Understanding Atomic Structure

• Atomic structure refers to an atom's internal makeup, including its components and arrangement.
• Atoms comprise protons, neutrons, and electrons.
• Protons and neutrons are in the nucleus, while electrons are in orbitals around it.
• The atomic number (Z) is the number of protons, defining the element.
• Neutron number variations lead to isotopes of the same element.
• Electrons are negatively charged, orbiting the nucleus in energy levels or shells.
• Electron arrangement dictates an atom's chemical properties.

Subatomic Particles

• Protons are positively charged particles in the nucleus.
• The number of protons determines the element's atomic number.
• Neutrons are neutral particles in the nucleus.
• Neutrons contribute to an atom's mass without affecting its charge.
• Isotopes are atoms of the same element differing in neutron number.
• Electrons are negatively charged, orbiting the nucleus.
• Electrons reside in specific energy levels or shells.
• Electron arrangement determines the chemical properties of an atom.

Nucleus

• The nucleus, the atom's core, contains protons and neutrons.
• The nucleus constitutes most of the atom's mass.
• Protons and neutrons are held together by the strong nuclear force.
• The nucleus carries a positive charge due to protons.
• The identity of the element is based on the number of protons.
• The isotope of the element is affected by the number of neutrons.

Electron Configuration

• Electron configuration is how electrons arrange in an atom's energy levels and sublevels.
• Electrons occupy energy levels or shells around the nucleus, indicated by principal quantum numbers (n = 1, 2, 3,...).
• Each energy level includes sublevels or orbitals (s, p, d, f).
• The s sublevel has 1 orbital, p has 3, d has 5, and f has 7.
• Each orbital holds up to two electrons, per the Pauli Exclusion Principle.
• The Aufbau principle determines the order of filling energy levels and sublevels.
• Hund's rule specifies that electrons occupy each orbital within a sublevel before pairing.
• Shorthand notation can express electron configurations, like [He]2s²2p⁴ for oxygen.
• Valence electrons are in the outermost energy level, responsible for chemical bonding.
• Core electrons reside in inner energy levels, not involved in bonding.

Atomic Number and Mass Number

• Atomic number (Z) is the count of protons in an atom's nucleus.
• The atomic number identifies the element.
• A given element's atoms all share the same atomic number.
• The atomic number is written as a subscript to the left of the element symbol (e.g., ₁H).
• Mass number (A) is the total of protons and neutrons in an atom's nucleus.
• The mass number is written as a superscript to the left of the element symbol (e.g., ¹H).
• Neutrons are found by subtracting the atomic number from the mass number (A - Z).
• Isotopes are atoms of the same element but differ in mass number due to varying neutron numbers.
• Atomic mass is the weighted average of an element's isotopes' masses.
• It is expressed in atomic mass units (amu).
• Atomic weight is often used in place of atomic mass.

Isotopes and Average Atomic Mass

• Isotopes are atoms of the same element, with differing neutron numbers.
• Isotopes share an atomic number but have different mass numbers.
• Some isotopes are stable, while others undergo radioactive decay.
• Radioactive isotopes decay over time, releasing particles and energy.
• Average atomic mass is the weighted average mass of an element's isotopes.
• It considers each isotope's natural abundance.
• Average atomic mass is calculated as: (mass of isotope 1 × abundance of isotope 1) + (mass of isotope 2 × abundance of isotope 2) + ...
• Abundances are commonly percentages.
• The periodic table lists average atomic mass values.

Ions

• Ions are atoms or molecules with a net electrical charge due to electron gain or loss.
• Cations are positive ions formed by losing electrons.
• Metals typically form cations.
• Anions are negative ions formed by gaining electrons.
• Nonmetals typically form anions.
• An ion's charge is a superscript to the right of the element symbol (e.g., Na⁺, Cl⁻).
• The number of electrons gained or lost determines the charge's magnitude.
• Ionic compounds arise from electrostatic attraction between cations and anions.
• Polyatomic ions consist of two or more bonded atoms (e.g., SO₄²⁻, NH₄⁺).

Quantum Numbers

• Quantum numbers describe atomic orbitals and their electrons.
• Four quantum numbers exist: principal (n), angular momentum (l), magnetic (ml), and spin (ms).
• The principal quantum number (n) describes the electron's energy level or shell (n = 1, 2, 3,...).
• Higher n values denote higher energy levels.
• The angular momentum quantum number (l) defines orbital shape (l = 0, 1, 2, ..., n-1).
• l = 0 is an s orbital (spherical).
• l = 1 is a p orbital (dumbbell-shaped).
• l = 2 is a d orbital (more complex).
• l = 3 is an f orbital (even more complex).
• The magnetic quantum number (ml) indicates orbital orientation in space (ml = -l, -l+1,..., 0,..., l-1, l).
• A p orbital (l = 1) has three orientations (ml = -1, 0, 1).
• The spin quantum number (ms) describes intrinsic angular momentum, or spin.
• Electrons act as if spinning, creating a magnetic dipole moment.
• The spin quantum number is +1/2 (spin up) or -1/2 (spin down).
• Pauli Exclusion Principle: No two electrons in an atom share the same set of four quantum numbers.
• Orbitals hold a maximum of two electrons with opposing spins.

Atomic Orbitals

• High electron probability regions around the nucleus are known as atomic orbitals.
• Orbitals are defined by shape and energy.
• The shape of s orbitals is spherical.
• p orbitals are dumbbell-shaped and align along the x, y, and z axes.
• d orbitals have more intricate shapes and orientations.
• f orbitals possess even more complex shapes and orientations.
• Each orbital accommodates a maximum of two electrons with opposing spins.
• An orbital's energy depends on its principal (n) and angular momentum (l) quantum numbers.
• Orbitals with lower n and l values have lower energies.
• Orbital filling follows the Aufbau principle and Hund's rule.
• Atom bonding is determined by the shape and orientation of orbitals.

Chemical Bonding

• Chemical bonding involves atoms sharing or transferring electrons to achieve electron configuration stability.
• Ionic bonds arise from electron transfer between atoms, resulting in ion formation.
• Covalent bonds occur when atoms share electrons.
• Metallic bonds involve electron delocalization within a metal atom lattice.
• The type of bonding influences the resulting compound's properties.
• Lewis structures depict atom and electron arrangement in molecules.
• The octet rule states atoms gain, lose, or share electrons to complete their outer shell with eight electrons.
• Exceptions to the octet rule include hydrogen needing only two electrons and elements capable of expanded octets, like sulfur and phosphorus.
• Bond polarity is unequal electron sharing in a covalent bond, causing partial charges on atoms.
• Electronegativity measures an atom's capacity to attract electrons in a chemical bond.