HONORS CHEMISTRY - CHAPTER 19 ACIDS, BASES, AND SALTS OBJECTIVES AND NOTES - V14 NAME: DATE: PAGE: THE BIG IDEA: REACTIONS Essential Questions 1. What are the different ways chemists define acids and bases? 2. What does the ph of a solution mean? 3. How do chemists use acid-base reactions? Chapter Objectives 1. Review nomenclature rules for acids and bases and the formation of acids and bases from anhydrides. (19.1) 2. Identify the common physical and chemical properties of acids and bases. (19.1) 3. Describe and identify the hydronium ion. (19.1) 4. Define, identify, and give examples of Arrhenius acids and bases. (19.1) 5. Define, identify, and give examples of Brønsted-Lowry acids and bases. (19.1) 6. Use the Brønsted-Lowry theory to classify substances as acids or bases, or as hydrogen ion donors or acceptors. (19.1) 7. Define, identify, and give examples of Lewis acids and bases. (19.1) 8. Describe and write the equation for the self-ionization of water. (19.2) 9. Given the hydrogen ion concentration, hydroxide ion concentration, ph, or poh of a solution, classify it as neutral, acidic, or basic. (19.2) 10. Describe the ph scale. (19.2) 11. Given the molarity, normality, hydrogen ion concentration, hydroxide ion concentration, ph, or poh of a solution calculate the other values. (19.2) 1
12. Given the hydrogen ion concentration, hydroxide ion concentration, ph, or poh of a solution at equilibrium, calculate the other values. (19.2) 13. Define acid-base indicators and ph meters. (19.2) 14. Describe the colors of phenolphthalein and litmus in acidic and basic solutions. (19.2) 15. Define and identify all strong acids and bases. (19.3) 16. Distinguish between strong and weak acids and bases. (19.3) 17. Use K a values to distinguish between weak acids. (19.3) 18. Use K b values to distinguish between weak bases. (19.3) 19. Demonstrate and be able to describe all aspects of laboratory safety rules and procedures. (Applicable every chapter) 19.1 Acid-Base Theories A. Arrhenius Acids and Bases (The "Classical" Theory) 1. hydronium ion: When the hydrogen ion [H + (aq)] comes in contact with water it immediately forms the hydronium ion [H 3 O + (aq)]. 2. Arrhenius acids: Compounds that contain hydrogen that ionize when placed in water to produce H + (aq) or H 3 O + (aq). a. Examples: HCl, H 2 SO 4. 3. Arrhenius bases: Compounds that when placed in water produce hydroxide ions, OH - (aq). Some authors use a more narrow definition that states bases are metallic hydroxides that dissociate when placed in water to produce hydroxide ions, OH - (aq). a. Examples: NaOH, KOH; most would include NH 3 in this list. 2 - HC - Chapter 19 - Objectives and Notes - V14
B. Brønsted-Lowry Acids and Bases 1. Brønsted-Lowry acid: A substance that can act as a proton donor. a. Examples: HC 2 H 3 O 2, H 2 O. 2. Brønsted-Lowry base: A substance that can act as a proton acceptor. a. Examples: NH 3, H 2 O. 3. This theory is not dependent on reactions occurring in a solution, they can occur anywhere, even in the air! 4. conjugate acid: It is a substance formed by the addition of a proton to a Brønsted base. 5. conjugate base: It is a substance formed by a loss of a proton from a Brønsted acid. 6. conjugate acid-base pair: An acid and base that differ by only the presence or absence of a proton. They are found on opposite sides of an equation at equilibrium that is demonstrating the Brønsted-Lowry theory of acids and bases. a. The stronger an acid, the weaker its conjugate base. The stronger a base, the weaker its conjugate acid. b. Examples: 1. HC 2 H 3 O 2 (aq) + H 2 O (l) H 3 O + (aq) + C 2 H 3 O - 2 (aq) acid + base conjugate acid + conjugate base 2. NH 3 (g) + H 2 O (l) NH 4 + (aq) + OH - (aq) base + acid conjugate acid + conjugate base 3. HCl (g) + NH 3 (g) NH 4 Cl (s) acid + base conjugate acid/base a. The reaction produces a white smoke, which often is found as a film on glassware in chem rooms. 7. amphoteric/amphiprotic: A substance that can act as either an acid or a base. a. Examples: H 2 O (see above); a weak acid can act as a base when it is combined with a strong acid; a weak base can act as an acid when it is combined with a strong base. 3 - HC - Chapter 19 - Objectives and Notes - V14
