General Chemistry IA: Atomic Structure, Bonding, and Molecular Properties – Study Guide

Atomic Structure and Electron Configuration

Quantum Numbers and Electron Arrangement

Understanding atomic structure requires knowledge of quantum numbers, electron configurations, and periodic trends. Quantum numbers describe the properties of atomic orbitals and the electrons in them.

• Principal Quantum Number (n): Indicates the energy level (shell) of an electron. Values: n = 1, 2, 3, ...

• Angular Momentum Quantum Number (l): Defines the shape of the orbital. Values: l = 0 (s), 1 (p), 2 (d), 3 (f)

• Magnetic Quantum Number (ml): Specifies the orientation of the orbital. Values: -l to +l

• Spin Quantum Number (ms): Indicates the spin direction of the electron. Values: +1/2 or -1/2

Example: The set n=5, l=2, ml=1, ms=+1/2 describes an electron in a 5d orbital.

Electron Configuration: The arrangement of electrons in an atom's orbitals, often abbreviated using noble gas notation.

• Example: Se: [Ar] 4s2 3d10 4p4

Maximum Number of Electrons in a Shell: $2n^2$

Periodic Table Trends: Elements in the same group have similar chemical properties due to similar valence electron configurations.

Periodic Table and Periodic Trends

Ionization Energy, Electron Affinity, and Atomic Radius

The periodic table organizes elements by increasing atomic number and groups elements with similar properties. Key periodic trends include:

• Ionization Energy: The energy required to remove an electron from an atom. Increases across a period, decreases down a group.

• Electron Affinity: The energy change when an atom gains an electron. Generally becomes more negative across a period.

• Atomic Radius: The size of an atom. Decreases across a period, increases down a group.

Example: First ionization energy order for Na, Mg, P, Cl: Cl > P > Mg > Na

Electromagnetic Radiation and Photons

Energy of Photons and Atomic Spectra

Atoms absorb and emit electromagnetic radiation, which can be described by the energy of photons.

• Energy of a Photon: $E = h \nu = \frac{hc}{\lambda}$

• Speed of Light (c): $3.00 \times 10^8$ m/s

• Planck's Constant (h): $6.63 \times 10^{-34}$ J·s

Example: Calculate the energy of a photon with wavelength 650 nm:

• $E = \frac{6.63 \times 10^{-34} \times 3.00 \times 10^8}{650 \times 10^{-9}}$

Relationship between Wavelength and Energy: As wavelength increases, energy decreases.

Chemical Bonding and Molecular Structure

Bond Types, Polarity, and Molecular Geometry

Chemical bonds form between atoms to create molecules. The type and strength of bonds affect molecular properties.

• Covalent Bonds: Atoms share electrons. Can be single, double, or triple bonds.

• Ionic Bonds: Atoms transfer electrons, forming ions.

• Bond Polarity: Determined by the difference in electronegativity between atoms. The most polar bond has the greatest difference.

• Bond Energy: The energy required to break a bond. Higher bond order generally means higher bond energy and shorter bond length.

Example: The C–F bond is more polar than C–H or C–O due to the high electronegativity of fluorine.

Octet Rule: Most atoms form bonds to achieve eight electrons in their valence shell. Exceptions include molecules like BF3 and SF6.

Molecular Geometry and VSEPR Theory

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

• Electronic Shape: Based on all electron domains (bonding and lone pairs).

• Molecular Shape: Based only on atoms (ignoring lone pairs).

• Bond Angles: Determined by the number of electron domains.

Example Table:

Molecule	Lewis Structure	Electronic Shape Name	Molecular Shape Name	Bond Angle(s)	Polarity
CCl4	Tetrahedral, all single bonds	Tetrahedral	Tetrahedral	109.5°	Nonpolar
NH3	Trigonal pyramidal, one lone pair	Tetrahedral	Trigonal pyramidal	107°	Polar
SeF4	See-saw, one lone pair	Trigonal bipyramidal	See-saw	~90°, ~120°, ~180°	Polar
XeF2	Linear, three lone pairs	Trigonal bipyramidal	Linear	180°	Nonpolar

Additional info: Table entries inferred from standard VSEPR theory and molecular geometry.

