Properties, ions in water, strong and weak acids and bases, the pH scale, indicators, neutralization, titration, buffers, and acids and bases in life and the environment
You have encountered acids and bases many times without thinking much about the chemistry behind them. The tart taste of lemon juice, the sting of vinegar on a cut, the burn of stomach acid — these are the sensations of acids at work. The slippery feel of soapy water, the bitterness of baking soda, the chemical that unclogs drains — these are signatures of bases.
Acids and bases are not exotic laboratory substances. They are among the most important classes of compounds in chemistry, biology, agriculture, medicine, industry, and environmental science. Your blood is kept at an almost perfectly constant acid-base balance by chemical systems that operate every second without your awareness. Farmers adjust the acid content of soil so crops can absorb nutrients. Water treatment plants monitor acid-base chemistry of drinking water carefully. Doctors use acid-base information to diagnose diseases.
In this chapter, you will move from the familiar experience of sour and slippery to the chemical explanation underneath. You will learn what makes a substance an acid or a base at the particle level, how chemists measure acidity using the pH scale, how acids and bases react with each other, and how these ideas connect to human health, agriculture, water quality, and the environment.
This chapter builds directly on ions in water (Chapter 17), chemical reactions and equations (Chapters 13 and 14), molecular polarity and intermolecular forces (Chapter 12), and equilibrium (Chapter 22). Acids and bases behave the way they do because of how molecules interact with water, release or capture ions, and reach dynamic balance in solution.
Long before chemists understood atoms and ions, they recognized acids and bases by their behavior. Over centuries of observation, a consistent set of properties emerged.
Acids share a recognizable set of behaviors. They have a sour taste — lemon juice, vinegar, and grapefruit juice are all acidic. They react with certain metals, such as zinc and iron, releasing hydrogen gas. They react with carbonates and bicarbonates, producing carbon dioxide bubbles. They change the color of certain plant-based dyes called indicators — blue litmus paper turns red in the presence of an acid. Acids conduct electricity when dissolved in water, which means they produce ions in solution.
Bases share their own recognizable properties. They taste bitter — unsweetened cocoa, coffee, and baking soda are mildly basic. They feel slippery or soapy to the touch. They do not react with metals the same way acids do. They also change indicator colors, but in the opposite direction — red litmus paper turns blue in a base. Like acids, bases conduct electricity in water solution.
When acids and bases are mixed together, they react in a process called neutralization. The characteristic properties of both tend to disappear. A neutralized acid-base mixture often produces a salt dissolved in water.
| Property | Acids | Bases |
|---|---|---|
| Taste | Sour (citric, acetic) | Bitter (caffeine, baking soda) |
| Feel | Normal to skin | Slippery, soapy |
| Litmus | Turns blue litmus red | Turns red litmus blue |
| Metals | React with many metals, releasing H₂ | Generally do not react with metals the same way |
| Carbonates | React with carbonates, releasing CO₂ | React differently; may dissolve carbonates slowly |
| Electrical conductivity | Conduct electricity in water solution | Conduct electricity in water solution |
| Mixed together | Neutralization: produce salt and (usually) water; properties diminish | |
Many acids and bases encountered in laboratories and households are corrosive. Concentrated acids like sulfuric acid and hydrochloric acid can cause severe chemical burns. Strong bases like sodium hydroxide are also highly corrosive. Even dilute solutions require proper eye protection and careful handling. Never taste laboratory chemicals. The observable properties described here are for scientific understanding only.
In the late 1800s, Swedish chemist Svante Arrhenius proposed definitions that connected acid-base behavior to ions in water. His definitions were among the first to explain why acids and bases have the properties they do.
Arrhenius Acid — A substance that produces hydrogen ions (H⁺) when dissolved in water.
Arrhenius Base — A substance that produces hydroxide ions (OH⁻) when dissolved in water.
Example — HCl in water: HCl → H⁺ + Cl⁻
Example — NaOH in water: NaOH → Na⁺ + OH⁻
The hydrogen ion, H⁺, is essentially just a bare proton. In water, bare protons immediately attach to water molecules (H₂O), forming hydronium ions (H₃O⁺). Chemists sometimes write H⁺ and H₃O⁺ interchangeably in acid-base chemistry, understanding that H⁺ in water is always hydrated. Throughout this chapter, we may use H⁺ as a shorthand for the hydronium ion.
