pH Scale Explained: Acids, Bases and Neutrality Simplified
Your stomach digests food with a liquid acidic enough to dissolve metal. Your blood must stay within a pH range so narrow that a shift of 0.4 units in either direction is fatal. The ocean has become measurably more acidic over the last century due to dissolved CO₂. All of these facts are described using one simple number — pH.
In this guide you will learn exactly what pH means, how the scale from 0 to 14 works, why it is logarithmic, and how to calculate it yourself.
What Is pH?
pH stands for potential of hydrogen (from the German Potenz, meaning power, and H, the symbol for hydrogen). It measures the concentration of hydrogen ions (H⁺) in a solution — specifically, how many H⁺ ions are floating around per litre.
The formal definition is:
The negative logarithm does two things: it turns a very small, awkward number into a manageable one, and it flips the scale so that higher acidity gives a lower pH number. A solution with [H⁺] = 0.001 mol/L has pH = −log(0.001) = −(−3) = 3.
pH was introduced by Danish chemist Søren Peder Lauritz Sørensen in 1909 while working at the Carlsberg Laboratory in Copenhagen. He was studying the effect of ion concentration on enzymes in beer brewing and needed a simple way to express very small hydrogen ion concentrations without writing numbers like 0.000001 mol/L every time.
How the pH Scale Works — From 0 to 14
The pH scale runs from 0 to 14 under standard conditions (25°C), with three regions:
More H⁺ than OH⁻
Neutral
More OH⁻ than H⁺
This means:
- pH 2 is 10× more acidic than pH 3
- pH 2 is 100× more acidic than pH 4
- pH 2 is 10,000× more acidic than pH 6
This logarithmic nature is why a small change in pH can have enormous biological consequences. Blood pH must stay between 7.35 and 7.45 — a window of just 0.1 units — because even that tiny shift represents a measurable change in hydrogen ion concentration that disrupts enzyme function throughout the body.
The pH Scale Chart — Common Substances
| pH | [H⁺] mol/L | Substance | Classification |
|---|---|---|---|
| 0 | 1.0 | Battery acid (H₂SO₄) | Strongly acidic |
| 1 | 0.1 | Stomach acid (HCl) | Strongly acidic |
| 2 | 0.01 | Lemon juice, vinegar | Acidic |
| 3 | 0.001 | Cola drinks, orange juice | Acidic |
| 4 | 0.0001 | Tomato juice, wine | Weakly acidic |
| 5 | 0.00001 | Black coffee, acid rain | Weakly acidic |
| 6 | 0.000001 | Urine, milk | Slightly acidic |
| 7 | 0.0000001 | Pure water ⭐ | Neutral |
| 8 | 0.00000001 | Seawater, baking soda | Slightly basic |
| 9 | 0.000000001 | Toothpaste, soap | Weakly basic |
| 10 | 10⁻¹⁰ | Milk of magnesia | Weakly basic |
| 11 | 10⁻¹¹ | Ammonia solution | Basic |
| 12 | 10⁻¹² | Soapy water | Strongly basic |
| 13 | 10⁻¹³ | Bleach, oven cleaner | Strongly basic |
| 14 | 10⁻¹⁴ | Liquid drain cleaner (NaOH) | Strongly basic |
Why pH Is Logarithmic — The Chemistry Behind It
Pure water undergoes a process called autoionisation — a tiny fraction of water molecules spontaneously split into hydrogen ions (H⁺) and hydroxide ions (OH⁻):
At 25°C, the product of [H⁺] and [OH⁻] concentrations in any aqueous solution is always the same constant:
This is the water dissociation constant (Kw). In pure water, [H⁺] = [OH⁻] = 1.0×10⁻⁷ mol/L. So pH = −log(10⁻⁷) = 7. This is why neutral pH is 7 at 25°C.
When you add an acid, it releases more H⁺ ions into the solution, pushing [H⁺] above 10⁻⁷ and pH below 7. When you add a base, it releases OH⁻ ions which react with H⁺ ions and reduce their concentration, pushing pH above 7.
The relationship between pH and pOH (the equivalent scale for hydroxide ions) is:
How to Calculate pH — Step-by-Step
Method 1: From Hydrogen Ion Concentration
Formula: pH = −log₁₀[H⁺]
Problem: A solution has [H⁺] = 0.01 mol/L. Find the pH.
2. Substitute: pH = −log₁₀(0.01)
3. Calculate: log₁₀(0.01) = log₁₀(10⁻²) = −2
4. Apply negative: pH = −(−2) = 2
Problem: A solution has [H⁺] = 3.5×10⁻⁵ mol/L. Find the pH.
2. log₁₀(3.5×10⁻⁵) = log₁₀(3.5) + log₁₀(10⁻⁵)
= 0.544 + (−5) = −4.456
3. pH = −(−4.456) = 4.46
Use our pH Calculator to solve this instantly for any hydrogen ion concentration.
Method 2: From Hydroxide Ion Concentration
When you are given [OH⁻] instead of [H⁺], find pOH first then convert.
Formula: pOH = −log₁₀[OH⁻], then pH = 14 − pOH
Problem: A solution has [OH⁻] = 0.001 mol/L. Find the pH.
2. pH = 14 − pOH = 14 − 3 = 11
Method 3: From a Strong Acid Concentration
For strong acids (HCl, H₂SO₄, HNO₃) that fully dissociate, [H⁺] = concentration of acid.
Problem: What is the pH of 0.05 mol/L HCl?