C. Lewis Acids and Bases 1. Lewis acid: A substance that can act as an electron pair acceptor. Example: BF 3. 2. Lewis base: A substance that can act as an electron pair donor. Example: NH 3. 3. As the Brønsted-Lowry theory expanded on the Arrhenius theory of acids, the Lewis theory expands on the Brønsted-Lowry theory. a. If a substance is an acid or base under the Arrhenius theory, it functions as an acid or base under the Brønsted-Lowry theory. However, the reverse is not necessarily true. b. If a substance is an acid or base under the Brønsted-Lowry theory, it functions as an acid or base under the Lewis theory. However, the reverse is not necessary true. c. In mathematical terms, Lewis acids and bases are the largest set. Brønsted-Lowry acids and bases are a subset of the Lewis acid-base set. Arrhenius acids and bases are a subset of the Brønsted-Lowry acid base set. 4. The Lewis theory is therefore not limited to only H + 's, OH - 's, and aqueous solutions, it can be applied to any compounds that are donating/accepting electron pairs. 5. In a reaction involving Brønsted-Lowry acids and bases a proton moves from the Brønsted-Lowry acid to the Brønsted-Lowry base. In a reaction involving Lewis acids and bases an electron pair does not move from the base to the acid. Rather the Lewis acid and base form a coordinate covalent bond and the resulting compound is called an adduct. a. adduct: The product of a Lewis acid and base reaction; it is a single substance that contains the new coordinate covalent bond. 1. It is the Lewis base that provides the electron pair for the coordinate covalent bond that is formed when producing the adduct. 4 - HC - Chapter 19 - Objectives and Notes - V14
5. Examples of Lewis acids and bases: a. molecule with incomplete + molecule with unshared adduct octet (acid) pair electrons (base) BF 3 + :NH 3 H 3 N:BF 3 b. small cation + molecule with unshared adduct (acid) pair electrons (base) Fe 3+ + 6 :C N: - [Fe(C N:) 6 ] 3- c. small cation (acid) + water (base) adduct + hydrogen ions Al 3+ + 6 :OH 2 Al(OH) 2 (H 2 O) 5 - (aq) + H + (aq) 19.2 Hydrogen Ions and Acidity A. Hydrogen Ions from Water 1. autoionization/self-ionization: A process in which a substance acts as both a proton donor and acceptor with itself at the same time. a. Example: H 2 O (l) + H 2 O (l) H 3 O + (aq) + OH - (aq) base + acid conjugate acid + conjugate base 1. The above equation can be simplified to yield the following: H 2 O (l) H + (aq) + OH - (aq) a. H + (aq) is really just a proton that is attracted to anything negative, like all of the negative ends of the water molecules around them. Thus, while writing H + (aq) in an equation is simple, the ion that exists in water is more like H 3 O + (aq) (If you want the truth, ask your chemistry instructor.) 5 - HC - Chapter 19 - Objectives and Notes - V14
2. The equilibrium constant for the ionization of water is: a. at 25 o C: K w = [H + ][OH - ] or K w = [H 3 O + ][OH - ] K w = [H + ][OH - ] = 1.00 x 10-14 or K w = [H 3 O + ][OH - ] = 1.00 x 10-14 B. & C. The ph Concept and Measuring ph 1. p is a convention that is used to signify the negative log of very small numbers. 2. ph: The power of the hydronium ion; it is defined as the negative logarithm (common logarithm) of the molar concentration of the hydronium ions. The ph scale ranges from 0 (most acidic) through 7 (neutral solution) to 14 (most basic). a. This system was created by S. P. L. Sorensen, a Danish biochemist, while he was working on the brewing of beer. b. The general forms for the ph equation is: ph = - log[h + ] or ph = - log[h 3 O + ] 6 - HC - Chapter 19 - Objectives and Notes - V14
