Atoms, Elements, and Chemical Bonding: Study Guide for GOB Chemistry

Atoms and Elements

Atomic Structure and Subatomic Particles

The structure of the atom is fundamental to understanding chemistry. Atoms consist of three main subatomic particles: protons, neutrons, and electrons.

• Protons: Positively charged particles found in the nucleus. The number of protons determines the atomic number and the identity of the element.

• Neutrons: Neutral particles also located in the nucleus. The number of neutrons, together with protons, determines the atomic mass.

• Electrons: Negatively charged particles that orbit the nucleus in energy levels (shells).

• Atomic Number (Z): The number of protons in an atom.

• Atomic Mass: The sum of protons and neutrons in the nucleus.

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

Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Equation:

$\text{Atomic Mass} = \text{Number of Protons} + \text{Number of Neutrons}$

Early Experiments and Modern Atomic Theory

Key experiments shaped our understanding of atomic structure:

• Cathode-Ray Tube: Demonstrated the existence of electrons.

• Gold Foil Experiment: Revealed the nucleus and led to the nuclear model of the atom.

Modern atomic theory describes atoms as having a dense nucleus surrounded by electrons in defined energy levels.

Periodic Table and Periodic Trends

Classification of Elements

The periodic table organizes elements by increasing atomic number and groups them based on similar properties.

• Metals: Conduct electricity, malleable, ductile, shiny.

• Nonmetals: Poor conductors, brittle, dull.

• Metalloids: Properties intermediate between metals and nonmetals.

Example: Silicon is a metalloid.

Periodic Trends

• Atomic Radius: Increases down a group, decreases across a period.

• Ionization Energy: Decreases down a group, increases across a period.

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

Electron Configuration and Quantum Numbers

Electron Configuration

Electrons occupy energy levels and subshells according to the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Shell (Principal Quantum Number, n): Indicates the main energy level.

• Subshell: s, p, d, f types within each shell.

• Orbital Diagrams: Visual representation of electron arrangement, including electron spin.

Example: Electron configuration for oxygen: $1s^2\ 2s^2\ 2p^4$

Valence Electrons and the Octet Rule

Valence electrons are the outermost electrons involved in chemical bonding. The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Exceptions: Hydrogen (2), Boron (6), expanded octets for elements in period 3 and beyond.

Chemical Bonding

Ionic and Molecular Compounds

Chemical compounds are formed by the combination of atoms through ionic or covalent bonds.

• Ionic Compounds: Formed by transfer of electrons from metals to nonmetals. Consist of cations and anions.

• Molecular (Covalent) Compounds: Formed by sharing electrons between nonmetals.

Example: Sodium chloride (NaCl) is ionic; water (H2O) is molecular.

Naming and Writing Formulas

• Naming Ionic Compounds: Name the cation first, then the anion. For transition metals, indicate charge with Roman numerals.

• Naming Molecular Compounds: Use prefixes to indicate the number of atoms (mono-, di-, tri-, etc.).

• Writing Formulas: Use element symbols and subscripts to indicate the number of each atom.

Example: Carbon dioxide: CO2; Iron(III) chloride: FeCl3

Polyatomic Ions

Polyatomic ions are groups of atoms with a charge that act as a single unit in compounds.

• Common Polyatomic Ions: Nitrate (NO3-), Sulfate (SO42-), Ammonium (NH4+), etc.

Ion Name	Formula
Nitrate	NO3-
Sulfate	SO42-
Phosphate	PO43-
Ammonium	NH4+
Hydroxide	OH-
Additional info: Other common ions include carbonate (CO32-), acetate (C2H3O2-).

Electromagnetic Radiation and Atomic Spectra

Electromagnetic Spectrum

Electromagnetic radiation includes a range of wavelengths and frequencies, from radio waves to gamma rays.

• Wavelength (λ): Distance between two peaks, measured in meters (m) or nanometers (nm).

• Frequency (ν): Number of cycles per second, measured in Hertz (Hz).

• Energy (E): Related to frequency and wavelength.

Equation:

$E = h\nu$

$c = \lambda\nu$

Where $E$ is energy, $h$ is Planck's constant, $\nu$ is frequency, $c$ is speed of light, and $\lambda$ is wavelength.

Absorption and Emission of Energy

Electrons absorb energy to move to higher energy levels and emit energy as photons when returning to lower levels. Atomic spectra are generated by these transitions.

• Photon: A quantum of electromagnetic energy.

• Atomic Spectra: Unique patterns of light emitted or absorbed by elements.

Lewis Structures and Molecular Geometry

Drawing Lewis Structures

Lewis structures represent the arrangement of electrons in molecules, showing bonds and lone pairs.

• Steps: Count valence electrons, arrange atoms, distribute electrons to satisfy the octet rule.

3D Geometry and Bond Angles

Molecular geometry describes the three-dimensional arrangement of atoms. Wedge-and-dash projections are used to represent 3D structures.

• Electron Domain Geometry: Based on the number of electron domains (bonding and lone pairs).

• Molecular Geometry: Actual shape of the molecule.

• Bond Angles: Determined by geometry (e.g., tetrahedral: 109.5°).

Geometry	Bond Angle	Example
Linear	180°	CO2
Trigonal Planar	120°	BF3
Tetrahedral	109.5°	CH4
Additional info: Bent and trigonal pyramidal geometries have bond angles less than 120° and 109.5°, respectively.

Intermolecular Forces and Polarity

Intermolecular Forces

Intermolecular forces are attractions between molecules, affecting physical properties.

• Types: London dispersion (weakest), dipole-dipole, hydrogen bonding (strongest).

Polarity

Polarity depends on the difference in electronegativity and molecular geometry.

• Polar Bonds: Unequal sharing of electrons.

• Polar Molecules: Molecules with an uneven distribution of charge.

Example: Water is a polar molecule; carbon dioxide is nonpolar.

Calculations and Applications

Average Atomic Mass

Average atomic mass is calculated using the masses and abundances of isotopes.

Equation:

$\text{Average Atomic Mass} = \sum (\text{Isotope Mass} \times \text{Fractional Abundance})$

Predicting Ionic Charges

Ionic charges can be predicted based on group number in the periodic table.

• Group 1: +1

• Group 2: +2

• Group 17: -1

• Transition Metals: Variable charges, indicated in compound names.

Units for Wavelength, Frequency, and Energy

• Wavelength: Meters (m), nanometers (nm)

• Frequency: Hertz (Hz)

• Energy: Joules (J)

Additional info:

• Lewis structures and orbital diagrams are essential for visualizing electron arrangements.

• Wedge-and-dash projections help represent 3D molecular shapes.

• Understanding periodic trends aids in predicting chemical behavior.