Acids, Bases & Titration — Print · Physical Science: Chemistry, Grade 11

Acids & Bases

Learning goals

• I can define acids and bases using the Arrhenius and Brønsted-Lowry models.
• I can distinguish strong from weak acids and bases, and separate 'strong' from 'concentrated.'

Readings for this unit

• A pH number is the beginning of an explanation

Acids and bases are two of the most important families of compounds in all of chemistry, and people had sorted substances into 'sour' and 'bitter' groups for centuries before anyone understood why. The first modern explanation came from Svante Arrhenius, who defined acids and bases by the ions they release in water. An Arrhenius acid produces hydrogen ions, H⁺ (which immediately grab onto water molecules to form hydronium, H₃O⁺), when dissolved in water. An Arrhenius base produces hydroxide ions, OH⁻. Lemon juice, vinegar, and stomach acid fit the first group; lye, ammonia solution, and dissolved baking soda fit the second.

Arrhenius's picture is powerful but a little narrow — it describes aqueous hydronium/hydroxide behavior; a base can increase OH⁻ through reaction with water without initially containing that ion. A broader definition came from Brønsted and Lowry, who shifted the focus from ions to protons. (H⁺ denotes a proton, but in water the proton is hydrated rather than remaining an isolated bare particle.) A Brønsted-Lowry acid is any proton donor, and a Brønsted-Lowry base is any proton acceptor. When HCl dissolves, it *donates* a proton to a water molecule: HCl + H₂O → H₃O⁺ + Cl⁻. Here HCl is the acid and water is the base — a general way of thinking that covers reactions Arrhenius's model can't.

One fascinating consequence is that water can play *either* role, depending on its partner. With hydrochloric acid, water accepts a proton and acts as a base; but with ammonia (NH₃ + H₂O → NH₄⁺ + OH⁻), water donates a proton and acts as an acid. A substance that can behave as either an acid or a base is called amphoteric, and water is the most famous example — a quiet hint that being an acid or a base is less a fixed label than a description of how two substances behave toward each other.

Acids and bases also announce themselves through characteristic properties. Acids taste sour (think of lemons and vinegar), react with many metals to release hydrogen gas, and turn blue litmus paper red. Bases taste bitter, feel slippery or soapy, and turn red litmus paper blue. Both kinds conduct electricity in solution, because both release ions that can carry charge. One firm safety rule, though: you never taste or touch laboratory chemicals to identify them — strong acids and strong bases are both corrosive enough to cause serious burns.

Strong and weak describe proton-transfer equilibrium, separate from concentration. A strong monoprotic acid such as HCl is essentially fully ionized in the usual dilute aqueous account; a weak acid has an equilibrium involving appreciable un-ionized species under appropriate conditions. Polyprotic acids require separate steps: sulfuric acid’s first ionization is strong, while its second step must be treated with its own equilibrium. A soluble strong hydroxide base such as NaOH dissociates into ions, while NH₃ produces hydroxide through a weak-base reaction with water.

Use proton-transfer and ionization models to interpret acids and bases rather than product labels alone. Strength, concentration, temperature, and the reacting species determine behavior. A weak acid can still be hazardous, and a stronger acid is not automatically more concentrated. Work with supplied equations and known values for quantitative comparisons.

Model check: Acid strength and concentration are independent; neither alone fixes pH. It accepts a proton from water, forming NH₄⁺ and OH⁻; Brønsted bases need not contain hydroxide initially.

Worked examples

Avoid these mistakes

• Acid strength and concentration are independent; neither alone fixes pH.

Words to know

Arrhenius acid

A substance that releases hydrogen ions (H⁺, which form H₃O⁺) when dissolved in water.

Arrhenius base

A substance that releases hydroxide ions (OH⁻) when dissolved in water.

Brønsted-Lowry acid

A proton (H⁺) donor in a chemical reaction.

Brønsted-Lowry base

A proton (H⁺) acceptor in a chemical reaction.

strong acid

An acid that ionizes completely in water, releasing nearly all of its hydrogen ions.

weak acid

An acid that ionizes only partially in water, so most of its molecules stay intact at any moment.

A third definition: Lewis acids and bases

Arrhenius focused on ions in water and Brønsted-Lowry on proton transfer, but the broadest view came from Gilbert Lewis, who shifted attention to electrons. A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor. Because a proton (H⁺) is an electron-pair acceptor, every Brønsted acid is also a Lewis acid — but the Lewis definition reaches further, covering reactions with no hydrogen at all, such as metal ions bonding to ammonia. Each definition is a wider lens: Arrhenius fits inside Brønsted-Lowry, which fits inside Lewis, and chemists pick whichever level of generality a problem needs.

