# Acid-Base Equilibrium

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Acid-Base Equilibrium
Learning Goal: I will understand the Bronsted-Lowry Theory of acids and bases, be able to identify conjugate acid/base pairs, understand what the autoionization of water means, and understand the difference between strong and weak acids/bases. I will be able to perform calculations to solve for the equilibrium constant (Kw) of the autoionization of water. I will understand what the pH scale really means, and use this knowledge to convert between pH, pOH and the concentration of [H+] and [OH-] ions in solution.

Acids and Bases Arrhenius acids: generate H+ in water
bases: generate OH- in water Brønsted-Lowry acids: H+ donors bases: H+ acceptors HCl H2O  Cl H3O+ acid base

H3O+ = H+(aq) = proton in water

Conjugate acid-base pairs
Conjugate base: remains after H+ is lost acid: HCl conj. base: Cl- Conjugate acid: remains after H+ is gained base: NH3 conj. acid: NH4+ Conjugate Pair HA B- + A- HB + Conjugate Pair

Conjugate Pair HA B- + A- HB + Conjugate Pair

Label the Conjugate Acid Base Pairs for Each of the Following:

Strong and Weak Acids Strong: 100% dissociation good H+ donor equilibrium lies far to right (HNO3) generates weak base (NO3-) Weak: <100% dissociation not-as-good H+ donor equilibrium lies far to left (CH3COOH) generates strong base (CH3COO-)

Strong Acids and Bases

Relative Acid Strength Relative Conj. Base Strength
Very weak Very strong Strong Weak Weak Strong Very weak Very strong The stronger the acid, the weaker its conjugate base, and conversely, the weaker the acid, the stronger its conjugate base

Acid Base Conj. Acid Conj. Base
Water is amphoteric/amphiprotic) meaning it can act as an acid (donates a proton)or a base (accepts a proton) Acid Base Conj. Acid Conj. Base Even pure water conducts electricity, which means it exists as an equilibrium between H2O (l), H3O+ (aq) and OH-(aq)

Autoionization of Water:
Equilibrium Equation for Autoionization of Water: Scientists often omit the water molecule that carries the H+ ion, so the autoionization of water can be changed from: to: Water equilibrium obeys equilibrium law, and can be written as (H2O is excluded because it’s in liquid state):

[H+(aq)] vs [OH-(aq)] Since: Therefore: neutral solutions: [H+(aq)] = [OH-(aq)] =1.0 x10-7 mol/L (pure water) acidic solutions: [H+(aq)] > [OH-(aq)] >1.0 x10-7 mol/L basic solutions: [H+(aq)] < [OH-(aq)] <1.0 x10-7 mol/L

Strength of Acids/Bases
Strong: 100% dissociation Ex: Strong Base: NaOH(s)  Na+(aq) + OH-(aq) Weak: <100% dissociation Ex: Weak Base CH3NH2(aq) + H2O(l) CH3NH2(aq) + OH-(aq)

Example #1: Calculating Unknown Concentrations of [H+(aq)] or [OH-(aq)]
Calculate the hydrogen ion concentration in a 0.25mol/L aqueous barium hydroxide (a strong base) solution.

You Try #1: Calculating Unknown Concentrations of [H+(aq)] or [OH-(aq)]
A 0.15M solution of hydrochloric acid (a strong acid) at SATP has a hydrogen ion concentration of 0.15M. Calculate the concentration of hydroxide ions. Assume STP conditions.

Let’s Try Together: Calculating Unknown Concentrations of [H+(aq)] or [OH-(aq)]
Determine the hydrogen ion and hydroxide ion concentration in 500mL of an aqueous solution containing 2.6g of dissolved NaOH.

What Does pH Really Mean? A measure of the hydronium/hydrogen ion
The hydrogen concentration can range anywhere from 10mol/L to 1x10-15mol/L. The pH scale must be able to cover a large range, so a logarithmic scale that covers factors of 10 is used. The “p” in pH stands for log. A solution with a pH of 1 has [H3O+] or [H+] of 0.1 mol/L or 10-1 a pH of 3 has [H3O+]/ [H+] of mol/L or 10-3 a pH of 7 has [H3O+]/ [H+] of mol/L or 10-7 pH = - log [H3O+] Or pH = - log [H+]

The pH scale pH = - log [H+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14
The pH scale ranges from 1 to 1x mol/L or from 1 to 14. pH = - log [H+] acid neutral base

pH + pOH = 14 ; the entire pH range!
Manipulating pH Algebraic manipulation of: pH = - log [H3O+] allows for: [H3O+] = 10-pH If pH is a measure of the hydronium ion concentration then the same equations could be used to describe the hydroxide (base) concentration. [OH-] = 10-pOH pOH = - log [OH-] thus: pH + pOH = 14 ; the entire pH range!

pH and pOH Formula Summary
pH = - log [H+] [H+] = 10-pH pOH = - log [OH-] [OH-] = 10-pOH pH + pOH = 14 (at SATP)

Example #1 Calculate the pH of a solution with a hydrogen ion concentration of 4.7 x10-11 mol/L

You Try #1 Calculate the pOH of a solution with a hydroxide ion concentration of 3.0x10-6 mol/L

Example # 2 What is the hydrogen ion concentration of a solution with a pH of 10.33?

What is the hydroxide ion concentration of a solution with a pOH 6.2?
You Try # 2 What is the hydroxide ion concentration of a solution with a pOH 6.2?

What is the pOH of a solution whose pH is 8.4?
Let’s Try Together What is the pOH of a solution whose pH is 8.4?

Let’s Throw Another Equation into the Mix
Recap: In neutral solutions: [H+(aq)] = [OH-(aq)] =1.0 x10-7 mol/L That means, at Kw= 1.0 x10-7 mol/L Therefore, pKw=-logKw And pH + pOH =pKw =14 (at STP)

Calculate the pH, pOH and [OH-] of a 0.042 mol/L HNO3 (aq) solution.
Example #3 Calculate the pH, pOH and [OH-] of a mol/L HNO3 (aq) solution.

Calculate the pH, pOH and [OH-]
You Try #3 Calculate the pH, pOH and [OH-] of a M HBr (aq) solution.

Acid-Base Equilibrium
Learning Goal: I will understand the Bronsted-Lowry Theory of acids and bases, be able to identify conjugate acid/base pairs, understand what the autoionization of water means, and understand the difference between strong and weak acids/bases. I will be able to perform calculations to solve for the equilibrium constant (Kw) of the autoionization of water. I will understand what the pH scale really means, and use this knowledge to convert between pH, pOH and the concentration of [H+] and [OH-] ions in solution.