Protein Chemistry

Lecture 1 Video 1 pKa

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🔬 Acid–Base Basics: The Big Picture

We start with the fundamental acid–base equilibrium:

mathrm{HA + H_2O ightleftharpoons H_3O^+ + A^-}

HA = the acid
A⁻ = the conjugate base
This equilibrium tells us how much the acid dissociates in water.

⚖️ Ka – The Acid Dissociation Constant

Ka measures how strongly an acid dissociates.
It is assigned to the acid (HA).
Higher Ka → more dissociation → stronger acid

Key idea:

Acid strength = how much HA wants to become A⁻

Ka values are:

Fixed for a given acid
Only change with temperature (which we ignore here)

🔢 pKa – A Logarithmic Convenience

Instead of using Ka directly, we usually use pKa:

oxed{mathrm{pKa = -log(Ka)}}

🔁 Inverse Relationship

Low pKa → High Ka → Strong acid
High pKa → Low Ka → Weak acid

This inverse relationship is critical and comes up constantly in biochemistry exams.

🧪 Conjugate Bases: Kb and pKb

For the conjugate base (A⁻):

You can define Kb and pKb
These describe how likely the base is to accept a proton

🔗 Acid–Base Pair Rule (Very Important)

For a conjugate acid–base pair in water:

oxed{mathrm{pKa + pKb = 14}}

This means:

Strong acid → weak conjugate base
Weak acid → strong conjugate base

You cannot have both strong at the same time.

🧠 What pKa vs pH Really Mean

pKa → intrinsic property of the molecule (“How much does this acid want to dissociate?”)
pH → property of the environment (“How many protons are actually around?”)

The balance between pH and pKa determines whether the molecule is mostly:

Protonated (HA)
Deprotonated (A⁻)

📐 Henderson–Hasselbalch Equation

This equation connects everything:

oxed{mathrm{pH = pKa + logleft(rac{A^-}{HA} ight)}}

🔑 Rules to remember

Base (A⁻) goes on top
Acid (HA) goes on bottom
pKa is fixed
pH depends on the environment

🔄 Three Key Scenarios (Must Know)

1️⃣ pH < pKa → Acid form dominates

Environment is acidic
Lots of H⁺ available
Protonation favored
(mathrm{HA > A^-})

🧠 Think: “Low pH = hold onto protons”

2️⃣ pH > pKa → Base form dominates

Environment is basic
Fewer protons available
Deprotonation favored
(mathrm{A^- > HA})

3️⃣ pH = pKa → 50/50 mixture

mathrm{HA = A^-}

This point is special:

Exactly half protonated
Exactly half deprotonated

🏥 Real Example: Drugs in the Stomach

Let’s apply this to a physiological setting.

Given:

Stomach pH ≈ 1.5
Two drugs with different pKa values:
- Drug A: pKa = 2.5
- Drug B: pKa = 4.5

We rearrange Henderson–Hasselbalch:

mathrm{pH - pKa = logleft(rac{A^-}{HA} ight)}

🧮 Drug A (pKa = 2.5)

1.5 - 2.5 = -1

rac{A^-}{HA} = 10^{-1} = 0.1

➡️ Some dissociation ➡️ Moderate amount of A⁻

🧮 Drug B (pKa = 4.5)

1.5 - 4.5 = -3

rac{A^-}{HA} = 10^{-3} = 0.001

➡️ Almost no dissociation ➡️ Mostly protonated (HA)

🧠 Final Takeaway from the Drug Example

Lower pKa (stronger acid) → more dissociation at the same pH
Higher pKa (weaker acid) → much less dissociation

Even a 2-unit difference in pKa leads to a 100-fold difference in protonation state.

That’s enormous in biochemistry and pharmacology.

✅ Core Concepts to Lock In

Ka and pKa describe acid strength
pKa is fixed, pH is environment-dependent
Protonation state depends on pH vs pKa
Henderson–Hasselbalch lets you quantify ratios
Small pKa differences → huge biological effects

Quiz

Score: 0/30 (0%)

Q0. What does the acid dissociation constant (Ka) quantitatively describe?

