3
11
0.001
M
1.0000e-11
M
3
11
0.001
M
1.0000e-11
M
The pH Calculator performs conversions between pH, pOH, hydrogen ion concentration [H⁺], and hydroxide ion concentration [OH⁻]. The pH scale is the universal measure of acidity and basicity in aqueous solutions, ranging from 0 (strongly acidic) to 14 (strongly basic), with 7 being neutral at 25°C. Developed by Soren Sorensen in 1909, pH is defined as the negative base-10 logarithm of the hydrogen ion activity: pH = -log₁₀[H⁺]. This calculator supports three modes: converting a hydrogen ion concentration to pH, converting a pH value back to hydrogen ion concentration, and converting hydroxide ion concentration to pOH (then pH). It provides all four related values simultaneously, making it an essential tool for chemistry students, laboratory technicians, environmental scientists, and anyone working with aqueous solutions.
The pH system is built on the autoionization of water: H₂O ⇌ H⁺ + OH⁻, with the equilibrium constant K_w = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C. This leads to the fundamental relationships:
pH = -log₁₀[H⁺]
pOH = -log₁₀[OH⁻]
pH + pOH = 14 (at 25°C)
[H⁺] = 10⁻ᵖᴴ
The logarithmic scale means each pH unit represents a tenfold change in hydrogen ion concentration. A solution at pH 3 has ten times more H⁺ ions than one at pH 4, and one hundred times more than pH 5. This compression is necessary because H⁺ concentrations in real solutions span many orders of magnitude — from about 10 M in concentrated acid to 10⁻¹⁵ M in concentrated base.
In the [H⁺] → pH mode, the calculator takes a hydrogen ion concentration and computes pH = -log₁₀(concentration). In pH → [H⁺] mode, it reverses the calculation: [H⁺] = 10⁻ᵖᴴ. In [OH⁻] → pOH mode, it first computes pOH = -log₁₀[OH⁻], then derives pH = 14 - pOH. All four values (pH, pOH, [H⁺], [OH⁻]) are always displayed regardless of the input mode.
A pH below 7 indicates an acidic solution, pH 7 is neutral, and pH above 7 is basic (alkaline). The further from 7, the stronger the acid or base. Remember that pH is logarithmic: the difference between pH 1 and pH 2 represents a 10× change in acidity. When comparing [H⁺] and [OH⁻], their product must always equal 10⁻¹⁴ at 25°C. At temperatures other than 25°C, the value of K_w changes, and the neutral pH shifts slightly from 7.
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0.001 M HCl fully dissociates, giving [H⁺] = 10⁻³ M, so pH = 3.0. The solution is acidic with pOH = 11 and [OH⁻] = 10⁻¹¹ M.
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At pH 8.5, [H⁺] = 10⁻⁸·⁵ ≈ 3.16 × 10⁻⁹ M and [OH⁻] = 10⁻⁵·⁵ ≈ 3.16 × 10⁻⁶ M. This is a mildly basic solution.
Yes. pH below 0 occurs in very concentrated strong acids (e.g., 10 M HCl has pH ≈ -1), and pH above 14 occurs in very concentrated strong bases. The 0-14 range is a practical guideline for most aqueous solutions, not a hard limit.
Hydrogen ion concentrations in practical solutions span roughly 15 orders of magnitude (from ~10 M to ~10⁻¹⁵ M). A logarithmic scale compresses this enormous range into a manageable 0-14 scale, making comparisons intuitive. Each unit change represents a tenfold change in [H⁺].
Yes. The water autoionization constant K_w increases with temperature. At 25°C, K_w = 10⁻¹⁴ and neutral pH = 7. At 37°C (body temperature), K_w ≈ 2.4 × 10⁻¹⁴ and neutral pH ≈ 6.8. At 60°C, neutral pH ≈ 6.5. The relationship pH + pOH = pK_w adjusts accordingly.
pH measures hydrogen ion concentration (-log₁₀[H⁺]) while pOH measures hydroxide ion concentration (-log₁₀[OH⁻]). They are complementary: pH + pOH = 14 at 25°C. Low pH means acidic (high [H⁺]), while low pOH means basic (high [OH⁻]).
Modern pH meters have accuracy of ±0.01 pH units when properly calibrated. pH indicator papers are less precise (±0.5 units). For solutions with very low ionic strength or very high acid/base concentration, activity corrections may be needed for accurate pH determination.
Pure water autoionizes: H₂O ⇌ H⁺ + OH⁻. At 25°C, [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M. Therefore pH = -log₁₀(10⁻⁷) = 7. This is the definition of neutral pH at 25°C.
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