Module 6: Covalent Bonds & Assessment Strategies

Assessment Methods and Final Exam Suggestions

• Instructor's Inquiry: The instructor asked students for their preferred assessment methods beyond traditional exams and for suggestions for the final exam format.

• Extra Credit Opportunity: Providing suggestions for both assessment methods and the final exam format earns $4$ extra credit points on Exam Two.

• Past Final Exam Formats:

  • Group Exam: Students self-sectioned into groups of three to complete a challenging exam collaboratively. This was described as a tough, not "patty cake," final.

  • Individual Cumulative Final: Covered chapters $1$ through $12$. Students generally disliked this format.

  • Critical Chapters Final: Focused on chapters $5$, $6$, $10$, and $11$ (as Exam Three covers chapters $7$, $8$, and $9$). This format was deemed more acceptable by students.

• Student Suggestion (Presentation): One student suggested a presentation/slideshow where they pick a topic and teach the class, similar to the instructor's teaching style. While the student felt they could perform better with this, they acknowledged that many people dislike presentations due to the potential for unequal team contribution.

• Instructor's Alternative Assessments (Past): The instructor has previously used creative methods like making chemistry-themed TikTok dances (e.g., demonstrating stoichiometry conversions like grams-to-moles) for their enjoyment, although only a few students participated.

Exam Two Information

• Challenge Level: Exam Two is considered the most challenging exam.

• Reason for Challenge: Primarily due to its mathematical nature, involving conversions.

• Conversion Types: Moves beyond simple conversions (e.g., milliliters to centimeters) to more complex ones (e.g., grams to moles, moles to moles).

• Content Focus: Chapter $5$ and $6$ will be a major focus, in addition to chapter $4$. Practice questions for chapter $4$ are available to give an idea of exam questions.

• Prompt for Extra Credit: The extra credit assignment requires a nice, well-thought-out response in essay format regarding preferred assessment methods and final exam format suggestions.

Chapter 4: Molecular Compounds and Covalent Bonds

Introduction to Intermolecular Forces

• Recall from Previous Chapters: We have already discussed ionic bonds, which involve the transferring of electrons.

• New Topic: Covalent Bonds: This chapter focuses on covalent bonds, which involve the sharing of electrons between atoms.

  • Analogy: Covalent bonds are like "soul mates" because the atoms are "stuck together" sharing electrons to complete their octet.

Key Concepts and Definitions

• Molecular Compounds: Compounds that consist of molecules rather than ions. Example: Polyatomic ions are held together by covalent bonds internally, even though the overall species carries a charge.

• Covalent Bond: A bond formed by sharing electrons between atoms.

• Molecule: A group of atoms held together by covalent bonds.

• Octet Rule: Main group elements undergo reactions to achieve $8$ valence electrons, mimicking a noble gas configuration. Nonmetals achieve this by sharing electrons in covalent bonds.

  • Goal: To have $8$ valence electrons (with the exception of hydrogen, which needs $2$ electrons).

• Valence Electrons: Electrons in the outermost shell of an atom. Can be determined by the column number on the periodic table (e.g., Group $5$A has $5$ valence electrons).

• Lewis Dot Structure: A representation that shows the valence electrons on the outer shell of an atom or in a molecule.

  • Finding Valence Electrons: The column number of an element in the periodic table indicates the number of valence electrons.

  • Drawing Lewis Structures: Lab $6$ will involve drawing these structures.

• Types of Covalent Bonds:

  • Single Bond: Involves the sharing of $2$ electrons ($1$ pair).

  • Double Bond: Involves the sharing of $4$ electrons ($2$ pairs).

  • Triple Bond: Involves the sharing of $6$ electrons ($3$ pairs).

• Bond Length: The optimum distance between nuclei in a bond. Generally, more electrons involved in a bond (e.g., double/triple) lead to shorter bond lengths due to stronger attraction.

• Diatomic Molecules: Seven elements that exist naturally in a double state (e.g., $H_2$) due to overlapping orbitals and the need to complete their octet. These must be memorized:

  • Hydrogen ($H_2$)

  • Nitrogen ($N_2$)

  • Oxygen ($O_2$)

  • Fluorine ($F_2$)

  • Chlorine ($Cl_2$)

  • Bromine ($Br_2$)

  • Iodine ($I_2$)

• Coordinate Covalent Bonds: Covalent bonds formed when both electrons are donated by the same atom (not extensively tested).

• Molecular Formulas and Lewis Structures: Different ways to represent molecular compounds (Lewis dot structure, structural formulas, condensed structure).

• Polar vs. Nonpolar Molecules: Will be discussed.

• Low Melting and Boiling Points: A characteristic of molecular compounds (briefly discussed).

Sections NOT Required for Memorization/Testing

• Section 4.4 (Coordinate Covalent Bonds): Not tested; just for awareness of their place in bonding.

• Section 4.8 (Shapes of Molecules): Not required for memorization on exams, but the equation sheet provides a good reference for geometric shapes (e.g., linear, trigonal planar) useful for Lab $6$.

• Section 4.11 (Naming Binary Molecular Compounds): Not required.

Detailed Examples of Covalent Bonding

1. Hydrogen ($H_2$) Bonding:

   • Hydrogen has $1$ valence electron.

   • Two hydrogen atoms combine, each sharing its single electron, forming a single covalent bond.

   • This results in $2$ electrons between the nuclei, achieving the electron configuration of helium ($1s^2$), completing hydrogen's need for $2$ electrons to be stable.

