Kb and pKb Explained

Ammonia, methylamine, aniline and pyridine all contain a nitrogen atom with a lone pair, and all four are weak bases. Yet at the same concentration, a methylamine solution contains roughly a thousand times more hydroxide ions than an aniline solution. Something about the rest of the molecule changes how willing that nitrogen is to pick up a proton.

To talk about this precisely, chemists use the base dissociation constant, Kb, and its logarithmic form, pKb. This article explains what they mean and, just as importantly, why they take the values they do.

Defining Kb

A weak base B reacts partly with water:

B + H₂O ⇌ BH⁺ + OH⁻

The equilibrium constant for this reaction, leaving out water (whose concentration barely changes), is:

Kb = [BH⁺][OH⁻] ÷ [B]

• Large Kb → the equilibrium lies further right → more OH⁻ → a stronger base.
• Small Kb → most of the base stays unprotonated → a weaker base.

Strong bases such as sodium hydroxide aren’t described with Kb at all, because they’re completely dissociated. Kb is a tool for weak bases.

Defining pKb

Just as with pH and pKa, the negative logarithm makes the numbers manageable:

pKb = −log₁₀ Kb and Kb = 10⁻ᵖᴷᵇ

The lower the pKb, the stronger the base. Each unit lower is ten times stronger.

Base	Formula	Kb	pKb
Dimethylamine	(CH₃)₂NH	5.4 × 10⁻⁴	3.27
Methylamine	CH₃NH₂	4.4 × 10⁻⁴	3.36
Carbonate ion	CO₃²⁻	2.1 × 10⁻⁴	3.67
Ammonia	NH₃	1.8 × 10⁻⁵	4.75
Hydrazine	N₂H₄	1.3 × 10⁻⁶	5.89
Hydroxylamine	NH₂OH	8.7 × 10⁻⁹	8.06
Pyridine	C₅H₅N	1.7 × 10⁻⁹	8.77
Ethanoate ion	CH₃COO⁻	5.6 × 10⁻¹⁰	9.26
Aniline	C₆H₅NH₂	4.3 × 10⁻¹⁰	9.37

Using Kb in calculations

For a weak base at concentration c, as long as only a small fraction reacts:

[OH⁻] ≈ √(Kb × c)

Then pOH = −log[OH⁻] and pH = 14.00 − pOH at 25 °C. Full worked examples, including when to use the quadratic, are in pH of weak bases. Here we’ll focus on what Kb tells you about molecules.

What makes a base stronger or weaker?

A base is strong when its lone pair is available and when the protonated form BH⁺ is stable. Four factors matter most.

1. Electron-donating groups make bases stronger

Alkyl groups such as methyl (–CH₃) push electron density towards the atom they’re attached to. In methylamine, the methyl group makes the nitrogen’s lone pair richer and more eager to bond to a proton, and it helps stabilise the positive charge on CH₃NH₃⁺. That’s why methylamine (pKb 3.36) is a stronger base than ammonia (4.75).

Adding a second methyl group (dimethylamine) strengthens it a little more. Adding a third (trimethylamine, pKb about 4.2) actually makes it slightly weaker in water, because the bulky groups get in the way of water molecules that would otherwise stabilise the cation by hydrogen bonding. Solvation matters as well as electronics.

2. Delocalisation makes bases weaker

In aniline, the nitrogen is bonded directly to a benzene ring. Its lone pair spreads (delocalises) partly into the ring’s π system, so it’s less available to grab a proton. Aniline is about 40,000 times weaker than ammonia as a base (a difference of about 4.6 pKb units).

The same effect explains why amides, such as ethanamide, are barely basic at all: the nitrogen’s lone pair is delocalised onto the neighbouring C=O group.

3. Hybridisation of the lone pair

A lone pair in an orbital with more s character is held closer to the nucleus and is less available. Nitrogen’s lone pair is:

• sp³ in ammonia and amines → most available → strongest bases
• sp² in pyridine → less available → weaker (pKb 8.77)
• sp in nitriles (R–C≡N) → hardly available → essentially not basic in water

4. Electronegative neighbours make bases weaker

In hydroxylamine, NH₂OH, the electronegative oxygen pulls electron density away from nitrogen. Its pKb of about 8 makes it far weaker than ammonia.

Anions as bases

Not every base is a neutral molecule. The anion left behind when a weak acid loses its proton is a weak base, and its Kb follows from the acid’s Ka:

Ka × Kb = Kw so pKa + pKb = 14.00 at 25 °C

This is why the ethanoate and carbonate ions appear in the table above. The weaker the parent acid, the stronger the anion is as a base. Full details are in the Ka × Kb = Kw relationship.

The modern convention: pKaH

Open a modern organic chemistry textbook or database and you may not find pKb values at all. Instead, base strength is given as the pKa of the conjugate acid, often written pKaH.

For ammonia, the conjugate acid is NH₄⁺, with pKa 9.25. So ammonia’s pKaH is 9.25.

The two scales carry the same information:

pKaH = 14.00 − pKb

But the rule flips: a higher pKaH means a stronger base. A base whose conjugate acid holds its proton tightly (high pKa) is a base that grabs protons eagerly.

Base	pKb	pKaH
Methylamine	3.36	10.64
Ammonia	4.75	9.25
Pyridine	8.77	5.23
Aniline	9.37	4.63

Chemists like pKaH because it puts acids and bases on one single scale. To predict whether a proton transfer will happen, you just compare two pKa values.

Where Kb matters

• Medicine. Many drugs are amines. Their pKaH decides what fraction is protonated in the stomach, blood or brain, which affects how well they’re absorbed and where they go.
• Water treatment. Ammonia and amines are used to adjust pH in boiler water, chosen by their base strength and volatility.
• Carbon capture. Amine solutions absorb CO₂ from power-station flue gases. Engineers pick amines with a Kb high enough to capture CO₂ well but low enough that heating releases it again for storage.
• Buffers. The ammonia/ammonium pair buffers solutions around pH 9.25, useful in analytical chemistry, for instance in EDTA titrations for water hardness.

Quick check

1. Which is the stronger base: hydrazine (pKb 5.89) or ammonia (pKb 4.75)? By roughly what factor?
2. The conjugate acid of a base has pKa 7.00. What is the base’s pKb?

Answers: (1) Ammonia, by about 10¹·¹⁴, roughly 14 times. (2) 14.00 − 7.00 = 7.00.

Common mistakes

• Mixing up the direction. Low pKb = strong base, but high pKaH = strong base.
• Using Kb for strong bases. NaOH and KOH are fully dissociated; just use their concentration.
• Forgetting water isn’t in the expression. Kb = [BH⁺][OH⁻] ÷ [B], with no [H₂O].
• Assuming more alkyl groups always means stronger. In water, tertiary amines are often slightly weaker than secondary amines because of solvation.

Key takeaways

• Kb measures how far a weak base reacts with water; pKb = −log Kb.
• Lower pKb means a stronger base; each unit is a factor of ten.
• Electron-donating groups and sp³ lone pairs strengthen bases; delocalisation, sp² or sp lone pairs and electronegative neighbours weaken them.
• pKaH (the pKa of the conjugate acid) is an equivalent scale where higher means more basic.
• For the matching acid side, read Ka and pKa.