The tendency of a base to accept protons

The extent to which an acid dissociates into H3O+ and A− in water

The pH of a solution containing the acid

The buffering capacity of an acid–base pair

Q1. Which statement correctly describes the relationship between Ka and acid strength?

Lower Ka means a stronger acid

Higher Ka means a stronger acid

Ka is unrelated to acid strength

Ka only applies to bases

Q2. What is the mathematical definition of pKa?

pKa = log(Ka)

pKa = Ka / 14

pKa = −log(Ka)

pKa = log(1/Ka)

Q3. Why is pKa often used instead of Ka in biochemistry?

It removes temperature dependence

It converts Ka into a linear scale

It compresses very large and small Ka values into a manageable range

It applies only to strong acids

Q4. Which relationship between pKa and acid strength is correct?

Higher pKa means stronger acid

Lower pKa means stronger acid

pKa does not correlate with acid strength

Only Ka, not pKa, reflects strength

Q5. For a conjugate acid–base pair in aqueous solution, what is the relationship between pKa and pKb?

pKa − pKb = 14

pKa + pKb = 7

pKa + pKb = 14

pKa × pKb = 14

Q6. If an acid is very strong, how would you describe its conjugate base?

Also very strong

Neutral

Very weak

Independent of the acid

Q7. What does pH primarily measure?

The intrinsic strength of an acid

The equilibrium constant of dissociation

The concentration of protons in the environment

The ratio of acid to base forms only

Q8. Which term is an intrinsic property of an acid molecule?

pKa

Proton concentration

Buffer capacity

Q9. What is the correct form of the Henderson–Hasselbalch equation?

pH = pKa − log(HA/A−)

pH = pKa + log(A−/HA)

pH = pKb + log(HA/A−)

pH = log(A−/HA)

Q10. In the Henderson–Hasselbalch equation, which species is placed in the numerator of the log term?

The acid form (HA)

The proton (H+)

The conjugate base form (A−)

Water

Q11. What happens when pH is lower than pKa?

The base form dominates

The acid and base forms are equal

The acid form dominates

The acid completely dissociates

Q12. What is true when pH equals pKa?

The acid is fully protonated

The acid is fully deprotonated

The concentrations of HA and A− are equal

The Ka value changes

Q13. If pH increases well above pKa, which form is favored?

A−

H3O+

Both equally

Q14. In the stomach (pH ≈ 1.5), which drug will be more protonated if Drug A has pKa 2.5 and Drug B has pKa 4.5?

Drug A

Drug B

Both equally

Neither

Q15. Ka and pKa are fixed values for a given acid unless temperature changes.

True

False

Q16. A lower pKa always corresponds to a higher Ka.

True

False

Q17. pH is an intrinsic property of a molecule.

True

False

Q18. Strong acids have strong conjugate bases.

True

False

Q19. For conjugate acid–base pairs in water, pKa + pKb is always 14.

True

False

Q20. In the Henderson–Hasselbalch equation, the base form is placed in the denominator.

True

False

Q21. When pH < pKa, protonation of the acid is favored.

True

False

Q22. When pH > pKa, the deprotonated form predominates.

True

False

Q23. At pH = pKa, the ratio of A− to HA is 1.

True

False

Q24. A difference of 1 pKa unit corresponds to a 10-fold difference in the A−/HA ratio.

True

False

Q25. At very low pH, most weak acids will be predominantly deprotonated.

True

False

Q26. The Henderson–Hasselbalch equation can be rearranged to calculate the ratio of A− to HA.

True

False

Q27. A drug with a higher pKa will dissociate more in a very acidic environment than a drug with a lower pKa.

True

False

Q28. pKa values alone determine the pH of a solution.

True

False

Q29. The ratio A−/HA becomes very small when pH is far below pKa.

True

False