Atoms, Ions, Compounds, and Nomenclature: Exam 2 Study Guide

Atoms and Atomic Structure

Size, Mass, and Structure of Atoms

• Atoms are the basic units of matter, composed of a nucleus (protons and neutrons) and electrons.

• The nucleus contains most of the atom's mass but occupies very little volume.

• The electron cloud takes up most of the atom's volume, where chemical interactions occur.

• Chemistry primarily happens in the outer regions (valence shell) where electrons are found.

Atomic Number, Mass Number, and Atomic Mass

• Atomic number (Z): Number of protons in the nucleus; defines the element.

• Mass number (A): Total number of protons and neutrons in the nucleus.

• Atomic mass: Weighted average mass of all isotopes of an element (measured in atomic mass units, amu).

• Atomic mass is not always a whole number, unlike mass number.

Isotopes

• Isotopes are atoms of the same element with different numbers of neutrons.

• Isotopic notation: $^A_Z\text{X}$ where X is the element symbol, A is the mass number, and Z is the atomic number.

• Number of neutrons = A - Z.

• Number of electrons = number of protons (for neutral atoms).

• Example: $^{14}_6\text{C}$ has 6 protons, 8 neutrons, and 6 electrons.

Valence Electrons and Shells

• Valence electrons: Electrons in the outermost shell; involved in chemical bonding.

• Valence shell: The highest energy level containing electrons in an atom.

• Number of valence electrons can often be determined from the group number in the periodic table (for main group elements).

Ions and Ionic Compounds

Ions: Cations and Anions

• Ions are charged particles formed when atoms gain or lose electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Example: Na loses one electron to become Na+ (cation); Cl gains one electron to become Cl- (anion).

Polyatomic Ions to Memorize

Name	Formula
Acetate	CH3COO-
Hydroxide	OH-
Ammonium	NH4+
Nitrate	NO3-
Carbonate	CO32-
Peroxide	O22-
Chromate	CrO42-
Phosphate	PO43-
Cyanide	CN-
Sulfate	SO42-
Hydrogen carbonate (bicarbonate)	HCO3-
Perchlorate	ClO4-

Ionization Energy

• Ionization energy: The energy required to remove an electron from a gaseous atom.

• Related to the attractive force between the nucleus (protons) and electrons.

• Factors influencing attraction: nuclear charge, distance from nucleus, electron shielding.

Ionic Bonds and Compounds

• Ionic bond: Electrostatic attraction between oppositely charged ions (cations and anions).

• Ionic compounds: Composed of a lattice of ions; have high melting points and conduct electricity when molten or dissolved in water.

• Formula unit: The simplest ratio of ions in an ionic compound (e.g., NaCl).

• To write formulas: balance charges so the total positive and negative charges are equal.

• Example: For Mg2+ and Cl-, the formula is MgCl2.

Covalent Compounds and Molecular Structure

Covalent Bonds and Compounds

• Covalent bond: Sharing of electron pairs between atoms.

• Covalent compounds: Exist as discrete molecules (e.g., H2O, CO2).

• Physical properties differ from ionic compounds (lower melting/boiling points, do not conduct electricity).

Comparison: Ionic vs. Covalent Compounds

• Ionic compounds: Extended lattice structure, formed by transfer of electrons, held by electrostatic forces.

• Covalent compounds: Discrete molecules, formed by sharing electrons, held by shared electron pairs.

Nomenclature (Naming Compounds)

Naming Ionic Compounds

• Type I cations: Metals with only one possible charge (e.g., Na+, Ca2+); name is just the metal name.

• Type II cations: Metals with more than one possible charge (e.g., Fe2+, Fe3+); use Roman numerals to indicate charge (e.g., iron(II) chloride).

• No prefixes are used in ionic compound names.

• Order: cation name first, then anion name (with -ide ending for monatomic anions).

• For polyatomic ions, use the ion name as is (e.g., sodium sulfate).

Naming Covalent Compounds

• Use prefixes to indicate the number of each atom (mono-, di-, tri-, tetra-, etc.).

• First element keeps its name; second element gets -ide ending.

• Example: CO2 is carbon dioxide; PCl5 is phosphorus pentachloride.

Writing Formulas from Names and Vice Versa

• For ionic compounds: determine the ions and their charges, balance to get a neutral formula.

• For covalent compounds: use prefixes to determine the number of each atom.

• Know the common ions and polyatomic ions for formula writing.

Classification of Matter

Pure Substances vs. Mixtures

• Pure substance: Matter with a fixed composition (elements or compounds).

• Mixture: Physical blend of two or more substances; composition can vary.

• Homogeneous mixture: Uniform composition throughout (solution).

• Heterogeneous mixture: Non-uniform composition (e.g., salad, sand in water).

Elements, Compounds, and Mixtures

• Element: Substance made of one type of atom; represented by symbols (e.g., H, O, Na).

• Compound: Substance made of two or more elements chemically combined; represented by formulas (e.g., H2O).

• Mixture: Combination of elements and/or compounds not chemically bonded.

Atoms, Molecules, and Ions

• Atom: Smallest unit of an element.

• Molecule: Two or more atoms bonded together (can be same or different elements).

• Ion: Atom or molecule with a net charge due to loss or gain of electrons.

The Periodic Table

Organization and Groups

• Elements are arranged by increasing atomic number.

• Groups (columns): Elements with similar properties; main groups include alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18).

• Periods (rows): Horizontal rows; properties change progressively across a period.

• Metals, nonmetals, metalloids: Metals are on the left, nonmetals on the right, metalloids border the staircase line.

• Transition elements: Found in the center block (Groups 3-12).

• Main group elements: Groups 1, 2, and 13-18.

The Seven Diatomic Elements

• Elements that exist as molecules of two atoms: H2, N2, O2, F2, Cl2, Br2, I2.

Dimensional Analysis and Units

Dimensional Analysis

• Method for converting between units using conversion factors.

• Set up conversion so units cancel appropriately.

• Example: To convert 10 cm to meters: $10\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.10\ \text{m}$

Typical Units

• Mass: grams (g), kilograms (kg)

• Length: meters (m), centimeters (cm)

• Volume: liters (L), milliliters (mL)

Significant Figures

• Rules for determining which digits are significant in a measurement.

• Important for reporting calculated results with correct precision.

Additional info: This guide covers foundational concepts in atomic structure, ions, compounds, nomenclature, classification of matter, periodic table organization, and essential calculation skills for general chemistry. Refer to your class notes and textbook for more detailed examples and practice problems.