The Arrhenius model has limits. Ammonia (NH₃) behaves as a base, but it contains no OH⁻ and cannot produce OH⁻ directly. In the 1920s, Johannes Brønsted and Thomas Lowry independently proposed broader definitions that solved this problem.
Brønsted-Lowry Acid — A proton donor: a substance that gives H⁺ to another substance.
Brønsted-Lowry Base — A proton acceptor: a substance that takes H⁺ from another substance.
Example — Ammonia in water: NH₃ + H₂O → NH₄⁺ + OH⁻
Water donates a proton (acts as acid); ammonia accepts it (acts as base). OH⁻ produced makes the solution basic.
Water can act as either an acid or a base — it is amphoteric. Water also ionizes very slightly on its own (autoionization): H₂O + H₂O → H₃O⁺ + OH⁻. In pure water at 25°C, [H₃O⁺] = [OH⁻] = 1 × 10⁻⁷ mol/L. This equal, tiny concentration of both ions is why pure water is neutral.
Not all acids and bases behave with the same intensity. Some completely release their ions when dissolved in water. Others only partially ionize, reaching an equilibrium between the ionized and un-ionized forms. This difference is one of the most important distinctions in acid-base chemistry.
Strong acid or base — Ionizes completely (≈ 100%) in water. No equilibrium arrow needed.
Weak acid or base — Only partially ionizes. An equilibrium (⇌) is established between the ionized and un-ionized forms.
Strength describes the degree of ionization — NOT the concentration. A strong acid can be dilute; a weak acid can be concentrated.
| Acid/Base | Name | Strength | Ionization in Water |
|---|---|---|---|
HCl | Hydrochloric acid | Strong acid | HCl → H⁺ + Cl⁻ (complete) |
HNO₃ | Nitric acid | Strong acid | HNO₃ → H⁺ + NO₃⁻ (complete) |
H₂SO₄ | Sulfuric acid | Strong acid (1st H) | H₂SO₄ → H⁺ + HSO₄⁻ (complete) |
HClO₄ | Perchloric acid | Strong acid | HClO₄ → H⁺ + ClO₄⁻ (complete) |
HBr, HI | Hydrobromic / hydroiodic | Strong acid | Complete ionization |
CH₃COOH | Acetic acid (vinegar) | Weak acid | CH₃COOH ⇌ H⁺ + CH₃COO⁻ (partial) |
H₂CO₃ | Carbonic acid | Weak acid | H₂CO₃ ⇌ H⁺ + HCO₃⁻ (partial) |
NaOH | Sodium hydroxide | Strong base | NaOH → Na⁺ + OH⁻ (complete) |
KOH | Potassium hydroxide | Strong base | KOH → K⁺ + OH⁻ (complete) |
NH₃ | Ammonia | Weak base | NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (partial) |
Strong and weak describe the degree of ionization, not the amount of acid. A 0.001 mol/L HCl solution is very dilute but HCl is still a strong acid — it still ionizes 100%. A 10 mol/L acetic acid solution is highly concentrated but acetic acid is still a weak acid. These properties are independent.
The concentration of H⁺ ions in aqueous solutions spans an enormous range — from 1 mol/L in strongly acidic solutions to 1 × 10⁻¹⁴ mol/L in strongly basic solutions. Working directly with such numbers is awkward. Chemists use a compressed, logarithmic scale called the pH scale to represent these values conveniently.
pH = −log[H⁺]
[H⁺] = 10⁻ᵖᴴ
Where [H⁺] = molar concentration of hydrogen ions (mol/L). The brackets indicate concentration. Because of the negative sign, higher [H⁺] gives a lower pH number. Lower pH = more acidic.
[H⁺] × [OH⁻] = 1.0 × 10⁻¹⁴ (at 25°C)
This product is constant for all dilute aqueous solutions at 25°C. When [H⁺] increases, [OH⁻] decreases by the same factor, and vice versa. In pure water: [H⁺] = [OH⁻] = 1 × 10⁻⁷ mol/L, giving pH = 7.
