HESI A2 · Chemistry · 6 min read
Chemical Reactions, Moles and Stoichiometry
- Identify types of chemical reactions and explain the principles of balancing chemical equations.
- Perform conversions between mass, moles, and representative particles using molar mass and Avogadro's number.
- Apply stoichiometric mole ratios to determine limiting reactants, theoretical yields, percent yields, and chemical formulas.
Chemical Equations and the Law of Conservation of Mass
A chemical equation represents a chemical reaction using chemical formulas and symbols. The starting substances entering the reaction are called reactants and are written on the left side of the equation. The substances formed as a result of the change are called products and appear on the right side. An arrow points from reactants to products, indicating the direction in which the chemical change proceeds.
The fundamental rule governing every chemical equation is the law of conservation of mass, which states that matter cannot be created or destroyed in a chemical reaction. Because matter is conserved, a balanced chemical equation must have the same number of atoms of each element on both sides of the reaction arrow. Consequently, in a balanced chemical equation, the total mass of the reactants equals the total mass of the products.
To balance an equation, coefficients are whole numbers placed before chemical formulas to balance the number of atoms of each element. In contrast, subscripts indicate the number of atoms of each element in a molecule or formula unit. Changing subscripts would alter the chemical identity of the substance, so subscripts cannot be changed when balancing equations. Equations also include state symbols to indicate physical states: solid (s), liquid (l), gas (g), and aqueous solution (aq).
Common mistakes
- Changing chemical subscripts instead of coefficients when balancing equations, which incorrectly alters the identity of the chemical substances.
- Using Avogadro's number when asked to convert between mass and moles, rather than using the substance's molar mass.
- Assuming that the reactant with the smaller starting mass is automatically the limiting reactant, without first converting masses to moles and applying the balanced mole ratio.
- Confusing empirical and molecular formulas by forgetting that the empirical formula must always represent the lowest whole-number ratio of atoms.
Key points
- Chemical equations represent reactions with reactants on the left and products on the right; the law of conservation of mass requires balancing equations using coefficients without altering subscripts.
- Reactions are classified into synthesis, decomposition, single replacement, double replacement, and combustion, with oxidation-reduction governing electron transfers.
- A mole is the SI unit for the amount of substance, containing Avogadro's number (approximately 6.022 x 10^23) of representative particles.
- Molar mass (in g/mol) is numerically equal to the formula or atomic mass and serves as the conversion factor between mass in grams and moles.
- Stoichiometry utilizes mole ratios from balanced equations to calculate amounts of reactants and products, identify limiting and excess reactants, and find theoretical yield.
- Percent yield compares the actual yield obtained in an experiment to the theoretical yield via the formula (actual yield / theoretical yield) x 100 percent.
- An empirical formula is the simplest whole-number ratio of elements in a compound, while a molecular formula shows the actual number of atoms in a molecule and is a whole-number multiple of the empirical formula.
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