7.3 Properties of Acids and Bases

Acids and Bases Definitions

Acids and bases show up everywhere in chemistry, from the vinegar in your kitchen to the reactions happening inside your cells. Two major definitions describe what makes something an acid or a base, and knowing the difference between them matters for this unit.

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Fundamental Acid and Base Concepts

The Arrhenius definition is the simpler, older model. It only works for reactions in water (aqueous solutions):

• An Arrhenius acid increases the concentration of $H^+$ ions when dissolved in water (examples: $HCl$, $H_2SO_4$)
• An Arrhenius base increases the concentration of $OH^-$ ions when dissolved in water (examples: $NaOH$, $KOH$)

The Brønsted-Lowry definition is broader. It doesn't require water at all, so it covers more types of reactions:

• A Brønsted-Lowry acid is any substance that donates a proton ($H^+$) to another substance
• A Brønsted-Lowry base is any substance that accepts a proton ($H^+$) from another substance

Comparing the Two Definitions

The key difference is scope. Arrhenius only works in water, while Brønsted-Lowry applies to non-aqueous solutions and even gas-phase reactions.

For example, ammonia ($NH_3$) acts as a Brønsted-Lowry base by accepting a proton, but it doesn't produce $OH^-$ ions directly, so it doesn't fit the Arrhenius definition as neatly. Similarly, $NH_4^+$ can act as a Brønsted-Lowry acid by donating a proton, even though it isn't a classic Arrhenius acid.

Acid and Base Strength

Strength refers to how completely an acid or base breaks apart (dissociates) in water. This is different from concentration, which is about how much acid or base you dissolved in the first place.

Strong vs. Weak Acids

A strong acid completely dissociates in water. Virtually every molecule splits into ions. Common strong acids include $HCl$, $HNO_3$, and $H_2SO_4$.

A weak acid only partially dissociates. Most of the molecules stay intact, and an equilibrium forms between the molecular form and the ionic form. Common weak acids include acetic acid ($CH_3COOH$), hydrofluoric acid ($HF$), and carbonic acid ($H_2CO_3$).

• Strong acids have lower $pK_a$ values, meaning they give up protons more readily
• Weak acids establish an equilibrium in solution, so only a fraction of the molecules are ionized at any given time

Strong vs. Weak Bases

The same logic applies to bases.

A strong base completely dissociates in water. Examples: $NaOH$, $KOH$, $Ca(OH)_2$.

A weak base only partially dissociates. Examples: ammonia ($NH_3$), methylamine ($CH_3NH_2$).

• Strong bases have lower $pK_b$ values (not higher), meaning they accept protons more readily
• Weak bases, like weak acids, establish an equilibrium between their molecular and ionic forms

Factors Affecting Acid and Base Strength

Several things determine how strong an acid or base is:

• Bond strength: The weaker the bond between the hydrogen and the rest of the molecule, the easier it is for the acid to release that proton. For example, $HI$ is a stronger acid than $HF$ partly because the $H-I$ bond is weaker.
• Electronegativity: More electronegative atoms pull electron density away from the hydrogen, making it easier to release as $H^+$.
• Conjugate stability: If the conjugate base (what's left after the acid donates its proton) is stable, the acid is stronger. A stable conjugate base "doesn't want" the proton back.
• Temperature: Higher temperatures generally increase dissociation for weak acids and bases.

Acid-Base Reactions

Conjugate Acid-Base Pairs

When an acid donates a proton, it becomes its conjugate base. When a base accepts a proton, it becomes its conjugate acid. A conjugate pair always differs by exactly one proton ($H^+$).

Here's the important pattern: the strength of an acid is inversely related to the strength of its conjugate base.

• Strong acids have weak conjugate bases. For example, $HCl$ is a strong acid, and its conjugate base $Cl^-$ is very weak (it barely attracts protons back).
• Weak acids have relatively strong conjugate bases. For example, $CH_3COOH$ is a weak acid, and its conjugate base $CH_3COO^-$ has a stronger tendency to accept a proton.

Conjugate pairs are also central to how buffer solutions work. Buffers resist changes in pH because the conjugate acid and base in the pair can neutralize small amounts of added acid or base.

Acid-Base Indicators and Their Applications

Indicators are substances that change color depending on the pH of a solution. They're actually weak acids or bases themselves, and their protonated and deprotonated forms have different colors.

Common indicators you should know:

A universal indicator is a mixture that produces a whole spectrum of colors across the pH range, giving you an approximate pH reading rather than just "acidic or basic."

In titrations, indicators help you find the endpoint, which is the point where the acid and base have fully reacted. Choosing the right indicator matters because different indicators change color at different pH ranges. You want one that changes color close to the expected pH at the equivalence point of your specific titration.