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Acids and bases - What are they?
Acids and bases - What are they?
For Level 2 Chemistry the definition of acids and bases we use and need to know is the Bronsted-Lowry definition.
For Level 2 Chemistry the definition of acids and bases we use and need to know is the Bronsted-Lowry definition.
This means that an acid can only act as an acid if another substance acts as a base at the same time.
This means that an acid can only act as an acid if another substance acts as a base at the same time.
In this way they form conjugate pairs; the difference in a pair is a proton (H⁺)
In this way they form conjugate pairs; the difference in a pair is a proton (H⁺)
an acid forms a conjugate base
an acid forms a conjugate base
a base forms a conjugate acid
a base forms a conjugate acid
As acids donate H⁺ acidic solutions have an excess of H⁺ in them. Water hydrolyses into H⁺ and OH⁻ equally so a neutral solution has [equal] of H⁺ and OH⁻and the dissociation of water , in equilibrium, is:
As acids donate H⁺ acidic solutions have an excess of H⁺ in them. Water hydrolyses into H⁺ and OH⁻ equally so a neutral solution has [equal] of H⁺ and OH⁻and the dissociation of water , in equilibrium, is:
H₂O ⇄ H⁺ + OH⁻
H₂O ⇄ H⁺ + OH⁻
H⁺ reacts with H₂O to form the hydronium ion, H₃O⁺ (H⁺ and H₃O⁺ are interchangeable), so the dissociation of water can also be written as:
H⁺ reacts with H₂O to form the hydronium ion, H₃O⁺ (H⁺ and H₃O⁺ are interchangeable), so the dissociation of water can also be written as:
2H₂O ⇄ H₃O⁺ + OH⁻
2H₂O ⇄ H₃O⁺ + OH⁻
From this we can see water is in equilibrium and the dissociation constant for water, Kw = 1 x10⁻¹⁴ . As [H₃O⁺] and [OH⁻ ] are equal, each has a concentration of 1 x10⁻⁷ molL⁻¹.
From this we can see water is in equilibrium and the dissociation constant for water, Kw = 1 x10⁻¹⁴ . As [H₃O⁺] and [OH⁻ ] are equal, each has a concentration of 1 x10⁻⁷ molL⁻¹.
Acidic solutions have a greater concentration of H⁺ (H₃O⁺) than OH⁻
Acidic solutions have a greater concentration of H⁺ (H₃O⁺) than OH⁻
Basic solution have a lesser concentration of H⁺ (H₃O⁺) than OH⁻
Basic solution have a lesser concentration of H⁺ (H₃O⁺) than OH⁻
Neutral solutions have an equal concentration of H⁺ (H₃O⁺) than OH⁻
Neutral solutions have an equal concentration of H⁺ (H₃O⁺) than OH⁻
From this it should be apparent that the more H⁺ (H₃O⁺) that is produced in a solution the stronger the acid; the opposite is also true - solutions with a greater amount of OH⁻ are stronger bases.
From this it should be apparent that the more H⁺ (H₃O⁺) that is produced in a solution the stronger the acid; the opposite is also true - solutions with a greater amount of OH⁻ are stronger bases.
The scale that measures acidity is the pH scale, where pH is a measure of the H⁺ ion concentration - the more H⁺ in solution the lower the pH, with the scale going from 1-14 on a logarithmic basis. We always use the pH scale and calculate pH's. We could use the pOH scale, which is a measure of the OH⁻ concentration - it is the converse of the pH scale.
The scale that measures acidity is the pH scale, where pH is a measure of the H⁺ ion concentration - the more H⁺ in solution the lower the pH, with the scale going from 1-14 on a logarithmic basis. We always use the pH scale and calculate pH's. We could use the pOH scale, which is a measure of the OH⁻ concentration - it is the converse of the pH scale.
