Acids and Bases Flashcards
What is a Brønsted acid and base in terms of protons and hydrogen ions
Acid = proton donor, release H+
Base = proton acceptor, attract H+
What does Brønsted acid-base equilibria involve
Transfer of protons
Equation that related pH and H+
pH = -log10[H+]
[H+] = 10^-pH
What is Ka
An equilibrium constant for the dissociation of an acid
What is the equation for Kw in pure water
Kw = [H+]^2
Equation for pKa
pKa = -log10Ka
What does it mean when acids and bases dissociate in water
Break up into positively and negatively charged ions
What dissociates the most and why
Strong acids = releasing all H+ ions
Strong bases = ionise completely
Weaker bonds dissociate more
What happens to the equilibrium if you add more acid or base in acid + base = salt + water
Position of equilibrium shifts right
(opposite happens if you add more salt or water)
What does the dissociation of water equation look like
H2O <–> H+ + OH-
What is a conjugate acid-base pair
A conjugate acid-base pair is two species that are different from each other by an H+ ion
what is Kc
the equilibrium constant, measured using products and reactants in either the gaseous or aqueous state
What is Kw
Kw i= autoprolysis constant of water
(at 25 degrees Celsius)
Is always = 1.0 x 10^-14.
What is the Kc and Kw equation for the dissociation of water
Kc = [H+][OH-] / [H2O]
Kw = [H+][OH-]
Why is [H2O] not shown in the Kw expression
[H2O] is almost constant
What is a pH scale a measure of
Hydrogen ion concentration in a solution
measured in mol dm-3
What is the equation for strong acid ionisation
HA (aq) -> H+ (aq) + A- (aq)
What is the equivalence point
where the amount of acid neutralises the alkali or vice versa
Kc equations
Kc x H2O = [H2O][OH-]
Strong base ionisation equation
BOH (aq) → B+ (aq) + OH- (aq)
Strong base equations:
[H+] = Kw/[OH-}
[H+] = 1 x 10^-14 / [OH-]
pH = -log[H+] = -log (Kw (1 x 10^=14) / [OH-])
Calculate the pH of 0.15 mol dm-3 sodium hydroxide, NaOH
[H+] = Kw ÷ [OH-]
[H+] = (1 x 10-14) / 0.15 = 6.66 x 10-14
pH = -log[H+]
= -log 6.66 x 10-14 = 13.17
Calculate the hydroxide concentration of a solution of sodium hydroxide when the pH is 10.50
Step 1: Calculate H concentration
[H+]= 10-pH
[H+]= 10-10.50
[H+]= 3.16 x 10-11 mol dm-3
Step 2: Rearrange the ionic product of water
Kw = [H+] [OH-]
[OH-]= Kw ÷ [H+]
Step 3: Find the concentration of hydroxide ions
Kw is 1 x 10-14 mol2 dm-6,
[OH-]= (1 x 10-14) ÷ (3.16 x 10-11)
[OH-]= 3.16 x 10-4 mol dm-3
What is Ka
Dissociation constant for a weak acid
Equation for Pka
pKa = -log10Ka
Ka equation
Ka = [H+][A-] / [HA]
What are you looking for when you are calculating:
pH
Ionic product of water
Weak acid dissociation constant
Weak acid pH
pH = -log10[H+]
Ionic product of water = Kw = [H+][OH-]
Equation for pKa = -log10 Ka
Weak acid pH = Ka = [H+][A-] / [HA]
What does the higher or lower the value of Ka mean
higher the value of Ka the more dissociated the acid and the stronger it is
The lower the value of Ka the weaker the acid
How to find the pH of weak acids step by step
1) Write Ka for weak acid
( Ka = [H+]^2 / [HA] )
2) Rearrange equation and substitute to find [H+]^2
3) Square root the number to find [H+]
4) Substitute [H+] into pH
( pH = -log10 [H+] )
How to find concentration of weak acids step by step
1) Substitute pH into pH equation to calculate [H+]
( [H+]=10^-pH )
2) Write an expression for Ka
3) Rerrange to give concentration
( [H+]^2 / Ka )
4) Substitute Ka and [H+] into equation
Find Ka of weak acids step by step
1) Use pH of acid to find [H+]
2) Write an expression for Ka
3) Substitute values for [H+] and [A-]
Step by step calculating Ka
1) Concert pKa to Ka
( Ka = 10^-pKa )
2) Write Ka expression
3) Rerrange to give [H+]^2
4) Square root it
5) Substitute [H+] into pH equation and solve
How to calibrate pH meter
1) Rinse probe with distilled water, shake off excess, place in standard buffer solution, record pH reading in table
2) Repeat solutions of other concentrations, rinsing the probe between each step
3) Plot graph of recorded pH against pH of buffer
What is Kw When going from [OH-] to [H+]
Kw = [H+][OH-]
As temperature increases, what happens to pH?
pH drops because water dissociated more so there are more H+
How to find the equivalence point in a titration graph?
In the middle of the graph where the line is completely straight going up
How to find the half equivalence point?
Half neutralised
So distance between straight line going up, half of it
What is the half equivalence point
PH = pKa
Pure weak acid equation by itself
Ka = [H+]^2 / [HA]
describe how the student would carry out a titration to obtain a pH curve
1) Use pH probe to measure the initial pH of ammonia solution
2) Add Hcl acid 2cm^3 at a time
3) Stir mixture
4) Record pH after each addition
5) Reduce size portions when close to the end point
6) Repeat until acid in excess
7) Plot a graph of the results
Explain why [H2O] is not shown in the Kw expression
[H2O] is very large compared to [H+
] and [OH-]
Explain why the value of Kw increases as the temperature increases
Equilibrium is endothermic
so shifts to the right to oppose temperature increase
Explain why water is neutral at 50 °C
[H+] = [OH–]
Suggest why the pH probe is washed with distilled water between each of
the calibration measurements.
To wash off any residual solution/substance (which could
interfere with the reading)
Explain why the volume of sodium hydroxide solution added between each
pH measurement is smaller as the end point of the titration is approached.
To wash off any residual solution/substance (which could
interfere with the reading)
How do you know if the indicators are suitable in the titration
All will have a colour change/pH range within the steep/vertical part
of the titration curve
State the meaning of the term weak acid.
partially dissociates (in water to form H+ ions)
