths of Acids and Bases n OBJECTIVES: –Define strong acids and weak acids.

Section 20.4 Strengths of Acids and Bases n OBJECTIVES: –Calculate an acid dissociation constant (K a ) from concentration and pH measurements.

Section 20.4 Strengths of Acids and Bases n OBJECTIVES: –Arrange acids by strength according to their acid dissociation constants (K a ).

Section 20.4 Strengths of Acids and Bases n OBJECTIVES: –Arrange bases by strength according to their base dissociation constants (K b ).

Strength n Strong acids and bases are strong electrolytes –They fall apart (ionize) completely. –Weak acids don’t completely ionize. n Strength different from concentration n Strong-forms many ions when dissolved n Mg(OH) 2 is a strong base- it falls completely apart when dissolved. –But, not much dissolves- not concentrated

Measuring strength n Ionization is reversible. n HAH + + A - n This makes an equilibrium n Acid dissociation constant = K a n K a = [H + ][A - ] (water is constant) [HA] n Stronger acid = more products (ions), thus a larger K a (Table 20.8, p.602)

What about bases? n Strong bases dissociate completely. n B + H 2 O BH + + OH - n Base dissociation constant = K b n K b = [BH + ][OH - ] [B] (we ignore the water) n Stronger base = more dissociated, thus a larger K b.

Strength vs. Concentration n The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the number of moles of acid or base in a given volume n The words strong and weak refer to the extent of ionization of an acid or base n Is concentrated weak acid possible?