3. poh: The power of the hydroxide ion; it is defined as the negative logarithm (common logarithm) of the molar concentration of the hydroxide ions. The poh scale ranges from 0 (most basic) through 7 (neutral solution) to 14 (most acidic). It is not nearly as commonly used as the ph scale. a. The general form for the poh equation is: poh = - log[oh - ] 4. Remember!: A change of one on the ph or poh scale is equal to a tenfold change in [H + ] or [OH - ] concentration. 5. The negative log of the K w equation produces a relationship between ph and poh. K w = [H + ][OH - ] = 1.00 x 10-14 - log K w = (-log[h + ])(-log[oh - ]) = - log(1.00 x 10-14 ) pk w = ph + poh = 14.00 a. The table below demonstrates the relationship among types of solutions, [H + ] and [OH - ] concentrations, and ph and poh. Types of Solutions [H + ] [OH - ] ph poh Taste Feel Acidic > 1.00 x 10-7 < 1.00 x 10-7 < 7.00 > 7.00 Sour - Basic < 1.00 x 10-7 > 1.00 x 10-7 > 7.00 < 7.00 Bitter Slippery Neutral 1.00 x 10-7 1.00 x 10-7 7.00 7.00 - - 7 - HC - Chapter 19 - Objectives and Notes - V14
8. indicators: Weak organic acids and bases whose colors differ from the colors of their conjugate bases and acids. a. They are not effective if added to a solution that is already colored. b. litmus: A very common indicator; litmus is blue in basic solutions and red in acidic solutions (Remember Don Showalter stating that fact). c. phenolphthalein: A very common indicator; phenolphthalein is red (magenta) in basic solutions and colorless in acidic solutions. d. universal indicators: Indicators that change various colors depending on the ph of the solution. 9. ph meter: A device that measures the concentration of [H + ] in a solution by determining the voltage between two electrodes. 19.3 Strengths of Acids and Bases A. Strong and Weak Acids and Bases 1. strong acid: An acid that ionizes completely in water. a. Since the reaction goes to completion, the yield sign is usually just a one way arrow: b. Examples: HCl, HBr, HI, HNO 3, HClO 4, H 2 SO 4. 1. In H 2 SO 4 only the first proton ionizes completely. 2. weak acid: An acid that does not ionize completely in water. a. Since the reaction reaches an equilibrium, lying somewhere to the left, the yield sign is a two-way arrow: b. Examples: HC 2 H 3 O 2, HF. 1. In H 2 SO 4 only the second proton does not ionize completely. 8 - HC - Chapter 19 - Objectives and Notes - V14
3. acid-dissociation constant/acid-ionization constant (K a ): The equilibrium constant that is produced when weak acids ionize. Since most acids are in reality, weak acids; these acids have small K a 's. The derivation of K a is seen below: weak acid (aq) + H 2 O (l) H 3 O + (aq) + conjugate base (aq) HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) K c = [H 3 O+ ][A - ] [HA][H 2 O] [H 2 O]K c = [H 3 O+ ][A - ] [HA][H 2 O] [H 2 O] a. The larger the K a, the stronger the weak acid. K a = [H 3 O+ ][A - ] [HA] 4. Acids react with metals more active than hydrogen (see the activity series) to produce hydrogen gas. 5. strong base: A metallic hydroxide that dissociates completely in water. The most common are the hydroxides of IA and IIA columns (except H and Be). a. Since the reaction goes to completion, the yield sign is usually just a one way arrow: b. Examples containing OH - : LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH) 2, Sr(OH) 2, and Ba(OH) 2. 6. weak base: A substance that does not produce many OH - (aq). a. Since the reaction reaches an equilibrium, lying somewhere to the left, the yield sign is a two-way arrow: b. Examples: NH 3 (aq) and CH 3 COO - (aq). c. While strong bases contain hydroxides and dissociate in aqueous solutions to produce OH - (aq), weak bases remove H + from the water (acting as Brønsted proton acceptors) to produce OH - (aq). 9 - HC - Chapter 19 - Objectives and Notes - V14
7. base-dissociation constant (K b ): The equilibrium constant that is produced when weak bases cause water to undergo hydrolysis; the value for K b 's are small. 1. The larger the K b, the stronger the weak base. 10 - HC - Chapter 19 - Objectives and Notes - V14