5: Covalent Bonding and Simple Molecular Compounds: Questions

Sacramento City College
CHEM 309: General, Organic and Biochemistry
Prof. Dianne Bennett

Agenda (Fall)  •   Agenda (Spring)  •  Homework Problems

1. Chapter 2: The Classical Wave Equation
2. 2.1: The One-Dimensional Wave Equation
3. 2.2: The Method of Separation of Variables
4. 2.3: Oscillatory Solutions to Differential Equations
5. 2.4: The General Solution is a Superposition of Normal Modes
6. 2.5: A Vibrating Membrane
7. 2.E: The Classical Wave Equation (Exercises)

 •  Homework Solutions

1. Chapter 11: Computational Quantum Chemistry
2. 11.1: Gaussian Basis Sets
3. 11.2: Extended Basis Sets
4. 11.3: Orbital Polarization Terms in Basis Sets
5. 11.4: The Ground-State Energy of H2H2
6. 11.5: Quantum Calculations
7. 11.E: Computational Quantum Chemistry (Exercises)

Chapters

1. Chapter 12: Group Theory: The Exploitation of Symmetry
2. 12.1: The Exploitation of Symmetry
3. 12.2: Symmetry Elements
4. 12.3: Symmetry Operations Define Groups
5. 12.4: Symmetry Operations as Matrices
6. 12.5: The C3VC3V Point Group
7. 12.6: Character Tables
8. 12.7: Characters of Irreducible Representations
9. 12.8: Using Symmetry to Solve Secular Determinants
10. 12.9: Generating Operators
11. 12.E: Group Theory: The Exploitation of Symmetry (Exercises)

  •  Video Tutorials

1. Chapter 13: Molecular Spectroscopy
2. 13.1: The Electromagnetic Spectrum
3. 13.2: Rotations Accompany Vibrational Transitions
4. 13.3: Unequal Spacings in Vibration-Rotation Spectra
5. 13.4: Unequal Spacings in Pure Rotational Spectra
6. 13.5: Vibrational Overtones
7. 13.6: Electronic Spectra Contain Electronic, Vibrational, and Rotational Information
8. 13.7: The Franck-Condon Principle
9. 13.8: Rotational Spectra of Polyatomic Molecules
10. 13.9: Normal Modes in Polyatomic Molecules
11. 13.10: Irreducible Representation of Point Groups
12. 13.11: Time-Dependent Perturbation Theory
13. 13.12: The Selection Rule for the Rigid Rotator
14. 13.13: The Harmonic Oscillator Selection Rule
15. 13.14: Group Theory Determines Infrared Activity
16. 13.E: Molecular Spectroscopy (Exercises)

 •  Basic Knowledge Practice Quizzes

Covalent Compounds

(1) How do atoms form covalent bonds? What is the easiest way to distinguish a covalent bond from an ionic bond?

(2) For each of the following elements, draw their Lewis dot structures.

1. Sulfur
2. Boron
3. Iodine
4. Silicon
5. ​Phosphorus

(3) Study the given Lewis structure and provide the following information:

1. Draw in any lone pairs that are needed.
2. How many bonding electrons does the carbon have? How many does Oxygen have?
3. Is the octet of each atom in the molecule satisfied?
4. How many single bonds are used in this molecule? How many double bonds?
5. ​How many bonds are shared between the Carbon and Oxygen atoms? How many are shared between the Carbon and each Hydrogen?

(4) Answer each question regarding the following structure:

1. Fill in all missing lone pairs
2. How many bonding electrons does the Nitrogen have? How many non-bonding electrons does it have?
3. How many bonding electrons do each Chlorine have? How many non-bonding electrons do they each have?
4. How many electrons, in total, does Nitrogen share with all three Chlorines? Does each atom have an octet?
5. ​Are the bonds formed in this molecule ionic or covalent?

(5) Draw the Lewis dot structures for each of the 7 diatomic elements in their diatomic states. Are these diatomic elements, when one atom is covalently bonded to another of the same element, considered molecules or compounds? Explain using the definition of whichever answer you choose.

(6) For each family or group number, describe the covalent bonding pattern found in organic compounds.

1. Halogens
2. Group 4A
3. Group 6A
4. Group 5A
5. Noble Gases

(7) Sketch the Lewis configurations for the following molecules. Include all lone pairs in the sketch and determine how many total electrons are in each molecule.

1. CH₂Cl₂
2. HCN
3. CH₂O
4. CH₃CH₂CH₃
5. ​CH₃COOH (Hint: Both oxygens bond to the second carbon)

(8) What is an expanded octet? Which elements cannot form an expanded octet? Where on the periodic table are the elements that can form expanded octets located?

(9) Give the names of the following covalent compounds.

