Stoichiometry

• (pronounced stoy-key-OM-uh-tree) is the area of study that examines the quantities of substances consumed and produced in chemical reactions.

• Stoichiometry (Greek stoicheion, “element,” and metron, “measure”) provides an essential set of tools widely used in chemistry.

• Stoichiometry is built on an understanding of atomic masses, chemical formulas, and the law of conservation of mass.

Antoine Lavoisier((1734–1794)

• French nobleman and scientist.

• Discovered this important chemical law during the late 1700s.

• He is generally considered the father of modern chemistry because he conducted carefully controlled experiments and used quantitative measurements.

3.1 ∣ Chemical Equations

The chemical formulas to the left of the arrow represent the starting substances, called reactants. The chemical formulas to the right of the arrow represent substances produced in the reaction, called products. The numbers in front of the formulas, called coefficients, indicate the relative numbers of molecules of each kind involved in the reaction.

Balancing Equations

• Chemists construct unbalanced equations to identify reactants and products. The chemical equation must be balanced to determine product or reactant amounts.

• In balancing an equation, you need to understand the difference between coefficients and subscripts.

Indicating the States of Reactants and Products

Symbols indicating the physical state of each reactant and product are often shown in chemical equations. We use the symbols (g ), (l), (s), and (aq) for substances that are gasses, liquids, solids, and dissolved in aqueous (water) solution, respectively.

3.2 ∣ Simple Patterns of Chemical Reactivity

Three types of reactions: combination reactions, decomposition reactions, and combustion reactions.

Combination and Decomposition Reactions

• In combination reactions, two or more substances react to form one product.

• A combination reaction between a metal and a nonmetal, produces an ionic solid.

• A single reactant breaks apart to form two or more substances. Many compounds react this way when heated.

• In a decomposition reaction one substance undergoes a reaction to produce two or more other substances.

Combustion Reactions

Combustion reactions are rapid reactions that produce a flame.

• Most combustion reactions we observe involve O2 from air as a reactant.

• Combustion of oxygen-containing derivatives of hydrocarbons, such as CH3OH, also produces CO2 and H2O.

• In our bodies, however, the reactions take place in a series of intermediate steps that occur at body temperature. These reactions that involve intermediate steps are described as oxidation reactions instead of combustion reactions.

3.3 ∣ Formula Weights

Chemical formulas and chemical equations both have a quantitative significance in that the subscripts in formulas and the coefficients in equations represent precise quantities.

Formula and Molecular Weights

• The formula weight (FW) of a substance is the sum of the atomic weights (AW) of the atoms in the chemical formula of the substance.

• If the chemical formula is the chemical symbol of an element, such as Na, the formula weight equals the atomic weight of the element, in this case 23.0 amu. If the chemical formula is that of a molecule, the formula weight is also called the molecular weight (MW).

• It's wrong to call ionic compounds molecules because they're three-dimensional arrays of ions. The formula weight of an ionic substance is calculated by adding the atomic weights of the atoms in the empirical formula.

Percentage Composition from Chemical Formulas

• Chemists must sometimes calculate the percentage composition of a compound—that is, the percentage by mass contributed by each element in the substance.

• Calculating the percentage composition of any element in a substance (sometimes called the elemental composition of a substance) is straightforward if the chemical formula is known.

Chemical formula:

% mass composition of element = (number of atoms of element)(atomic weight of element)/ formula weight of substance x 100%

3.4 ∣ Avogadro’s Number and the Mole

• Even the smallest samples we deal with in the laboratory contain enormous numbers of atoms, ions, or molecules.

• Chemists therefore have devised a counting unit for describing large numbers of atoms or molecules.

• In chemistry the counting unit for numbers of atoms, ions, or molecules in a laboratory-size sample is the mole, abbreviated mol.

• One mole is the amount of matter that contains as many objects.

• Avogadro’s number, NA, in honor of the Italian scientist Amedeo Avogadro (1776–1856), and it is often cited with units of reciprocal moles, 6.02 * 1023 mol-1.

• Avogadro’s number is so large that it is difficult to imagine.

Molar Mass

• Amole is always the same number, but 1-mol samples of different substances have different masses.

• The mass in grams of one mole, often abbreviated as 1 mol, of a substance (that is, the mass in grams per mole) is called the molar mass of the substance.

• The molar mass in grams per mole of any substance is numerically equal to its formula weight in atomic mass units.

Interconverting Masses and Numbers of Particles

• The molar mass and Avogadro’s number are used as conversion factors to convert grams to moles and then moles to atoms.

• Any time you calculate the number of atoms, molecules, or ions in an ordinary sample of matter, you can expect the answer to be very large.

• Procedure for interconverting mass and number of formula units. The number of moles of the substance is central to the calculation.

3.5 ∣ Empirical Formulas from Analyses

• The empirical formula for a substance tells us the relative number of atoms of each element in the substance.

• The ratio of the numbers of moles of all elements in a compound gives the subscripts in the compound’s empirical formula.

Molecular Formulas from Empirical Formulas

The subscripts in the molecular formula of a substance are always whole-number multiples of the subscripts in its empirical formula.

Formula:

Whole-number multiple= molecular weight/empirical formula weight

Combustion Analysis

One technique for determining empirical formulas in the laboratory is combustion analysis, commonly used for compounds containing principally carbon and hydrogen.

3.6 ∣ Quantitative Information from Balanced Equations

The coefficients in a chemical equation represent the relative numbers of molecules in a reaction. The mole concept allows us to convert this information to the masses of the substances in the reaction.

Stoichiometrically equivalent

• A specific amount of one substance reacts with a specific amount of every other chemical listed in the equation.

• Procedure for calculating amounts of reactants consumed or products formed in a reaction. The number of grams of a reactant consumed or product formed can be calculated in three steps, starting with the number of grams of any reactant or product.

3.7 ∣ Limiting Reactants

• The reactant that is completely consumed in a reaction is called the limiting reactant because it determines, or limits, the amount of product formed.

• The other reactants are sometimes called excess reactants.

Theoretical and Percent Yields

• The quantity of product calculated to form when all of a limiting reactant is consumed is called the theoretical yield.

• The amount of product actually obtained, called the actual yield, is almost always less than (and can never be greater than) the theoretical yield.

The percent yield of a reaction relates actual and theoretical yields:

Percent yield = actual yield/theoretical yield x 100%