General Chemistry Core Concepts: Atomic Structure, Bonding, and Molecular Geometry

Atomic Structure and Properties

Density

Density is a fundamental physical property defined as mass per unit volume. It is commonly used to characterize substances and can be calculated using the formula:

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Units: g/cm3, kg/m3

• Example: Water has a density of 1.00 g/cm3 at 4°C.

Conversions

Unit conversions are essential in chemistry for expressing measurements in different units.

• Key conversions: mass (g, kg), volume (mL, L), length (cm, m), energy (J, kJ)

• Example: 1 L = 1000 mL

Element Symbols and Atomic Structure

Each element is represented by a unique symbol and is characterized by its atomic number (number of protons), electrons, and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Ions: Atoms or molecules with a net electric charge due to loss or gain of electrons.

• Neutrals: Atoms with equal numbers of protons and electrons.

• Example: $^{12}\text{C}$ and $^{13}\text{C}$ are isotopes of carbon.

Common/Stable Ions of Elements

Many elements form ions with characteristic charges.

• Example: Na+, Cl-, Mg2+

Basic Data: Mass, Charge, Distances

Fundamental constants and properties are crucial for calculations.

• Electron mass: $9.11 \times 10^{-31}$ kg

• Proton mass: $1.67 \times 10^{-27}$ kg

• Elementary charge: $1.60 \times 10^{-19}$ C

• Atomic radii: Typically 0.1–0.3 nm

Basic Experiments: Atom and Ion Sizes Using the Periodic Table (PT)

Atomic and ionic sizes can be inferred from periodic trends.

• Atomic radius: Decreases across a period, increases down a group.

• Ion size: Cations are smaller, anions are larger than parent atoms.

Electronegativity Using PT

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond.

• Trend: Increases across a period, decreases down a group.

• Example: Fluorine is the most electronegative element.

Ionization Potentials (1st, 2nd...) Trends Using PT

Ionization energy is the energy required to remove an electron from an atom.

• Trend: Increases across a period, decreases down a group.

• Successive ionization energies: Each subsequent electron requires more energy to remove.

Reading Data from PT: Atomic Mass, Mass Number, Mole, Molar Mass

The periodic table provides atomic mass and other key data for calculations.

• Atomic mass: Weighted average of isotopes.

• Mass number: Sum of protons and neutrons.

• Mole: $1 \text{ mole} = 6.022 \times 10^{23}$ particles (Avogadro's number)

• Molar mass: Mass of one mole of a substance (g/mol)

Atom/Object Mass → Mole → Mass and Similar Calculations

Conversions between mass, moles, and number of particles are fundamental in stoichiometry.

• Example: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

Electronic Structure and Bonding

Valence Electrons

Valence electrons are the outermost electrons involved in chemical bonding.

• Determines: Reactivity and bonding behavior

• Example: Oxygen has 6 valence electrons.

Effective Charge (Screening)

Effective nuclear charge ($Z_{eff}$) is the net positive charge experienced by valence electrons.

• Formula: $Z_{eff} = Z - S$ (where $Z$ is atomic number, $S$ is shielding constant)

Electron Configuration (Condensed)

Electron configuration describes the arrangement of electrons in an atom.

• Condensed notation: Uses noble gas core (e.g., [Ne] 3s2 3p1)

Isoelectronic Atoms and Ions

Isoelectronic species have the same number of electrons.

• Example: Na+ and Ne are isoelectronic.

Unpaired Electrons (Spin)

Unpaired electrons contribute to magnetic properties.

• Paramagnetic: Atoms with unpaired electrons

• Diamagnetic: All electrons paired

Quantum Numbers (n, l, ml, s)

Quantum numbers describe the properties of atomic orbitals and electrons.

• n: Principal quantum number (energy level)

• l: Angular momentum quantum number (orbital shape)

• ml: Magnetic quantum number (orbital orientation)

• s: Spin quantum number (+1/2 or -1/2)

Electron Configuration Boxes: Hund's, Pauli, Aufbau Rules

Rules for filling electron orbitals:

• Hund's Rule: Maximize unpaired electrons in degenerate orbitals.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Aufbau Principle: Electrons fill lowest energy orbitals first.

Chemical Bonding and Molecular Structure

Ionic and Covalent Bonds

Chemical bonds form between atoms to create compounds.

• Ionic bond: Transfer of electrons between metals and nonmetals.

• Covalent bond: Sharing of electrons between nonmetals.

Energy of Photon

The energy of a photon is related to its frequency and wavelength.

• Formula: $E = h\nu$

• Where: $h$ is Planck's constant, $\nu$ is frequency

Compounds and Molecules: Naming and Formulas

Compounds are named according to systematic rules.

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

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

Lattice Energy

Lattice energy is the energy released when ions form a crystalline lattice.

• Formula: $U = k \frac{Q_1 Q_2}{r}$ (where $Q_1$, $Q_2$ are ion charges, $r$ is distance)

Lewis Rules and Structures

Lewis structures represent valence electrons and bonding in molecules.

• Octet rule: Atoms tend to have 8 electrons in their valence shell.

• Exceptions: Less than 8 (e.g., H, B), more than 8 (expanded octet)

Resonance

Some molecules have multiple valid Lewis structures (resonance forms).

• Example: Ozone (O3)

Electron Domain and VSEPR Rule

Electron domains determine molecular geometry using the Valence Shell Electron Pair Repulsion (VSEPR) model.

• Electron domain: Region where electrons are likely to be found (bonding or lone pairs)

• VSEPR: Predicts shapes based on repulsion between electron domains

Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms.

• Examples: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral

Polarity: Polar and Nonpolar Bonds/Molecules

Polarity depends on electronegativity differences and molecular geometry.

• Polar bond: Unequal sharing of electrons

• Nonpolar bond: Equal sharing of electrons

• Polar molecule: Has a net dipole moment

Formal Charges and Best Lewis Structure

Formal charge helps determine the most stable Lewis structure.

• Formula: $\text{Formal charge} = \text{Valence electrons} - (\text{Lone electrons} + \frac{1}{2}\text{Bonding electrons})$

Hybridization (sp, sp2, sp3)

Hybridization explains molecular shapes and bond angles.

• sp: Linear (180°)

• sp2: Trigonal planar (120°)

• sp3: Tetrahedral (109.5°)

Bond Angles

Bond angles are determined by molecular geometry and hybridization.

• Example: Tetrahedral: 109.5°, Trigonal planar: 120°, Linear: 180°

Sigma and Pi Orbitals

Sigma (σ) and pi (π) bonds are types of covalent bonds formed by orbital overlap.

• Sigma bond: End-to-end overlap, single bonds

• Pi bond: Side-to-side overlap, found in double/triple bonds

Molecular Orbitals (MO)

Molecular orbital theory describes bonding using orbitals that extend over the entire molecule.

• Bonding MO: Lower energy, increased electron density between nuclei

• Antibonding MO: Higher energy, decreased electron density between nuclei

Additional info:

• Some topics were inferred and expanded for completeness, such as quantum numbers and molecular orbital theory.