O - Acids, Bases, pH and buffers Flashcards
Give the Bronsted-Lowry definition of an acid
Proton donors
Give the Bronsted-Lowry definition of a base
Proton acceptors
Give the equation for the reaction between an acid and water
HA (aq) + H₂O (l) → H₃O⁺ (aq) + A⁻ (aq)
Give the equation for the reaction between a base and water
B (aq) + H₂O (l) → BH⁺ (aq) + OH⁻ (aq)
Define a strong acid and give an example
Strong acids ionise almost completely in water (nearly all H⁺ released)
Example: Hydrochloric acid
HCl (g) + H₂O (l) → H⁺ (aq) + Cl⁻ (aq)
(this reaction is reversible but the equilibrium is well to the right)
Define a strong base and give an example
Strong bases ionise almost completely in water
Example: Sodium Hydroxide
NaOH (s) + H₂O (l) → Na⁺ (aq) + OH⁻ (aq)
(this reaction is reversible but the equilibrium is well to the right)
Define a weak acid and give an example
Weak acids only ionise very slightly in water, so small numbers of H⁺ ions are formed. An equilibrium is set up which is well over to the left
Example: Ethanoic acid
CH₃OOH (aq) + H₂O (l) ⇌ CH₃OO⁻ (aq) + H⁺ (aq)
Define a weak base and give an example
Weak bases only ionise slightly in water. An equilibrium is set up which is well over to the left
Example: Ammonia
NH₃ (aq) + H₂O (l) ⇌ NH₄⁺ (aq) + OH⁻ (aq)
Describe how acids and bases react
Acids donate their protons to the base
HA (aq) + B (aq) ⇌ BH⁺ (aq) + A⁻ (aq)
Describe what happens when you add an acid to water
The water acts as a base and accepts the protons:
HA (aq) + H₂O (l) → H₃O⁺ (aq) + A⁻ (aq)
Define conjugate pairs
A pair made between an acid and a base either side of the equilibrium which transfer an Hydrogen ion between them
State the conjugate pairs in the following reaction:
HCl + H₂O ⇌ H₃O⁺ + Cl⁻
- HCl and Cl⁻ form a conjugate pair where HCl is the acid
- H₂O and H₃O⁺ form a conjugate pair where H₃O⁺ is the acid
State the conjugate pairs in the following reaction:
B + H₂O ⇌ BH⁺ + OH⁻
where B is a base
- B and BH⁺ form a conjugate pair where BH⁺ is the acid
- H₂O and OH⁻ form conjugate pair where H₂O is the acid
Describe how water can behave as both an acid and a base
It can act as an acid by donating a proton or as a base by accepting a proton
Both hydroxonium ions and hydroxide ions will always be present at the same time due to the equilibrium:
H₂O ⇌ H⁺ + OH⁻
Describe why the equilibrium lies to the left in the following reaction:
H₂O ⇌ H⁺ + OH⁻
Give an expression for the equilibrium constant Kc
Because water only dissociates a small amount
Kc = [H⁺][OH⁻] / [H₂O]
Define Ionic product of water
The Ionic Product of Water, Kw, is the equilibrium constant for the reaction in which water undergoes an acid-base reaction with itself.
Kw = Kc x [H₂O] = [H⁺][OH⁻]
What makes a solution neutral?
An equal concentration of H⁺ and OH⁻ ions
[H⁺] = [OH⁻]
Why is the pH logarithmic?
Because the concentration of hydrogen ions can vary enormously
Give the equation to find the pH of a solution
pH = -log₁₀[H⁺]
Name 2 strong monoprotic acids
Hydrochloric Acid (HCl) Nitric Acid (HNO₃)
Describe the relationship between the concentration of hydrogen ions and the acid concentration for a strong monoprotic acid
Each mole of acid produces 1 mole of hydrogen ions
[HA] = [H⁺]
Describe how do you find the pH of a strong base
[H⁺] is linked to [OH⁻] by the ionic product of water Since Kw = [H⁺][OH⁻] [H⁺] = Kw / [OH⁻] And thus the pH of a base is: pH = -log₁₀[H⁺] = -log₁₀(Kw / [OH⁻])
Give the equation for the equilibrium constant Ka for the reaction:
HA ⇌ H⁺ + A⁻
Ka = [H⁺]² / [HA]
Define pKa
pKa is the negative base-10 logarithm of the acid dissociation constant of a solution
pKa = -log₁₀Ka
Ka = 10⁻ᵖᴷᵃ
The lower the pKa, the stronger the acid
What is Ka?
The acid dissociation constant
Give 5 examples for the use of buffers
1) Shampoo
2) Biological washing powders
3) Regulation of cell pH
4) Regulation of blood pH
5) Food products to control pH change due to bacteria and fungi
State the equation for the process of coupling of a buffer in the blood to the respiratory system
2H₂O + CO₂ ⇌ H₂CO₃ + H₂O ⇌ H₃O⁺ + HCO₃⁻
