Atoms, Periodic Table, Subatomic Particles

• Atoms are composed of positively charged protons, neutral neutrons, and negatively charged electrons.
• Isotopes are atoms of the same element that contain different numbers of neutrons.
• Ions are atoms that have gained or lost electrons, resulting in a net electrical charge.
• Cations are positively charged ions formed when atoms lose electrons.
• Anions are negatively charged ions formed when atoms gain electrons.

The Periodic Table

• Elements in the same group (column) of the periodic table exhibit similar chemical properties.
• Elements in the same period (row) of the periodic table have the same number of electron shells.
• Alkali metals (Group 1) are highly reactive, particularly with water.
• Alkaline earth metals (Group 2) are reactive, but less so than Group 1 elements.
• Halogens (Group 17) are very reactive nonmetals.
• Noble gases (Group 18) are generally inert and rarely participate in chemical reactions.

Subatomic Particles, Structures, and Configurations

• A proton has a mass of 1 atomic mass unit (amu) and a charge of +1.
• A neutron has a mass of 1 amu and a charge of 0.
• An electron has a negligible mass (approximately 0 amu) and a charge of -1.
• Electron configuration describes the arrangement of electrons in different energy levels within an atom (e.g., 2-8-8).

Periodic Trends

• Atomic radius increases as you move down a group and decreases as you move across a period.
• Ionization energy is the energy required to remove an electron from an atom.
• Ionization energy increases across a period and decreases down a group.
• Electronegativity is the ability of an atom to attract electrons in a chemical bond.
• Electronegativity increases across a period and decreases down a group.

Ionic Bonding, Structure, Names, and Formulas

• Ionic bonds form between metals and nonmetals due to the transfer of electrons.
• When naming ionic compounds, the metal name is followed by the nonmetal name with the suffix "-ide" (e.g., NaCl = Sodium Chloride).
• Roman numerals are used to indicate the charge of transition metals in ionic compounds (e.g., FeCl₃ = Iron(III) Chloride).
• Chemical formulas for ionic compounds are determined using the crisscross method to balance charges.

Metallic Bonding and Properties

• Metallic bonding involves the delocalization of electrons, creating a "sea of electrons" around metal atoms.
• Metals typically have high melting and boiling points due to strong metallic bonds.
• Metals are good conductors of heat and electricity because of the mobility of their electrons.
• Metals are malleable (can be hammered into sheets) and ductile (can be drawn into wires) due to the ability of metal atoms to slide past each other.

Covalent Bonding, Structure, Names, and Formulas

• Covalent bonds involve the sharing of electrons between nonmetal atoms.
• Types of covalent bonds include single bonds (one shared pair of electrons), double bonds, and triple bonds.
• Prefixes are used in the naming of covalent compounds to indicate the number of atoms of each element (e.g., CO₂ = Carbon dioxide).
• Lewis structures are diagrams that show the bonding and lone pairs of electrons in a molecule.

Molecular Polarity

• Nonpolar molecules have an even distribution of electron density (e.g., CO₂).
• Polar molecules have an uneven distribution of electron density (e.g., H₂O).
• Molecular polarity can be determined by examining electronegativity differences between atoms and the symmetry of the molecule. Asymmetrical molecules are typically polar.

Intermolecular Forces (IMFs)

• London dispersion forces are weak forces present in all molecules.
• Dipole-dipole forces occur between polar molecules.
• Hydrogen bonding is a relatively strong intermolecular force that occurs in molecules containing hydrogen bonded to fluorine (F), oxygen (O), or nitrogen (N).