pH of Acids and Bases Study Guide (page 2)
The applications of acid-base chemistry extend beyond what was learned in Lesson 6, "Aqueous Reactions." The acidity of an acid solution can be measured by its pH, and equilibria can be established between an acid and a base to create a buffer.
Acid and Base Definitions
Acids are proton donors (according to Brönsted) or electron acceptors (according to Lewis, this is a more general concept). Strong acids are completely dissociated in water, releasing protons (H +) and anionic conjugate bases. Acids have a sour taste.
Bases are proton acceptors (Brönsted) or electron donors (Lewis).When dissolved in water, strong bases such as NaOH dissociate to release hydroxide ions and sodium cation. Bases have a bitter taste and feel slippery like soap. (See also Lesson 6, Table 6.2, for more on acids and bases.)
Reactions of Acids
When acids react with another substance, the products can be predicted based on solubility rules and the four principal acid reactions (see Table 15.1)
Show the reaction when HCl reacts with Mg(OH)2 (the neutralization of stomach acid with milk of magnesia).
2HCl + Mg(OH)2 MgCl2 + H2O
Show the reaction when magnesium carbonate reacts with nitric acid.
Autoionization of Water
In pure water, H2O dissociates to H+ ions (protons) and OH– ions (hydroxide):
H2O H+ + OH–
The molar concentration of H+ equals the molar concentration of OH–:
[H+] = [OH–] = 1 * 10–7 M In turn, the ion product of water is Kw = [H+] [OH–] = 1 * 10–14
The pH measures the negative logarithm (for the presentation of a very small number in a large scale) of the hydrogen ion concentration (in mol/L):
pH = –log [H+]
If pure water has a hydrogen ion concentration of 1 * 10–7 M, the pH of neutral water is 7. A pH of 7 defines a neutral solution. The pH scale ranges from 0 to 14 with acids in the lower end of the scale (smaller than pH 7), whereas bases are at the higher end (greater than pH 7). Some strong acids will result in a negative pH value.
pH of Strong Acids and Bases
Strong acids and bases completely dissociate (ionize) in solution. Therefore, the concentration of the solute is the same as the concentration of the [H+] for acid or [OH–] for base. A 6.0 M HCl solution produces 6.0 M H+ ions, and a 3.5 M solution of sodium hydroxide produces 3.5 M of OH– ions.
Calculate the pH of a 0.050 M HCl solution.
Because HCl is a strong acid, the [H+] = 0.05 M; pH = – log[H+] = – log(0.050) =
Calculate the pH of a 0.0030 M NaOH solution.
Because NaOH is a strong base, [OH–] = 0.0030 M;
1.0 * 10–14 = [OH–] [H+] [H+] = pH = –log[H+] = –log (3.33333 * 10–12) = 11.477 = 11
pH of Weak Acids
Weak acids are weak electrolytes and do not dissociate completely. An equilibrium exists between the reactants and the products, and the equilibrium constant must be taken into account to solve for the pH value. When a weak acid (HA) is dissolved in water, the conjugate base (A–) and conjugate acid (H+) are formed. The equilibrium constant for an acid is called the acid dissociation constant.
HA A– + H+
[Alternatively written HA + H2O A– + H3O+]
and pKa = –log [Ka]
What is the pH of a 1.0 M solution of acetic acid?
CH3CO2H CH3CO– + H+ so,
|Equilibrium||1.0 – x||+x||+x|
|x = 0.0042 = [H+]||(1.0 – x ≈ 1.0 if x is small)|
pH = –log [0.0042] =
Diprotic and Triprotic Acids
Substances containing more than one acidic proton are called polyprotic acids. Diprotic acids contain two acidic protons, and triprotic acids contain three acidic protons. Acid protons dissociate one at a time and have different Ka and pKa constants. Carbonic acid (H2CO3) is a diprotic acid.
pH of Weak Bases
Problems involving weak bases are treated similarly to the problems with weak acids. Weak bases (B) accept a proton from water (H2O) to produce a conjugate acid (HB+) and hydroxide ions:
B + H2O + HB+ + OH–
A buffer is a solution of a weak base and its conjugate acid (also weak) that prevents drastic changes in pH. The weak base reacts with any H+ ions that could increase acidity, and the weak conjugate acid reacts with OH– ions that may increase the basicity of the solution.
Carbonic Acid/Bicarbonate Buffer: H2CO3/HCO3–
Blood pH must be maintained at a pH of 7.40 by a buffer system consisting of the couple H2CO3/HCO3–:
H2CO3 HCO3– + H+
Neutralization of acid: HCO3– + H+ → H2CO3
Neutralization of base: H2CO3 + NaOH → NaHCO3 + H2O
Phosphate Buffer: H2PO4–/HPO4–2
The principal buffer system inside cells in blood consists of the couple H2PO4–/HPO4–2:
Neutralization of acid: HPO42– + H+ → H2PO4–
Neutralization of base:H2PO4– + OH – → HPO42– + H2O
The pH of a buffer solution can be calculated by the Henderson-Hasselbalch equation, which states that the pH of a buffer solution has a value close to the value of the weak acid (pKa):
What is the pH of a carbonate buffer solution with 0.65 M H2CO3 and 0.25 M HCO3–?
H2CO3 HCO3– + H+
Practice problems for these concepts can be found at - pH of Acids and Bases Practice Questions
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