Acids and Bases

Quick introduction to acids and bases.

Acid-base chemistry is everywhere: in the sour tang of citrus, the bitterness of alkalis, the colors of hydrangeas, the fizz of baking, the balance of blood, and the heart of heavy industry. At its core the rule is simple—acids donate protons and bases accept them—but the consequences shape both our daily lives and the world economy.

Acids and Bases

Acid-base chemistry is about the transfer of protons (H⁺). The word acid comes from the Latin acidus (“sour”), while alkali traces back through Arabic to mean “ashes of saltwort,” since early bases were made from plant ash. At the simplest level, acids donate protons and bases accept them. When an acid and a base meet, the proton is transferred and a salt is formed.

The classic definition is due to Brønsted and Lowry:

• Acid: proton donor
• Base: proton acceptor

Hydrochloric acid, HCl, is a strong acid: in water it dissociates completely, \[ \mathrm{HCl \;\to\; H^+ + Cl^-}. \] Sodium hydroxide, NaOH, is a strong base: it releases hydroxide ions, \[ \mathrm{NaOH \;\to\; Na^+ + OH^-}. \] Mix the two and the H⁺ and OH⁻ combine to form water, leaving Na⁺ and Cl⁻ as a neutral salt: \[ \mathrm{HCl + NaOH \;\to\; NaCl + H_2O}. \]

Why H⁺ really means H₃O⁺

In water, free protons do not exist on their own. They immediately attach to water molecules to form hydronium, H₃O⁺. So when we write “H⁺,” it is shorthand for a proton transferred through water as H₃O⁺. Chains of hydrogen bonds allow this proton to “hop” rapidly from one water molecule to another, a unique feature of aqueous acid-base chemistry.

A weak acid only partly dissociates. Acetic acid (CH₃COOH), the acid in vinegar, sets up an equilibrium: \[ \mathrm{CH_3COOH \;\rightleftharpoons\; H^+ + CH_3COO^-}. \] This incomplete ionization is why vinegar tastes sour but does not burn like hydrochloric acid. Weak bases are similar: ammonia, NH₃, accepts a proton from water to make ammonium, NH₄⁺, but only partially.

Every acid has a conjugate base: the molecule left behind after giving up its proton. For example, when hydrochloric acid loses H⁺, the leftover Cl⁻ is its conjugate base. Likewise, bases have conjugate acids: when ammonia (NH₃) accepts H⁺, the product NH₄⁺ is its conjugate acid. Acids and bases thus come in pairs.

Taste

Acids typically taste sour—think lemon juice or vinegar. Bases tend to taste bitter and feel slippery or soapy on the skin. Sodium hydroxide (NaOH) and ammonia (NH₃) are obvious examples, though both are too caustic to taste directly. Everyday safe bases include bicarbonate (baking soda) and certain vegetables, which have a mild alkaline bite.

Soap

Soap is made by reacting fats or oils with a strong base such as sodium hydroxide (NaOH) in a process called saponification. The reaction breaks triglycerides into glycerol and the sodium (or potassium) salts of long-chain fatty acids—these salts are what we call soap. Chemically, soap is therefore a salt, not a pure base. However, when dissolved in water the fatty acid salts partially react with water, creating a mildly alkaline solution. This is why soap feels slippery and can irritate skin: it is just basic enough to change the surface chemistry of oils and dirt, letting them wash away.

Salts

In chemistry, a salt is the product of an acid-base reaction: the positive ion (cation) from the base and the negative ion (anion) from the acid combine. Sodium chloride (NaCl) is the most familiar, but many salts exist: potassium nitrate (KNO₃), calcium carbonate (CaCO₃), ammonium sulfate ((NH₄)₂SO₄). The key point is that a salt is neutral overall, even if its parent acid and base were strong and reactive.

Measuring acidity: pH and pKₐ

Chemists use two related but distinct measures:

First, pH, a property of the solution. It measures the concentration of free protons (H⁺, really H₃O⁺) in water \[ \mathrm{pH = -\log_{10}[H^+]}. \] The concentration of free protons is written [H⁺], with square brackets meaning “molar concentration.” Pure water has [H⁺] = 1×10⁻⁷ mol/L, which corresponds to pH 7.

Neutral water has pH ≈ 7. Thus, the scale runs roughly 0–14, with 7 neutral. pH < 7 means acidic, pH > 7 means basic. Hydrochloric acid solution has low pH, sodium hydroxide solution has high pH. Because it is logarithmic, each unit represents a 10-fold change in [H⁺].

Second is pKₐ, a property of the molecule. It measures how willing a molecule is to give up its proton. Low pKₐ = strong acid (readily donates H⁺), high pKₐ = weak acid. Acetic acid has pKₐ ≈ 4.8 (weak), hydrochloric acid pKₐ < 0 (very strong).

