Chapter 3 – Periodic Properties of the Elements

Electron Configurations and Orbital Diagrams for Ions

Electron configurations describe the arrangement of electrons in an atom or ion. Orbital diagrams visually represent these configurations using arrows for electrons and boxes for orbitals.

• Electron Configuration: The notation showing the distribution of electrons among atomic orbitals (e.g., 1s2 2s2 2p6).

• For Ions: Remove electrons from the highest energy orbital (usually the outermost shell) for cations; add electrons for anions.

• Example: For Na+, remove one electron from Na: Na: 1s2 2s2 2p6 3s1 → Na+: 1s2 2s2 2p6

Magnetic Properties: Diamagnetic vs. Paramagnetic

Atoms and ions can be classified based on their magnetic properties, which depend on the presence of unpaired electrons.

• Diamagnetic: All electrons are paired; not attracted to a magnetic field.

• Paramagnetic: At least one unpaired electron; attracted to a magnetic field.

• Example: O2 is paramagnetic due to two unpaired electrons.

Shielding and Effective Nuclear Charge (Zeff)

Shielding occurs when inner electrons block the attraction between the nucleus and outer electrons, reducing the effective nuclear charge experienced by valence electrons.

• Effective Nuclear Charge (Zeff): The net positive charge experienced by valence electrons.

• Formula: where is the atomic number and is the number of shielding (core) electrons.

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

Periodic Trends: Atomic Radii, Ionic Radii, and First Ionization Energy

Periodic trends describe how certain properties of elements change across periods and down groups in 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.

• First Ionization Energy: Energy required to remove the first electron; increases across a period, decreases down a group.

• Explanation: Trends are due to changes in Zeff and electron shielding.

Identifying Elements from Ionization Energies

Successive ionization energies can be used to identify elements by observing large jumps, which indicate removal of core electrons.

• Example: A large increase between the second and third ionization energies suggests the element has two valence electrons (e.g., Mg).

Chapter 4 – Molecules and Compounds

Types of Chemical Bonds

Chemical bonds are the forces holding atoms together in compounds. The type of bond can often be predicted from the periodic table.

• Ionic Bonds: Formed between metals and nonmetals; involve transfer of electrons.

• Covalent Bonds: Formed between nonmetals; involve sharing of electrons.

• Metallic Bonds: Occur between metal atoms; electrons are delocalized.

• Example: NaCl is ionic; H2O is covalent.

Lewis Theory and the Octet Rule

Lewis theory uses dot structures to represent valence electrons and predict bonding. The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Octet Rule: Atoms are most stable with eight valence electrons (exceptions: H, He, Li, Be, B).

• Limitations: Does not explain all bonding situations (e.g., expanded octets, odd-electron species).

• Predicting Ionic Compounds: Combine ions in ratios that result in a neutral compound.

• Example: Mg2+ and Cl- combine to form MgCl2.

Naming (Nomenclature) of Compounds

Nomenclature rules allow systematic naming of ionic and molecular compounds.

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

• Molecular Compounds: Use prefixes to indicate number of atoms (e.g., CO2: carbon dioxide).

• Generating Formulas: Use the names to determine the elements and their ratios.

Polyatomic Ions

Polyatomic ions are charged species composed of two or more atoms covalently bonded.

• Common Polyatomic Ions: NO3- (nitrate), SO42- (sulfate), NH4+ (ammonium).

• Use in Naming: Name compounds containing polyatomic ions by naming the cation and then the polyatomic ion (e.g., NaNO3: sodium nitrate).

Formula and Molar Mass

The molar mass is the mass of one mole of a substance, calculated from the atomic masses of its constituent elements.

• Formula Mass: Sum of atomic masses in a formula unit.

• Molar Mass:

• Conversions:

• Example: Molar mass of H2O = 2(1.01) + 16.00 = 18.02 g/mol.

Mass Percent Composition

Mass percent composition expresses the mass of each element as a percentage of the total mass of the compound.

• Formula:

• Use: Can be used as a conversion factor in stoichiometric calculations.

