AP CHAPTER 15 OUTLINE
ACIDS AND BASES
I. Bronsted Acids and Bases
     A.
Arrhenius Definition -- an acid and a base combine to form water
     B. Bronsted Concept
          1.
Bronsted Acid -- a proton donor
          2.
Bronsted Base -- a proton acceptor
     C. Conjugate Acids and Bases
          1.
Conjugate Acid-Bae pair -- two substances that only differ by one proton
          2.
Conjugate acid -- proton donor
          3.
Conjugate base -- proton acceptor
          4. In all Bronsted acid-base equilibrium, there is always two conjugate acid-base pairs
     D.
Amphoteric substances
          1. a substance that can be either acid or base depending on the other substance
II. Strengths of Bronsted Acids and Bases and Periodic Trends
     A. Comparing the Acid-Base Strengths of Conjugate Pairs
          1.
The position of an acid-base equilibrium favors the weaker acid and base
          2. Stronger acids and bases tend to react with each other to produce their weaker conjugates
     B. Reciprocal Relationships within Conjugate Acid-Base pairs
          1.
The stronger a Bronsted Acid is, the weaker is its conjugate base
          2.
The weaker a Bronsted acid is, the stronger is its conjugate base
     C. Periodic Trends in the Strengths of Binary Acids
          1.
The sstrengths of the binary acids increase from left to right within the same period
               a. increased electronegativity causes the corresponding H-X bonds to become more polar
          2.
The strengths of binary acids increases from top to bottom within the same group
               a. increased atomic size decreases the H-X bond strengths
     D.
Oxoacids
          1. Acids with Hydrogen, Oxygen, and some other element
          2. How easy a hydrogen is released is determined by how the group of atoms attached to the oxygen affect the polarity of the O-H bond
     E. The Effect of the Electronegativity
          1.
When the central atoms hold the same number of oxygen atoms, the acid strength increaes from bottom to top within the group and from left to right within a period
               a. follows electronegativity trends
     F. The Effect of the Number of Oxygens and the acidity of Oxaoacids
          1.
For a given central atom, the acid strength of an oxoacid increases with the number of oxygens held by the central atom
              a. The lone oxygens pull the electron density away from the O-H bond, making them more polar (easily removed)
     G. Oxoacids and the Effect of the Dispersal of Negative Charge to the Lone Oxygens
          1.
Acids with smaller negative charge per each lone oxygen, are less able to attract H+ ions, weaker base
               a. Acids with more lone oxygens will more fully ionize than acids with fewer lone oxygens

III. Lewis Acids and Baes
     A. Definitions
          1.
Lewis Acid -- any ionic or molecular species that can accept a pair of electrons in the formation of a coordinate covalent bond
          2.
Lewis Base -- any ionic or molecular species that can donate a pair of electrons in the formation of a coordinate covalent bond
          3.
Neutralization -- the formation of a coordinate covalent bond betwen the donor (base) and the acceptor (acid)
     B. examples of Lewis Acid-Base Reactions
          1. A Lewis base is a substance that has completed valence shells and unshared pairs of electrons
          2. A Lewis acid is a substance that has incomplete valence shells
IV. Acid-Base Properties of the Elements and Their Oxides
     A.
Most metal oxids react with water to form bases, and nonmetal oxides react with water to form acids
     B. Hydrated Metal Ions as Weak Acids
          1. tend to be proton donors in water
          2. degree to which the metal ions produce acidic solutions
               a.
amount of charge on the cation
                    1. The higher the charge (more able to draw electrons) on the metal ion the more acidic the solution
               b.
size of the cation
                    1. The smaller the cation (more able to pull electrons from O-H bond) the more acidic the solution
          3. Charge Density = ionic charge/ionic volume
               a. the higher the charge density, the more effective the metal ion is at drawing electrons --> more acidic solutions
     C. Periodic Trends in the Acidity of Metal Ions
          1.
Hydrated metal ions at the top of a group in the P.T. are the most acidic within a group
               a. Same charge with smaller cation size
          2.
Hydrated metal ions increae in acidity left to right within a period
               a. Increaed charge with smaller cation size
     D. Nonmetal Oxides as Acids
          1.
Acid anhydrides
          2. React with water to give acid solutions (acid rain, pollution)
V. Ionization of Water and the pH Concept
     A.
Self-Ionization (Autoionization) of Water
          1. Two water molecules react to form H3O+ (H+) and OH-
          2.
Ion product constant of water
               a. Kw = [H3O+][OH-]
               b. Since [H2O] is essentially always 55.6 M, it can be made part of the constant (only ions)
               c. Simplify to
Kw = [H+][OH-]
          3. In pure water, autoionization is equal between H+ and OH-
               a.
[H+] = [OH-] = 1.0 x 10^7 mol/L
               b.
Kw = 1.0 x 10^14
     B. Autoionization of Water when Solutes are Present
          1.
In any aqueous solution, the product of [H+] and [OH-] equals Kw, although these two molar concentrations may not actually equal each other
     C. Criteria for Acidic, Basic, and Neutral Solutions
          1.
Neutral solution --> [H+] = [OH-]
          2.
Acids solution --> [H+] >[OH-]
          3.
Basic solution --> [H+] < [OH-]
     D. the pH Concept
          1. In most solutions of weak acids and baes, the molar concentrations of H= and OH- are very small
          2. pH scale was developed by Sorenson to alleviate this problem with small numbers
               a.
pH = -log[H+]
               b. [H+] = 10^-pH
          3. Logarithm relationship will work for any concentration based number
               a.
pOH = -log[OH-]
               b.
pKw = -logKw
                    1.
pKw = 14.00
          4. Using the equation from "V. A. 2. c." we can find a new relationship
               a.
pH + pOH = pKw = 14.00
     E. Definitions of Acidic, Basic, and Neutral solutions According to pH
          1.
Acidic solution --> pH < 7.00
          2.
Basic solution --> pH > 7.00
          3.
Neutral solution --> pH = 7.00
          4. Solutions of 1.0 M or higher give negative pH values are usually never used
     F. Measuring pH
          1. pH meters
          2. Acid-Base indicators --> change different colors in the presence of acids and bases
          3. Litmus paper --> red in acid and blue in bse
          4. Universal indicator (paper or solution) --> mixture of many indicators
VI. solutions of Strong Acids and Bases
     A. pH of dilute Solutions of Strong Acids and Bases
          1. Assume 100% ionize, so the ion concentrations are the acid/base concentrations times the molar coefficient from the equation
     B. Suppression of the Ionization of Water by Acids and Bases
          1.
Assume that all the H+ in the solution of an acid comes from the solute
          2.
Assume that all the OH- in the solution of a base comes from the solute
          3. Since the amounts of H+ and OH- from the water are equal


Outline based upon:
     Brady, J. E., Holum, J. R., Russell, J. W. (2000)
. Chemistry: The Study of Matter and Its Changes. (3rd ed.). New York: John Wiley & Sons, Inc. pp. 667-694.
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