3.9 Acid-Base Equilbiria Flashcards
Describe the Lowry-Bronsted theory
- acid base equilibria involve the transfer of protons between substances
- a Lowry-Bronsted acid is a proton donor
- a Lowry-Bronsted base is a proton acceptor
Strong acid vs weak acid
Strong: completely dissociates into its ions
Weak: partially dissociates
What is the equilibrium dissociation constant?
Ka
Describe the relationship between Ka and acid strength
- stronger acid = higher Ka value
- strong acids dissociate to ions more, meaning the concentration of ions at equilibria is much greater making Ka increase in value
When is methyl orange used?
- reactions with a more acidic neutralisation point
- red in acids and turns yellow at neutralisation point
When is phenolpthaliein used?
- reactions with more basic neutralisation point
- pink in alkali and turns colourless at neutralisation point
pH equation
pH = -log[H+]
[H+] conc equation
10^-pH
Differentiate between monoprotic, diprotic and triprotic acids
Mono: conc of H+ ions same as conc of acid
Di: conc of H+ ions is double acid conc
Tri: conc of H+ ions is triple acid conc
What is Ka?
Acid dissociation constant
What does the size of Ka tell you?
Larger = more dissociation
Smaller = less dissociation
What is Kw?
The ionic product of water
Kw equation at 25°C
Kw = [H+][OH-]
Kw = 1x10^-14
Define a buffer
Any system that minimises pH changes by addition of small amounts of acid or base
—> can’t prevent completely but maintained for as long as buffer solution remains
Difference between acidic and basic buffer
Acidic: made of weak acid and salt of the acid
Basic: mass of weak base and salt of base
Describe how addition of acid and alkali impact an acidic buffer
- if small amounts of acid added to buffer, equilibrium shifts lefts removing nearly all excess H+
- if small amounts of alkali is added to the buffer, the OH- ions will react to the H+ ions to form water
—> equilibrium right and more H+ ions produced
Describe how adding acidic and alkali impacts basic buffed
- adding acid shifts equilibrium to right to remove H+ ions from solution
- adding base shifts equilibrium to left as reacting OH- ions with H+ to form water
pH of buffer equations
pKa + log([A-]/[HA])
pKa = -logKa
Ka = 10^-pKa
What is meant by the equivalence point?
The point at which the volume of one solution has reacted exactly with the volume of the second solution
Indicator for strong acid, strong base
Phenolphathlein or methyl orange
Indicator for strong acid, weak base
Methyl orange
Indicator for weak acid, strong base
Phenolphthalein
Graphs for strong acid and strong base
Graphs for weak acid and weak base
Graph for strong acid and weak base
Graph for weak acid and strong base
pH of bases equations
[H+] = Kw.
——
OH-
pH = -log Kw
——
OH-
Difference between neutral salts, acidic salts and basic salts
Neutral: salt produced from strong acid and strong alkali
Acidic: salt produced from a strong acid and a weak alkali
Basic: salt produced from a weak acid and strong alkali
Ammonia equilibria equation
NH4+ —> NH3 + H+
Ethanoate equilibria equation
CH3COO- + H+ —> CH3COOH
What is meant by conjugate acid-base pairs?
2 species with differing H+
What factors determine the pH of a buffer solution?
- temperature
- Ka
How do you find pKa on a titration curve?
pKa = pH at half neutralisation point
Uses for buffer solutions
- storing/using enzymes
- fermentation
- dyeing
Difference between base and alkali
Base: species that accepts a proton
Alkali: a soluble base that releases OH- ions