Remember: pH + pOH = 14
Remember: pH + pOH = 14
See further down this post for the section on calculations for more information on pH / pOH and how they relate to [H₃O⁺] and [OH⁻]
Strong acids and bases compared to weak acids and bases
Strong acids and bases compared to weak acids and bases
This means that:
This means that:
Strong acids and bases are much better conductors of electricity due to the greater amount of dissociated ions in solution
Strong acids and bases are much better conductors of electricity due to the greater amount of dissociated ions in solution
Strong acids and bases undergo faster rates of reactions due to the greater amount of dissociated ions in solution - there are more particles
Strong acids and bases undergo faster rates of reactions due to the greater amount of dissociated ions in solution - there are more particles
Strong acids due to a greater concentration of H₃O⁺ have lower pH's, whilst strong bases with a greater concentration of OH⁻ have higher pH's
Strong acids due to a greater concentration of H₃O⁺ have lower pH's, whilst strong bases with a greater concentration of OH⁻ have higher pH's
Strong bases have a lower pOH, whilst strong acids have a higher pOH
Strong bases have a lower pOH, whilst strong acids have a higher pOH
You will need to be able to answer questions on the differences between strong acids and bases with regards to rates of reaction (see the separate post) and conductivity.
Reactions of acids and bases
Reactions of acids and bases
Metals + acids → Salt + Hydrogen gas
Metals + acids → Salt + Hydrogen gas
Acid + base → Salt + water
Acid + base → Salt + water
Acid + carbonate → Salt + water + carbon dioxide
Acid + carbonate → Salt + water + carbon dioxide
Conjugate pairs
Conjugate pairs
calculations
calculations
As mentioned acids and bases are measured on the pH scale which is the logarithm (to base 10) of the [H⁺] ion. (H₃O⁺) - this gives us a more manageable number. The pOH scale is the converse of the pH scale and is the logarithm (to base 10) of the [OH⁻]. If we know the pOH we can find the pH by subtracting pOH from 14. Here are the 2 key equations:
As mentioned acids and bases are measured on the pH scale which is the logarithm (to base 10) of the [H⁺] ion. (H₃O⁺) - this gives us a more manageable number. The pOH scale is the converse of the pH scale and is the logarithm (to base 10) of the [OH⁻]. If we know the pOH we can find the pH by subtracting pOH from 14. Here are the 2 key equations:
Recap on logs and anti-logs
Recap on logs and anti-logs
It is a straightforward exercise to convert pH back in [H⁺] and pOH back into [OH⁻] by using the anti-log.
It is a straightforward exercise to convert pH back in [H⁺] and pOH back into [OH⁻] by using the anti-log.
Alternatively we can use the dissociation of water, Kw, which is 1 x10⁻¹⁴
Alternatively we can use the dissociation of water, Kw, which is 1 x10⁻¹⁴
Remember the dissociation of water is: H₂O ⇄ H⁺ + OH⁻
Remember the dissociation of water is: H₂O ⇄ H⁺ + OH⁻
From the example to the right, if we wanted to use Kw to find [OH⁻] assuming we were given a pH of 3.21 we would do the following:
From the example to the right, if we wanted to use Kw to find [OH⁻] assuming we were given a pH of 3.21 we would do the following:
Find [H⁺] by using antilog of the pH. This gives us a hydrogen ion concentration, [H⁺] of 6.17 x10⁻⁴
Find [H⁺] by using antilog of the pH. This gives us a hydrogen ion concentration, [H⁺] of 6.17 x10⁻⁴
Write out the Kw expression: Kw = [H⁺] [OH⁻]
Write out the Kw expression: Kw = [H⁺] [OH⁻]
Rearrange to find the hydroxide ion concentration: [OH⁻] = Kw / [H⁺]
Rearrange to find the hydroxide ion concentration: [OH⁻] = Kw / [H⁺]
Enter the numbers: [OH⁻] = 1 x10⁻¹⁴ / 6.17 x10⁻⁴
Enter the numbers: [OH⁻] = 1 x10⁻¹⁴ / 6.17 x10⁻⁴
Solve: [OH⁻] = 1.63 x10⁻¹¹
Solve: [OH⁻] = 1.63 x10⁻¹¹
Be aware of rounding issues - try not to round up until the end, when answers should be to 3 significant figures
Calculations examples
Calculations examples
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