pH and pKa
Introduced| English |
|---|
| pKa/piː kɑː/ |
A number for acid strength
- Comparing tiny $K_a$ values full of exponents is awkward.
- A logarithm turns each into a friendly small number.
- Now stronger acids get lower scores.
- One glance ranks any list of acids.
Defining pKa
- pKa 酸度系数 measures acid strength: $\text{p}K_a = -\log K_a$.
- A smaller $\text{p}K_a$ means a stronger acid.
- It compresses the huge range of $K_a$ values.
$\text{p}K_a$ is defined as...
$\text{p}K_a = -\log K_a$.
Lower pKa, stronger acid
- Because of the negative log, a big $K_a$ gives a small $\text{p}K_a$.
- So the strongest acids have the lowest $\text{p}K_a$.
- Rank acids by lining up their $\text{p}K_a$ values.
Which acid is stronger?
A lower $\text{p}K_a$ means a stronger acid.
A larger $K_a$ corresponds to a smaller $\text{p}K_a$.
The negative log inverts the order.
Comparing pKa to pH
- If a solution's pH is below the acid's $\text{p}K_a$, the acid form dominates.
- If pH is above $\text{p}K_a$, the base form dominates.
- They are present equally when pH equals $\text{p}K_a$.
Read strength from pKa
Sort each acid by its pKa. A smaller pKa means a stronger acid.
When a solution's pH equals the acid's $\text{p}K_a$, the acid and conjugate base are...
At $\text{pH} = \text{p}K_a$, the two forms are equal.
When pH is below $\text{p}K_a$, the ____ form of the species dominates.
A more acidic solution favours the protonated (acid) form.
Acid A has $\text{p}K_a = 3$; acid B has $\text{p}K_a = 5$. Which is stronger?
- A lower $\text{p}K_a$ means a stronger acid.
- So acid A ($\text{p}K_a = 3$) is the stronger acid.
$\text{p}K_a$ is a fixed property of the acid, while pH depends on the solution.
$\text{p}K_a$ describes the acid; pH describes the mixture.
A smaller $\text{p}K_a$ is a stronger acid -- the negative log flips the order, a common mix-up. $\text{p}K_a$ is a property of the acid; pH is a property of the solution. And when $\text{pH} = \text{p}K_a$, the acid and its conjugate base are present in equal amounts.
pKa $= -\log K_a$ ranks acid strength on a friendly scale: a smaller $\text{p}K_a$ is a stronger acid. Comparing it to a solution's pH tells you which form dominates -- the acid form below $\text{p}K_a$, the base form above, and equal amounts when $\text{pH} = \text{p}K_a$.