Section 20.1Describing Acids and Bases • OBJECTIVES: • List the properties of acids and bases.
Section 20.1Describing Acids and Bases • OBJECTIVES: • Name an acid or base, when given the formula.
Properties of acids • Taste sour (don’t try this at home). • Conduct electricity. • Some are strong, others are weak electrolytes. • React with metals to form hydrogen gas. • Change indicators (blue litmus to red). • React with hydroxides to form water and a salt.
Properties of bases • React with acids to form water and a salt. • Taste bitter. • Feel slippery (don’t try this either). • Can be strong or weak electrolytes. • Change indicators (red litmus turns blue).
Names and Formulas of Acids • An acid is a chemical that produces hydrogen ions (H1+) when dissolved in water • Thus, general formula = HA, where A is a monatomic or polyatomic anion • HCl(g) is hydrogen monochloride • HCl(aq) is named as an acid • Name focuses on the anion present
Names and Formulas of Acids 1. BINARY - 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. TERNARY - When anion ends with -ite, the anion has the suffix -ous, then acid 3. TERNARY - When anion ends with -ate, the anion suffix is -ic and then acid
Names and Formulas of Bases • A base produces hydroxide ions (OH1-) when dissolved in water. • Named the same way as any other ionic compound • name the cation, followed by anion • To write the formula: write symbols; write charges; then cross (if needed) • Sample Problem 20-1, p. 579
Section 20.2Hydrogen Ions and Acidity • OBJECTIVES: • Given the hydrogen-ion or hydroxide-ion concentration, classify a solution as neutral, acidic, or basic.
Section 20.2Hydrogen Ions and Acidity • OBJECTIVES: • Convert hydrogen-ion concentrations into values of pH, and hydroxide-ion concentrations into values of pOH.
Hydrogen Ions from Water • Water ionizes, or falls apart into ions: H2O ® H1+ + OH1- • Called the “self ionization” of water • Occurs to a very small extent: [H1+ ] = [OH1-] = 1 x 10-7 M • Since they are equal, a neutral solution results from water • Kw = [H1+ ] x [OH1-] = 1 x 10-14 • Kw is called the “ion product constant”
Ion Product Constant • H2O H+ + OH- • Kw is constant in every aqueous solution: [H+] x [OH-] = 1 x 10-14 • If [H+] > 10-7 then [OH-] < 10-7 • If [H+] < 10-7 then [OH-] > 10-7 • If we know one, other can be determined • If [H+] > 10-7, it is acidic and [OH-] < 10-7 • If [H+] < 10-7, it is basic and [OH-] > 10-7 • Basic solutions also called “alkaline” • Sample problem 20-2, p. 582
Logarithms and the pH concept • Logarithms are powers of ten. • Review from earlier lessons, and p. 585 • definition: pH = -log[H+] • in neutral pH = -log(1 x 10-7) = 7 • in acidic solution [H+] > 10-7 • pH < -log(10-7) • pH < 7 (from 0 to 7 is the acid range) • in base, pH > 7 (7 to 14 is base range)
pH and pOH • pH = -log[H+] • pOH = -log [OH-] • Kw = [H+] x [OH-] = 1 x 10-14 • pH + pOH = 14 • Thus, a solution with a pH less than 7 is an acid; a pH greater than 7 is a base; 7 is neutral
100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 14 13 11 9 7 5 3 1 0 10-14 10-13 10-11 10-9 Basic 10-7 10-5 10-3 10-1 100 [H+] pH 0 1 3 5 7 9 11 13 14 Acidic Neutral Basic pOH [OH-]
Examples: • Sample 20-3, p.586 • Sample 20-4, p.586 • Sample 20-5, p.587 • Sample 20-6, p.588
Measuring pH • Why measure pH? • Everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc. • Sometimes we can use indicators, other times we might need a pH meter
Acid-Base Indicators • 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) • Examples: litmus, phenolphthalein, bromthymol blue: Fig 20.8, p.590
Acid-Base Indicators • Although useful, there are limitations to indicators: • usually given for a certain temperature (25 oC), thus may change at different temperatures • what if the solution already has color? • ability of human eye to distinguish colors
Acid-Base Indicators • 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. 20.10, p.591
Section 20.3Acid-Base Theories • OBJECTIVES: • Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis
Section 20.3Acid-Base Theories • OBJECTIVES: • Identify conjugate acid-base pairs in acid-base reactions.
