Download Acid-Base Equilibria: Understanding the Nature of Hydrated Protons and Acid-Base Reactions and more Lecture notes Chemistry in PDF only on Docsity!
Learning goals and key skills: Understand the nature of the hydrated proton, represented as either H+(aq) or H 3 O+(aq) Define and identify Arrhenuis acids and bases. Define and identify Bronsted-Lowry acids and bases, and identify conjugate acid-base pairs. Relate the strength of an acid to the strength of its conjugate base. Understand how the equilibrium position of a proton transfer reaction relates the strengths of acids and bases involved. Describe the autoionization of water and understand how [H 3 O+] and [OH-] are related Calculate the pH of a solution given [H 3 O+] or [OH-] Calculate the pH of a strong acid or strong base given its concentration Calculate Ka or Kb for a weak acid or weak base given its concentration and the pH of the solution Calculate pH of a weak acid or weak base or its percent ionization given its concentration and Ka or Kb. Calculate Kb for a weak base given Ka of its conjugate acid, and similarly calculate Ka from Kb. Predict whether and aqueous solution of a salt will be acidic, basic, or neutral Predict the relative strength of a series of acids from their molecular structures Define and identify Lewis acids and bases.
Chapter 16
Acid-Base Equilibria
Acids and Bases
Arrhenius -An acid is a substance that, when dissolved in water, increases the concentration of hydrogen ions. -A base is a substance that, when dissolved in water, increases the concentration of hydroxide ions. Brønsted-Lowry -An acid is a proton donor. -A base is a proton acceptor. Acids and bases may be inorganic (7 strong acids, 8 strong bases) or organic (acids have –COOH group):
- HCl, HNO 3 , H 2 SO 4 , HBr, HI, HClO 3 , HClO 4
- AOH (A = Li, Na, K, Rb, Cs); A(OH) 2 (A = Ca, Sr, Ba)
What happens when an acid
dissolves in water?
- Water acts as a Brønsted-Lowry base and abstracts a proton (H+) from the acid.
- As a result, the conjugate base of the acid and a hydronium ion are formed.
Acid and base strength
Strong acids are completely dissociated in water. Their conjugate bases are weak. Weak acids only dissociate partially in water. Their conjugate bases are strong. In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base.
A Strong acid
HCl(aq) + H 2 O(l) → H 3 O+(aq) + Cl(aq)
A Weak acid
HF(aq) + H 2 O(l) ⇌ H 3 O+(aq) + F–(aq)
Autoionization of water
Kw = [H 3 O+][OH−]
- This equilibrium constant is referred to as
the ion-product constant for water, Kw.
Water is amphoteric. In pure water, some
molecules act as bases and some as acids.
This is referred to as autoionization.
pH and pOH scale
pH = -log [H 3 O
]
pOH = -log [OH
]
- At 25 °C in pure water, Kw = [H 3 O+] [OH−] = 1.0 10 − [H 3 O+] = 1.0 10 -14^ = 1.0 10 -7^ M Since in pure water [H 3 O+] = [OH-], pH + pOH = pKw = 14. Neutral pH is 7.00. Acidic pH is below 7.00. Basic pH is above 7.00.
Three ways to measure pH
red-to-blue:
basic, pH > 8
blue-to-red:
acidic, pH < 5
Strong acids completely ionize.
HA + H 2 O → H 3 O+^ + A
For the monoprotic strong acids,
[H 3 O+] = [acid]
- HCl, HNO 3 , H 2 SO 4 , HBr, HI, HClO 3 , HClO 4
Strong bases completely ionize.
MOH(aq) → M+(aq) + OH
(aq) or
M(OH) 2 (aq) → M2+(aq) + 2 OH
(aq)
- AOH (A = Li, Na, K, Rb, Cs); A(OH) 2 (A = Ca, Sr, Ba)
Calculating Ka from the pH
The pH of a 0.100 M solution of formic
acid, HCOOH, at 25C is 2.38. Calculate
Ka for formic acid at this temperature.
We know that
[H 3 O+] [HCOO-]
[HCOOH]
Ka =
To calculate Ka, we need the equilibrium
concentrations of all three things.