Conjugate acid-base pairs: the partners in every proton transfer

Every Brønsted-Lowry reaction actually involves two acid-base pairs. When an acid donates a proton, what's left behind is a conjugate base; when a base accepts a proton, it becomes a conjugate acid. In HCl + H₂O → H₃O⁺ + Cl⁻, HCl and Cl⁻ are one conjugate pair, and H₂O and H₃O⁺ are the other. A useful rule follows: the stronger an acid, the weaker its conjugate base, because a substance that lets go of its proton easily has little tendency to grab it back. Recognizing conjugate pairs turns a confusing equation into a simple story of a proton changing hands.

Strong vs. weak, concentrated vs. dilute: untangling four words

These four words trip up nearly every chemistry student because two of them describe ionization and two describe amount. Strong and weak describe how completely an acid or base breaks apart into ions: a strong acid ionizes almost 100%, a weak acid only a few percent. Concentrated and dilute describe how much acid is dissolved per liter — the molarity. The two axes are independent, so all four combinations exist: dilute strong acid, concentrated strong acid, dilute weak acid, and concentrated weak acid. Vinegar is a dilute weak acid; the acid in a car battery is a concentrated strong acid. Keeping the axes separate is the key to the whole topic.

How your tongue tastes an acid

The sour taste of acids isn't a coincidence — it's your biology detecting hydrogen ions directly. Specialized taste receptor cells on your tongue respond to the H⁺ that acids release, which is why lemon juice (citric acid) and vinegar (acetic acid) taste sharp and sour. Evolution likely favored sour detection to warn against spoiled or unripe food, both of which tend to be acidic. Bitterness, the hallmark of many bases and alkaloids, is detected by a completely different family of receptors and probably evolved to flag potential toxins. Your sense of taste is, quite literally, a portable acid-base indicator — though, unlike litmus, one you should never use on lab chemicals.

The acids and bases in your home — and one dangerous mix

Household chemistry is full of acids and bases: vinegar and lemon juice are mild acids, while baking soda, ammonia, oven cleaner, and bleach are bases. Mixing an acid and a base intentionally can be handy — baking soda neutralizes an acid spill or a bee sting. But some combinations are genuinely dangerous. Mixing bleach with ammonia (or with acidic cleaners) releases toxic chloramine and chlorine gases that can severely damage the lungs. This is why cleaning-product labels warn 'do not mix,' and it's a real-world reminder that acid-base reactions are powerful chemistry, not just textbook color changes.

In the real world

Your stomach makes strong hydrochloric acid to digest food, while the weak citric and acetic acids in fruit and vinegar are mild enough to eat every day.

Exit check

Explain why NH₃ is a base without containing OH in its formula.

pH & Neutralization

Learning goals

• I can use the pH scale to classify a solution and connect its pH to the concentration of hydrogen ions.
• I can explain a neutralization reaction and predict the products when an acid reacts with a base.

Readings for this unit

• A pH number is the beginning of an explanation

pH describes hydrogen-ion activity. For dilute aqueous solutions near 25°C, a familiar chart commonly spans 0–14, with neutral water near pH 7, acidic solutions below neutral, and basic solutions above it. That range is not an absolute bound, and neutral pH changes with temperature because water’s ionization constant changes. Compare pH under stated conditions rather than treating exactly 7 as universal. A pH value alone does not identify a chemical or establish safety.

The exact definition uses dimensionless hydrogen-ion activity: pH = −log₁₀(aH⁺). In the introductory dilute-solution approximation, activity is replaced by the numerical molar concentration relative to the 1 M standard, giving the familiar pH ≈ −log₁₀[H⁺]. A concentration of 1 × 10⁻³ M then gives about pH 3. Concentrated or strongly interacting solutions may depart from this approximation, so state the model before using concentration in a logarithm.

A difference of one pH unit corresponds to a tenfold change in hydrogen-ion activity. A solution at pH 3 has 100 times the activity of one at pH 5; the dilute model gives the same concentration ratio approximately. This comparison does not say that pH 3 contains 100 times as much total acid or requires 100 times as much base to titrate: volume, acid identity, and dissociation matter. Keep the logarithmic activity comparison distinct from total reacting amounts.

Indicators change color over characteristic pH intervals rather than at one exact value. Litmus distinguishes broad acidic and basic behavior, while an appropriate indicator can help locate a titration endpoint. A pH meter needs calibration, suitable electrodes, temperature control or compensation, and correct handling. A display with two decimal places does not guarantee matching accuracy. Colored samples, contamination, and calibration problems can affect results, so record uncertainty and method.

Neutralization is an acid-base reaction. For a strong acid reacting with hydroxide, the net ionic equation is H⁺(aq) + OH⁻(aq) → H₂O(l). Spectator ions remain in aqueous solution; they do not automatically assemble into solid salt. Calculate reacting moles to identify excess acid or base. At equivalence, the pH is near 7 for an ideal strong-acid/strong-base system at 25°C, but weak-acid or weak-base salts can hydrolyze and give a different equivalence pH.

