General Chemistry: Atomic Structure, Bonding, and Periodic Trends – Study Guide

Atomic Structure and Electron Configuration

Electron Configuration of Ions

Understanding electron configurations is essential for predicting chemical behavior and properties of elements and ions.

• Electron Configuration: The arrangement of electrons in an atom or ion, typically written using the order of subshell filling (e.g., 1s2 2s2 2p6).

• Isoelectronic Species: Atoms or ions with the same number of electrons. For example, N3−, O2−, F−, and Ne are isoelectronic.

• Example: The electron configuration for Mg2+ is the same as that for Ne: $1s^2 2s^2 2p^6$.

Periodic Trends: Ionization Energy and Electron Affinity

Periodic trends help predict the reactivity and properties of elements.

• Ionization Energy (IE): The energy required to remove an electron from a gaseous atom or ion. IE increases across a period and decreases down a group.

• Electron Affinity (EA): The energy change when an electron is added to a neutral atom. EA generally becomes more negative across a period.

• Example: $\text{Na} \rightarrow \text{Na}^+ + e^-$ (IE process)

Ionic Radii

Ionic radii vary depending on the charge and the number of electrons relative to protons.

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

• Isoelectronic Series: For ions with the same number of electrons, the more positive the charge, the smaller the ion.

• Example: $\text{Mg}^{2+} < \text{Na}^+ < \text{F}^- < \text{O}^{2-}$ (increasing ionic radius)

Chemical Bonding

Ionic and Covalent Bonds

Chemical bonds form by the transfer or sharing of electrons between atoms.

• Ionic Bond: Formed by the transfer of electrons from a metal to a nonmetal, resulting in oppositely charged ions.

• Covalent Bond: Formed by the sharing of electrons between two nonmetals.

• Example: NaCl is ionic; H2O is covalent.

Electronegativity and Bond Polarity

Electronegativity differences determine bond polarity and ionic character.

• Electronegativity: The ability of an atom to attract electrons in a bond. Increases across a period, decreases down a group.

• Polar Covalent Bond: Electrons are shared unequally due to differences in electronegativity.

• Nonpolar Covalent Bond: Electrons are shared equally.

• Example: H2O has polar covalent bonds; Cl2 has nonpolar covalent bonds.

Lewis Structures and Resonance

Lewis structures represent the arrangement of electrons in molecules and ions.

• Octet Rule: Atoms tend to have eight electrons in their valence shell (exceptions exist).

• Lone Pairs: Non-bonding pairs of electrons on an atom.

• Resonance: Some molecules have more than one valid Lewis structure; the actual structure is a hybrid.

• Formal Charge: Calculated as: $\text{Formal charge} = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2}(\text{bonding electrons})$

• Example: SO2 has resonance structures with different S–O bond arrangements.

VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

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

• Common Geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Example: CO2 is linear; CH4 is tetrahedral.

Properties of Elements and Compounds

Periodic Table Trends

The periodic table organizes elements by increasing atomic number and similar properties.

• Groups: Vertical columns; elements in a group have similar chemical properties.

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

• Alkali Metals (Group 1A): Highly reactive, form +1 ions.

• Alkaline Earth Metals (Group 2A): Less reactive than alkali metals, form +2 ions.

Reactivity and Chemical Reactions

Reactivity trends can be predicted using periodic trends and knowledge of chemical bonding.

• Alkali metals react vigorously with water to form hydroxides and hydrogen gas.

• Example: $2\text{Na} + 2\text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{H}_2$

Ionic Compounds and Nomenclature

Naming Ionic Compounds

Ionic compounds are named by stating the cation first, followed by the anion.

• Monatomic Ions: Named by the element (cation) or element root + -ide (anion).

• Polyatomic Ions: Have specific names (e.g., SO42− is sulfate, SO32− is sulfite).

• Example: NaCl is sodium chloride; Na2SO4 is sodium sulfate.

Common Polyatomic Ions

Ion	Formula	Name
Thiosulfate	SO32−	Thiosulfate
Acetate	C2H3O2−	Acetate
Carbonate	CO32−	Carbonate
Chlorate	ClO3−	Chlorate

Lattice Energy and Born-Haber Cycle

Lattice Energy

Lattice energy is the energy released when gaseous ions form an ionic solid.

• Born-Haber Cycle: A thermochemical cycle used to calculate lattice energy using enthalpy changes for each step in the formation of an ionic compound.

• Example Equation: $\text{Na}(s) + \frac{1}{2}\text{Cl}_2(g) \rightarrow \text{NaCl}(s)$

• Steps include: sublimation, ionization, bond dissociation, electron affinity, and lattice formation.

Sample Table: Electronegativity Values

Element	Electronegativity
F	4.0
O	3.5
N	3.0
Cl	3.0
Br	2.8
I	2.5

Lewis Dot Structures and Formal Charge

Drawing Lewis Structures

Lewis structures show the arrangement of valence electrons around atoms in a molecule.

• Count total valence electrons.

• Arrange atoms, connect with single bonds, complete octets, and assign lone pairs.

• Check for resonance and calculate formal charges to find the most stable structure.

Formal Charge Calculation

• Use the formula: $\text{Formal charge} = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2}(\text{bonding electrons})$

• Structures with formal charges closest to zero are generally most stable.

Summary Table: Types of Chemical Bonds

Bond Type	Formation	Example
Ionic	Transfer of electrons	NaCl
Covalent	Sharing of electrons	H2O
Metallic	Delocalized electrons	Fe, Cu
Hydrogen	Attraction between H and electronegative atom	H2O (between molecules)

Additional info:

• Some questions reference specific sections (e.g., "Section 3.3 Electron Configuration of Ions"). These correspond to standard textbook topics in General Chemistry.

• Where tables or values were referenced but not fully visible, standard values and examples were provided for completeness.