CHEM 1770 General Chemistry I Study Notes

Contact Information:
- Help Forum or chem1770bq@iastate.edu (use "CHEM 1770" in subject line)
- Chemistry Help Center located at 1761 Gilman Hall
- Hours: MTWR 9 AM - 6 PM; F 9 AM - 1 PM
- Office Hours: MW at 12 PM - 12:50 PM in Troxel or by appointment.
- Supplemental Instruction (SI) Sessions:
- Monday 7:10 PM - 8:00 PM in Gilman 1051
- Wednesday 6:10 PM - 7:00 PM in Carver 0174
- Friday 2:15 PM - 3:05 PM in Gilman 1051
Homework: Chapter 5 due on Sunday, 10/12 at 11:59 PM.
Recitation Worksheet: Scheduled for 10/16 covering topics from the next exam, specifically Chapter 6.1-2 (Monday) and Chapter 6.2-3 (Wednesday).
Lecture Materials: Lecture notes available on Canvas, exams will be recorded.

Exam 2: Scheduled for October 15th from 8:15 PM to 9:15 PM.
- Important notes for the exam:
- Read Canvas Announcements.
- Refer to the Module Exams Information and Instructions section.
- Read the Exam Information and Instructions document thoroughly.
- Topics to review: Chapters 3, 4, and 5 (NO Chapter 6).
- Practice exams are available; note that content may vary.
- Exam Duration: 60 minutes, students should arrive early.
Supplemental Instruction (SI) exam review sessions:
- Monday, October 13 in Troxel 1001 from 7:10 PM to 9:10 PM.
Calculator Requirement: Please bring a calculator to the exam.
Academic Accommodations: Register via the Canvas Navigation Menu under the Exam Accommodation Center.

Use valence bond (VB) theory to describe molecular bonding.
Explain hybridization and its reconciliation with observed molecular shapes.
Understand multiple bond formation via hybridized and non-hybridized atomic orbitals.
Identify sigma (σ) and pi (π) bonds and their differences.
Utilize molecular orbital (MO) theory to explain covalent bonding in diatomic molecules.
Draw molecular orbital energy diagrams for diatomic molecules.
Predict bond order and magnetic properties in diatomic molecules.
Recognize molecular orbital diagrams of simple polyatomic molecules.
Awareness of evolving bonding theories.

Valence Bond Theory (VB):
- Bonds result from the overlap of atomic orbitals (AOs).
- Conditions for overlap:
  - Two half-filled AOs overlap, causing spin pairing of valence electrons.
  - One full AO overlaps an empty AO.
- Each bond is localized between bonding atoms and involves two electrons.
Molecular Orbital Theory:
- A quantum mechanical model where overlapping atomic orbitals produce new molecular orbitals.
- Electrons in this model are delocalized across the molecule.
- Energy and symmetry of orbitals are critical in both techniques.

Driving Force for Bond Creation:
- Energy released upon bond formation (example: $2 H <br>ightarrow H_2$ : $E = -458$ kJ).
- Energy consumed upon bond breakage ( $H_2 <br>ightarrow 2H$ : $E = 458$ kJ).
Bond Energy (D):
- Defined as the energy required to break a bond, always a positive value.
- Measured in kJ per mole of substance.

Bonds arise from overlapping atomic orbitals, necessitating:
- Correct spatial orientation and similar energy states of orbitals.
- VB theory asserts that bonding electrons are localized between nuclear centers.

Example for H₂:
- Each hydrogen has one valence electron in a 1s orbital.
- Overlap of two 1s orbitals results in a covalent bond.
- Characteristically, s orbitals are spherical with no preferred direction of approach.

Formed by end-to-end overlaps:
- Two s orbitals (non-directional).
- One s and one p orbital (non-directional).
- Two p orbitals aligned along the same axis.

Formed from side-to-side overlaps of p orbitals:
- Can form two π bonds between two nuclei.
- Electron density exists above and below the axis joining the atoms.

Describes how atomic orbitals combine to form new, equivalent, hybrid orbitals.
The concept reconciles observed molecular geometries with the overlap model.
- Trends in Hybridization:
- sp³: Tetrahedral geometry, 4 identical orbitals; angle of 109.5°.
- sp²: Trigonal planar geometry; angle of 120°, involves one non-hybridized p orbital.
- sp: Linear geometry, 180° angle; forms two π bonds.

Molecular orbitals are generated through the linear combination of atomic orbitals (LCAO).
- By this method, bonding orbitals (lower energy) and antibonding orbitals (higher energy) are created.

Bond order calculation:
- Given by the formula:
$ext{Bond Order} = rac{ ext{Number of bonding electrons} - ext{Number of antibonding electrons}}{2}$

Bond characteristics (Example: H₂, He₂):
- H₂: Bond order of 1.
- He₂: Bond order of 0; does not exist.

O2 is paramagnetic due to unpaired electrons, despite VB theory suggesting otherwise.

Common hybrid orbitals include sp³, sp², sp, each applicable to various molecular geometries based on the number of atomic orbitals mixed.
Hybridization Diagram Summary:
- sp³ (tetrahedral), sp² (trigonal planar), sp (linear).

Draw Lewis structures (including resonance) and assign formal charges to various ions and molecules.
Examples include: HCO2– (formate), CH3CO2– (acetate), CH2O (formaldehyde).

The strong foundational concepts such as VB Theory and MO Theory remain relevant in explaining chemical bonding, with notable distinctions between localized bonds in the VB approach and delocalization in the MO methodology.
Continuously emerging theories and experimental evidence through techniques like Photoelectron Spectroscopy (PES) serve to validate and refine existing bonding models.

Understanding chemical bonding through these theories is crucial for further explorations in chemistry and various applications across the field.

For further questions or clarifications, reach out to provided academic resources or utilize SI sessions for collaborative learning.