A titration uses a measured volume of a standard solution and reaction stoichiometry to determine an unknown amount or concentration. The equivalence point is the stoichiometric condition; an indicator endpoint is the observed color change used to approximate it. Select an indicator whose transition range fits the equivalence region, and account for overshoot or volume uncertainty. Equal acid and base volumes are not required. Use supplied titration records to calculate concentration and judge whether the evidence supports the reported precision.

Model check: pH differences are logarithmic; neutrality near 7 assumes about 25°C. pH3 has 10^(5−3)=100 times the hydronium concentration in the simple concentration model.

pH

pH = −log[H⁺]

[H⁺] is the hydrogen ion concentration in mol/L. Because it is a base-10 logarithm, each whole pH unit is a tenfold change in acidity; pH < 7 is acidic, 7 is neutral, and pH > 7 is basic.

Worked examples

Avoid these mistakes

• pH differences are logarithmic; neutrality near 7 assumes about 25°C.

Words to know

pH scale

A scale from 0 to 14 that measures how acidic or basic a solution is; below 7 is acidic, exactly 7 is neutral, and above 7 is basic.

pH

Negative base-ten logarithm of hydronium activity, approximated using suitably normalized concentration in dilute solutions; 0–14 is not an absolute limit.

neutralization

A reaction between an acid and a base that produces water and a salt, driving the pH of the mixture toward 7.

salt

An ionic compound formed from the positive ion of a base and the negative ion of an acid during neutralization — table salt (NaCl) is just one example.

indicator

A substance, such as litmus or phenolphthalein, that changes color depending on the pH of the solution it is added to.

titration

A lab procedure that slowly adds a solution of known concentration to another until the reaction is exactly complete, used to find an unknown concentration.

pH is a logarithm: why one number hides a factor of ten

The definition pH = −log[H⁺] makes pH a logarithmic scale, and this is the idea students most often miss. Because it's a base-10 log, each whole pH unit corresponds to a tenfold change in hydrogen ion concentration. A solution of pH 3 has ten times more H⁺ than pH 4 and ten thousand times more than pH 7. This 'compression' of huge ranges into small numbers is exactly why scientists use logarithms elsewhere too — the Richter scale for earthquakes and the decibel scale for sound work the same way, so a magnitude-6 quake releases about ten times the ground motion of a magnitude-5.

Søren Sørensen and the brewery where pH was born

The pH scale was invented in 1909 by the Danish biochemist Søren Sørensen, who was not studying stomachs or oceans but beer. Working at the Carlsberg brewery's laboratory, he needed a compact way to describe the acidity that affected the enzymes in brewing, rather than writing out tiny hydrogen-ion concentrations in scientific notation. His solution — take the negative logarithm and call it 'pH,' loosely the 'power of hydrogen' — was so convenient that it spread far beyond brewing. More than a century later, a shorthand born in a brewery is standard vocabulary in medicine, agriculture, oceanography, and every chemistry classroom.

Buffers: how your blood defends its pH

Human blood must stay in an astonishingly narrow window, pH 7.35 to 7.45, or enzymes stop working and cells are damaged. It manages this with a buffer — a mixture of a weak acid and its conjugate base that neutralizes small additions of acid or base, resisting pH change. Your blood's main buffer is the carbonic-acid/bicarbonate system, and it constantly mops up the acids your metabolism produces. When buffering fails, the conditions acidosis (blood too acidic) or alkalosis (too basic) are medical emergencies. Every breath you take, by controlling carbon dioxide, is quietly adjusting this acid-base balance in real time.

Titration: finding a hidden concentration drop by drop

Titration turns neutralization into a precision measuring tool. A solution of known concentration (the titrant) is slowly dripped from a burette into a measured volume of the unknown solution, with an indicator like phenolphthalein added. As long as excess acid remains, the indicator stays one color; the instant just enough base has been added to neutralize it — the equivalence point — a single drop flips the color. From the volume and concentration of titrant used, the unknown concentration is calculated. Titration is a workhorse of analytical chemistry, used to check everything from the acidity of food and wine to the purity of medicines.

Ocean acidification: acids and bases at planetary scale

The surface ocean has shifted from an average pH near 8.2 before the Industrial Revolution to about 8.1 today, as seawater absorbs extra atmospheric carbon dioxide and forms carbonic acid. A 0.1 drop sounds trivial, but because pH is logarithmic it represents roughly a 30% rise in hydrogen-ion concentration in only two centuries. More acidic water makes it harder for corals, oysters, and other shell-builders to form their calcium carbonate skeletons, threatening entire marine food webs. It is a sobering, planet-sized example of a single number — pH — linking chemistry, the carbon cycle, and the future of life in the sea.

