Difficulty: Introductory | Prerequisites: Atoms and Molecules study notes (Part A)
Water is the solvent of life. This section covers how to prepare solutions at specific concentrations (a skill tested both in exams and in lab), how pH works mathematically, and how pH changes affect living organisms. If you are comfortable with exponents and logarithms, the pH calculations will be straightforward. If not, work through the examples below until the pattern clicks.
Solution preparation uses dilution formulas (C1V1 = C2V2) and mass/volume calculations. pH is the negative log of hydrogen ion concentration, and pH + pOH always equals 14 in water at 25 degrees C. Acid rain lowers pH in lakes, and different organisms have different tolerance thresholds, so small pH changes can wipe out entire populations.
Molarity (M)
The number of moles of solute per litre of solution (mol/L). Think of it as the "strength" of a solution: 1 M glucose means one mole of glucose dissolved in enough water to make one litre total.
Millimolar (mM)
1/1000 of a molar concentration. 100 mM = 0.1 M.
Percent solution (w/v)
Grams of solute per 100 mL of solution. A 10% sucrose solution contains 10 g of sucrose in every 100 mL.
Dilution
Reducing the concentration of a solution by adding more solvent. The key relationship: C1 x V1 = C2 x V2 (concentration times volume before = concentration times volume after). The amount of solute stays the same; only the volume changes.
pH
The negative base-10 logarithm of the hydrogen ion concentration: pH = -log[H+]. A low pH means high [H+] (acidic); a high pH means low [H+] (basic).
pOH
The negative base-10 logarithm of the hydroxide ion concentration: pOH = -log[OH-]. In water at 25 degrees C, pH + pOH = 14.
Kw (ion product of water)
The equilibrium constant for the self-ionisation of water: Kw = [H+] x [OH-] = 1 x 10^-14 at 25 degrees C. This value is constant regardless of whether the solution is acidic, basic, or neutral.
Buffer
A solution that resists changes in pH when small amounts of acid or base are added. Buffers consist of a weak acid and its conjugate base (or a weak base and its conjugate acid). In simple terms, a buffer acts as a chemical sponge for H+ or OH- ions.
Acid rain
Precipitation with a pH below about 5.6 (the pH of normal rain, which is slightly acidic due to dissolved CO2). Caused primarily by sulphur dioxide and nitrogen oxides reacting with water vapour in the atmosphere.
a. 50 mL of 100 mM NaCl from a 0.5 M stock
First, convert units: 100 mM = 0.1 M.
Apply C1 x V1 = C2 x V2:
C1 = 0.5 M, V1 = ?, C2 = 0.1 M, V2 = 50 mL
V1 = (C2 x V2) / C1 = (0.1 M x 50 mL) / 0.5 M = 10 mL
Procedure: measure 10 mL of the 0.5 M NaCl stock solution and add 40 mL of water (distilled/deionised) to bring the total volume to 50 mL.
b. 300 mL of 10% sucrose from solid
10% (w/v) means 10 g per 100 mL.
For 300 mL: (10 g / 100 mL) x 300 mL = 30 g
Procedure: weigh 30 g of solid sucrose, dissolve it in some water, then add water to reach a final volume of 300 mL.
c. 200 mL of 0.1 M glucose from 1.0 M glucose
Apply C1 x V1 = C2 x V2:
C1 = 1.0 M, V1 = ?, C2 = 0.1 M, V2 = 200 mL
V1 = (0.1 M x 200 mL) / 1.0 M = 20 mL
Procedure: measure 20 mL of the 1.0 M glucose stock and add 180 mL of water to bring the total to 200 mL.
The completed table from the worksheet:
pH | [H+] | [OH-] | Kw | Acidic/Basic/Neutral |
|---|---|---|---|---|
7 | 10^-7 M | 10^-7 M | 10^-14 | Neutral |
10 | 10^-10 M | 10^-4 M | 10^-14 | Basic |
5 | 10^-5 M | 10^-9 M | 10^-14 | Acidic |
Row-by-row reasoning:
Row 1: [OH-] = 10^-7 M is given. Since Kw = [H+][OH-] = 10^-14, then [H+] = 10^-14 / 10^-7 = 10^-7 M. pH = -log(10^-7) = 7. Equal [H+] and [OH-] means neutral.
Row 2: pH 10 is given. [H+] = 10^-10 M. [OH-] = 10^-14 / 10^-10 = 10^-4 M. More OH- than H+, so basic.
Row 3: [H+] = 10^-5 M is given. pH = -log(10^-5) = 5. [OH-] = 10^-14 / 10^-5 = 10^-9 M. More H+ than OH-, so acidic.
a. pH 6.0 lake with 100-fold increase in [H+]:
A 100-fold increase in [H+] means the pH drops by 2 units (since pH is a log scale). pH 6.0 becomes pH 4.0.