Key differences: pH changes with concentration, while pKₐ is fixed for each molecule. pH is measured in a solution, pKₐ is looked up in a table. Together they describe how acidic a solution is and why.

Litmus and flowers

Litmus paper is a simple pH indicator: red in acid, blue in base. It comes from dyes extracted from lichens. Many natural pigments behave similarly. The blue–pink color change of hydrangea flowers really is linked to soil acidity—acidic soils (low pH) produce blue blooms, while alkaline soils give pink. The mechanism involves aluminum ions, which are soluble in acidic soils and alter the pigment chemistry in the petals.

How much does a mole of protons weigh?

A single proton has a mass of about \(1.673 \times 10^{-27}\) kg. Multiply by Avogadro’s number (\(6.022 \times 10^{23}\)) and you get about 1.007 g per mole of H⁺. This is why the atomic weight of hydrogen is close to 1.

What does this mean for water? Pure water at 25 °C has pH 7, so [H⁺] = \(1 \times 10^{-7}\) mol/L. That works out to: \[ 10^{-7} \ \text{mol/L} \times 1.007 \ \text{g/mol} \approx 1 \times 10^{-7} \ \text{g per liter}. \] That is just 0.1 micrograms of protons in a liter of neutral water—utterly negligible compared to the liter’s total mass of ~1000 g.

So while pH is conceptually about counting protons, in mass terms the protons are vanishingly small compared to the water itself.

What do pH and pKₐ mean?

Both terms come from German chemical notation:

pH – “p” stands for potenz (power), and H is hydrogen. Together it means power of hydrogen, i.e. the negative logarithm of the hydrogen ion concentration. It is a solution property specific to water.

pKₐ – “p” again means potenz (negative log), “K” is the equilibrium constant, and the subscript “a” stands for acid. It is the negative logarithm of the acid dissociation constant. It is meaningful in any solvent.

pH of water

At 25 °C (room temperature), pure water is empirically observed to have \([H^+]= [OH^-] \approx 1 \times 10^{-7}\) mol/L, which corresponds to pH = 7. This comes from the autoionization of water: \[ \mathrm{2H_2O \;\rightleftharpoons\; H_3O^+ + OH^-}. \] The equilibrium constant for this process is called the ionic product of water, \[ K_w = [H^+][OH^-]. \] At 25 °C, \(K_w \approx 1.0 \times 10^{-14}\). Taking the square root gives \([H^+]= [OH^-] = 1 \times 10^{-7}\) mol/L, so pH = 7.

But this is temperature dependent. At 0 °C, \(K_w \approx 0.11 \times 10^{-14}\), so pH is about 7.5. At 50 °C, \(K_w \approx 5.5 \times 10^{-14}\), so pH is about 6.6.

How many protons at extreme pH?

Using [H⁺] = \(\mathrm{10^{-\mathrm{pH}}}\) and molar mass of H⁺ ≈ 1.007 g/mol at 25 °C:

pH	[H⁺] (mol·L⁻¹)	H⁺ mass per L (g)	\([OH^-]\) (mol·L⁻¹)*
−1	\(10\)	\(10.07\)	\(10^{-15}\)
0	\(1\)	\(1.007\)	\(10^{-14}\)
1	\(10^{-1}\)	\(1.01\times10^{-1}\)	\(10^{-13}\)
7	\(10^{-7}\)	\(1.01\times10^{-7}\)	\(10^{-7}\)
13	\(10^{-13}\)	\(1.01\times10^{-13}\)	\(10^{-1}\)
14	\(10^{-14}\)	\(1.01\times10^{-14}\)	\(1\)

Values of pH < 0 and > 14 occur for very concentrated strong acids/bases (e.g., ~12 M HCl gives pH ≈ −1). The numbers above use “ideal” behavior. In real concentrated solutions, activity ≠ concentration, so pH is reported on an activity scale and the simple back-calculation gives an approximation. Even at pH 0, the liter only contains ~1 g of free protons—tiny versus ~1000 g of water.

The classic “fountain” experiment

Strong acids like hydrochloric (HCl) or nitric acid (HNO₃) have such a strong affinity for water that they can pull it up a tube. In a famous classroom experiment, concentrated HCl gas is dissolved in water inside a flask fitted with a tube. As the gas dissolves, the volume contracts and water is literally sucked up the tube into the flask, creating a dramatic “fountain” effect. It’s a vivid demonstration of how powerfully some acids hydrate.

Buffers and how they work

A buffer is a mixture of a weak acid and its conjugate base. It resists swings in pH because any added H⁺ is absorbed by the base, while any added OH⁻ is neutralized by the acid. For example, in an acetic acid/acetate buffer:

• Add H⁺ → acetate (CH₃COO⁻) grabs it, forming CH₃COOH.
• Add OH⁻ → acetic acid donates a proton to neutralize it, forming water plus acetate.