Empirical and Molecular Formulas

The empirical formula gives the simplest whole-number ratio of elements in a compound, while the molecular formula gives the actual number of atoms.

• Finding Empirical Formula: Convert mass percent to moles, divide by smallest number of moles, and round to nearest whole number.

• Finding Molecular Formula: where

• Example: If empirical formula is CH2O and molar mass is 180 g/mol, molecular formula is C6H12O6.

Chapter 5 – Chemical Bonding I: Lewis Structures and Molecular Shape

Electronegativity and Bond Polarity

Electronegativity (EN) is the ability of an atom to attract electrons in a bond. Differences in EN determine bond polarity.

• EN Trend: Increases across a period, decreases down a group.

• Bond Types:

  • Non-polar covalent: EN difference < 0.5

  • Polar covalent: EN difference 0.5–1.9

  • Ionic: EN difference > 1.9

• Dipole Moment: Direction points toward more electronegative atom.

Lewis Structures

Lewis structures represent the arrangement of valence electrons in molecules. They help predict molecular shape and reactivity.

• Drawing Steps: Count valence electrons, connect atoms with single bonds, complete octets, place remaining electrons as lone pairs.

• Exceptions: Incomplete octets (e.g., B, Be), odd-electron species (radicals), expanded octets (period 3 and beyond).

Resonance Structures

Some molecules can be represented by two or more valid Lewis structures, called resonance structures. The actual molecule is a hybrid of these forms.

• Example: Ozone (O3) has two resonance structures with different locations for double bonds.

• Significance: Resonance structures show delocalization of electrons.

Formal Charge

Formal charge helps determine the most stable Lewis structure.

• Formula:

• Best Structure: The one with formal charges closest to zero and negative charges on the most electronegative atoms.

Evaluating Lewis Structures

Use formal charge and the octet rule to select the most plausible Lewis structure.

• Criteria: Minimize formal charges, obey the octet rule, place negative charges on electronegative atoms.

Bond Energy and Bond Length

Bond energy is the energy required to break a bond; bond length is the distance between nuclei of bonded atoms.

• Trends: Single bonds are longest and weakest; triple bonds are shortest and strongest.

• Example: C–C single bond > C=C double bond > C≡C triple bond (in terms of length: single > double > triple).

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the 3D shape of molecules based on electron group repulsions.

• Electron Groups: Regions of electron density (bonds or lone pairs) around a central atom.

• Common Geometries:

  • 2 groups: linear (180°)

  • 3 groups: trigonal planar (120°)

  • 4 groups: tetrahedral (109.5°)

• Drawing: Use wedge-dash notation for 3D representation.

Polarity of Molecules and Dipole Moments

The overall polarity of a molecule depends on both bond polarity and molecular geometry.

• Polar Molecule: Has a net dipole moment due to asymmetrical shape or polar bonds.

• Nonpolar Molecule: Symmetrical shape causes dipoles to cancel.

• Example: CO2 is nonpolar (linear), H2O is polar (bent).

Chapter 6 – Chemical Bonding II (Partial)

Valence Bond (VB) Theory

Valence Bond theory explains bonding as the overlap of atomic orbitals to form a covalent bond.

• Key Idea: A bond forms when two atomic orbitals overlap and a pair of electrons is shared.

• Example: H2 forms by overlap of two 1s orbitals.

Hybridization and Electron Geometry

Hybridization is the mixing of atomic orbitals to form new, equivalent hybrid orbitals that determine molecular geometry.

• Types: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral), etc.

• Prediction: Count electron groups around the atom to determine hybridization.

• Example: CH4 (methane) has four electron groups → sp3 hybridization.

Sigma (σ) and Pi (π) Bonds

Sigma and pi bonds are types of covalent bonds formed by different orbital overlaps.

• Sigma (σ) Bond: Formed by head-on overlap; all single bonds are sigma bonds.

• Pi (π) Bond: Formed by side-to-side overlap; present in double and triple bonds (in addition to a sigma bond).

• Assignment: In a double bond, one is sigma, one is pi; in a triple bond, one is sigma, two are pi.