pH is minus the log of the hydrogen-ion concentration in mol/L, so each unit is a factor of ten: pH 3 is a hundred times more acidic than pH 5. A 0.1 mol/L solution of a strong acid like hydrochloric acid is pH 1.00.
A strong acid splits completely into ions; a weak one only a little. At the same 0.1 mol/L, hydrochloric acid is pH 1.00 and ethanoic acid pH 2.88, with 1.3 percent of its molecules ionised.
Run a base into an acid and the pH rises slowly, jumps at the equivalence point and levels off. The pH at the jump is 7 for a strong acid with a strong base, 8.73 for ethanoic acid with sodium hydroxide, and 5.28 for hydrochloric acid with ammonia.
A weak acid with its salt soaks up added acid or base. Ten millimoles of hydrochloric acid in a litre of plain water takes the pH from 7 to 2.00; in a litre of 0.1 mol/L ethanoate buffer it moves from 4.76 to 4.58.
Hydrochloric acid ionises completely, so [H+] is the concentration, 0.1 mol/L, and the pH is its negative log. Ten times more dilute is one unit higher: 0.01 mol/L is 2.00, 0.001 mol/L 3.00, 0.0001 mol/L 4.00.
A weak acid HA has Ka = [H+][A−]/[HA]. Very little ionises, so [H+] = √(Kac) and pH = ½(pKa − log c). For ethanoic acid, pKa 4.76, at 0.1 mol/L that is ½(4.76 + 1) = 2.88, and the exact solution is 2.88, with 1.3 percent of the molecules ionised. Ten times more dilute it is 3.39: a weak acid moves only about half a unit for each tenfold dilution, because more of it ionises as it thins.
A hundred-millionth of a mole per litre of hydrochloric acid is not pH 8. Water supplies its own 10−7 mol/L of hydrogen ions, so the acid only adds to it: the pH of 10−8 mol/L is 6.98, and no acid, however dilute, makes a solution alkaline. The same applies to bases from the other side.
Sodium hydroxide is a strong base: 0.1 mol/L is pH 13.00, and 1 mol/L is pH 14.00. Ammonia is a weak base, taking a hydrogen ion from water to make ammonium (pKa 9.25, so pKb 4.75); 0.1 mol/L is pH 11.12 with 1.3 percent of the molecules converted. At 25 °C, pH + pOH = 14.
With 0.1 mol/L on both sides, the pH at the equivalence point is 7.00 for hydrochloric acid with sodium hydroxide, 8.73 for ethanoic acid with sodium hydroxide (the ethanoate that is left is a weak base) and 5.28 for hydrochloric acid with ammonia (the ammonium left is a weak acid). Ethanoic acid with ammonia gets to 7.00 with a jump of only 0.4 pH units: with two weak partners there is no sharp end.
An indicator is a weak acid whose two forms have different colours, and it changes over about two pH units centred on its own pKa. It must change inside the steep part of the curve. For ethanoic acid with sodium hydroxide that part runs from pH 7.1 to 10.4 in the half percent of volume either side of the end point (0.1 mL in 25), so phenolphthalein (8.3 to 10.0) works and methyl orange (3.1 to 4.4) would turn long before. For hydrochloric acid with ammonia the steep part is 3.6 to 6.9, and it is the other way round.
Half-way to the equivalence point of a weak acid, half the acid has been turned into its salt, so [HA] = [A−] and pH = pKa: for ethanoic acid 4.76, which is how a pKa is measured. The region between a tenth and nine tenths of the way is a buffer, flat because the weak acid and its salt are both present.
The Henderson–Hasselbalch equation, pH = pKa + log([A−]/[HA]), gives the pH of a buffer from the ratio of its two forms. In 0.1 mol/L of ethanoic acid and ethanoate in equal parts it is 4.76, and 10 mmol of hydrochloric acid in a litre moves it only to 4.58, against a fall from 7 to 2.00 in water. The capacity, the strong acid or base needed for one pH unit, is greatest at pH = pKa and equals 2.3 c/4, which is 57.6 mmol per litre for this one.
