Chapter 3: Periodic Properties of the Elements

Effective Nuclear Charge

The effective nuclear charge () is the net positive charge experienced by an electron in a multi-electron atom. It accounts for both the attraction to the nucleus and the repulsion from other electrons.

• Definition: , where is the atomic number and is the shielding constant.

• Trend: increases across a period and slightly increases down a group.

• Application: Explains atomic size and ionization energy trends.

Trends in Atomic and Ionic Radius

Atomic and ionic radii describe the size of atoms and ions, which vary predictably across the periodic table.

• Atomic Radius: Decreases across a period (left to right), increases down a group.

• Ionic Radius: Cations are smaller than their parent atoms; anions are larger.

• Example: is smaller than , is larger than .

Trends in Ionization Energy and Electron Affinity

Ionization energy is the energy required to remove an electron from an atom. Electron affinity is the energy change when an atom gains an electron.

• Ionization Energy: Increases across a period, decreases down a group.

• Electron Affinity: Becomes more negative across a period (especially for nonmetals).

• Application: Used to predict chemical reactivity.

Electron Configuration of Atoms and Ions

Electron configuration describes the arrangement of electrons in an atom or ion.

• Notation: Use the Aufbau principle, Pauli exclusion principle, and Hund’s rule.

• Example: ;

Hund’s Rule and Pauli Exclusion Principle

These principles govern electron arrangement in orbitals.

• Hund’s Rule: Electrons fill degenerate orbitals singly before pairing.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

Magnetic Properties: Paramagnetic vs. Diamagnetic

Magnetic properties depend on electron configuration.

• Paramagnetic: Atoms/ions with unpaired electrons; attracted to magnetic fields.

• Diamagnetic: All electrons paired; weakly repelled by magnetic fields.

• Example: is paramagnetic; is diamagnetic.

Valence Electrons

Valence electrons are the outermost electrons involved in chemical bonding.

• Distinguishing: Valence electrons are those in the highest principal energy level.

• Application: Used to predict chemical reactivity and bonding.

Chapter 4: Molecules and Compounds

Ionic and Covalent Compounds

Compounds are classified as ionic or covalent based on the nature of the bonding between atoms.

• Ionic Compounds: Formed from metals and nonmetals; involve transfer of electrons.

• Covalent Compounds: Formed from nonmetals; involve sharing of electrons.

• Example: (ionic), (covalent)

Lewis Dot Structures

Lewis dot structures represent valence electrons as dots around chemical symbols.

• Purpose: Visualize bonding and lone pairs in molecules.

• Example: : O has two lone pairs and forms two bonds with H.

Naming Ionic and Binary Molecular Compounds

Naming conventions differ for ionic and molecular compounds.

• Ionic Compounds: Name cation first, then anion (e.g., = sodium chloride).

• Binary Molecular Compounds: Use prefixes to indicate number of atoms (e.g., = carbon dioxide).

Polyatomic Ions: Formulas and Charges

Polyatomic ions are groups of atoms with a net charge.

• Common Polyatomic Ions:

Molar Mass and Percent Composition

Molar mass is the mass of one mole of a substance. Percent composition is the percentage by mass of each element in a compound.

• Calculate Molar Mass: Sum atomic masses of all atoms in the formula.

• Percent Composition:

• Example: For , molar mass = g/mol.

Conversions: Moles, Grams, Molecules

Stoichiometric calculations often require conversion between moles, grams, and number of molecules.

• Moles to Grams:

• Moles to Molecules:

Hydrates: Number of Moles of Water

Hydrates are compounds that contain water molecules within their crystal structure.

• Determination: Use mass loss upon heating to calculate moles of water.

• Example: contains 5 moles of water per mole of salt.

Empirical Formula from Combustion Analysis

Combustion analysis is used to determine the empirical formula of a compound.

• Steps:

  1. Measure mass of and produced.

  2. Calculate moles of C and H.

  3. Determine moles of O by difference.

  4. Find simplest whole-number ratio.

• Example: A sample produces 2.20 g and 0.90 g ; calculate empirical formula.