Svante Arrhenius • Swedish chemist (1859-1927) - Nobel prize winner in chemistry (1903) • one of the first chemists to explain the chemical theory of the behavior of acids and bases • Dr. Hubert Alyea-last graduate student of Arrhenius. (link below) http://www.woodrow.org/teachers/ci/1992/Arrhenius.html
1. Arrhenius Definition • Acids produce hydrogen ions (H1+) in aqueous solution. • Bases produce hydroxide ions (OH1-) when dissolved in water. • Limited to aqueous solutions. • Only one kind of base (hydroxides) • NH3 (ammonia) could not be an Arrhenius base.
Polyprotic Acids • Some compounds have more than 1 ionizable hydrogen. • HNO3 nitric acid - monoprotic • H2SO4 sulfuric acid - diprotic - 2 H+ • H3PO4 phosphoric acid - triprotic - 3 H+ • Having more than one ionizable hydrogen does not mean stronger!
Polyprotic Acids • However, not all compounds that have hydrogen are acids • 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... • Consider HCl • What about CH4 (methane)? • CH3COOH (ethanoic acid, or acetic acid) - it has 4 hydrogens like methane does…? • Table 20.4, p. 595 for bases
2. Brønsted-Lowry Definitions • Broader definition than Arrhenius • Acid is hydrogen-ion donor (H+ or proton); base is hydrogen-ion acceptor. • Acids and bases always come in pairs. • HCl is an acid. • When it dissolves in water, it gives it’s proton to water. • HCl(g) + H2O(l) H3O+ + Cl- • Water is a base; makes hydronium ion.
Johannes Bronsted / Thomas Lowry (1879-1947) (1874-1936)
Acids and bases come in pairs... • A conjugate base is the remainder of the original acid, after it donates it’s hydrogen ion • A conjugate acid is the particle formed when the original base gains a hydrogen ion • Indicators are weak acids or bases that have a different color from their original acid and base
Acids and bases come in pairs... • General equation is: • HA(aq) + H2O(l) H3O+(aq) + A-(aq) • Acid + Base Conjugate acid + Conjugate base • NH3 + H2O NH41+ + OH1- base acid c.a. c.b. • HCl + H2O H3O1++ Cl1- • acid base c.a. c.b. • Amphoteric - acts as acid or base
3. Lewis Acids and Bases • Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction • Lewis Acid - electron pair acceptor • Lewis Base - electron pair donor • Most general of all 3 definitions; acids don’t even need hydrogen! • Sample Problem 20-7, p.599
Section 20.4Strengths of Acids and Bases • OBJECTIVES: • Define strong acids and weak acids.
Section 20.4Strengths of Acids and Bases • OBJECTIVES: • Calculate an acid dissociation constant (Ka) from concentration and pH measurements.
Section 20.4Strengths of Acids and Bases • OBJECTIVES: • Arrange acids by strength according to their acid dissociation constants (Ka).
Section 20.4Strengths of Acids and Bases • OBJECTIVES: • Arrange bases by strength according to their base dissociation constants (Kb).
Strength • Strong acids and bases are strong electrolytes • They fall apart (ionize) completely. • Weak acids don’t completely ionize. • Strength different from concentration • Strong-forms many ions when dissolved • Mg(OH)2 is a strong base- it falls completely apart when dissolved. • But, not much dissolves- not concentrated
Measuring strength • Ionization is reversible. • HA H+ + A- • This makes an equilibrium • Acid dissociation constant = Ka • Ka = [H+ ][A- ] (water is constant) [HA] • Stronger acid = more products (ions), thus a larger Ka (Table 20.8, p.602)
What about bases? • Strong bases dissociate completely. • B + H2O BH+ + OH- • Base dissociation constant = Kb • Kb = [BH+ ][OH-] [B] (we ignore the water) • Stronger base = more dissociated, thus a larger Kb.
Strength vs. Concentration • 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 • The words strong and weak refer to the extent of ionization of an acid or base • Is concentrated weak acid possible?
Practice • Write the expression for HNO2 • Write the Kb for NH3 • Sample 20-8, p. 604 • Carefully study Key Terms and equations, p. 608 • Be sure to do the ChemASAP programs, and take all the self-tests that are available!