Compare Arrhenius and Bronsted-Lowry acids and bases, conjugate pairs, strong and weak acids, pH, pOH, neutralization, and titration basics.
A lesson of Chemistry, in Equilibrium, Acids and Bases, and Redox, about 35 minutes.
By the end of this lesson, you will be able to define an acid three ways, calculate pH, tell strength from concentration, and explain what a neutralization and a titration actually do.
Three Definitions, Each Wider Than the Last
Arrhenius. An acid releases $\mathrm{H^+}$ in water; a base releases
$\mathrm{OH^-}$. Simple, and it only works in water.
Brønsted-Lowry. An acid donates a proton; a base accepts one. This
frees the idea from water entirely, and it is the definition to use by default.
Lewis. An acid accepts an electron pair; a base donates one. The widest of the
three, and it covers substances with no hydrogen at all, such as
$\mathrm{BF_3}$.
Each definition contains the previous one. Every Arrhenius acid is a
Brønsted-Lowry acid, and every Brønsted-Lowry acid is a Lewis acid, but not the
reverse. That nesting is the reason all three survive: each is the most convenient
tool at its own level.
A free proton does not exist in water. $\mathrm{H^+}$ is a bare nucleus, and it
attaches immediately to a water molecule to form the hydronium ion
$\mathrm{H_3O^+}$. Writing $\mathrm{H^+}$ is a shorthand for that, and both forms
mean the same thing.
Conjugate pairs differ by exactly one proton. $\mathrm{HCl}$ and
$\mathrm{Cl^-}$ are a pair; so are $\mathrm{H_2O}$ and $\mathrm{OH^-}$. Water
appears in pairs on both sides, which is why it can act as an acid or a base
depending on what it meets.
Sort It: Acid or Base
By the Brønsted-Lowry definition.
Strength Is Not Concentration
This is the distinction the whole topic rests on, and the two words are used
interchangeably in ordinary speech, which is why the confusion is universal.
Strength is about how completely it dissociates, and it is a property of the
substance.
Concentration is about how much is dissolved, and it is a property of your
particular solution.
The weak acid uses an equilibrium arrow, and that is the whole difference. Most
of the acetic acid molecules in vinegar are sitting there intact at any moment.
So a dilute strong acid can be less hazardous than a concentrated weak one.
Very dilute hydrochloric acid is harmless, and glacial acetic acid causes serious
burns. Strong does not mean dangerous and weak does not mean safe.
Each pH unit is a factor of TEN in concentration, and each two units is a
factor of a hundred. A pH of $3$ has a thousand times more $\mathrm{H^+}$ than a
pH of $6$, so "slightly more acidic" can mean an enormous chemical difference.
That is why ocean acidification matters at a change of $0.1$ pH units, and why
lemon juice at pH $2$ is not merely twice as sour as coffee at pH $5$.
In water the two concentrations are locked together:
Adding acid does not only raise $[\mathrm{H^+}]$; it forces $[\mathrm{OH^-}]$ down
to compensate, because their product is fixed. Neither is ever zero.
Easy values worth recognizing: $[\mathrm{H^+}] = 10^{-3}$ gives pH $3$, and
$10^{-11}$ gives pH $11$. When the concentration is a clean power of ten, the pH is
just the exponent without its sign.
Example 1: Calculating pH
Here is every step, worked all the way through.
$0.001$ M $\mathrm{HCl}$.
Is HCl strong or weak? Type one word strong
A strong acid dissociates fully, so the H+ concentration equals the acid concentration. Write 0.001 as a power of ten 1e-3
pH is minus the log of that. What is the pH? 3
What is the pOH? 11
Worked Example 1: pH and pOH
A solution has $[\mathrm{OH^-}] = 1 \times 10^{-4}$ M.
The sodium and chloride ions were dissolved before and are dissolved after. They
are spectator ions, and stripping them out gives the net ionic equation,
which is the same for every strong acid and strong base pair.
Neutralization is strongly exothermic, releasing about $57$ kJ per mole of
water formed, and that value is nearly identical for all strong acid-base pairs
precisely because the same net reaction is occurring.
The product is not always neutral, which surprises people given the name. A
strong acid with a weak base gives a slightly acidic salt; a weak acid with a
strong base gives a slightly basic one. Sodium acetate solution is basic even
though it was made by neutralization.
Titration
A titration finds an unknown concentration by reacting it with a known one.
Add the known solution from a burette until an indicator changes color. That
color change is the endpoint, and the underlying event is the equivalence
point, where the moles of acid and base exactly match.
$$M_aV_a \times (\text{ratio}) = M_bV_b$$
The ratio comes from the balanced equation. $\mathrm{H_2SO_4}$ supplies two
protons per molecule, so it needs twice as much $\mathrm{NaOH}$ as $\mathrm{HCl}$
would, and forgetting that halves the answer.
Choose an indicator that changes near the equivalence pH. Phenolphthalein
turns at about pH $9$, which suits a weak acid with a strong base. Methyl orange
turns at about pH $4$, which suits a strong acid with a weak base. Using the wrong
one gives a color change at the wrong volume.
Rinse the burette with the solution it will hold, not with water, since
residual water dilutes the titrant and inflates the volume delivered.
Worked Example 2: A Titration Calculation
$25.0$ mL of $\mathrm{HCl}$ needs $20.0$ mL of $0.100$ M $\mathrm{NaOH}$.
Strong means fully dissociated. Concentrated means a lot of it.
Mistake 2: Reading pH as linear
Each unit is a factor of ten.
Mistake 3: Treating a weak acid's $[\mathrm{H^+}]$ as its concentration
Only strong acids dissociate fully.
Mistake 4: Forgetting the ratio in a titration
$\mathrm{H_2SO_4}$ needs twice the base.
Mistake 5: Assuming every neutralization gives pH 7
Only strong with strong does.
Where This Goes Next
Buffers are where equilibrium and acids meet, and they keep you alive.
A buffer is a weak acid together with its conjugate base, and it resists pH
change. Add acid, and the base half consumes it; add base, and the acid half
consumes that. Le Chatelier's principle, deployed deliberately.
Blood is buffered by that pair, and it is held between pH $7.35$ and $7.45$.
Below $6.8$ or above $7.8$ is fatal, and the buffer is what keeps the margin.
Hyperventilation blows off $\mathrm{CO_2}$, shifting that equilibrium left, raising
the pH and causing the dizziness and tingling that follow, and breathing into a
bag works by restoring the $\mathrm{CO_2}$.
Ocean acidification is the same chemistry at planetary scale. Dissolved
$\mathrm{CO_2}$ forms carbonic acid, and ocean pH has fallen about $0.1$ units
since industrialization. On a logarithmic scale that is a $30%$ increase in
hydrogen ion concentration, and it makes it harder for corals and shellfish to
build calcium carbonate skeletons.
Soil pH decides what will grow, and blueberries at pH $4.5$ and asparagus at
pH $8$ cannot share a bed. Agriculture manages pH deliberately, with lime to raise
it and sulfur to lower it.
The pH scale looks like a laboratory abstraction and turns out to be one of the
most consequential numbers in biology, medicine and the environment.
Check Your Understanding
Five of them, no instructions this time.
Summary
Three definitions, each wider: releases $\mathrm{H^+}$, donates a proton,
accepts an electron pair.
Strength is dissociation; concentration is amount. They are independent.
Each pH unit is a factor of ten, so small-looking shifts are large chemical
changes.
Neutralization is really $\mathrm{H^+} + \mathrm{OH^-} \rightarrow \mathrm{H_2O}$,
and a titration finds an unknown concentration using the balanced ratio.