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Basic Answer Key Chapter 9: Covalent Bonding, Study notes of Geometry

Faculté Polytechnique de MonsGeometry

A covalent bond is formed when two atoms share one or more electrons. 2. Covalent bonds involve electron sharing, while ionic bonds are formed as a result of ...

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Chemistry Student Edition - Basic
Answer Key
Chapter 9: Covalent Bonding
Lewis Electron Dot Structures
Questions
1. How is a covalent bond formed?
2. What is the major difference between a covalent bond and an ionic bond?
3. What orbitals are used in covalent bonding?
4. What types of elements generally form covalent bonds?
5. How do double and triple covalent bonds differ from single covalent bonds?
6. Predict the number of covalent bonds the following atoms will make:”
a. N (nitrogen)
b. S (sulfur)
c. Br (bromine)
d. F (fluorine)
7. Draw Lewis structures for the following molecules:
a. HBr
b. CO2
c. NI3
d. PBr5
8. Determine how many lone pairs are present in each of the molecules from the problem
above.
Answers
1. A covalent bond is formed when two atoms share one or more electrons.
2. Covalent bonds involve electron sharing, while ionic bonds are formed as a result of charge
attractions between two ions.
3. Covalent bonds use s and p orbitals
4. Covalent bonds are usually formed by compounds with O, N, C, S, and P
5. Double and triple covalent bonds share two and three pairs of electrons, respectively
6.
a. 3
b. 2
c. 1
d. 1
7. Draw Lewis structures for the following molecules:
a.
b. O::C::O
c.
pf3
pf4
pf5

Partial preview of the text

Download Basic Answer Key Chapter 9: Covalent Bonding and more Study notes Geometry in PDF only on Docsity!

Chemistry Student Edition - Basic

Answer Key

Chapter 9: Covalent Bonding

Lewis Electron Dot Structures

Questions

  1. How is a covalent bond formed?
  2. What is the major difference between a covalent bond and an ionic bond?
  3. What orbitals are used in covalent bonding?
  4. What types of elements generally form covalent bonds?
  5. How do double and triple covalent bonds differ from single covalent bonds?
  6. Predict the number of covalent bonds the following atoms will make:” a. N (nitrogen) b. S (sulfur) c. Br (bromine) d. F (fluorine)
  7. Draw Lewis structures for the following molecules: a. HBr b. CO 2 c. NI 3 d. PBr 5
  8. Determine how many lone pairs are present in each of the molecules from the problem above.

Answers

  1. A covalent bond is formed when two atoms share one or more electrons.
  2. Covalent bonds involve electron sharing, while ionic bonds are formed as a result of charge attractions between two ions.
  3. Covalent bonds use s and p orbitals
  4. Covalent bonds are usually formed by compounds with O, N, C, S, and P
  5. Double and triple covalent bonds share two and three pairs of electrons, respectively
  6. a. 3 b. 2 c. 1 d. 1
  7. Draw Lewis structures for the following molecules: a. b. O::C::O c.

d.

Molecular Geometry

Questions

  1. What is the basic principle of VSEPR theory?
  2. How many covalent bonds would there be attached to a central atom in the following configurations: a. tetrahedral b. octahedral c. trigonal planar d. linear
  3. What is the difference between the electron domain geometry and the molecular geometry?
  4. How do lone pairs act differently than bonding pairs in terms of electron repulsion?
  5. Using the VSEPR method, predict the molecular geometries (including bond angles) for each of the following molecules: a. SF 2 b. PBr 3 c. AlCl 3 d. TeCl 6 e. HCN Answers
  6. Electron pairs in the valence shell of an atom repel one another. Repulsion between electrons in different bonding pairs causes them to remain as far apart as possible.

a. 4 b. 6 c. 3 d. 2

  1. Molecular geometry only considers atoms, while electron domain geometry includes lone pairs
  2. Lone pairs take up more space than bonding pairs

a. Bent, <109.5 degrees b. Trigonal pyramidal, <109.5 degrees c. Trigonal planar, 120 degrees d. Octahedral, 90, 180 degrees e. Linear, 180 degrees

Intermolecular Forces

Questions

  1. Explain what gives a molecule a molecular dipole.
  2. What is a dipole-dipole interaction?
  3. Define the term hydrogen bond.
  4. Explain what dispersion forces are.
  5. What would be the strongest intermolecular force holding together collections of each of the following molecules? a. CH 3 CH 2 CH 2 CH 3 b. CH 3 OH c. PCl 5 d. O 2
  6. Draw the hydrogen-bond interactions that can occur between molecules of CH 3 NH 2.
  7. How would molecules containing N-F bonds interact with one another? Answers
  8. A molecule becomes a dipole when there is sufficient difference in electronegativity between atoms to create centers of partial positive and partial negative charges in the molecule.
  9. Interactions of partial positive and partial negative charge centers between molecules.
  10. A polar interaction between a H attached to a N, O, or F with a N, O, or F on another molecule.
  11. Weak, temporary polar interactions between molecules.
  12. What would be the strongest intermolecular force holding together collections of each of the following molecules? a. CH 3 CH 2 CH 2 CH 3 : London Dispersion Forces b. CH 3 OH: Hydrogen bonding c. PCl 5 : London Dispersion Forces d. O 2 : Dispersion forces caused by external temporary polarization of the electron cloud between the two atoms.
  13. CH 3 - N-H │ H····N-H │ CH 3
  14. The electronegativity difference between N and F is 1.0, so the N would have a partial positive charge that cold interact with the partial negative charge of a F atom on another molecule to form a dipole-dipole interaction.

Hybridization and Molecular Orbitals

Questions

  1. Why do atoms “promote” electrons to higher orbitals in order to form bonds?
  2. Describe the type of hybrid orbitals formed corresponding to each of the following electron domain geometries:

a. tetrahedral b. trigonal planar c. linear

  1. In sp^3 hybridization, how many hybrid orbitals are there? How do the hybrid orbitals compare to one another?
  2. What is the hybridization state of Si in SiH 4?
  3. What hybrid orbitals are used by the nitrogen atoms in H 2 N-NH 2?
  4. What types of bonds make up single, double, and triple bonds?
  5. What hybrid orbitals are used by the carbon atoms in each of the following: a. H 3 C-CH=CH 2 b. CO 2
  6. Describe the covalent bonds in each of the molecules from the question above. Answers
  7. Promotion allows atoms to hybridize orbitals and create new bonds with identical energy

a. sp^3 b. sp^2 c. sp

  1. There are four hybrid orbitals in sp^3 hybridization. All four hybrid orbitals are identical
  2. The hybridization state is sp^3 d
  3. Sp^3 for both N atoms
  4. 1 sigma bond is found in single bonds; 1 sigma and 1 pi bond in double bonds, and 1 sigma and 2 pi bonds in triple bonds

a. Sp^3 for the single-bonded carbon and sp^2 for the carbon atoms involved in the double bond. b. Carbon dioxide is a linear molecule, so there would be sp 2 hybridization in all three atoms.

  1. Describe the covalent bonds in each of the molecules from the question above.