Chemistry Chapter: Atomic Notation and Ions

Standard Atomic Notation

• Mass Number (A) = number of protons (Z) + number of neutrons (N)
• Symbol (X) = chemical symbol for the element
• Atomic Number (Z) = number of protons in the atom's nucleus
• Atomic Mass Unit (u) = 1/12 the mass of one carbon-12 atom (1.660 540 2 x 10^-27 kg)

Ions and Octet Rule

• Elements with a full valence shell (octet) are stable.
• Atoms that don't have a full octet tend to combine with other atoms to achieve a stable octet by sharing, losing, or gaining electrons.
• Ions are formed when atoms lose or gain electrons, creating a positive or negative charge.
• Positive ions (cations) are formed when metals lose electrons.
• Negative ions (anions) are formed when nonmetals gain electrons.

Multivalent Elements

• Some elements can form more than one type of stable ion.
• Transition metals are often multivalent.
• Different charges are often represented using Roman numerals (e.g., copper(I) and copper(II)).

Polyatomic Ions

• Ions containing more than one atom that act as a single unit.
• Students should memorize common polyatomic ions.
• Mnemonic device provided for memorization (NICK the CAMEL ate a CLAM for SUPPER in PHOENIX).

Ions in the Human Body

• About 99% of the human body is composed of oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus.
• Other elements like sulfur, chlorine, sodium, magnesium, iodine, and iron are present in smaller amounts, playing key roles.
• Many elements are dissolved in water within the body.

Isotopes, Radioisotopes, and Atomic Mass

• Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
• Isotopic abundance is the percentage of a particular isotope in a sample.
• Atomic mass is the average mass of all isotopes of an element in a naturally occurring sample.

Average Atomic Mass

• The average mass of an element considering the prevalence of the isotopes naturally.
• Relative atomic mass is the average weight of all isotopes in a normal environment on the Earth's crust.

Periodic Table and Periodic Law

• Elements in the same vertical column (group) have similar characteristics.
• Groups IA (group 1) alkali metals, IIA (Group 2) alkaline earth metals, VIIA (Group 17) halogens, and VIIIA (Group 18) noble gases are notable periodic table groups.
• Transition metals are in groups 3-12.
• Inner transition metals are lanthanides (57 -70) and actinides (89-102).

Electron Configuration and Trends

• Electron shielding influences the attractive force between electrons and the nucleus.
• More orbits (larger atoms) = less effective nuclear charge.
• More protons (smaller atoms) = stronger effective nuclear charge.
• Atomic radius increases going down a group and decreases going across a period.
• Metal reactivity increases going down a group and decreasing across a period.
• Nonmetal reactivity increases going up a group and increasing across a period.

Ionic Radius

• Cations are always smaller than the corresponding atom because the loss in electrons reduces the electron-electron repulsion.
• Anions are always larger because the gain of electrons leads to increased repulsion among electrons.

Ionization Energy

• Energy required to remove an electron from an atom or ion.
• First ionization energy (IE₁), second ionization energy (IE₂), etc.
• Removal becomes increasingly difficult for successive electrons.
• Ionization energy increases going across a period and decreases going down a group.

Electron Affinity

• Energy change that occurs when an electron is added to a neutral atom.
• Higher electron affinity means more energy is released when an electron is added.
• Electron Affinity increases across a period and decreases down a group.

Electronegativity

• Ability of an atom to attract electrons in a covalent bond.
• Electronegativity increases across a period.
• Electronegativity decreases down a group.

Polar Bonds and Molecules

• Differences in electronegativity result in bonds with unequal sharing of electrons (polar covalent).
• Polar molecules have distinct positive and negative ends due to asymmetrical charge distribution.
• Non-polar molecules have equal sharing of electrons (nonpolar covalent) and symmetrical distribution of charge.

Intermolecular Forces

• Forces of attraction between molecules.
• London dispersion forces (temporary dipoles), dipole-dipole forces (permanent dipoles).
• Hydrogen bonds (strong dipole-dipole forces when hydrogen is bonded to nitrogen, oxygen, or fluorine).

Chemical Formulas

• Binary molecular compounds - molecular formulas, and prefixes (mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca).
• Binary ionic compounds - include metal and nonmetal, find charge of each using periodic table and crossover to find subscript.

Nomenclature of Acids

• Naming Binary acids (hydro ____ acid).
• Naming Oxyacids (____ ic acid or ____ ous acid).
• Naming Oxyacids with polyatomic ions (use naming rules for polyatomic ions).

Hydrates

• Ionic compounds that contain bound water molecules in their crystal structure.
• Hydrates are often referred to as the 'water of hydration'.

Dive into the fundamentals of atomic notation, ions, and the octet rule in this quiz. Explore mass numbers, atomic symbols, and the behavior of multivalent elements. Perfect for students looking to solidify their understanding of chemical principles.