1. N₂O₄
2. SCl₆
3. CBr₄
4. P₂O₅
5. ClF₇

Three-Dimensional Shapes of Molecules

(10) Include the Lewis dot structure for each given molecule. Give both the electron geometry and molecular shape of the central atom(s).

1. CHF₃
2. H₂O
3. NH₂I
4. CH₂S

(11) If a central atom has three bonding groups surrounding it, how can you differentiate between the possible molecular geometries it can take on? What electron geometry, molecular shape and bond angles would such a molecule have?

(12) For each of the following molecules, tell whether they are two-dimensional or three-dimensional based on their electron geometry.

1. NH₃
2. H₂S
3. CH₂Cl₂
4. CH₂S
5. CO₂

(13) What are the approximate bond angles in the given molecules?

1. PBr₃
2. HNO
3. CH₃F
4. CH₂O
5. CS₂

(14) In each of the following molecules, give the central atom’s molecular shape and the approximate bond angles.

a. O=C=O

b.

c. NH₃

d.

e.

(15) Use your knowledge of what you’ve learned so far to provide the Lewis dot structures, electron geometry, molecular shape, bond angles, number of lone pair electrons, and the number of shared pair electrons for each molecule given.

1. CI₂Br₂
2. NOF
3. ClCN
4. CH₂O

(16) For each of the three carbons in this molecule, provide the requested information.

1. The molecular shape of each carbon
2. Is this molecule more likely to be two-dimensional or three dimensional?
3. What is the bond angle of the CH₂?
4. What is the CH-C bond angle?
5. What is the CH-C-N bond angle?

Modeling Exercises

(17) Construction of acetic acid, CH₃COOH

1. Obtain two carbons, four hydrogens, and two oxygens with six single-bond sticks and two double-bond sticks
2. Sketch the Lewis dot structure to get an idea of how to put it together
3. Construct the model to get a three-dimensional visual of acetic acid
4. For each carbon atom, how many different electron groups, both bonding and non-bonding, surround them?
5. Of those groups and for each carbon, respectively, how many are bonding and how many are non-bonding?
6. What is the electron geometry around each carbon? Molecular geometry?
7. Is the majority of the molecule two-dimensional, or three-dimensional? Justify. (Hint: Try rotating some of the central atoms and looking at it from different angles.)
8. What are the bond angles on each central atom (remember to include the oxygen in the -OH as one of the central atoms.)
9. Why does one of the two oxygen atoms get treated as a central atom and not the other?

(18) Construction of acetonitrile, CH₃CN

1. Obtain two carbons, three hydrogens and a nitrogen, four single-bond sticks and three double-bond sticks.
2. Sketch the Lewis dot structure of the molecule to determine its structure.
3. Construct the model of the molecule for a visual representation.
4. How many non-bonding electrons does each carbon have? Bonding electrons?
5. Determine the electron geometry and molecular geometry of each carbon.
6. Is the nitrogen a central atom or not? Explain.
7. What is the electron geometry of the nitrogen? What is the molecular geometry?

Molecular Polarity

(19) Between the following pairs, which has the higher electronegativity?

1. Oxygen and Nitrogen
2. Sulfur and Oxygen
3. Chlorine and Bromine
4. Carbon and Silicon
5. Phosphorus and Nitrogen

(20) Describe the trend in electronegativity as you move left and right across the periodic table. Does electronegativity increase to the right or the left? Now do the same for moving up and down; does electronegativity increase as you go up or down?

(21) Sketch each of the following molecules to help you determine whether they are polar or non-polar. Draw dipole arrows to indicate polarity in polar molecules.

1. CCl₄
2. PH₃
3. HOCl
4. CS₂
5. HCOOH

(22) Electronegativity is not the sole decider of polarity within a molecule. Keeping this in mind, determine if the following molecules are polar or not. If not, explain why.

1. CH₂F₂
2. CO₂
3. CH₃CH₃
4. H₂O
5. CH₃OCH₃

Intermolecular Forces of Attraction

(23) Are intermolecular forces considered to be bonds? Explain. Is an H-bond an exception?

(24) Which of the three intermolecular forces is the most common? List two other names that can be used to describe this interaction. What is its relative strength compared to the others?

(25) Explain why a dipole-dipole attraction is a stronger force than London Dispersion forces. What is the critical difference that makes this so?

(26) What is the dominant intermolecular forces in each of the following molecules?

1. CH₂CH₂
2. CH₃OH
3. HCN
4. CH₃OCH₃
5. NH₃

(27) Water (H₂O) is one of the most perfect H-bonding molecules we know of. Use your knowledge of H-bonding to explain why solid H₂O (ice) floats in liquid H₂O while in most other substances, the density of the solids is greater than their respective liquids.

(28) Intermolecular forces are significant for allowing molecules to do things that bonding alone will not allow for. Explain how something like DNA utilizes intermolecular forces to take on its helical shape without relying solely on covalent bonding.