Which pH values are acids and which are bases
The pH scale runs from 0 to 14, and the number answers one question: is this water-based solution an acid, a base, or neither?
- Acids are low. pH 0 up to just below 7. Battery acid, lemon juice, vinegar, black coffee.
- Neutral is exactly 7. Pure water, and only pure water sits there.
- Bases are high. Just above 7 up to 14. Baking soda, hand soap, ammonia, bleach. A base dissolved in water is also called an alkali.
So “are acids high or low pH” has a one-word answer: low. It reads backwards from what most people expect, and the next section explains why.
You will see this same scale called the acid-base scale, the acidity scale, the alkalinity scale, or just pH levels. Those are all names for the one scale below.
| pH range | What it is | Acid or base | Ions |
|---|---|---|---|
| 0 up to 3 | Strongly acidic | acid | far more H⁺ than OH⁻ |
| 3 up to 6 | Acidic | acid | more H⁺ than OH⁻ |
| 6 up to 7 | Weakly acidic | acid | slightly more H⁺ than OH⁻ |
| exactly 7 | Neutral | neither | [H⁺] = [OH⁻] |
| above 7 up to 8 | Weakly basic | base (alkali) | slightly more OH⁻ than H⁺ |
| above 8 up to 11 | Basic (alkaline) | base (alkali) | more OH⁻ than H⁺ |
| above 11 up to 14 | Strongly basic | base (alkali) | far more OH⁻ than H⁺ |
The boundaries are worth being exact about. Neutral means pH 7 and nothing else. A solution at pH 6.9 is acidic, not neutral, and pH 7.1 is basic. Human blood at pH 7.4 is slightly basic, even though people often call it neutral.
What every number on the scale means
Each whole number is a fixed hydrogen-ion concentration, and the two ion columns below always multiply to the same thing.
| pH | [H⁺] mol/L | [OH⁻] mol/L | Acid or base | Close to |
|---|---|---|---|---|
| 0 | 1 × 10⁰ | 1 × 10⁻¹⁴ | strongly acidic | 1 M hydrochloric acid |
| 1 | 1 × 10⁻¹ | 1 × 10⁻¹³ | strongly acidic | battery acid |
| 2 | 1 × 10⁻² | 1 × 10⁻¹² | strongly acidic | lemon juice |
| 3 | 1 × 10⁻³ | 1 × 10⁻¹¹ | acidic | vinegar |
| 4 | 1 × 10⁻⁴ | 1 × 10⁻¹⁰ | acidic | orange juice |
| 5 | 1 × 10⁻⁵ | 1 × 10⁻⁹ | acidic | black coffee |
| 6 | 1 × 10⁻⁶ | 1 × 10⁻⁸ | weakly acidic | saliva |
| 7 | 1 × 10⁻⁷ | 1 × 10⁻⁷ | neutral | pure water |
| 8 | 1 × 10⁻⁸ | 1 × 10⁻⁶ | weakly basic | seawater |
| 9 | 1 × 10⁻⁹ | 1 × 10⁻⁵ | basic (alkaline) | hand soap |
| 10 | 1 × 10⁻¹⁰ | 1 × 10⁻⁴ | basic (alkaline) | hand soap |
| 11 | 1 × 10⁻¹¹ | 1 × 10⁻³ | basic (alkaline) | household ammonia |
| 12 | 1 × 10⁻¹² | 1 × 10⁻² | strongly basic | limewater |
| 13 | 1 × 10⁻¹³ | 1 × 10⁻¹ | strongly basic | household bleach |
| 14 | 1 × 10⁻¹⁴ | 1 × 10⁰ | strongly basic | 1 M sodium hydroxide |
Read a row across and the pattern is the whole scale in one line: as [H⁺] falls by a factor of ten, [OH⁻] rises by a factor of ten, and the two always multiply to 1 × 10⁻¹⁴.
