Understanding pH: The Measure of Acidity
pH is a logarithmic scale that measures the hydrogen ion concentration in a solution. The scale runs from 0 to 14, with 7 being neutral. Values below 7 indicate acidic solutions, while values above 7 indicate basic (alkaline) solutions. This simple number encapsulates a millionfold range of hydrogen ion concentrations.
The pH calculator above handles calculations for both strong and weak acids and bases, converting between pH, pOH, and ion concentrations. Understanding these relationships is fundamental to chemistry, biochemistry, environmental science, and many industrial processes.
The pH Definition
pH = -log10[H+]
Where [H+] is the hydrogen ion concentration in moles per liter (M)
Key Relationships in Acid-Base Chemistry
Fundamental Equations
| pH + pOH = 14 | (at 25C) |
| [H+][OH-] = Kw = 10^-14 | (ion product of water) |
| pOH = -log10[OH-] | (hydroxide definition) |
| pKa + pKb = 14 | (conjugate pair) |
The pH Scale Explained
Because pH is logarithmic, each unit change represents a 10-fold change in hydrogen ion concentration:
| pH | [H+] (M) | [OH-] (M) | Character |
|---|---|---|---|
| 0 | 1 | 10^-14 | Strongly Acidic |
| 3 | 10^-3 | 10^-11 | Acidic |
| 5 | 10^-5 | 10^-9 | Weakly Acidic |
| 7 | 10^-7 | 10^-7 | Neutral |
| 9 | 10^-9 | 10^-5 | Weakly Basic |
| 11 | 10^-11 | 10^-3 | Basic |
| 14 | 10^-14 | 1 | Strongly Basic |
Strong vs. Weak Acids and Bases
Strong Acids and Bases
Strong acids and bases dissociate completely in water. For strong acids, [H+] equals the acid concentration. For strong bases like NaOH, [OH-] equals the base concentration.
Strong Acid Example: 0.01 M HCl
HCl -> H+ + Cl- (complete dissociation)
[H+] = 0.01 M
pH = -log(0.01) = -log(10^-2) = 2
Strong Base Example: 0.001 M NaOH
NaOH -> Na+ + OH- (complete dissociation)
[OH-] = 0.001 M
pOH = -log(0.001) = 3
pH = 14 - 3 = 11
Weak Acids and Bases
Weak acids and bases only partially dissociate. The extent of dissociation is characterized by the equilibrium constant Ka (for acids) or Kb (for bases).
Weak Acid Equilibrium
HA <-> H+ + A-
Ka = [H+][A-] / [HA]
For weak acid calculations:
[H+] = sqrt(Ka x C)
or more precisely: [H+] = (-Ka + sqrt(Ka^2 + 4KaC)) / 2
Common pKa Values Reference
The pKa value indicates acid strength - lower pKa means stronger acid:
| Acid | Formula | pKa | Classification |
|---|---|---|---|
| Hydrochloric | HCl | -7 | Strong |
| Sulfuric | H2SO4 | -3 | Strong |
| Phosphoric | H3PO4 | 2.15 | Weak |
| Hydrofluoric | HF | 3.17 | Weak |
| Acetic | CH3COOH | 4.76 | Weak |
| Carbonic | H2CO3 | 6.35 | Weak |
| Hydrogen sulfide | H2S | 7.0 | Weak |
| Hydrogen cyanide | HCN | 9.21 | Very Weak |
| Ammonium ion | NH4+ | 9.25 | Very Weak |
pH Calculation Examples
Example 1: Weak Acid pH
Problem: Find the pH of 0.1 M acetic acid (pKa = 4.76)
Step 1: Find Ka from pKa
Ka = 10^-4.76 = 1.74 x 10^-5
Step 2: Calculate [H+]
[H+] = sqrt(Ka x C) = sqrt(1.74 x 10^-5 x 0.1) = 1.32 x 10^-3 M
Step 3: Calculate pH
pH = -log(1.32 x 10^-3) = 2.88
Percent dissociation: (1.32 x 10^-3 / 0.1) x 100 = 1.32%
Example 2: Converting Between pH and [H+]
Given: A solution has pH = 5.4
Find [H+]:
[H+] = 10^-5.4 = 3.98 x 10^-6 M
Find [OH-]:
pOH = 14 - 5.4 = 8.6
[OH-] = 10^-8.6 = 2.51 x 10^-9 M
Example 3: Weak Base pH
Problem: Find the pH of 0.05 M ammonia (pKb = 4.75)
Step 1: Find Kb
Kb = 10^-4.75 = 1.78 x 10^-5
Step 2: Calculate [OH-]
[OH-] = sqrt(Kb x C) = sqrt(1.78 x 10^-5 x 0.05) = 9.43 x 10^-4 M
Step 3: Calculate pOH and pH
pOH = -log(9.43 x 10^-4) = 3.03
pH = 14 - 3.03 = 10.97
pH of Common Substances
| Substance | Typical pH | Category |
|---|---|---|
| Battery acid | 0-1 | Strongly acidic |
| Stomach acid | 1-2 | Strongly acidic |
| Lemon juice | 2-3 | Acidic |
| Vinegar | 2.5-3 | Acidic |
| Orange juice | 3-4 | Acidic |
| Coffee | 4-5 | Mildly acidic |
| Pure water | 7 | Neutral |
| Blood | 7.35-7.45 | Slightly basic |
| Seawater | 7.5-8.4 | Basic |
| Baking soda | 8.3-8.4 | Basic |
| Household ammonia | 11-12 | Strongly basic |
| Bleach | 12-13 | Strongly basic |
| Drain cleaner | 13-14 | Strongly basic |
Temperature Effects on pH
The ion product of water (Kw) changes with temperature, affecting neutral pH:
- At 0C: Kw = 0.114 x 10^-14, neutral pH = 7.47
- At 25C: Kw = 1.0 x 10^-14, neutral pH = 7.00
- At 37C: Kw = 2.4 x 10^-14, neutral pH = 6.81
- At 100C: Kw = 51 x 10^-14, neutral pH = 6.14
This is important for biochemical measurements at body temperature (37C) and industrial processes at elevated temperatures.
Frequently Asked Questions
Yes, concentrated strong acids can have pH below 0, and concentrated strong bases can exceed pH 14. For example, 10 M HCl has a theoretical pH of -1. However, the pH scale becomes less meaningful at extreme concentrations where activity coefficients deviate significantly from 1.
Enzymes and proteins have optimal pH ranges for function. Blood pH outside 7.0-7.8 can denature proteins and disrupt enzyme activity. The body uses multiple buffer systems (carbonic acid/bicarbonate, phosphate, proteins) to maintain pH around 7.4.
Ka is the acid dissociation constant, a measure of acid strength. pKa = -log(Ka), so lower pKa means stronger acid. pKa values are easier to work with because they avoid dealing with very small numbers and fit conveniently on the pH scale.
For polyprotic acids like H3PO4, each proton has its own pKa. The pH is primarily determined by the first dissociation if pKa values are well-separated. For accurate calculations, all equilibria must be considered simultaneously.
The simplified formula [H+] = sqrt(Ka x C) assumes negligible dissociation. This works when C >> Ka (less than 5% dissociation). For dilute solutions or moderate Ka values, the quadratic formula gives more accurate results.