Introduction: Matter, Energy, and Measurement

Definitions and Properties of Matter

Matter is defined as anything that occupies space and has mass. Understanding the basic properties and classifications of matter is fundamental to chemistry.

• Atom: The smallest unit into which an element can be divided without losing its identity.

• Element: A substance that cannot be broken down into simpler components.

• Compound: A substance composed of two or more elements chemically combined.

• Physical Change: Alters the state or appearance of matter without changing its composition.

• Chemical Change: Alters the composition of matter, resulting in the rearrangement of atoms.

• Example: Flammability is a chemical change.

Measurement and Significant Figures

Accurate measurement and proper use of significant figures are essential in chemical calculations.

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Example: The number 1.050 × 109 has 4 significant figures.

• Example: 0.00253 has 3 significant figures.

• Unit Conversion: 6.00 inches = 15.2 cm.

• Example Calculation: (1.3 × 103)(5.724 × 104) = 7.4 × 107 (rounded to correct sig figs).

Accuracy vs. Precision

Accuracy refers to how close a measurement is to the true value, while precision refers to how close repeated measurements are to each other.

• Example: The image below shows a set of measurements that are precise but not necessarily accurate.

Atoms, Molecules, and Ions

Atomic Structure and Mass

Atoms consist of protons, neutrons, and electrons. The atomic mass is calculated as the sum of protons and neutrons.

• Atomic Mass Equation:

• Example: An oxide ion (O2–) has 8 protons and 10 electrons.

• Neutrons: Have no charge.

• Example: One atom of 79Br has 44 neutrons.

Diatomic Elements and Molecular Compounds

Certain elements exist naturally as diatomic molecules, and molecular compounds are typically composed of nonmetals.

• Diatomic Elements: Br2, I2, N2, Cl2, H2, O2, F2

• Molecular Compounds: Usually comprised of nonmetals.

• States of Nonmetals: Solid, liquid, or gas.

Structural and Empirical Formulas

Formulas represent the composition and structure of compounds.

• Molecular Formula: Shows the exact number of atoms of each element (e.g., dinitrogen tetroxide is N2O4).

• Empirical Formula: Shows the simplest whole-number ratio (e.g., C6H12O6 has empirical formula CH2O).

• Structural Formula: Shows the bonding arrangement of atoms.

Common Ions and Nomenclature

Chemical nomenclature is used to name ions and compounds.

• Example: Calcium Nitrate is Ca(NO3)2.

• Example: Acetic acid is CH3COOH.

• Example: Iron(III) Chlorate is Fe(ClO3)3.

• Oxoanion of Sulfur with -2 charge: Sulfite ion.

Chemical Reactions and Reaction Stoichiometry

Types of Chemical Reactions

Chemical reactions are classified based on the changes occurring in reactants and products.

• Combination Reaction: Two or more substances combine to form one product.

• Decomposition Reaction: A single compound breaks down into two or more products.

• Combustion Reaction: A substance reacts with oxygen, releasing energy and forming products like CO2 and H2O.

Balancing Chemical Equations

Balancing equations ensures the conservation of mass and atoms.

• Example: 6 Li(s) + 1 N2(g) → 2 Li3N(s)

• Example: 2 N2H4 + N2O4 → 3 N2 + 4 H2O

• Example: 2 C3H10(g) + 11 O2(g) → 6 CO2(g) + 10 H2O(g)

Stoichiometry: Mole Calculations

Stoichiometry involves calculating the quantities of reactants and products in chemical reactions using balanced equations.

• Subscripts: Indicate the number of atoms of each element.

• Superscripts: Indicate the charge of ions.

Example: KClO3 Decomposition

• 1.65 mol KClO3 produces 2.48 mol O2

• 3.50 mol KCl needed requires 3.50 mol KClO3

• 2.73 mol KClO3 produces 2.73 mol KCl

Example: Molecule and Atom Calculations

• Calculate molecules in 1.058 g H2O: molecules

Example: Comparing Molecules in Samples

• 10.0 g O2 contains more molecules than 50.0 g I2

Molar Mass Calculations

Molar mass is the mass of one mole of a substance, used for converting between grams and moles.

• Example: Molar mass of PbSO4 is 303.3 g/mol

Stoichiometry: Mass Calculations in Reactions

Mass calculations use balanced equations and molar masses to determine the amount of product or reactant.

• Example: 4 Fe + 3 O2 → 2 Fe2O3

• 61.0 g Fe2O3 produced from 42.7 g Fe

• 56.6 g Fe2O3 produced from 17.0 g O2

• 53.7 g O2 needed to react with 125 g Fe

Example: Combustion of Butane

• 303 g CO2 produced from 100 g butane

• 358 g O2 needed for 100 g butane

• 2.33 g H2O produced from 5.38 g O2

Limiting Reactant and Yield

The limiting reactant determines the maximum amount of product that can be formed in a reaction.

• Example: In the reaction 2 Al(s) + Fe2O3(s) → 2 Fe(l) + Al2O3(s), aluminum is the limiting reactant and 5.4 moles of iron can be formed.

Additional info:

• Chalcogens are found in Group 6 (also known as Group 16).

• Noble gases are inert due to their full valence electron shells.