Acids And Bases Flashcards
Def of acid
A proton donor
Def of base
A proton acceptor
Def of salt
When a metal replaces the hydrogen ions on an acid in a neutralisation reaction
pH formula
pH = -log[H+]
Def of pH
The concentration of hydrogen ions in solution
How does the pH scale work?
- what is the logarithmic scale?
- pH 1 has 10x more hydrogen ions than in pH 2
- Therefore the scale goes up in multiples of 10
What type of reaction is dissociation?
- A reversible equation
- Therefore the reaction is in equilibrium
Different indicators
- Universal Indicator
- Phenolphthalein
- Methyl Orange
- Bromophenol Blue
Universal Indicator Colour Changes
- Acid: green to red
- Alkali: green to purple
Phenolphthalein colour changes
- Acid: colourless
- Alkali: pink/dark pink
Methyl Orange Colour changes
- Acid: 1-4 - red
- Alkali: above 4 is yellow
Bromophenol blue colour changes
- Acid: 1-3 - yellow
- Alkali: above 3 - blue
Conjugate acid-base pairs def
Conjugate acid-base pairs differ by the presence or absence of a transferable proton
What word describes water as an acid or base
Amphoteric
- It can act as both an acid and a base
- It can donate and accept protons
What are mono, di and tribasic acids
Monobasic - can donate 1 protons to a base/1 hydrogen ions
Dibasic - can donate 2 protons/hydrogen ions to a base
Tribasic - can donate 3 protons/hydrogen ions to a base
How to work out the pH of an acid when given its concentration in moles
- Write out equation for dissociation of the acid
- Work out if acid is monobasic, dibasic or tribasic
- Work out ratio of moles of H+ ions to base ions
- Work out moles of the number of protons/H+
- N = c x v
- Put concentration of H+ into: pH = -log[H+]
Why can’t pH = - log[H+] be used to work out pH of weak acids
They do not fully dissociate
What is the name given to Kw?
Ionic product
What is the expression for Kw?
Kw = [H+(aq)] x [OH-(aq)]
HA + H2O =(reversible) H3O+ + A-
What is Kc?
why is Kc always going to be very small?
- Kc is always going to be very small, as the no. of moles is always going to be close to 55 in 1dm3 of solution/water
- Therefore the value of Kc will always be very low
What substance is not included when calculating Ka
- Why?
- Water
- The concentration of water in solution is always going to be very high
- Therefore the concentration of water is effectively CONSTANT
What is Ka
Acid dissociation constant
Equation for Ka
Ka = [H+][A-] / [HA]
Times when approximations for Ka equation could be wrong
- If Ka is particularly high
- For very weak acids, pH > 6, the dissociation is no longer negligible
- Therefore [H+]equilibrium = [A-]equilibrium does not apply
- [HA]equilibrium = [HA]start only holds true for acids with a vet small Ka
- does not hold true for stringer weak acids or very dilute solutions
Assumptions made in calculations for Ka of weak acids
Kemdndhhdbd
Calculation for pKa
pKa = -log(Ka)
- Ka = 10^-pKa
Method for working out the pH of a strong base
> Work out the concentration of the hydroxide ions.
Use Kw to work out the hydrogen ion concentration.
Convert the hydrogen ion concentration to a pH.
What is the value of Kw at 298K?
1.00 x 10^-14 mol^2dm^-6
Equation for Kw - for strong bases
Kw = [H+][OH-]
Buffer def
A system that minimises pH changes on addition of small amounts of an acid or a base
What is in a buffer?
Buffered contain a weak acid and a salt of that acid, e.g. ethanoic acid and sodium ethanoate
In a buffer solution containing ethanoic acid and sodium ethanoate, what reactions occur
- CH3COOH = CH3COO- + H+
2. H2O = H+ + OH-
What will happen to the equilibria if acid is added to the buffer solution?
- H+ ion conc will increase
- H+ ions react with ethanoate ions removing them
- Equilbrium reaction 1 (acid dissociation) will shift left
- Minimal change in pH
What happens to the equilibria if an alkali is added to a buffer solution?
- OH- ion concentration will increase. This will react with and remove H+ ions
- H+ ion conc. decreases
- Equilbrium 2 reaction (water dissociation) will shift right
- More of the water will dissociate to replace the missing protons
- Minimal change in pH
How could you make a buffer solution with sodium hydroxide and excess ethanoic acid?