In both cases, the ratio of acid to base shifts slightly, but the overall [H⁺] stays nearly constant. Human blood relies on the carbonic acid/bicarbonate buffer (H₂CO₃/HCO₃⁻) to maintain pH ~7.4—without it, even small amounts of acid from metabolism would be fatal.

Extremes

The strongest common acids are hydrochloric, nitric, and sulfuric acid, which at high concentration drive pH close to 0. Superacids like fluoroantimonic acid (HSbF₆) go far beyond, with effective pH values below –20. On the base side, sodium hydroxide and potassium hydroxide reach pH ~14; still stronger are superbases such as butyllithium, which react violently with water and must be handled in special solvents.

Beyond water

Water is the most familiar solvent for acid–base chemistry, but it is not the only one. The Brønsted–Lowry idea—acids donate protons (H⁺), bases accept them—works in other liquids too. What changes is which ions take the roles of “acid form” and “base form.”

• Liquid ammonia (NH₃): Here, an acid is anything that produces the ammonium ion (NH₄⁺), while a base is anything that produces the amide ion (NH₂⁻). For example, dissolving ammonium chloride (NH₄Cl) in liquid ammonia releases NH₄⁺, showing its acidic behavior.
• Sulfuric acid (H₂SO₄) as a solvent: In pure sulfuric acid, acids generate H₃SO₄⁺ ions, and bases generate HSO₄⁻ ions. So sulfuric acid can act as both solvent and participant in acid–base reactions.
• Other solvents: In hydrogen fluoride (HF), the key ions are H₂F⁺ and F⁻. In molten salts or acetic acid, similar conjugate pairs define acidity and basicity.

Water is special because it is amphoteric—it can act as either acid (giving H⁺ to become OH⁻) or base (accepting H⁺ to become H₃O⁺). But the underlying idea is universal: in any solvent, acids and bases are defined by proton transfer, with the solvent itself carrying those protons back and forth.

Industry and everyday life

Acid-base reactions underpin much of modern life. Sulfuric acid is the world’s most produced chemical, used in fertilizer, petroleum refining, and chemical manufacturing. Hydrochloric acid is vital for food processing and metal treatment. Sodium hydroxide drives paper making, soap production, and water treatment. Ammonia is the base behind nitrogen fertilizers that sustain global agriculture. In the kitchen, baking soda reacting with vinegar shows the same chemistry: an acid and a base react, producing carbon dioxide bubbles that make dough rise.

Common acids, bases, and salts

Chemical	Type	Common name / notes	Everyday / industrial use	Typical pH (aq)
HCl	Acid	Hydrochloric acid	Stomach acid; metal cleaning; food processing	0–1 (strong)
H₂SO₄	Acid	Sulfuric acid	Fertilizers; batteries; petroleum refining	0 (strong)
HNO₃	Acid	Nitric acid	Explosives (TNT); fertilizers; etching metals	0–1 (strong)
CH₃COOH	Acid	Acetic acid	Vinegar; food preservative; plastics	2–3 (weak)
H₂CO₃	Acid	Carbonic acid	Sparkling drinks; blood buffer system	3–4 (weak)
HF	Acid	Hydrofluoric acid	Glass etching; semiconductor industry	1–2 (strong, toxic)
NaOH	Base	Sodium hydroxide (caustic soda)	Soap making; paper; water treatment	13–14 (strong)
KOH	Base	Potassium hydroxide	Alkaline batteries; biodiesel production	13–14 (strong)
NH₃	Base	Ammonia	Fertilizers; cleaning products	11–12 (weak)
Ca(OH)₂	Base	Slaked lime	Mortar; plaster; water treatment	12 (moderate)
NaHCO₃	Base	Baking soda	Baking; antacids; fire extinguishers	8–9 (mild)
Mg(OH)₂	Base	Milk of magnesia	Antacid; laxative	10.5 (mild)
NaCl	Salt	Sodium chloride	Table salt; de-icing; food preservation	~7 (neutral)
KNO₃	Salt	Potassium nitrate (saltpeter)	Fertilizers; fireworks; gunpowder	~7 (neutral)
CaCO₃	Salt	Calcium carbonate	Limestone; chalk; shells; construction	~9 (slightly basic, sparingly soluble)
(NH₄)₂SO₄	Salt	Ammonium sulfate	Fertilizer; soil treatment	~5 (slightly acidic)
Na₂CO₃	Salt	Washing soda	Glass making; cleaning agents	11 (basic)
NaOCl	Salt	Sodium hypochlorite	Household bleach; disinfectant	11–12 (basic)