Acids and Bases in Organic Chemistry: pKa, Stability, and Predicting Proton Transfers
Acid-base chemistry is the single highest-leverage topic in the first semester, because proton transfers are everywhere. Activating an alcohol before substitution, protonating a carbonyl, generating an enolate, choosing a base that eliminates instead of substituting, working up a reaction and separating products in a funnel: all of it is acid-base reasoning wearing different clothes.
The good news is that one number does most of the work. If you can read a pKa correctly, explain where it comes from, and use it to predict which way a proton will move, you can answer a large share of exam questions without memorizing separate rules for each situation. This guide covers what pKa really tells you, how to rank conjugate base stability from structure, the values worth knowing cold, and how to turn all of it into a prediction.
Two Definitions You Use for Different Jobs
The Bronsted-Lowry definition is about protons: an acid donates one, a base accepts one. This is the definition you use whenever you compare strengths, because pKa is a Bronsted measurement. Every Bronsted acid has a conjugate base, the species left behind after the proton leaves, and the two are locked together. A strong acid has a stable, weak conjugate base; a weak acid has an unstable, strong conjugate base.
The Lewis definition is broader and about electrons: an acid accepts an electron pair, a base donates one. It covers species with no proton at all, such as AlCl3 or BF3, and it is the definition hiding behind arrow pushing. Every curved arrow in a mechanism runs from a Lewis base to a Lewis acid. Use Bronsted language when you are ranking acidity and Lewis language when you are explaining why a particular arrow is drawn.
Reading pKa Without Getting the Direction Backwards
pKa is a logarithmic measure of how readily a proton comes off, and the direction catches people out constantly: a lower pKa means a stronger acid. Because the scale is logarithmic, each unit is a factor of ten, so the gap between acetic acid at about 4.76 and ethanol at about 16 is more than eleven orders of magnitude, not a modest difference. Negative values are ordinary; hydronium sits near -1.7 and HCl near -7.
The other half of the skill is reading pKa as a statement about the conjugate base. A low pKa means the conjugate base is stable and comfortable holding the electrons left behind. A high pKa means the conjugate base is unstable, which makes it a strong, reactive base. So when you are asked to compare two acids, do not stare at the acids. Draw both conjugate bases and ask which one handles its electrons better.
- Lower pKa = stronger acid. One pKa unit is a factor of ten.
- Strong acid, stable and weak conjugate base; weak acid, unstable and strong conjugate base.
- To compare acidity, draw the conjugate bases and rank their stability.
- In water the strongest base that can exist is hydroxide, which is why stronger bases are made in other solvents.
Ranking Conjugate Base Stability: Charge, Atom, Resonance, Induction, Orbital
Five structural factors explain nearly every acidity comparison, and they are usually applied in this order. Charge comes first: a neutral species is more stable than an anion, so the protonated form of something is always the stronger acid. Atom comes next, and it has two halves. Across a row of the periodic table electronegativity dominates, which is why the trend runs methane near 50, ammonia near 38, water at 15.7 and HF near 3.2. Down a column size dominates, because a larger atom spreads the charge over more volume, which is why a thiol near 10 is far more acidic than an alcohol near 16.
Resonance is the third and often the decisive factor: delocalizing the negative charge over several atoms stabilizes it strongly, which is why acetic acid at 4.76 beats ethanol at 16 and phenol at about 10 beats cyclohexanol. Induction is the fourth, a weaker through-bond pull from electronegative atoms that fades with distance; chloroacetic acid sits near 2.9 and trifluoroacetic acid near 0.5, both well below acetic acid. Orbital hybridization is the last: more s character holds electrons closer to the nucleus, so a terminal alkyne C-H near 25 is far more acidic than an alkene C-H near 44 or an alkane C-H near 50.
The pKa Values Worth Knowing Cold
You do not need a full table in your head. A short anchor set covering the common families lets you place almost anything by interpolation, and every value below is approximate because solvent and substitution shift it. What matters is the neighborhood and the ordering, not the second decimal place.
Notice how much of the list is carbonyl chemistry. The alpha C-H of a ketone near 19 to 20 is millions of times more acidic than an ordinary alkane, because the resulting enolate delocalizes onto oxygen, while an ester alpha C-H near 25 is less acidic because the ester oxygen already donates into the carbonyl. Those two numbers are why enolate chemistry uses the bases it does.
- Strong mineral acids: HCl about -7, sulfuric acid about -3, hydronium -1.7.
- Carboxylic acid about 4 to 5 (acetic acid 4.76); ammonium about 9 to 10; phenol about 10.
- Water 15.7; alcohols about 16 to 18 (ethanol 16); amide N-H about 17.
- Ketone alpha C-H about 19 to 20; ester alpha C-H about 25; terminal alkyne about 25.
- Ammonia N-H about 38; alkene C-H about 44; alkane C-H about 50.
Predicting Which Way the Proton Goes
A proton transfer favors the side with the weaker acid, which is the side with the higher pKa. That one sentence answers most questions of the form 'will this base deprotonate that?'. You can even estimate the equilibrium constant: it is roughly ten raised to the difference between the pKa of the acid you form and the pKa of the acid you started with, so a gap of a few units already means the reaction runs essentially to completion.
Work an example. Can hydroxide deprotonate a terminal alkyne? The alkyne is about 25 and the water you would form is 15.7, so the stronger acid ends up on the product side and the equilibrium sits far to the left. Sodium amide works instead, because the ammonia formed is about 38 and the gap now runs the right way. The same arithmetic explains a standard separation: a carboxylic acid at 4.76 dissolves in aqueous sodium bicarbonate, since carbonic acid at about 6.4 is the weaker acid, while phenol at about 10 does not and needs hydroxide instead. Pick a base whose conjugate acid has a pKa comfortably above the proton you are trying to remove, and you will choose correctly every time. Octet's reference tables keep the pKa values one tap away while you practise, and the quiz lets you test the direction rule until it is reflex.
Frequently asked questions
Does a lower pKa mean a stronger or weaker acid?
A lower pKa means a stronger acid. The scale is logarithmic, so each unit is a factor of ten. Acetic acid at 4.76 is over eleven orders of magnitude more acidic than ethanol at about 16, and negative values such as -7 for HCl are perfectly normal.
How do I tell which of two acids is stronger from structure?
Draw both conjugate bases and rank their stability using charge, then the atom holding the charge, then resonance, then induction, then hybridization. The more stable the conjugate base, the stronger the acid. Resonance is usually the factor that decides realistic exam comparisons.
Which way does an acid-base reaction go?
Toward the weaker acid, meaning the side with the higher pKa. Estimate the equilibrium constant as ten raised to the pKa of the acid formed minus the pKa of the acid consumed. A positive difference of a few units means the reaction essentially goes to completion.
Why can't hydroxide deprotonate a terminal alkyne?
Because a terminal alkyne is around pKa 25 and water is 15.7, so hydroxide would generate the stronger acid and the equilibrium stays on the alkyne side. Sodium amide works, since the ammonia it forms is around 38 and the proton transfer becomes strongly favorable.
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