Chemical Equations and Stoichiometry

Introduction to Chemical Equations

Chemical equations are concise representations of chemical reactions, showing the reactants and products involved. They are fundamental to understanding how substances interact and transform in chemistry.

• Reactants appear on the left side of the equation.

• Products appear on the right side of the equation.

• Coefficients are used to balance the equation, ensuring the law of conservation of mass is followed.

• States of matter are indicated by (g) for gas, (l) for liquid, (s) for solid, and (aq) for aqueous solution.

Key Point: Coefficients are added to balance equations, not subscripts. Changing subscripts alters the identity and properties of the compound.

Example: Hydrogen and oxygen can form water or hydrogen peroxide, depending on the coefficients and subscripts:

Balancing Chemical Equations

Balancing chemical equations ensures that the number of atoms of each element is the same on both sides of the equation. The process involves adjusting coefficients, not subscripts.

• Write the correct formula for each substance.

• Balance the most complicated substance first.

• Adjust coefficients to balance each element.

• Check the balance by counting atoms.

Example: Complete combustion of pentane:

Types of Chemical Reactions

Combination Reactions

Combination reactions involve two or more substances reacting to form one product. Formation reactions are a subset, where a compound is formed directly from its elements.

Decomposition Reactions

In decomposition reactions, one substance breaks down into two or more substances. These reactions are often used in industrial and laboratory settings.

Example: Sodium azide decomposition is used to inflate airbags in cars.

Combustion Reactions

Combustion reactions are rapid reactions that produce a flame, typically involving oxygen as a reactant. They are important in energy production and everyday life.

Molecular Weight, Formula Weight, and Percent Composition

Molecular Weight (MW) and Formula Weight (FW)

The molecular weight is the sum of the atomic weights of the atoms in a molecule. Formula weight is used for ionic compounds, which do not exist as discrete molecules.

• Molecular weight (MW): For molecules, sum the atomic weights.

• Formula weight (FW): For ionic compounds, sum the atomic weights in the empirical formula.

Example: Ethane (C2H6): MW = amu

Example: Calcium chloride (CaCl2): FW = amu

Percent Composition

Percent composition is the mass percentage of each element in a compound, calculated from its formula.

Example: Oxalic acid (H2C2O4): H = 2.24%, C = 26.68%, O = 71.08%

The Mole and Avogadro's Number

The Mole Concept

The mole is a fundamental unit in chemistry, representing 6.02214076 × 1023 particles (Avogadro's number). It allows chemists to count atoms and molecules by weighing.

• 1 mole = 6.02214076 × 1023 particles

• The value is based on the number of atoms in exactly 12 grams of pure carbon-12.

Example: 1 mole of water (H2O) contains 6.022 × 1023 molecules and has a mass of 18.0 g.

Molar Mass

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol). For elements, it is the atomic weight from the periodic table. For compounds, it is the sum of the atomic weights in the formula.

• For diatomic gases, multiply the atomic weight by 2.

• Formula weight (in amu) equals molar mass (in g/mol).

Stoichiometric Calculations and Empirical Formulas

Stoichiometric Calculations

Stoichiometry involves using balanced chemical equations to calculate the quantities of reactants and products. The "mole bridge" connects grams, moles, and number of particles.

• Convert grams to moles using molar mass.

• Use mole ratios from the balanced equation.

• Convert moles to number of particles using Avogadro's number.

Determining Empirical Formulas

The empirical formula is the simplest whole-number ratio of atoms in a compound. It can be determined from percent composition or experimental data.

• Convert percentages to grams (assume 100 g sample).

• Convert grams to moles using molar mass.

• Write a pseudoformula using moles as subscripts.

• Divide by the smallest number of moles.

• Multiply ratios to obtain whole numbers if necessary.

Example: Chromium oxide (68.4% Cr, 31.6% O): Empirical formula is Cr2O3.

Example: Magnesium oxide (60.3% Mg): Empirical formula is MgO.

Determining Molecular Formulas

The molecular formula is a multiple of the empirical formula. If the molar mass is known, the molecular formula can be determined.

• Divide the molar mass by the empirical formula mass.

• Multiply the empirical formula by this factor.

Example: Empirical formula C4H8O, molar mass 216 g/mol: Molecular formula is C12H24O3.

Combustion Analysis

Combustion analysis is a technique for determining the empirical formula of organic compounds by burning a known mass and measuring the products.

• All original C forms CO2, H forms H2O, O is found by subtraction.

• Masses of constituent elements are determined, then empirical formula is calculated.

Example: A 0.1000 g sample produces 0.0930 g H2O and 0.227 g CO2. Empirical formula is C3H6O.

Redefining the Kilogram and Avogadro's Number

Modern Definition of Avogadro's Number

Avogadro's number was historically defined as the number of atoms in 12 grams of carbon-12. In 2018, the kilogram was redefined using precise measurements of silicon spheres, allowing Avogadro's number to be set as exactly 6.02214076 × 1023 particles.

Additional info: The precise definition of Avogadro's number is fundamental for accurate scientific measurements and standards.

Naming Inorganic Compounds: Common Cations and Anions

Common Cations

Cations are positively charged ions. The table below lists common cations encountered in general chemistry.

Common Anions

Anions are negatively charged ions. The table below lists common anions encountered in general chemistry.