Source: Comprehensive Guide to Chemical Principles, CHM 111
Tags: molarity, concentration, electrolyte, nonelectrolyte, strong electrolyte, weak electrolyte, dissociation, stoichiometry, balanced equation, limiting reagent, mole ratio, mass-to-mass conversion, solution, solute, solvent, CHM 111, general chemistry
Difficulty: Introductory to Intermediate Prerequisites: Mole concept and molar mass conversions (see "Mole Concept and Molar Mass, CHM 111 – Study Notes"). Comfort with algebraic rearrangement.
Solutions are everywhere in chemistry, from the acid in your stomach to the saline drip in a hospital. Understanding how to describe their concentration (molarity) and how dissolved substances behave (electrolytes vs. nonelectrolytes) is essential for predicting reactions in aqueous environments. Stoichiometry then gives you the quantitative toolkit: given a balanced equation, you can calculate exactly how much product forms from a given amount of reactant, or figure out which reactant runs out first. Together, these topics form the quantitative backbone of CHM 111.
Electrolytes dissociate into ions in water and conduct electricity; nonelectrolytes do not. Molarity measures concentration as moles of solute per litre of solution. Stoichiometry uses mole ratios from balanced equations to convert between amounts of reactants and products.
Solution
A homogeneous mixture of two or more substances. The substance present in the larger amount is the solvent; the substance dissolved in it is the solute. In simple terms, it is a thoroughly mixed combination where you cannot see the individual components.
Electrolyte
A substance that dissociates into ions when dissolved in water, making the solution conduct electricity. Think of it as a compound that "breaks apart" in water.
Strong electrolyte
An electrolyte that dissociates completely. Every formula unit splits into ions. Examples: NaCl, HCl, NaOH, and most soluble ionic compounds.
Weak electrolyte
An electrolyte that only partially dissociates. At any given moment, most of the substance remains as intact molecules, with only a small fraction existing as ions. Examples: acetic acid (CH₃COOH), ammonia (NH₃).
Nonelectrolyte
A substance that dissolves in water without forming ions. The solution does not conduct electricity. Examples: sugar (sucrose), ethanol.
Molarity (M)
Concentration expressed as moles of solute per litre of solution. In simple terms, it tells you how "crowded" the solution is with solute molecules or ions.
Stoichiometry
The area of chemistry that deals with the quantitative relationships between reactants and products in a chemical reaction. Think of it as the accounting system for chemical reactions.
Balanced chemical equation
An equation where the number of atoms of each element is the same on both sides. This reflects the law of conservation of mass: matter is neither created nor destroyed.
Mole ratio
The ratio of moles of one substance to moles of another in a balanced equation. It comes directly from the coefficients. In 2H₂ + O₂ → 2H₂O, the mole ratio of H₂ to O₂ is 2 : 1.
Limiting reagent (limiting reactant)
The reactant that is completely consumed first in a reaction, determining the maximum amount of product that can form. The other reactant is in excess.
When a substance dissolves in water, one of three things happens:
Complete dissociation (strong electrolyte): NaCl → Na⁺ + Cl⁻. The solution conducts electricity well. Strong acids (HCl, HNO₃, H₂SO₄), strong bases (NaOH, KOH), and most soluble salts fall here.
Partial dissociation (weak electrolyte): CH₃COOH ⇌ CH₃COO⁻ + H⁺. The solution conducts electricity poorly. Weak acids and weak bases are the main examples.
No dissociation (nonelectrolyte): Sugar dissolves but stays as intact molecules. No ions form, so no conductivity.
A quick practical test: if you put electrodes in the solution and the light bulb glows brightly, it is a strong electrolyte. Dimly, a weak electrolyte. Not at all, a nonelectrolyte.
The formula:
Molarity (M) = Moles of solute ÷ Volume of solution in litres
Rearranged forms:
Moles of solute = M × V (in litres)
Volume (L) = Moles of solute ÷ M
Worked example:
Dissolve 0.25 mol NaCl in enough water to make 0.50 L of solution.
M = 0.25 mol ÷ 0.50 L = 0.50 M
Important detail: Volume refers to the final volume of the solution, not the volume of solvent alone. If you dissolve 0.25 mol of NaCl in water and then add water until the total volume reaches 0.50 L, you have a 0.50 M solution.
The law of conservation of mass requires that atoms are neither created nor destroyed. Balancing means adjusting coefficients (the numbers in front of formulas) so each element has equal counts on both sides.
Practical approach:
Write the unbalanced equation with correct formulas.
Start by balancing elements that appear in only one reactant and one product.
Balance hydrogen and oxygen last (they tend to appear in multiple compounds).
Use whole-number coefficients. If you end up with fractions, multiply everything through to clear them.
Check your work: count every element on both sides.
Worked example:
Unbalanced: Fe + O₂ → Fe₂O₃
Balance Fe: 2 Fe on the right, so put 2 in front of Fe on the left. 2Fe + O₂ → Fe₂O₃
Balance O: 3 O on the right, but O₂ gives only even numbers. Use 3/2: 2Fe + 3/2 O₂ → Fe₂O₃
Clear the fraction by multiplying everything by 2: 4Fe + 3O₂ → 2Fe₂O₃
Check: Fe: 4 = 4. O: 6 = 6. Balanced.