2. pH = −log₁₀(0.05)
3. pH = −(−1.301) = 1.30
Method 4: Finding [H⁺] from pH
The reverse calculation — given pH, find [H⁺]:
Problem: A solution has pH 4.5. Find [H⁺].
2. [H⁺] = 3.162×10⁻⁵ mol/L
Strong Acids vs Weak Acids — Why It Matters for pH
Not all acids produce the same pH at the same concentration. The key difference is dissociation:
Strong acids fully dissociate in water — every molecule releases a H⁺ ion. HCl, HNO₃, and H₂SO₄ are strong acids. At 0.1 mol/L, HCl gives [H⁺] = 0.1 mol/L and pH = 1.
Weak acids only partially dissociate — most molecules stay intact. Acetic acid (vinegar, CH₃COOH) at 0.1 mol/L gives [H⁺] ≈ 0.00134 mol/L and pH ≈ 2.87, not pH 1. This requires the acid dissociation constant (Ka) to calculate precisely.
The same principle applies to bases: strong bases like NaOH fully dissociate while weak bases like ammonia (NH₃) only partially react with water.
Buffer Solutions — How pH Is Maintained in Living Systems
A buffer is a solution that resists changes in pH when small amounts of acid or base are added. Buffers contain a weak acid and its conjugate base (or a weak base and its conjugate acid) in roughly equal concentrations.
The pH of a buffer is calculated using the Henderson-Hasselbalch equation:
The most important buffer in human physiology is the carbonic acid–bicarbonate buffer:
The kidneys and lungs constantly adjust these concentrations to keep pH within the 7.35–7.45 range. Failure to maintain this balance causes acidosis (pH too low) or alkalosis (pH too high), both of which are medical emergencies.
pH in Everyday Life
Soil pH and agriculture: Most crops grow best between pH 6 and 7. Acidic soils (below 6) are treated with lime (calcium carbonate) to raise pH. Alkaline soils (above 7.5) are treated with sulfur or acidifying fertilisers. Blueberries are an exception — they thrive at pH 4 to 5.
Swimming pools: Pool water is maintained at pH 7.2 to 7.6. Too acidic (below 7.2) and the water irritates eyes and corrodes metal fittings. Too basic (above 7.8) and chlorine becomes ineffective at killing bacteria, creating a health hazard.
Cooking: Baking soda (NaHCO₃, pH ~8.3) and baking powder both rely on acid-base reactions to produce CO₂ gas that makes baked goods rise. Buttermilk (pH ~4.5) is acidic enough to react with baking soda and produce the necessary bubbles.
Ocean acidification: Since the Industrial Revolution, the ocean has absorbed approximately 30% of all CO₂ emitted by human activities. CO₂ dissolves in seawater to form carbonic acid, dropping ocean pH from a pre-industrial 8.2 to the current 8.1. That 0.1 unit drop represents a 26% increase in hydrogen ion concentration — enough to dissolve coral skeletons and disrupt marine ecosystems.
Common Mistakes When Working with pH
❌ Thinking the Scale Only Goes 0 to 14
The pH scale can technically go below 0 and above 14 for extremely concentrated strong acids and bases. Concentrated sulfuric acid can have a pH of −1 or lower. The 0–14 range covers most practical situations.
❌ Forgetting the Logarithm Is Base 10
pH uses log₁₀ (common logarithm), not ln (natural logarithm). Using ln instead of log₁₀ will give a wrong answer every time.
❌ Confusing pH and pOH
pH measures H⁺ concentration; pOH measures OH⁻ concentration. They add up to 14 but are not interchangeable. A basic solution has high pOH AND low pH.
❌ Assuming Lower pH Always Means More Dangerous
Context matters. Stomach acid at pH 1–2 is essential for digestion and harmless to your stomach lining. Battery acid at pH 0 is dangerous. Bleach at pH 12–13 is dangerous on skin. The chemical identity of the acid or base matters as much as the pH.
❌ Thinking Neutral Means Safe
pH 7 means neutral, not harmless. Distilled water is pH 7. So is a solution of many neutral-pH toxic compounds. pH tells you about hydrogen ion concentration only — not toxicity.
Frequently Asked Questions
pH stands for “potential of hydrogen” (from the German Potenz Wasserstoff). It was coined by Sørensen in 1909 to describe the power or potential of hydrogen ion concentration in a solution.
At 25°C, pure water has [H⁺] = [OH⁻] = 10⁻⁷ mol/L. Since pH = −log(10⁻⁷) = 7, pure water defines neutrality. Note that at higher temperatures water dissociates more, so neutral pH drops slightly — at 37°C (body temperature) neutral pH is approximately 6.8.
Yes. Extremely concentrated strong acids can have negative pH values. Concentrated hydrochloric acid (12 mol/L) has a pH of approximately −1.08. The 0–14 range covers typical dilute aqueous solutions but is not a physical limit.
Acidity is the qualitative concept (a substance is acidic if it donates H⁺). pH is the quantitative measure of how acidic. All acidic solutions have pH below 7, but the pH number tells you exactly how acidic — a solution at pH 2 is 100 times more acidic than one at pH 4.
pH strips contain dyes called acid-base indicators that change colour at specific pH values. Different dyes change at different pH ranges. Universal indicators contain a mixture of dyes to give a rainbow of colours across the full 0–14 range, matched to a colour chart.
🧪 Want to Calculate pH Instantly?
Rather than calculating pH by hand, our pH Calculator solves for pH, pOH, [H⁺], and [OH⁻] from any starting value — including strong acid or base concentration — with full step-by-step working. For concentration calculations involving moles in solution, the Molarity Calculator handles all mole-based chemistry instantly.