Acids, Bases & pH Cheat Sheet

A printable reference covering pH, pOH, Kw, acid-base definitions, neutralization, titration, and indicators for grades 10-11.

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Acids, bases, and pH explain how substances donate or accept hydrogen ions and how their solutions behave in water. This cheat sheet helps students quickly connect definitions, formulas, and common reaction patterns. It is especially useful for solving concentration, pH, neutralization, and titration problems. Grade 10-11 chemistry often depends on recognizing whether a substance is acidic, basic, neutral, strong, or weak. The most important ideas are the relationships among $[\mathrm{H_3O^+}]$ , $[\mathrm{OH^-}]$ , pH, pOH, and $K_w$ . Strong acids and bases dissociate completely, while weak acids and bases only partially ionize. Neutralization reactions form water and a salt when acids and bases react. Titration calculations usually rely on mole ratios from a balanced equation and the relationship $n = C V$ .

Key Facts

An Arrhenius acid increases $[\mathrm{H_3O^+}]$ in water, while an Arrhenius base increases $[\mathrm{OH^-}]$ in water.
A Brønsted-Lowry acid donates a proton, and a Brønsted-Lowry base accepts a proton.
The pH of a solution is calculated with $\mathrm{pH} = -\log[\mathrm{H_3O^+}]$ .
The pOH of a solution is calculated with $\mathrm{pOH} = -\log[\mathrm{OH^-}]$ .
At $25^{\circ}\mathrm{C}$ , water has $K_w = [\mathrm{H_3O^+}][\mathrm{OH^-}] = 1.0 \times 10^{-14}$ .
At $25^{\circ}\mathrm{C}$ , the relationship between pH and pOH is $\mathrm{pH} + \mathrm{pOH} = 14.00$ .
For a monoprotic strong acid such as $\mathrm{HCl}$ , the acid concentration equals the hydronium concentration, so $[\mathrm{H_3O^+}] = [\mathrm{HCl}]$ .
For a neutralization titration, use the balanced mole ratio with $n = C V$ , where $n$ is moles, $C$ is concentration, and $V$ is volume in liters.

Vocabulary

Acid: An acid is a substance that donates $\mathrm{H^+}$ or increases $[\mathrm{H_3O^+}]$ in water.
Base: A base is a substance that accepts $\mathrm{H^+}$ or increases $[\mathrm{OH^-}]$ in water.
pH: pH is a logarithmic measure of acidity defined by $\mathrm{pH} = -\log[\mathrm{H_3O^+}]$ .
Neutralization: Neutralization is a reaction between an acid and a base that usually produces water and a salt, such as $\mathrm{HCl} + \mathrm{NaOH} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}$ .
Indicator: An indicator is a substance that changes color over a specific pH range to show whether a solution is acidic or basic.
Buffer: A buffer is a solution that resists pH change because it contains a weak acid and its conjugate base or a weak base and its conjugate acid.

Common Mistakes to Avoid

Using $[\mathrm{OH^-}]$ in the pH formula, which is wrong because $\mathrm{pH} = -\log[\mathrm{H_3O^+}]$ and $\mathrm{pOH} = -\log[\mathrm{OH^-}]$ .
Forgetting that pH is logarithmic, which is wrong because a decrease of $1$ pH unit means $[\mathrm{H_3O^+}]$ increases by a factor of $10$ .
Treating weak acids as fully dissociated, which is wrong because weak acids only partially ionize and do not usually have $[\mathrm{H_3O^+}]$ equal to the original acid concentration.
Using milliliters directly in $n = C V$ , which is wrong because volume must be converted to liters when concentration is in $\mathrm{mol/L}$ .
Ignoring coefficients in neutralization reactions, which is wrong because titration calculations must use the mole ratio from the balanced equation.

Practice Questions

1 Calculate the pH of a solution with $[\mathrm{H_3O^+}] = 1.0 \times 10^{-3}\,\mathrm{mol/L}$ .
2 A solution has $\mathrm{pOH} = 4.25$ at $25^{\circ}\mathrm{C}$ . Find its pH and decide whether it is acidic, basic, or neutral.
3 What volume of $0.200\,\mathrm{mol/L}$ $\mathrm{NaOH}$ is needed to neutralize $25.0\,\mathrm{mL}$ of $0.100\,\mathrm{mol/L}$ $\mathrm{HCl}$ in the reaction $\mathrm{HCl} + \mathrm{NaOH} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}$ ?
4 Explain why a $0.100\,\mathrm{mol/L}$ weak acid usually has a higher pH than a $0.100\,\mathrm{mol/L}$ strong acid.