1.
Explain how buffers minimise change. Use carbonic acid (in blood) as an example of a buffer.
Dissociation:
H2CO3 = H+ + HCO3-
Add alkali:
- Increase concentration of OH- ions
- OH- + H+ —> H2O
- Conc. of H+ ions decreases
- equilibrium shifts right
- replacing H+
- so the change in pH is minimised
Add acid:
- Increase conc. of H+
- H+ + HCO3- —> H2CO3
- equilibrium shifts left
- minimise change in pH
Assumptions made when making buffer solutions
- Acid has not dissociated at all
2. Our salt has fully dissociated
What does pH of the buffer depend on?
- Ka of the acid
2. The concentration ratio of the acid:conjugate base
When doing titrations, would the titration values be different if a weak acid was used.
- Why?
- There would be no difference in titration values
- The same number of moles is still required if a weak acid is used
- As when H+ dissociates form a weak acid, more will start to dissociate
- Eventually, all the H+ will dissociate from the weak acid during titration
What factors determine the pH of a buffer solution?
2 Marks
- Ka - acid dissociation constant
- Temperature
- Concentration of weak acid and conjugate salt/base
Give one reason why the pH scale is a more convenient measurement for acid concentration than [H+].
(1 Mark)
- Deals with negative indices over a very wide range
- pH makes numbers manageable
- removes very small numbers
- any 1 of these reasons for 1 Mark
What information is provided by Ka values?
- The strength of the acid
Or - extent of acid dissociation in solution
Predict and explain the acid-base reaction that would take place if ethanoic acid were mixed with phenol. Include an equation in your answer.
(2 Marks)
- Ethanoic acid is the stronger acid/Ka of ethanoic acid is greater
C6H5OH + CH3COOH = C6H5OH2+ + CH3COO-
An excess of magnesium was added to 100 cm3 of 0.0450 mol dm–3 hydrochloric acid. The same mass of magnesium was added to 100 cm3 of 0.0450 mol dm–3 ethanoic
acid.
Both reactions produced 54 cm3 of hydrogen gas, measured at room temperature and pressure, but the reaction with ethanoic acid took much longer to produce this gas volume.
Explain why the reactions produced the same volume of a gas but at different rates. Use equations in your answer.
(4 Marks)
- Ethanoic acid is a weaker acid than Hydrochloric acid
- HCl and CH3COOH have the same number of moles
- So they release the same number of moles of H+
- Mg + 2HCl —> MgCl2 + H2
- Mg + 2CH3COOH —> (CH3COO-)2Mg2+ + H2
How to calculate concentration of a given volume of nitric acid (HNO3) diluted with the same amount of water (H2O)?
Concentration of HNO3 halves:
E.g. 0.015/2 = 0.0075moldm-3
A student measured the pH of water as 7.0 at 25 °C. The student then warmed the water to 40 °C and measured the pH as 6.7.
What do these results tell you about the tendency of water to ionise as it gets warmer? Explain your reasoning in terms of equilibrium.
(2 Marks)
- [H+] increases
- As H2O ionises more/dissociates more
- H2O = OH- + H+
- Equilibrium shifts to the right
Significance of equivalence point in different titration curves
- The point in titration at which the amount of titrant added is just enough to completely neutralize the analyte solution (in acid-base titration)
- Moles of base = moles of acid
- Solution only contains salt and water.
pH curve info for strong base, strong acid neutralisation
- start point: 2 pH
- end point: 12 pH
- equivalence point: 7 - 10/11 pH
- phenolphthalein suitable
- methyl orange suitable
pH curve info for weak base, weak acid neutralisation
start point: 3
end point: 9
equivalence point: 7
- no indicator is suitable
pH curve info for weak base, strong acid neutralisation
start point: 1.2
end point: 9 pH
equivalence point: 3 - 7 pH
- methyl orange suitable indicator
pH curve info for strong base, weak acid neutralisation
start point: 3 pH
end point: 12 pH
equivalence point: 7-11 pH
- phenolphtalein suitable indicator
reasons why pH scale is a more convenient measurement for measuring acid concentration than [H+].
(1 Mark)
makes (small or very large) numbers manageable
removes very small numbers
[H+] deals with negative indices over a very wide range
what information is provided by Ka values?
strength of acid
extent of acid dissociation
what two chemicals are needed to make a buffer solution?
weak acid and its conjugate base
