6.2 Solutions, Acids & Bases

Why Solutions, Acids & Bases Matter for the NAPT

Most of the chemistry a Sailor actually touches day to day isn't a dry powder — it's a solution. Lead-acid batteries that start diesel generators and submarines run on sulfuric acid dissolved in water. Boiler and reactor coolant water is chemically treated and constantly monitored for pH. Firefighting foam concentrate, cleaning compounds, and even seawater itself are solutions with measurable concentrations and acid-base properties. The NAPT's chemistry content tests the vocabulary and math behind all of this: what a solution is, how concentrated it is, and where it falls on the acid-base scale.

Solutions: Solute, Solvent & Concentration

A solution is a homogeneous mixture in which one substance (the solute) is dissolved uniformly throughout another (the solvent). In saltwater, salt (NaCl) is the solute and water is the solvent. Because water dissolves so many substances, it is often called the "universal solvent," and solutions with water as the solvent are called aqueous solutions (written with an (aq) label).

A solution's concentration describes how much solute is packed into a given amount of solvent or solution. A concentrated solution has a large amount of solute relative to solvent; a dilute solution has comparatively little. Chemists express concentration precisely using molarity (M):

M = moles of solute ÷ liters of solution

Worked Example 1: What is the molarity of a solution made by dissolving 2.00 mol of sodium chloride (NaCl) in enough water to make 4.00 L of solution?

• M = mol ÷ L = 2.00 mol ÷ 4.00 L = 0.500 M

Worked Example 2 (mass to moles to molarity): A technician dissolves 58.5 g of NaCl (molar mass = 58.5 g/mol) in enough water to make 2.00 L of solution. What is the molarity?

• Step 1 — convert mass to moles: moles = mass ÷ molar mass = 58.5 g ÷ 58.5 g/mol = 1.00 mol
• Step 2 — apply the molarity formula: M = 1.00 mol ÷ 2.00 L = 0.500 M

This two-step pattern — grams to moles using molar mass, then moles per liter — is the single most common calculation style tested for solution questions.

The pH Scale

The pH scale measures how acidic or basic (alkaline) a solution is, based on the concentration of hydrogen ions ([H+]) it contains:

pH = -log[H+]

The scale runs from 0 to 14 at 25°C:

• pH less than 7 — acidic (more H+ ions)
• pH equal to 7 — neutral (pure water)
• pH greater than 7 — basic/alkaline (fewer H+ ions, more OH- ions)

Because pH is a logarithmic scale, each whole-number step represents a tenfold (10×) change in hydrogen-ion concentration. A solution at pH 3 is 10 times more acidic than one at pH 4, and 100 times more acidic than one at pH 5.

A closely related measure is pOH, based on hydroxide ion concentration ([OH-]):

pOH = -log[OH-], and pH + pOH = 14 (at 25°C)

Strong vs. Weak Acids & Bases

"Strong" and "weak" describe how completely an acid or base dissociates (breaks apart into ions) in water — not how concentrated the solution is or how dangerous it is. A strong acid or strong base ionizes essentially 100% in solution. A weak acid or weak base only partially ionizes, leaving most of the molecules undissociated.

Strong acids (memorize these — there are only six commonly tested):

• Hydrochloric acid (HCl)
• Sulfuric acid (H2SO4)
• Nitric acid (HNO3)
• Hydrobromic acid (HBr)
• Hydroiodic acid (HI)
• Perchloric acid (HClO4)

Strong bases include sodium hydroxide (NaOH), potassium hydroxide (KOH), and the other Group 1 (alkali metal) and heavier Group 2 (alkaline earth) hydroxides.

Weak acids (such as acetic acid, CH3COOH, found in vinegar, and carbonic acid, H2CO3, formed when CO2 dissolves in water) and weak bases (such as ammonia, NH3) establish an equilibrium — some molecules ionize, but most remain intact. This is why a 1 M solution of acetic acid is far less corrosive and has a much higher (less extreme) pH than a 1 M solution of hydrochloric acid, even though both are described as "acids" at the same concentration.

Common trap: "Strong" does not mean "concentrated." A dilute strong acid (say, 0.01 M HCl) fully dissociates and can still have a very low pH, while a concentrated weak acid may dissociate only slightly and have a much higher pH than a strong acid at the same molarity.

Worked pH Calculations

Worked Example 3: What is the pH of a 0.010 M solution of hydrochloric acid (HCl)?

Because HCl is a strong acid, it dissociates essentially completely, so [H+] = 0.010 M = 1.0 × 10⁻² M.

• pH = -log(1.0 × 10⁻²) = 2.00

Worked Example 4: What is the pH of a 0.0010 M solution of sodium hydroxide (NaOH)?

NaOH is a strong base and dissociates completely, so [OH-] = 0.0010 M = 1.0 × 10⁻³ M.

• pOH = -log(1.0 × 10⁻³) = 3.00
• pH = 14.00 - pOH = 14.00 - 3.00 = 11.00

Notice the pattern: for strong acids and bases, the ion concentration equals the molarity directly (because dissociation is complete), which reduces the pH or pOH calculation to a single logarithm step.

Exam Strategy

1. Molarity is always moles of solute divided by liters of total solution — not liters of solvent added before mixing.
2. Convert grams to moles first using molar mass before applying the molarity formula.
3. Memorize the six strong acids and the common strong bases — anything else defaults to "weak" for exam purposes.
4. Remember pH + pOH = 14 at 25°C, and that pH is logarithmic — a 1-unit pH difference is a 10-fold concentration difference.