[OH⁻] = (1.0 × 10⁻¹⁴) ÷ [H⁺]
| pH Range | Classification | Relationship | Examples |
|---|---|---|---|
| 0 – 1 | Strongly acidic | [H⁺] ≫ [OH⁻] | Battery acid, concentrated HCl |
| 1 – 3 | Strongly acidic | [H⁺] ≫ [OH⁻] | Stomach acid (1.5–2.5), lemon juice (2–3) |
| 3 – 5 | Acidic | [H⁺] > [OH⁻] | Vinegar (2.5–3.5), coffee (4.5–5) |
| 5 – 7 | Mildly acidic | [H⁺] > [OH⁻] | Clean rainwater (5.6), milk (6.5–6.8) |
| 7 | Neutral | [H⁺] = [OH⁻] | Pure water at 25°C |
| 7 – 9 | Mildly basic | [H⁺] < [OH⁻] | Blood (7.35–7.45), seawater (7.9–8.3) |
| 9 – 11 | Basic | [H⁺] < [OH⁻] | Baking soda solution (8.5), antacid (10) |
| 11 – 14 | Strongly basic | [H⁺] ≪ [OH⁻] | Household ammonia (11–12), drain cleaner (13–14) |
Each one-unit change in pH represents a tenfold change in [H⁺]. A solution at pH 4 has 10× more H⁺ than pH 5, and 100× more than pH 6. A pH change of 0.1 units — small as it sounds — represents about a 26% change in H⁺ concentration. This is why even small pH shifts in blood or ocean water carry large consequences.
0.01 = 1 × 10⁻²
pH = −log(1 × 10⁻²) = −(−2) = 2
[H⁺] = 10⁻¹¹ = 1 × 10⁻¹¹ mol/L[H⁺] × [OH⁻] = 1 × 10⁻¹⁴
[OH⁻] = (1 × 10⁻¹⁴) ÷ (1 × 10⁻¹¹) = 1 × 10⁻³ mol/L
An indicator is a substance that changes color depending on the pH of the solution it is in. Indicators are typically weak acids or weak bases where the ionized and un-ionized forms have different colors. When pH changes, the balance between the two forms shifts, and the color changes. Long before pH was formally defined, chemists used natural plant pigments — from red cabbage, berries, turmeric, and flowers — that responded to acids and bases with color changes.
| Indicator | Color in Acid | Color in Base | Useful pH Range |
|---|---|---|---|
| Litmus | Red | Blue | pH ~5–8 (rough yes/no) |
| Methyl orange | Red (below 3.1) | Orange-yellow (above 4.4) | pH 3.1–4.4 |
| Bromothymol blue | Yellow (below 6) | Blue (above 7.6); green near 7 | pH 6–7.6 |
| Phenolphthalein | Colorless (below 8.2) | Pink/magenta (above 10) | pH 8.2–10 |
| Universal indicator | Red → orange → yellow | Blue → violet | Full pH 1–14 spectrum |
An indicator is a weak acid (HIn) where HIn is one color and In⁻ is another. The equilibrium HIn ⇌ H⁺ + In⁻ shifts depending on [H⁺]. In acidic solutions, extra H⁺ pushes the equilibrium left (HIn dominates, color A visible). In basic solutions, OH⁻ removes H⁺, pulling the equilibrium right (In⁻ dominates, color B visible). The indicator's color reports which direction equilibrium has been pushed — and that depends on pH.
Indicators give approximate pH values. For precise measurements, chemists use pH meters — instruments with electrodes that generate a voltage depending on [H⁺]. The meter converts voltage to a pH reading accurate to two decimal places. pH meters must be calibrated using buffer solutions of known pH before use.