Looking at the tolerance table at pH 4.0, only frogs can survive. All other groups (trout, bass, perch, salamanders, clams, crayfish, snails) would be eradicated.
b. [OH-] = 10^-9 M: which organisms survive?
If [OH-] = 10^-9 M, then [H+] = Kw / [OH-] = 10^-14 / 10^-9 = 10^-5 M, so pH = 5.
At pH 5.0, the surviving groups are: trout, perch, frogs, and salamanders. Bass, clams, crayfish, and snails cannot tolerate pH 5.0.
c. Buffer X and vulnerable animals:
The buffer graph shows the buffering range is approximately pH 5.5 to 6.5. Even if the buffer successfully holds pH in this range, animals that cannot tolerate pH 5.5 will still be vulnerable. Looking at the tolerance table at pH 5.5: clams and snails cannot survive pH 5.5 (they need pH 6.0 or above). So clams and snails remain vulnerable even with successful buffering.
Dilution formula: C1 x V1 = C2 x V2
Percent solution (w/v): mass of solute (g) / volume of solution (mL) x 100
pH: -log[H+]
pOH: -log[OH-]
pH + pOH = 14 (at 25 degrees C)
Kw = [H+] x [OH-] = 1 x 10^-14 (at 25 degrees C)
pH scale: each unit = 10-fold change in [H+]. A drop from pH 6 to pH 4 = 100-fold increase in [H+].
Acid rain is not a textbook abstraction. Regions downwind of coal-burning power plants have experienced lake acidification that eliminated fish populations, exactly as the tolerance table predicts. The buffering capacity of a lake depends on the local geology: limestone-bedded lakes resist pH changes because CaCO3 neutralises incoming acid, while granite-bedded lakes have almost no buffering capacity and are highly vulnerable.
Dilution calculations are a daily reality in any biology or chemistry lab. Preparing media, running assays, and diluting stock reagents all require C1V1 = C2V2.
Students often think a pH change from 6 to 4 is "just two units." On a log scale, that is a 100-fold increase in hydrogen ion concentration. Each pH unit represents a tenfold change.
Students sometimes confuse concentration with amount. Diluting a solution lowers the concentration, but the total amount of solute (in moles) stays the same.
Students often think Kw changes with pH. It does not (at a given temperature). Kw is constant at 10^-14 at 25 degrees C regardless of whether the solution is acidic or basic.
Students sometimes think buffers prevent all pH change. Buffers resist pH change within their buffering range, but they have a capacity limit and cannot hold pH against overwhelming additions of acid or base.
You will almost certainly be asked to perform a dilution calculation. Know C1V1 = C2V2 cold.
Be able to convert between pH, [H+], and [OH-] quickly. Practice until it is automatic.
Understand how to read a pH tolerance table and determine which organisms survive at a given pH.
The buffer question tests whether you understand the concept of buffering range and can identify vulnerable organisms even within that range.
True or False: A pH of 3 has 10 times more H+ ions than a pH of 4.
Fill in the blank: C1 x V1 = ______ x ______
True or False: Kw changes when you add acid to water.
Fill in the blank: If [H+] = 10^-9 M, the pH is ______ and the solution is ______.
True or False: A buffer can prevent all pH changes regardless of how much acid is added.
Answers: 1. True, 2. C2 x V2, 3. False, 4. 9, basic, 5. False
Q: You need 250 mL of 0.2 M NaCl from a 2.0 M stock. How much stock do you use, and how much water do you add?
A: V1 = (C2 x V2) / C1 = (0.2 x 250) / 2.0 = 25 mL of stock. Add 225 mL of water to bring the total to 250 mL.
Q: A lake has a pH of 6.0. A nearby factory begins releasing acidic waste that increases the hydrogen ion concentration by 1000-fold. What is the new pH, and which organisms from the tolerance table can survive?
A: A 1000-fold increase in [H+] drops the pH by 3 units: pH 6.0 becomes pH 3.0. At pH 3.0, no organisms in the table can survive (the most acid-tolerant, frogs, only tolerate down to pH 4.0).
Q: If the [OH-] in a solution is 10^-3 M, what are the pH, [H+], and whether the solution is acidic or basic?
A: [H+] = Kw / [OH-] = 10^-14 / 10^-3 = 10^-11 M. pH = 11. The solution is basic.
Q: Why do buffers have a limited effective range?
A: Buffers work by absorbing excess H+ or OH- using their weak acid/conjugate base pair. Once either component is depleted, the buffer can no longer absorb additional acid or base, and pH changes freely.
The pH concepts here connect directly to enzyme kinetics later in the course: enzymes have optimal pH ranges, and deviations denature them (see Part A study notes). Buffer chemistry is also essential for understanding blood pH regulation (carbonic acid/bicarbonate buffer system) in physiology courses.
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