Atoms and Elements

Atomic Structure and Isotopes

Atoms are the fundamental units of matter, composed of protons, neutrons, and electrons. Understanding their structure is essential for describing elements and their isotopes.

• Atom: The smallest unit of an element, consisting of a nucleus (protons and neutrons) and electrons in orbitals.

• Atomic Number (Z): The number of protons in the nucleus; defines the element.

• Mass Number (A): The total number of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element (same Z) with different numbers of neutrons (different A).

• Ions: Atoms or molecules that have gained or lost electrons, resulting in a net charge.

Example: Carbon-12 and Carbon-14 are isotopes of carbon. Both have 6 protons, but Carbon-12 has 6 neutrons, while Carbon-14 has 8 neutrons.

Isotope Symbol: (where X is the element symbol, A is mass number, Z is atomic number)

The Periodic Table: Arrangement and Classification

The periodic table organizes elements by increasing atomic number and similar chemical properties.

• Groups: Vertical columns; elements in the same group have similar properties.

• Periods: Horizontal rows; properties change progressively across a period.

• Metals, Nonmetals, Metalloids: Metals are typically on the left and center, nonmetals on the right, and metalloids border the staircase line.

• Main Group Elements: Groups 1, 2, and 13-18.

• Transition Elements: Groups 3-12.

• Special Groups:

  • Alkali Metals: Group 1

  • Alkaline Earth Metals: Group 2

  • Halogens: Group 17

  • Noble Gases: Group 18

Example: Sodium (Na) is an alkali metal in Group 1, Period 3.

Predicting Ion Charges

The charge of ions formed by elements can often be predicted from their group position.

• Group 1: +1 charge (e.g., Na+)

• Group 2: +2 charge (e.g., Ca2+)

• Group 17: -1 charge (e.g., Cl-)

• Group 16: -2 charge (e.g., O2-)

• Transition metals: Variable charges

Calculating Atomic Mass

The atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.

• Formula:

Example: If chlorine has two isotopes, Cl-35 (75.8%, mass = 34.97 u) and Cl-37 (24.2%, mass = 36.97 u):

• u

Mole Conversions

The mole is a counting unit in chemistry, relating mass, number of particles, and volume (for gases).

• Avogadro's Number: particles/mol

• Conversions:

  • Grams ↔ Moles:

  • Moles ↔ Number of Atoms:

Example: 12.0 g of carbon contains mol, or atoms.

Molecules, Compounds, and Chemical Equations

Chemical Bonds and Types of Compounds

Chemical bonds hold atoms together in compounds. Compounds can be classified as molecular or ionic.

• Ionic Compounds: Formed from metals and nonmetals; consist of cations and anions held by electrostatic forces.

• Molecular Compounds: Formed from nonmetals; consist of molecules held by covalent bonds.

Example: NaCl is ionic; H2O is molecular.

Writing Formulas and Names

• Ionic Compounds: Name cation first, then anion (e.g., NaCl: sodium chloride).

• Molecular Compounds: Use prefixes to indicate number of atoms (e.g., CO2: carbon dioxide).

• Writing Formulas: Use charges to balance ionic compounds; use prefixes for molecular compounds.

Formula Mass and Mole Calculations

• Formula Mass: Sum of atomic masses of all atoms in a formula unit.

• Formula:

• Grams ↔ Moles:

Example: Formula mass of H2O: g/mol

Mass Percent Composition

• Definition: The percent by mass of each element in a compound.

• Formula:

Example: In H2O, mass % of H:

Empirical and Molecular Formulas

• Empirical Formula: Simplest whole-number ratio of atoms in a compound.

• Molecular Formula: Actual number of atoms of each element in a molecule.

• Determination: Use mass or percent composition to find moles of each element, then ratio.

• Combustion Analysis: Used to determine empirical formula from combustion products (CO2, H2O).

Example: A compound with 40% C, 6.7% H, 53.3% O has empirical formula CH2O.

Writing and Balancing Chemical Equations

• Balanced Equation: Same number of each atom on both sides; obeys the law of conservation of mass.

• Steps:

  1. Write correct formulas for reactants and products.

  2. Balance atoms one element at a time using coefficients.

  3. Check that all atoms are balanced.

Example:

Chemical Quantities and Aqueous Reactions

Stoichiometry and Chemical Calculations

Stoichiometry involves quantitative relationships between reactants and products in a chemical reaction.

• Using Balanced Equations: Coefficients indicate mole ratios for reactants and products.

• Stoichiometric Calculations: Use mole ratios to convert between amounts of substances.

• General Steps:

  1. Convert given quantity to moles.

  2. Use mole ratio from balanced equation.

  3. Convert moles to desired units (grams, molecules, etc.).

Example: How many grams of CO2 are produced from 16.0 g CH4?

• 1. Convert 16.0 g CH4 to moles: mol

• 2. Use mole ratio (1:1): 1.00 mol CH4 produces 1.00 mol CO2

• 3. Convert moles CO2 to grams: g CO2

Additional info: Metric-non-metric conversions and periodic table usage are assumed as standard skills for these topics.