Common pH values
| Substance | Typical pH | Published range | [H⁺] mol/L | Acid or base |
|---|---|---|---|---|
| 1 M hydrochloric acid a strong acid at 1 mol/L | 0.0 | exact | 1.0 × 10⁰ | strongly acidic |
| battery acid | 0.5 | 0 to 1 | 3.2 × 10⁻¹ | strongly acidic |
| stomach acid | 1.5 | 1.5 to 3.5 | 3.2 × 10⁻² | strongly acidic |
| lemon juice | 2.2 | 2 to 2.6 | 6.3 × 10⁻³ | strongly acidic |
| cola | 2.5 | 2.3 to 2.8 | 3.2 × 10⁻³ | strongly acidic |
| vinegar | 2.9 | 2.4 to 3.4 | 1.3 × 10⁻³ | strongly acidic |
| apple juice | 3.3 | 3.3 to 4 | 5.0 × 10⁻⁴ | acidic |
| orange juice | 3.7 | 3.3 to 4.2 | 2.0 × 10⁻⁴ | acidic |
| tomato juice | 4.3 | 4.1 to 4.6 | 5.0 × 10⁻⁵ | acidic |
| black coffee | 5.0 | 4.8 to 5.1 | 1.0 × 10⁻⁵ | acidic |
| clean rainwater acidic from dissolved CO₂, with no pollution at all | 5.6 | exact | 2.5 × 10⁻⁶ | acidic |
| saliva | 6.2 | 6.2 to 7.6 | 6.3 × 10⁻⁷ | weakly acidic |
| milk | 6.7 | 6.5 to 6.8 | 2.0 × 10⁻⁷ | weakly acidic |
| pure water neutral at 25 °C | 7.0 | exact | 1.0 × 10⁻⁷ | neutral |
| human blood | 7.4 | 7.35 to 7.45 | 4.0 × 10⁻⁸ | weakly basic |
| seawater | 8.1 | 7.5 to 8.4 | 7.9 × 10⁻⁹ | basic (alkaline) |
| baking soda solution | 8.3 | exact | 5.0 × 10⁻⁹ | basic (alkaline) |
| hand soap | 9.5 | 9 to 10 | 3.2 × 10⁻¹⁰ | basic (alkaline) |
| milk of magnesia | 10.5 | exact | 3.2 × 10⁻¹¹ | basic (alkaline) |
| household ammonia | 11.3 | 11 to 11.5 | 5.0 × 10⁻¹² | strongly basic |
| limewater saturated calcium hydroxide | 12.4 | exact | 4.0 × 10⁻¹³ | strongly basic |
| household bleach | 12.6 | 11 to 13 | 2.5 × 10⁻¹³ | strongly basic |
| oven and drain cleaner | 13.5 | 13 to 14 | 3.2 × 10⁻¹⁴ | strongly basic |
| 1 M sodium hydroxide a strong base at 1 mol/L | 14.0 | exact | 1.0 × 10⁻¹⁴ | strongly basic |
These are typical published values, not measurements of one particular sample. Real pH moves with concentration, temperature and whatever else is dissolved, so the range column shows the published spread where there is one. Two rows are exact rather than typical: hydrochloric acid at 1 mol/L gives pH 0 and sodium hydroxide at 1 mol/L gives pH 14, because that is what the definition works out to at those concentrations.
That last point catches people out. A strong acid is not automatically pH 0. Hydrochloric acid is as strong as acids get, but its pH depends entirely on how concentrated it is: 1 mol/L is pH 0, 0.1 mol/L is pH 1, and 0.001 mol/L is pH 3. Strength and concentration are two different things.
Why one step means ten times
pH is defined as:
pH = −log₁₀[H⁺], which rearranges to [H⁺] = 10⁻ᵖᴴ mol/L
The logarithm is what makes the scale readable. Hydrogen-ion concentrations in ordinary solutions run from about 1 mol/L down to 0.00000000000001 mol/L, a span of fourteen powers of ten. Taking the log compresses that into the numbers 0 through 14.
It also means the scale is not linear. Every whole step is a ten-fold change in H⁺:
- pH 4 has 10 times more H⁺ than pH 5
- pH 4 has 100 times more H⁺ than pH 6
- pH 4 has 1,000 times more H⁺ than pH 7
That is why a lake dropping from pH 6 to pH 5 is a serious change and not a small one, and why the minus sign matters: more H⁺ gives a smaller pH number.
H⁺ and OH⁻ move together
Water is never only water. A tiny fraction of it splits into ions, and at 25 °C the product of the two concentrations is fixed:
[H⁺] × [OH⁻] = 1.0 × 10⁻¹⁴ (this constant is called Kw)
Take the negative log of both sides and it becomes the shortcut most students actually use:
pH + pOH = 14
So a solution at pH 3 has pOH 11, and a solution at pOH 2 has pH 12. The calculator above will start from any of the four quantities, because a problem is as likely to hand you a concentration and ask for the pH as the other way round.
At pH 7 the two concentrations are equal, both at 1 × 10⁻⁷ mol/L. That balance is what neutral means. If you are told a solution has a pH of 7, what you have been told is that it contains exactly as many H⁺ ions as OH⁻ ions.
One caveat worth knowing: the 14 comes from water at 25 °C. Warm the water and Kw rises, so neutral shifts slightly below 7 and the pair no longer adds to 14. At any temperature you meet in a school lab, 14 is the number to use.
Working backwards from a concentration
Given [H⁺], the pH is the negative log:
- [H⁺] = 1 × 10⁻³ mol/L gives pH 3
- [H⁺] = 4.0 × 10⁻⁸ mol/L gives pH 7.4, which is human blood
- [OH⁻] = 1 × 10⁻² mol/L gives pOH 2, so pH 12
Switch the calculator to [H⁺] or [OH⁻] and type the concentration in any
of the forms a textbook uses, including 1e-7, 0.0000001 and 1x10^-7.
What pH actually stands for
The H is the hydrogen ion. The p is a mathematical operator meaning “the negative base-10 logarithm of”, which is why you also see pOH, pKa and pKw written the same way.
pH is not an element and it is not a compound of phosphorus and hydrogen. It also has no units: it is the log of a ratio, and taking a log strips the units away. The concentration underneath it, [H⁺], is in moles per litre.
Using this with a class
Project the scale and ask students to predict the pH of a substance before dragging to check it. Then switch the calculator to [H⁺] and give them a concentration to work back from, which is the direction exam questions usually run. It is free to embed on a class site or LMS page.