When an acid and a base are combined in the right proportions, they react to form water and a salt. This process is called neutralization because the characteristic acidic and basic properties are reduced or eliminated as the H⁺ and OH⁻ ions combine.
acid + base → salt + water
The net ionic equation for any strong acid-strong base neutralization is always: H⁺ + OH⁻ → H₂O
A salt is any ionic compound formed from the cation of the base and the anion of the acid.
| Acid | Base | Salt Formed | Balanced Equation |
|---|---|---|---|
| HCl | NaOH | NaCl (sodium chloride) | HCl + NaOH → NaCl + H₂O |
| HNO₃ | KOH | KNO₃ (potassium nitrate) | HNO₃ + KOH → KNO₃ + H₂O |
| H₂SO₄ | 2 NaOH | Na₂SO₄ (sodium sulfate) | H₂SO₄ + 2 NaOH → Na₂SO₄ + 2 H₂O |
| CH₃COOH | NaOH | NaCH₃COO (sodium acetate) | CH₃COOH + NaOH → NaCH₃COO + H₂O |
H₂SO₄ + KOH → K₂SO₄ + H₂OH₂SO₄ + 2 KOH → K₂SO₄ + 2 H₂ONo. The word neutralization describes the reaction type, not that the result is exactly pH 7. Strong acid + strong base → pH 7. Weak acid + strong base → pH above 7 (basic). Strong acid + weak base → pH below 7 (acidic). The final pH depends on the relative strengths of the acid and base involved. The salt produced may partially react with water (hydrolysis), producing H⁺ or OH⁻ ions.
Titration is a laboratory procedure in which a solution of known concentration (the titrant) is carefully added to a solution of unknown concentration until the reaction between them is complete. The point at which the reaction is complete is called the equivalence point or endpoint. At the equivalence point, the moles of H⁺ from the acid exactly equal the moles of OH⁻ from the base.
A titration uses a burette — a long, graduated glass tube with a stopcock valve at the bottom. The burette allows the titrant to be delivered in precise, controlled amounts and read to two decimal places. The solution being analyzed (the analyte) is placed in an Erlenmeyer flask below the burette, with a few drops of indicator added.
M_acid × V_acid = M_base × V_base
Where M = molarity (mol/L) and V = volume in liters (or consistent units). At the equivalence point, moles of acid = moles of base. Since moles = M × V, the formula follows directly. Rearrange to solve for any one unknown.
M_acid × V_acid = M_base × V_base
M_acid = (M_base × V_base) ÷ V_acid
M_acid = (0.200 mol/L × 0.01850 L) ÷ 0.02500 L
M_acid = 0.003700 mol ÷ 0.02500 L
M_acid = 0.148 mol/L
For strong acid–strong base titrations, phenolphthalein or bromothymol blue both work well because the pH jumps sharply near the equivalence point. For weak acid–strong base titrations (equivalence point above 7), use phenolphthalein. For strong acid–weak base titrations (equivalence point below 7), use methyl orange or methyl red. The indicator should change color as close to the equivalence point as possible.
Living systems, water systems, and industrial processes often need to maintain a nearly constant pH even when small amounts of acid or base are added or produced. A buffer is a solution that resists changes in pH when small amounts of acid or base are added to it.
A buffer typically contains a weak acid and its conjugate base (the salt of the weak acid), or a weak base and its conjugate acid. The two components work as a team — one handles added acid, and the other handles added base.
H₂CO₃ ⇌ H⁺ + HCO₃⁻
When acid (H⁺) is added: HCO₃⁻ + H⁺ → H₂CO₃ (H⁺ is consumed; pH barely changes)
When base (OH⁻) is added: H₂CO₃ + OH⁻ → HCO₃⁻ + H₂O (OH⁻ is consumed; pH barely changes)
A buffer works only as long as the supply of both components is not exhausted. Exceeding buffer capacity causes the pH to change sharply.
Human blood is maintained at pH 7.35–7.45. If blood pH drops below 7.35, the condition is called acidosis. If it rises above 7.45, it is called alkalosis. Both can be life-threatening if untreated. The carbonic acid–bicarbonate buffer works together with the lungs (which control CO₂ by adjusting breathing rate) and the kidneys (which adjust HCO₃⁻ concentration) to keep blood pH in this narrow safe range even as the body constantly produces acids through metabolism.
Your stomach produces hydrochloric acid (HCl) with a pH of approximately 1.5 to 2.5 — strongly acidic. This acid denatures proteins, making them accessible to digestive enzymes, activates the enzyme pepsin, and kills most bacteria that enter with food. The small intestine operates at pH 6 to 7. The pancreas secretes bicarbonate (HCO₃⁻) into the small intestine to neutralize the arriving stomach acid, raising the pH to the range where intestinal enzymes can function.