Phosphoric acid gives up its three protons at pKa 2.16, 7.21 and 12.32, so each form is the main one between two of them. At pH 7.4 it is mostly hydrogenphosphate (61 percent) and dihydrogenphosphate (39 percent). Carbon dioxide dissolved in rain, 0.042 percent of the air, makes carbonic acid (pKa1 6.35) and a pH of 5.6; a 0.1 mol/L solution of sodium hydrogencarbonate (baking soda) sits between its two pKa values at pH 8.3.
Water ionises a little: Kw = [H+][OH−] is 10−14 at 25 °C and grows with temperature, so pure water is neutral at pH 7.0 only at 25 °C. At 100 °C, pKw is 12.25 and neutral water has a pH of about 6.1; at body temperature, 37 °C, about 6.8; near freezing, about 7.5. Water at 100 °C is not acidic: it has equal hydrogen and hydroxide ions.
Every solution here is ideal: the ions do not interact (no activity coefficients), so a measured pH above about 0.1 mol/L differs from the figure by up to a few tenths. The dissociation constants are the tabulated values at 25 °C and in dilute solution, which differ by about 0.1 between sources, and they are not corrected for temperature; only the pure water is. Calcium hydroxide is capped at what dissolves (0.02 mol/L), no solid or gas equilibrium is modelled, and carbonic acid stands for dissolved carbon dioxide and the acid together.
The concentration of hydrogen ions in a solution, as pH = −log10[H+] with the concentration in mol/L. Below 7 is acidic, 7 neutral, above 7 alkaline, at 25 °C. A pH meter measures the activity of the ions, which is close to the concentration in dilute solutions.
Because hydrogen-ion concentrations in everyday liquids span a factor of a hundred trillion, from strong acids to strong bases. A log scale squeezes that into 0 to 14, and each step is a factor of ten: pH 3 is ten times more acidic than pH 4 and a thousand times more than pH 6.
Strength is how completely each molecule ionises; concentration is how much is dissolved. A strong acid at 0.001 mol/L is dilute but fully ionised, pH 3.00; ethanoic acid at 10 mol/L is concentrated but weak, and its pH is 1.9, still less acidic than 0.1 mol/L hydrochloric acid at 1.00.
The negative log of the acid dissociation constant Ka. The smaller the pKa, the stronger the acid: nitric acid is about −1.4, ethanoic acid 4.76, hydrocyanic acid 9.21. At a pH equal to the pKa the acid and its conjugate base are present in equal amounts.
Yes. A concentrated strong acid has a hydrogen-ion concentration above 1 mol/L, so a negative log: ideally 10 mol/L of hydrochloric acid would be pH −1 (a real one reads lower, because the ions are not ideal), and 1 mol/L sodium hydroxide is pH 14.00.
Because neutral means equal hydrogen and hydroxide ions, and the amount of each depends on temperature: pKw falls from 14.0 at 25 °C to about 12.25 at 100 °C, so neutral water is pH about 6.1 there. Hot water is not acidic; the pH scale has shifted.
A solution that resists changes of pH, made from a weak acid and its conjugate base (or a weak base and its conjugate acid) in similar amounts. Added acid is soaked up by the base, added base by the acid. Blood is buffered by carbonic acid and hydrogencarbonate, which is how it stays within about 0.05 of pH 7.4.
By the equivalence point: pick the one whose colour change, about its pKa ± 1, lies in the steep part of the titration curve. Phenolphthalein for a weak acid with a strong base, methyl orange for a strong acid with a weak base, nearly any for strong with strong, and none for weak with weak.
The equivalence point is where exactly equal amounts of acid and base have been mixed; the end point is where the indicator changes colour. A well-chosen indicator makes them differ by less than a drop, and a poor one by millilitres.
Because at the equivalence point the acid has been turned into its salt, and the anion of a weak acid is itself a weak base, reacting with water to make hydroxide. 0.1 mol/L ethanoic acid with 0.1 mol/L sodium hydroxide ends at pH 8.73 for the same reason that a salt of a strong acid with a weak base ends below 7.
This page has changed in 3 archived releases, but each of those was a site-wide pass, so none is attributable to this tool on its own and none is listed here. That is not a claim that the tool never changed — a release that reworked many pages at once may well have altered this one too. The changelog has them.