   • A single line often represents a covalent bond, with each line symbolizing $2$ shared electrons.

2. Water ($H_2O$):

   • Example of a molecule held together by covalent bonds, where two hydrogen atoms share electrons with one oxygen atom.

3. Ammonia ($NH_3$):

   • Nitrogen (N): In Column $5$A, so it has $5$ valence electrons.

   • Hydrogen (H): Has $1$ valence electron and needs $1$ more to complete its octet ($2$ electrons).

   • Bonding: Each of the three hydrogen atoms forms a single covalent bond with the nitrogen atom, sharing $1$ electron from hydrogen and $1$ from nitrogen.

   • Octet Check:

     • Hydrogen: Each H atom participates in one single bond, giving it $2$ electrons (shared) and fulfilling its required octet.

     • Nitrogen: The N atom forms three single bonds (total of $3 imes 2 = 6$ shared electrons) and has one lone pair of $2$ electrons. Total electrons around N are $6 + 2 = 8$, completing its octet.

   • Real-world Connection (Cirrhosis): The liver converts ammonia to ammonium. In cirrhosis (scarring of the liver), the liver's function is impaired, leading to ammonia buildup. High ammonia levels cause hepatic encephalopathy (dizziness, confusion). Lactulose is a drug given to bind ammonia and facilitate its excretion via bowel movements.

4. Methane ($CH_4$):

   • Carbon (C): In Column $4$A, so it has $4$ valence electrons. Needs $4$ more.

   • Hydrogen (H): Has $1$ valence electron. Needs $1$ more.

   • Bonding: Carbon, being very stable, forms four single covalent bonds, each with a hydrogen atom.

   • Octet Check:

     • Hydrogen: Each H atom has $2$ shared electrons (from its single bond).

     • Carbon: Carbon participates in four single bonds, totaling $4 imes 2 = 8$ electrons, completing its octet.

5. Exceptions to the Octet Rule (FYI only, not tested):

   • Boron Trifluoride ($BF_3$): Boron only has $6$ valence electrons, demonstrating an exception where an atom does not achieve an octet.

6. Hydrogen Bromide ($HBr$):

   • Hydrogen (H): $1$ valence electron, needs $1$ more = $2$ total.

   • Bromine (Br): In Column $7$A, has $7$ valence electrons, needs $1$ more = $8$ total.

   • Bonding: Hydrogen shares its $1$ electron with bromine, and bromine shares one of its electrons with hydrogen, forming a single covalent bond.

   • Octet Check:

     • Hydrogen: Has $2$ shared electrons (happy).

     • Bromine: Has $6$ lone pair electrons and $2$ shared electrons from the bond with H, totaling $8$ electrons (happy).

   • Conclusion: HBr will exist as a single bond.

7. Carbon Dioxide ($CO_2$) - Demonstrating Double Bonds:

   • Carbon (C): $4$ valence electrons.

   • Oxygen (O): $6$ valence electrons. Since there are two oxygen atoms, total valence electrons from oxygen are $6 imes 2 = 12$.

   • Total Valence Electrons for $CO_2$: $4 ( ext{from C}) + 12 ( ext{from 2 O}) = 16$ total valence electrons.

   • Attempt with Single Bonds:

     • If C formed a single bond with each O, and O was left with $6$ lone electrons.

     • Oxygen: Would have $6$ lone electrons $+$ $2$ shared electrons (from single bond) $=$ $7$ electrons (not happy).

     • Carbon: Would have $2$ shared electrons from one O $+$ $2$ shared electrons from the other O $=$ $4$ electrons (not happy).

     • Conclusion: Single bonds do not work for $CO_2$; it fails to satisfy the octet rule for all atoms.

   • Solution: Double Bonds: To achieve the octet for C and O while adhering to the $16$ total valence electrons.

     • Carbon forms a double bond with each oxygen atom.

     • Octet Check with Double Bonds:

       • Each Oxygen: Has $2$ lone pairs ($4$ electrons) $+$ $2$ shared pairs (a double bond, $4$ electrons) $=$ $8$ electrons (happy).

       • Carbon: Has $2$ double bonds ($2 imes 4 = 8$ electrons) $=$ $8$ electrons (happy).

     • Total Electron Count Check: $4 ( ext{from lone pairs on 1st O}) + 4 ( ext{from double bond}) + 4 ( ext{from double bond}) + 4 ( ext{from lone pairs on 2nd O}) = 16$ total electrons. This matches the calculated total valence electrons, confirming the structure.

   • Nitrogen ($N<em>2$) as a Triple Bond Example: Nitrogen exists as a diatomic molecule ($N</em>2$) due to a triple bond, sharing $6$ electrons to complete the octet for each nitrogen atom.

• Representations: Both Lewis dot structures (with dots) and skeleton structures (just lines representing bonds) are valid. The instructor will ask which structure is correct, not require students to draw them for exams.

Practice and Review

• Practice Questions: Covalent bond practice questions are available under "Week 6, Chapter 4 PowerPoint" on Canvas. These provide an idea of exam questions for Chapter $4$.

• Polyatomic Ions: Regularly read the polyatomic ions on the conversion sheet (e.g., for $5$ seconds daily) to become familiar with them, rather than memorizing them.

Conclusion

• Covalent bonds are a smooth concept once the octet rule and electron sharing are understood.

• It is important to determine if a molecule is likely to exist by checking if all atoms achieve their octet (or duet for hydrogen) and if the total number of valence electrons is accounted for.