Antacids work by neutralizing excess stomach acid. They typically contain basic compounds such as calcium carbonate (CaCO₃), magnesium hydroxide (Mg(OH)₂), or aluminum hydroxide (Al(OH)₃). H₂ blockers and proton pump inhibitors reduce acid production upstream.
Soil pH is one of the most important factors controlling plant growth. Most crops grow best in soil with pH between 6.0 and 7.0. At this range, essential nutrients — nitrogen, phosphorus, potassium, calcium, magnesium, iron — are in chemical forms that plant roots can absorb.
When soil becomes too acidic (pH below 5.5), aluminum and manganese ions dissolve in concentrations toxic to plants, and phosphorus becomes locked in insoluble compounds. When soil is too basic (pH above 7.5), iron, manganese, and micronutrients become insoluble and unavailable. Farmers add lime (CaCO₃) to raise pH, or sulfur compounds to lower it. Soil pH testing is routine in professional agriculture.
Clean rainwater is slightly acidic — about pH 5.6 — because atmospheric CO₂ dissolves in it to form carbonic acid. Acid rain refers to precipitation with pH below about 5.0, forming when sulfur dioxide (SO₂) from burning sulfur-containing fuels and nitrogen oxides (NOₓ) from combustion react with water vapor:
CO₂(g) + H₂O(l) → H₂CO₃(aq) (clean rainwater, pH 5.6)
2 SO₂ + O₂ + 2 H₂O → 2 H₂SO₄ (sulfuric acid — acid rain)
NO₂ + H₂O → HNO₃ (nitric acid — acid rain)
Acid rain damages forests, acidifies lakes and streams, corrodes stone and metal structures, and dissolves limestone monuments (CaCO₃ + acid → soluble products + CO₂).
When CO₂ from burning fossil fuels dissolves into seawater, it forms carbonic acid, lowering ocean pH. The average ocean pH has decreased by about 0.1 pH units since pre-industrial times — representing approximately a 26% increase in H⁺ concentration due to the logarithmic scale. Ocean acidification threatens marine organisms that build calcium carbonate shells and skeletons — corals, oysters, clams, sea urchins, and organisms at the base of the marine food chain:
CaCO₃ + CO₂ + H₂O → Ca²⁺ + 2 HCO₃⁻
Shells and coral skeletons dissolve as ocean water becomes more acidic. This is an active area of scientific research and environmental monitoring.
Many pharmaceuticals are either weak acids or weak bases. The pH of the body environment affects how well drugs are absorbed. A weak-acid drug will be mostly un-ionized in the acidic stomach (where it can cross lipid membranes and be absorbed) and mostly ionized in the basic small intestine (where it may be less easily absorbed). Understanding acid-base chemistry is essential for designing drugs that will be absorbed in the right location and at the right rate.
| Substance | Type | Approximate pH | Key Chemical |
|---|---|---|---|
| Vinegar | Weak acid | 2.5–3.5 | Acetic acid (CH₃COOH) |
| Lemon juice | Weak acid | 2–3 | Citric acid |
| Baking soda | Weak base | 8.3 | Sodium bicarbonate (NaHCO₃) |
| Household ammonia cleaner | Weak base | 11–12 | Ammonia (NH₃) |
| Bleach | Base | 12–13 | Sodium hypochlorite (NaOCl) |
| Drain cleaner | Strong base | 13–14 | Sodium hydroxide (NaOH) |
Never mix bleach with ammonia-based cleaners or acidic cleaners. This produces dangerous chlorine gas and other toxic compounds. Chemical safety knowledge is not merely academic — it saves lives.
pH = −log[H⁺]
[H⁺] = 10⁻ᵖᴴ
[H⁺] × [OH⁻] = 1.0 × 10⁻¹⁴ (at 25°C)
[OH⁻] = (1.0 × 10⁻¹⁴) ÷ [H⁺]
All concentrations in mol/L (molarity, M). Brackets [] indicate concentration.
