Acids, bases, and pH
Acids and bases are everywhere, from lemon juice to soap. The pH scale measures how acidic or basic a solution is.
Definitions
Arrhenius acids produce \(\text{H}^+\) in water and bases produce \(\text{OH}^-\). The broader Bronsted-Lowry definition says acids donate protons (\(\text{H}^+\)) and bases accept them. In \(\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-\), water donates a proton, so ammonia acts as a base.
Conjugate pairs
When an acid donates a proton, what remains is its conjugate base: \(\text{HCl}\) and \(\text{Cl}^-\) are a conjugate pair. When a base accepts a proton, it forms its conjugate acid: \(\text{NH}_3\) and \(\text{NH}_4^+\). Strong acids have very weak conjugate bases.
The pH scale
\(\text{pH} = -\log[\text{H}_3\text{O}^+]\). Neutral water at \(25^\circ\)C has pH 7; acids have pH below 7 and bases above 7. Because the scale is logarithmic, each unit is a tenfold change: pH 3 is ten times more acidic than pH 4 and a hundred times more than pH 5.
\text{pH} = -\log[\text{H}_3\text{O}^+]The water equilibrium
Water ionises slightly: \([\text{H}_3\text{O}^+][\text{OH}^-] = K_w = 1.0 \times 10^{-14}\) at \(25^\circ\)C. So \(\text{pH} + \text{pOH} = 14\). A solution with \([\text{OH}^-] = 10^{-3}\) M has pOH 3 and pH 11.
K_w = [\text{H}_3\text{O}^+][\text{OH}^-] = 1.0\times 10^{-14}Strong and weak acids
Strong acids, such as HCl, \(\text{HNO}_3\), and \(\text{H}_2\text{SO}_4\), ionise completely in water. Weak acids, like acetic acid in vinegar, ionise only partly, establishing an equilibrium. Strength (how completely it ionises) is different from concentration (how much is dissolved).
Worked example: pH of a strong acid
What is the pH of 0.010 M HCl? It ionises completely, so \([\text{H}_3\text{O}^+] = 0.010 = 10^{-2}\) M and pH \(= 2\). Diluting tenfold to 0.0010 M raises the pH to 3.
Everyday pH
Stomach acid is about pH 1-2, lemon juice about 2, coffee about 5, pure water 7, blood about 7.4, baking soda solution about 8-9, and household bleach about 12-13. The body tightly controls blood pH; a drop below about 7.35 (acidosis) is a medical emergency.
Indicators
Acid-base indicators change colour over particular pH ranges. Litmus is red in acid and blue in base; phenolphthalein turns pink above about pH 8.2; universal indicator shows a rainbow of colours across the scale. Red cabbage juice is a homemade indicator.
Key takeaways
Bronsted-Lowry acids donate protons and bases accept them, forming conjugate pairs. pH \(= -\log[\text{H}_3\text{O}^+]\), with each unit a tenfold change, and \(\text{pH} + \text{pOH} = 14\) at \(25^\circ\)C. Strong acids ionise fully, weak acids partly; strength differs from concentration.
Flashcard
Proton donor
Check your understanding
What is the pH of a 0.001 M HCl solution?