In the real world

A chalky antacid tablet is a mild base that neutralizes the excess hydrochloric acid in your stomach, easing heartburn by pushing the pH back up toward neutral.

Exit check

Compare pH3 andpH5 hydronium concentrations.

Acid Strength, Conjugate Pairs & Ka

Learning goals

• I can identify conjugate pairs and distinguish strength from concentration.
• I can estimate weak-acid ionization and check a simplifying assumption.

Readings for this unit

• A pH number is the beginning of an explanation

A Brønsted–Lowry acid donates a proton and a base accepts one. After donating, an acid becomes its conjugate base; after accepting, a base becomes its conjugate acid. Members of a conjugate pair differ by one proton. Water can play either role depending on its reaction partner.

Strength describes the extent of ionization in water. Concentration describes amount per solution volume. A dilute strong acid can have lower hydrogen-ion concentration than a concentrated weak acid, or higher, depending on numbers. Neither label alone determines pH; use actual concentration and ionization.

For HA + H₂O ⇌ H₃O⁺ + A⁻, Ka=[H₃O⁺][A⁻]/[HA] in the introductory concentration approximation. Water is omitted because it is the solvent. Larger Ka generally means more extensive ionization at comparable concentration and temperature. pKa=−log 10 Ka provides a convenient scale.

For a weak monoprotic acid initially at concentration C, let x be the concentration ionized. Then Ka≈x²/(C−x), assuming no important added ions and negligible initial hydronium. If x is small relative to C, x≈√(KaC). Check the approximation by calculating x/C; a rough 5% criterion is common.

Strong-acid pH calculations often use [H₃O⁺]≈C for a fully ionized monoprotic acid. This approximation needs care at extremely low concentrations where water’s own ionization matters. Neutral pH is 7 at about 25°C, not a temperature-independent law.

Use a reaction and numbers to explain an acid claim. Do not treat a weak acid as harmless: concentration, exposure, and other chemical properties matter. These calculations are for supplied data, with household observations limited to safe, labeled food materials and no tasting or mixing cleaners.

Worked examples

Avoid these mistakes

• Do not infer actual pH from strength alone.

Words to know

conjugate pair

Species that differ by one proton.

Ka

Acid-ionization equilibrium constant for a specified temperature.

pKa

Negative base-ten logarithm of Ka.

In the real world

Use the chapter investigation to connect this unit to measured evidence and a practical decision.

Exit check

What distinguishes a conjugate base from a different unrelated base?

Titration, Buffers & Neutralization Curves

Learning goals

• I can distinguish endpoint from stoichiometric equivalence.
• I can calculate a buffer pH and explain finite buffer capacity.

Readings for this unit

• A pH number is the beginning of an explanation

Titration adds a known-concentration solution to a measured unknown to infer its amount through a reaction. Equivalence means stoichiometric amounts have reacted; endpoint is the observed indicator or instrument signal. A useful endpoint approximates equivalence but is not identical by definition.

For monoprotic strong acid and strong base, nH⁺=nOH⁻ at equivalence. Calculate moles from M×V and account for coefficients for other reactions. The familiar MAVA=MBVB equality applies directly only to a 1:1 stoichiometric case, not every acid-base titration.

A buffer contains appreciable amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid. Added small amounts of acid are consumed mainly by the conjugate base; added base is consumed mainly by the weak acid. Buffer capacity is finite; adding enough acid or base exhausts it.

The Henderson–Hasselbalch approximation is pH≈pKa+log 10([A⁻]/[HA]) for an appropriate buffer with both species present and activities approximated by concentrations. Equal acid and base amounts give pH≈pKa. At half-equivalence for a simple weak-acid/strong-base titration, the same relation often holds.

Equivalence pH depends on conjugate-ion hydrolysis. A strong-acid/strong-base equivalence is near neutral at 25°C; a weak-acid/strong-base equivalence is usually basic. An indicator must change over the steep relevant section of the curve; choosing a fixed pH 7 indicator for every titration is not justified.

Work from supplied titration data at home. Plot pH against added volume and distinguish the curve’s evidence from the calculation model. Record aliquot volume, titrant concentration, reading resolution, repeats, and possible overshoot before reporting an inferred concentration.

Worked examples

Avoid these mistakes

• An equivalence point is not always pH 7; a buffer cannot absorb unlimited additions.

Words to know

endpoint

Observed signal used to stop a titration.

equivalence point

Stoichiometric completion for the selected titration reaction.

buffer capacity

Amount of added acid/base a buffer can accommodate effectively.

In the real world

Use the chapter investigation to connect this unit to measured evidence and a practical decision.

Exit check

What evidence should accompany a titration concentration result?