Chemical Composition and Formulas

Empirical and Molecular Formulas

The empirical formula represents the simplest whole-number ratio of elements in a compound, while the molecular formula shows the actual number of each type of atom in a molecule.

• Empirical Formula: Simplest ratio of atoms (e.g., CH2O for glucose).

• Molecular Formula: Actual number of atoms (e.g., C6H12O6 for glucose).

• Relationship: The molecular formula is a whole-number multiple of the empirical formula.

• Example: Hydrogen peroxide has an empirical formula of HO and a molecular formula of H2O2.

Percent Composition

Percent composition indicates the mass percentage of each element in a compound.

• Calculation:

• Example: In H2O, percent H = 11.2%, percent O = 88.8%.

Chemical Reactions and Stoichiometry

Chemical Changes and Evidence

A chemical change results in the formation of new substances. Evidence includes color change, gas production, precipitate formation, and energy change.

• Examples of Evidence: Bubbling (gas), temperature change, color change, formation of a solid (precipitate).

Law of Conservation of Mass

The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction.

• Implication: The total mass of reactants equals the total mass of products.

• Equation:

Classifying Chemical Reactions

Chemical reactions can be classified into several types based on the rearrangement of atoms.

• Synthesis (Combination):

• Decomposition:

• Single Displacement:

• Double Displacement:

• Combustion: Hydrocarbon + O2 CO2 + H2O

Balancing Chemical Reactions

Balancing ensures the same number of each atom on both sides of the equation, in accordance with the Law of Conservation of Mass.

• Steps:

  1. Write the unbalanced equation.

  2. Count atoms of each element on both sides.

  3. Add coefficients to balance atoms.

  4. Check your work.

• Example:

Stoichiometry

Stoichiometry involves using balanced equations to calculate quantities of reactants and products.

• Mole Ratio: Derived from coefficients in the balanced equation.

• Example: For , 2 moles of H2 react with 1 mole of O2 to produce 2 moles of H2O.

Molar Mass and Moles

Molar mass is the mass of one mole of a substance (g/mol). The mole is a counting unit for atoms, molecules, or ions.

• Calculation:

• Example: 18 g of H2O is 1 mole (molar mass = 18 g/mol).

Chemical Bonding and Molecular Structure

Types of Bonds

Chemical bonds form when atoms share or transfer electrons. The main types are ionic, pure covalent, and polar covalent bonds.

• Ionic Bonds: Electrons are transferred from a metal to a nonmetal, forming ions (e.g., NaCl).

• Pure Covalent Bonds: Electrons are shared equally between atoms (e.g., Cl2).

• Polar Covalent Bonds: Electrons are shared unequally due to differences in electronegativity (e.g., H2O).

Electronegativity and Bond Polarity

Electronegativity is an atom's ability to attract shared electrons. The difference in electronegativity determines bond type and polarity.

• Nonpolar Bond: Electronegativity difference < 0.4 (e.g., H2).

• Polar Covalent Bond: Electronegativity difference between 0.4 and 2.0 (e.g., HCl).

• Ionic Bond: Electronegativity difference > 2.0 (e.g., NaCl).

• Dipole Moment: A measure of bond polarity; occurs when there is a separation of charge.

Lewis Structures

Lewis structures are diagrams showing the arrangement of valence electrons among atoms in a molecule.

• Steps:

  1. Count total valence electrons.

  2. Draw skeletal structure with single bonds.

  3. Distribute remaining electrons to satisfy the octet rule.

• Example: Lewis structure of H2O shows two lone pairs on O and single bonds to H atoms.

Electron Geometry and Molecular Geometry

Electron geometry describes the arrangement of electron groups around a central atom, while molecular geometry describes the arrangement of atoms (excluding lone pairs).

• Example: In H2O, electron geometry is tetrahedral, molecular geometry is bent.