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Chapter 20 Acids and Bases. Section 20.1 Describing Acids and Bases. OBJECTIVES: List the properties of acids and bases. Section 20.1 Describing Acids and Bases. OBJECTIVES: Name an acid or base, when given the formula. Properties of acids. Taste sour (don’t try this at home).

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Chapter 20 Acids and Bases

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Chapter 20 acids and bases l.jpg

Chapter 20Acids and Bases


Section 20 1 describing acids and bases l.jpg

Section 20.1Describing Acids and Bases

  • OBJECTIVES:

    • List the properties of acids and bases.


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Section 20.1Describing Acids and Bases

  • OBJECTIVES:

    • Name an acid or base, when given the formula.


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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.


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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).


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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


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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


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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


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Section 20.2Hydrogen Ions and Acidity

  • OBJECTIVES:

    • Given the hydrogen-ion or hydroxide-ion concentration, classify a solution as neutral, acidic, or basic.


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Section 20.2Hydrogen Ions and Acidity

  • OBJECTIVES:

    • Convert hydrogen-ion concentrations into values of pH, and hydroxide-ion concentrations into values of pOH.


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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”


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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


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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)


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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


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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-]


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Examples:

  • Sample 20-3, p.586

  • Sample 20-4, p.586

  • Sample 20-5, p.587

  • Sample 20-6, p.588


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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


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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


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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


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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


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Section 20.3Acid-Base Theories

  • OBJECTIVES:

    • Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis


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Section 20.3Acid-Base Theories

  • OBJECTIVES:

    • Identify conjugate acid-base pairs in acid-base reactions.


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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


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Svante Arrhenius (1859-1927)


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Hubert N. Alyea (1903-1996)


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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.


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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!


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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


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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


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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.


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Johannes Bronsted / Thomas Lowry (1879-1947) (1874-1936)


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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


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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


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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


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Gilbert Lewis (1875-1946)


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Section 20.4Strengths of Acids and Bases

  • OBJECTIVES:

    • Define strong acids and weak acids.


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Section 20.4Strengths of Acids and Bases

  • OBJECTIVES:

    • Calculate an acid dissociation constant (Ka) from concentration and pH measurements.


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Section 20.4Strengths of Acids and Bases

  • OBJECTIVES:

    • Arrange acids by strength according to their acid dissociation constants (Ka).


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Section 20.4Strengths of Acids and Bases

  • OBJECTIVES:

    • Arrange bases by strength according to their base dissociation constants (Kb).


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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


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Measuring strength

  • Ionization is reversible.

  • HAH+ + 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)


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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.


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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?


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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!


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