General Chemistry Practice Exam Study Guide

Atomic Structure and Electron Configuration

Electron Subshells and Capacity

The arrangement of electrons in atoms is described by subshells (s, p, d, f), each with a specific capacity.

• p Subshell: Can hold a maximum of 6 electrons.

• Electron configuration: Follows the Aufbau principle, filling lower energy subshells first.

• Example: The p subshell in the second energy level (2p) can hold up to 6 electrons.

Equation:

$\text{Maximum electrons in a subshell} = 2(2l + 1)$

For p subshell ($l=1$): $2(2 \times 1 + 1) = 6$

Ions and Electron Count

An ion is an atom or molecule with a net electric charge due to the loss or gain of electrons.

• Cations: Positive ions formed by losing electrons.

• Anions: Negative ions formed by gaining electrons.

• Example: An atom with 17 protons and 18 electrons is a chloride ion ($\text{Cl}^-$).

Valence Electrons

Valence electrons are the electrons in the outermost shell, crucial for chemical bonding.

• Nitrogen: Has 5 valence electrons (group 15 element).

• Determination: Group number for main group elements indicates valence electrons.

Chemical Bonding

Bond Polarity

Bond polarity arises from differences in electronegativity between bonded atoms.

• Most polar bond: O-H bond, due to large electronegativity difference between O and H.

• Polarity order: Greater difference in electronegativity = more polar bond.

Types of Chemical Bonds

• Ionic Bonds: Formed between metals and nonmetals (e.g., Na and F).

• Covalent Bonds: Shared electrons between nonmetals.

• Polar Covalent: Unequal sharing of electrons.

• Metallic Bonds: Delocalized electrons among metal atoms.

• Example: NaF is an ionic compound.

Octet Rule Exceptions

Most atoms aim for 8 electrons in their valence shell, but some molecules violate this rule.

• BF3: Boron has only 6 valence electrons in this molecule.

• Other exceptions: Molecules with odd numbers of electrons, or expanded octets (period 3 and beyond).

Oxidation Numbers and Redox Reactions

Oxidation Number Assignment

Oxidation numbers indicate the charge an atom would have if electrons were transferred completely.

• Mn in MnO2: Oxygen is -2, total for two oxygens is -4. Mn must be +4.

• Equation:

$\text{Oxidation number of Mn} + 2 \times (-2) = 0$

$\text{Oxidation number of Mn} = +4$

Redox Reactions

• Reduction: Gain of electrons.

• Oxidation: Loss of electrons.

• Species reduced: Gains electrons.

Nomenclature and Compound Naming

Transition Metal Compounds

Transition metals can have multiple oxidation states, indicated by Roman numerals in compound names.

• Fe2(SO4)3: Iron(III) sulfate.

• Systematic naming: Metal (oxidation state) + anion name.

Molecular Compounds

• Dinitrogen tetroxide: N2O4

• Prefix system: Used for nonmetal compounds (di-, tri-, tetra-, etc.).

Formal Charge and Molecular Geometry

Formal Charge Calculation

Formal charge helps determine the most stable Lewis structure.

• Formula:

$\text{Formal charge} = \text{Valence electrons} - (\text{Nonbonding electrons} + \frac{1}{2}\text{Bonding electrons})$

• NH4+: Nitrogen has a formal charge of +1.

Molecular Geometry

The shape of a molecule is determined by the arrangement of electron pairs around the central atom.

• SO3: Trigonal planar geometry.

• VSEPR Theory: Predicts molecular shapes based on electron pair repulsion.

Electrostatic Potential Maps

Electron Density Regions

Electrostatic potential maps show regions of high and low electron density in molecules.

• Highest electron density: Most negative region (often colored red or dark).

• Lowest electron density: Most positive region (often colored blue or light).

Summary Table: Key Concepts

Concept	Key Point	Example
p Subshell Capacity	Holds 6 electrons	2p6
Ion Formation	17 protons, 18 electrons = Cl-	Chloride ion
Valence Electrons	Nitrogen has 5	Group 15 element
Bond Polarity	O-H is most polar	Water molecule
Oxidation Number	Mn in MnO2 is +4	Manganese dioxide
Compound Naming	Fe2(SO4)3 is Iron(III) sulfate	Transition metal sulfate
Bond Type	Na and F form ionic bond	Sodium fluoride
Octet Rule Exception	BF3 violates octet rule	Boron trifluoride
Formal Charge	N in NH4+ is +1	Ammonium ion
Molecular Geometry	SO3 is trigonal planar	Sulfur trioxide
Electrostatic Map	Highest density = most negative region	Oxygen in water
Redox Reaction	Species reduced gains electrons	Fe3+ to Fe2+