Molecular Orbital Theory

Bond Order, Magnetism, and Bond Length

Molecular orbital theory explains bonding in molecules by combining atomic orbitals to form molecular orbitals.

• Bond Order: $(\text{Bond Order}) = \frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2}$

• Paramagnetic: Molecules with unpaired electrons are attracted to magnetic fields.

• Diamagnetic: Molecules with all electrons paired are repelled by magnetic fields.

• Bond Length: Higher bond order generally means shorter bond length.

Example: N2 has a bond order of 3 (triple bond) and is diamagnetic.

Bond Energies and Thermochemistry

Calculating Enthalpy Changes

The enthalpy change ($\Delta H$) of a reaction can be estimated using bond energies:

• $\Delta H = \sum \text{Bond energies of bonds broken} - \sum \text{Bond energies of bonds formed}$

Example: For the reaction $\text{H–C≡C–H} + 4\text{O}_2 \rightarrow 3\text{CO}_2 + 2\text{H}_2\text{O}$, use the bond energies from the provided table to calculate $\Delta H$.

Exothermic vs. Endothermic: If $\Delta H$ is negative, the reaction is exothermic (releases heat); if positive, endothermic (absorbs heat).

Lattice Energy and Born-Haber Cycle

Lattice energy is the energy released when ions form a crystalline solid. The Born-Haber cycle uses enthalpy changes to calculate lattice energy.

• Steps: Sublimation, ionization, electron affinity, bond energy, and formation enthalpy.

• Example Equation: $\text{Ca}(s) + \text{Cl}_2(g) \rightarrow \text{CaCl}_2(s)$

Additional info: The Born-Haber cycle is a standard method for calculating lattice energies in ionic compounds.

Lewis Structures, Resonance, and Formal Charge

Drawing and Evaluating Lewis Structures

Lewis structures represent the arrangement of electrons in a molecule. Resonance structures show delocalization of electrons, and formal charge helps identify the most stable structure.

• Formal Charge: $\text{FC} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$

• Resonance: Multiple valid Lewis structures for a molecule or ion.

Example: For CNO-, draw all resonance structures and calculate formal charges to identify the best structure.

Hybridization and Bonding in Organic Molecules

Hybrid Orbitals and Sigma/Pi Bonds

Atoms in molecules may hybridize their orbitals to form bonds. Sigma (σ) bonds are single bonds formed by head-on overlap; pi (π) bonds are formed by side-on overlap in double/triple bonds.

• sp3 Hybridization: Tetrahedral geometry, 4 sigma bonds.

• sp2 Hybridization: Trigonal planar geometry, 3 sigma bonds and 1 pi bond.

• sp Hybridization: Linear geometry, 2 sigma bonds and 2 pi bonds.

Example: In cysteine, identify the hybridization of C, O, N atoms and the overlapping orbitals forming sigma and pi bonds.

Bond Energies Table

Single and Multiple Bond Energies

Bond energies are used to estimate reaction enthalpies and compare bond strengths.

Bond	Bond Energy (kJ/mol)
C–H	413
O=O	495
C≡C	839
N≡N	941
O–H	463
F–F	155

Additional info: Table entries selected for relevance to common reactions and bond comparisons.

Molecular Orbital Diagrams

Second Period Diatomic Molecules

Molecular orbital diagrams show the relative energies of molecular orbitals formed from atomic orbitals. For second period diatomic molecules, the ordering of σ and π orbitals changes across the period.

• For B2, C2, N2: π2p orbitals are lower in energy than σ2p.

• For O2, F2: σ2p is lower than π2p.

Additional info: This affects bond order and magnetic properties.

Summary Table: Key Concepts

Concept	Definition	Example/Application
Quantum Numbers	Describe electron properties in atoms	n=2, l=1, ml=0, ms=+1/2 (2p electron)
Bond Order	Number of chemical bonds between atoms	N2 bond order = 3
Hybridization	Mixing of atomic orbitals to form new orbitals	sp3 in methane (CH4)
Resonance	Delocalization of electrons in molecules	NO3- ion