We can find [H 3 O+], which is the same as
[HCOO-], from the pH.
HCOOH (aq) + H 2 O (l) ⇌ H 3 O+^ (aq) + HCOO-^ (aq)
Calculating Ka from pH
Now we can set up an ICE table…
[HCOOH], M [H 3 O+], M [HCOO-], M Initially 0.100 0 0 Change - 4.17 10 -3^ + 4.17 10 -3^ + 4.17 10 - Equilibrium 0.10 - 4.17 10 - = 0. 4.17 10 -3^ 4.17 10 -
Calculating Percent Ionization
concentration ionized original concentration
In this example,
Percent Ionization = 100%
[H 3 O+]eq = 4.2 10 -3^ M
[HCOOH]initial = 0.10 M
[H 3 O+]eq [HA]initial
Percent Ionization = 100%
4.2 10 -
= 4.2%
Percent Ionization =^ ^ 100%
Method to Calculate pH Using Ka
- Write the chemical equation for the ionization equilibrium.
- Write the equilibrium constant expression.
- Set up a table for Initial/Change in/Equilibrium Concentration to determine equilibrium concentrations as a function of change (x).
- Substitute equilibrium concentrations into the equilibrium constant expression and solve for x. (Make assumptions if possible)
Strong vs. Weak Acids
Differences in conductivity and in rates of chemical reactions.
Polyprotic Acids…
…have more than one acidic proton Easier to remove the first proton than any successive proton. If the difference between the Ka for the first dissociation and subsequent Ka values is 10^4 or more, the pH generally depends only on the first dissociation.
Example (polyprotic acids)
H 3 PO 4 (aq) + H 2 O (l) ⇌ H 3 O+^ (aq) + H 2 PO 4 -^ (aq)
Ka1 = 7.5×10- H 2 PO 4 -^ (aq) + H 2 O (l) ⇌ H 3 O+^ (aq) + HPO 4 2-^ (aq) Ka2 = 6.2×10- HPO 4 2-^ (aq) + H 2 O (l) ⇌ H 3 O+^ (aq) + PO 4 3-^ (aq) Ka3 = 3.6×10- Successive Ka values are smaller; it is less favorable to remove H+^ from an increasingly negatively charged ion.
Weak Bases
Kb can be used to find [OH-] and, through it, pH.
pH of Basic Solutions
What is the pH of a 0.15 M solution of NH 3 at 25 °C? [NH 4 +] [OH-] [NH 3 ] Kb = (^) = 1.8 10 - NH 3 (aq) + H 2 O (l) (^) ⇌ NH 4 +^ (aq) + OH-^ (aq) Tabulate the data. [NH 3 ], M [NH 4 +], M [OH-], M Initial 0.15 0 0 Change -x +x +x Equilibrium 0.15 - x 0.15 x x
Types of Weak Bases
Two main categories
1) Neutral substances with an
Atom that has a nonbonding
pair of electrons that can
accept H+^ (e.g. ammonia
and the amines)
2) Anions of weak acids
Ka and Kb
Ka and Kb are related in this way: Ka Kb = Kw Therefore, if you know one of them, you can calculate the other. at 25 °C
Reactions of Cations with Water
- Attraction between nonbonding electrons on oxygen and the metal causes a shift of the electron density in water.
- This makes the O-H bond more polar and the water more acidic.
- Greater charge and smaller size make a cation more acidic.
Effect of Cations and Anions
- An anion that is the conjugate base of a strong acid will not affect the pH.
- An anion that is the conjugate base of a weak acid will increase the pH.
- A cation that is the conjugate acid of a weak base will decrease the pH.
- Cations of the strong Arrhenius bases will not affect the pH.
- Other metal ions will cause a decrease in pH.
- When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the affect on pH depends on the Ka and Kb values.
Factors Affecting Acid Strength
- The more polar the H-X bond and/or the weaker the H-X bond strength, the more acidic the compound.
- Acidity increases from left to right across a row and from top to bottom down a group.
Factors Affecting Acid Strength
In oxyacids, in which an -OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.