M_acid × V_acid = M_base × V_base
moles = molarity × volume (in liters)
Rearrange to solve for any unknown. Works when the acid-to-base mole ratio is 1:1.
pH = −log(1 × 10⁻⁴) = −(−4) = 4[H⁺] = 10⁻⁹ = 1 × 10⁻⁹ mol/L
[OH⁻] = (1 × 10⁻¹⁴) ÷ (1 × 10⁻⁹) = 1 × 10⁻⁵ mol/L
CH₃COOH(aq) + NaOH(aq) → NaCH₃COO(aq) + H₂O(l)
M_NaOH × 0.02000 L = 0.150 mol/L × 0.01500 L
M_NaOH = (0.150 × 0.01500) ÷ 0.02000
M_NaOH = 0.002250 ÷ 0.02000 = 0.113 mol/L
Answer these questions to test understanding before the practice problems.
Problems 1–4: Calculate pH from [H⁺]. Problems 5–8: Calculate [H⁺] from pH.
Write balanced neutralization equations and name the salt produced.
A student tests unknown solutions using indicators. Use the table to answer questions 33–36.
| Solution | Litmus | Phenolphthalein | Methyl Orange | Bromothymol Blue |
|---|---|---|---|---|
| A | Red | Colorless | Red | Yellow |
| B | Blue | Pink | Yellow | Blue |
| C | Purple | Colorless | Yellow | Green |
| D | Red | Colorless | Orange | Yellow |
Indicator reference — Litmus: red below pH 7, blue above pH 7. Phenolphthalein: colorless below pH 8, pink above pH 10. Methyl orange: red below pH 3.1, orange 3.1–4.4, yellow above 4.4. Bromothymol blue: yellow below pH 6, green near pH 7, blue above pH 7.6.
Each problem contains a mistake in reasoning, calculation, or concept. Identify and correct each error.
What is the error? What distinguishes strength from sensation intensity?
Who is correct? Explain the relationship between [H⁺] and pH scale direction.
Identify the error. What are the correct products of HCl + NaOH?
Identify the error and write the correct equation.
Is this correct? Explain when the equivalence point pH is above or below 7.
What is the flaw? What does buffer capacity mean?
Strong and weak describe the degree of ionization, not the amount of acid present. HCl is a strong acid whether it is 0.001 mol/L or 10 mol/L — it still ionizes 100% in water. Concentration tells you how much acid you have; strength tells you how completely it ionizes. A 0.001 mol/L HCl solution is dilute but still strong. A 10 mol/L acetic acid solution is concentrated but still weak.
Neutralization is a reaction type — acid reacts with base. The final pH depends on the relative strengths of the acid and base. Weak acid + strong base: the resulting salt hydrolyzes in water, giving pH above 7. Strong acid + weak base: pH below 7. Only strong acid + strong base gives pH exactly 7.
The relationship is inverse. Higher [H⁺] gives lower pH. The negative sign in pH = −log[H⁺] causes the inversion. Acidic solutions have pH below 7; basic solutions have pH above 7. A solution at pH 2 has more H⁺ than a solution at pH 4, not less.
Buffers resist pH changes — they do not prevent them entirely. Any buffer has a limited capacity. Adding enough acid or base will eventually exhaust the buffer components and cause significant pH changes. Buffers operate within a pH range close to the pKa of the weak acid component; outside this range, their effectiveness drops sharply.
The equivalence point is the theoretical point where moles of acid equal moles of base. The endpoint is where the indicator changes color. A well-chosen indicator makes these nearly identical, but they are not exactly the same — the indicator introduces a small discrepancy. This is one reason indicator choice matters and why pH meters give more precise results than indicators.
These questions draw on earlier chapters as well as Chapter 23.
In Chapter 23, you studied one of the most important classes of chemical reactions — the transfer of protons between acids and bases. Chapter 24 introduces another great class of reactions built on particle transfer: oxidation-reduction reactions, where electrons move from one substance to another. You will learn what oxidation numbers mean, how to identify which substance loses electrons and which gains them, and how this electron transfer powers batteries, enables electroplating, drives corrosion, and forms the basis for fuel cells and energy technology. The tools developed in this chapter — recognizing ions, writing reaction equations, understanding equilibrium — carry directly into Chapter 24.