Once you have a balanced equation, the coefficients give you mole ratios that let you convert between any two substances in the reaction.
Mole-to-mole conversion:
Using 2H₂ + O₂ → 2H₂O, if you have 3.0 mol of O₂:
Moles of H₂O = 3.0 mol O₂ × (2 mol H₂O / 1 mol O₂) = 6.0 mol H₂O
Mass-to-mass conversion (the full chain):
Convert grams of the given substance to moles (÷ molar mass).
Use the mole ratio from the balanced equation to find moles of the desired substance.
Convert moles of the desired substance to grams (× molar mass).
This three-step chain (grams → moles → mole ratio → moles → grams) is the standard pattern for nearly every stoichiometry problem.
Molarity:
M = n / V
Where n = moles of solute, V = volume of solution in litres.
Stoichiometry conversion chain:
Grams of A → (÷ Molar mass of A) → Moles of A
Moles of A → (× mole ratio B/A) → Moles of B
Moles of B → (× Molar mass of B) → Grams of B
Molarity calculations underpin every titration in an analytical chemistry lab. When a technician titrates an unknown acid with a standardised base, they use molarity and stoichiometry together to determine the acid's concentration.
Stoichiometry is how industrial chemists scale reactions from a beaker to a reactor. If you need 500 kg of product, stoichiometry tells you exactly how many kilograms of each reactant to order, minimising waste and cost.
Students often confuse "volume of solution" with "volume of solvent" when calculating molarity. Molarity uses the total volume of the final solution, which includes both solvent and solute.
When balancing equations, you may only change coefficients, never subscripts. Changing a subscript changes the substance itself (H₂O becomes H₂O₂, a completely different compound).
A common stoichiometry error is forgetting to check for a limiting reagent. If a problem gives you amounts of two reactants, one of them limits the amount of product. You must identify which one before calculating the product yield.
Students sometimes assume "strong electrolyte" means "very soluble." They are different properties. A strong electrolyte dissociates completely when it dissolves, but that says nothing about how much dissolves in the first place.
⚠️ Molarity is the default concentration unit throughout CHM 111. Expect to use it in dilution problems (M₁V₁ = M₂V₂), titration calculations, and equilibrium expressions.
⚠️ Stoichiometry problems are almost guaranteed on every exam. The mass-to-mass conversion chain is your bread and butter. Practise until the steps feel routine.
⚠️ Identifying electrolyte type (strong, weak, none) is a common multiple-choice question. Know the rules: strong acids, strong bases, and soluble salts are strong electrolytes. Weak acids and weak bases are weak electrolytes. Molecular compounds like sugar are nonelectrolytes.
⚠️ Balancing equations must be done correctly before stoichiometry can work. An unbalanced equation gives wrong mole ratios and therefore wrong answers for everything downstream.
True or false: A 1.0 M NaCl solution contains 1.0 mol of NaCl in 1.0 L of water.
Fill in the blank: In a balanced equation, the coefficients represent the ______ ratios of reactants and products.
True or false: Acetic acid (CH₃COOH) is a strong electrolyte.
Fill in the blank: To convert grams of reactant A to grams of product B, you must first convert grams of A to ______.
True or false: When balancing a chemical equation, you may change the subscripts in a chemical formula.
Answers: 1. False (it is 1.0 mol in 1.0 L of solution, not 1.0 L of water; the distinction matters). 2. Mole. 3. False (it is a weak electrolyte; it only partially dissociates). 4. Moles of A. 5. False (only coefficients may be changed).
Q: What is the molarity of a solution made by dissolving 5.85 g of NaCl (molar mass 58.44 g/mol) in enough water to make 250 mL of solution?
A: Moles of NaCl = 5.85 ÷ 58.44 = 0.1001 mol. Volume = 250 mL = 0.250 L. M = 0.1001 ÷ 0.250 = 0.400 M.
Q: Given the reaction 2H₂ + O₂ → 2H₂O, how many grams of water are produced from 4.0 g of H₂? (Molar masses: H₂ = 2.016 g/mol, H₂O = 18.016 g/mol)
A: Moles of H₂ = 4.0 ÷ 2.016 = 1.984 mol. Mole ratio: 2 mol H₂O per 2 mol H₂ = 1 : 1. So 1.984 mol H₂O is produced. Mass = 1.984 × 18.016 = 35.7 g.
Q: Classify each as a strong electrolyte, weak electrolyte, or nonelectrolyte: NaOH, CH₃COOH, glucose (C₆H₁₂O₆).
A: NaOH is a strong electrolyte (strong base, dissociates completely). CH₃COOH is a weak electrolyte (weak acid, partially dissociates). Glucose is a nonelectrolyte (molecular compound, no ions in solution).
Q: Balance the equation: Al + HCl → AlCl₃ + H₂
A: 2Al + 6HCl → 2AlCl₃ + 3H₂. Check: Al: 2 = 2. H: 6 = 6. Cl: 6 = 6. Balanced.
Molarity connects back to the mole concept: you cannot prepare a solution of known concentration without first converting a mass of solute to moles. Stoichiometry builds on both the mole concept and balanced equations, and later extends into limiting reagent and percent yield problems. In subsequent CHM 111 topics, these skills feed into acid-base chemistry, titrations, gas stoichiometry (PV = nRT), and equilibrium calculations.
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