CHCl3

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  1. CHCl3

a. Lewis structure

b. ???Electronic geometry.

???c.Molecular shape.

e. ???Bond angle?

f. draw the different bonds and label their polarity:

  1. SO4^2-

a. Lewis structure

b. ???Electronic geometry.

c. ???Molecular shape.

  1. ???Bond angle?

f. draw the different bonds and label their polarity:

  1. CO2

a. Lewis structure:

b. ???Electronic geometry.:

c. ???Molecular shape:

d. ???Bond angle?:

f. Draw the different bonds and label their polarity.

  1. HCN

a. Lewis structure

b. ???Electronic geometry.

???c.Molecular shape.

e. ???Bond angle?

f. draw the different bonds and label their polarity:

  1. SO3

a. Lewis structure

b. ???Electronic geometry.

c. ???Molecular shape.

  1. ???Bond angle?

f. draw the different bonds and label their polarity:

  1. NO2^-

a. Lewis structure:

b. ???Electronic geometry.:

c. ???Molecular shape:

d. ???Bond angle?:

f. Draw the different bonds and label their polarity.

  1. SF6

a. Lewis structure

b. ???Electronic geometry.

???c.Molecular shape.

e. ???Bond angle?

f. draw the different bonds and label their polarity:

  1. XeF4

a. Lewis structure

b. ???Electronic geometry.

c. ???Molecular shape.

  1. ???Bond angle?

f. draw the different bonds and label their polarity:

The Correct Answer and Explanation is :

1. CHCl3 (Chloroform)

a. Lewis Structure:

  • Central carbon (C) atom bonded to three chlorine (Cl) atoms and one hydrogen (H) atom.
   H
   |
Cl-C-Cl
   |
  Cl

b. Electronic Geometry: Tetrahedral

  • The four regions of electron density (three single bonds with chlorine atoms and one with hydrogen) around the central carbon.

c. Molecular Shape: Tetrahedral

  • Based on the electron geometry, but considering bond polarization, the shape is tetrahedral.

d. Bond Angle: Approximately 109.5°

  • The ideal angle for a tetrahedral molecular shape.

e. Polarity of Bonds:

  • C-H: Slightly polar (C is more electronegative than H).
  • C-Cl: Polar (Cl is more electronegative than C).

2. SO4^2- (Sulfate Ion)

a. Lewis Structure:

  • Central sulfur (S) atom surrounded by four oxygen (O) atoms, with double bonds to each oxygen.
  O=S=O
  |   |
  O   O

b. Electronic Geometry: Tetrahedral

  • Four regions of electron density (four O atoms).

c. Molecular Shape: Tetrahedral

  • The shape remains tetrahedral due to the symmetrical distribution of the bonds.

d. Bond Angle: 109.5°

  • For a tetrahedral shape, the bond angles are approximately 109.5°.

e. Polarity of Bonds:

  • S=O: Polar bonds due to the difference in electronegativity.

3. CO2 (Carbon Dioxide)

a. Lewis Structure:

  • Central carbon (C) atom with double bonds to two oxygen (O) atoms.
O=C=O

b. Electronic Geometry: Linear

  • There are two regions of electron density (two double bonds to oxygen).

c. Molecular Shape: Linear

  • With two bonding pairs and no lone pairs on carbon, the shape is linear.

d. Bond Angle: 180°

  • The ideal angle for a linear molecule.

e. Polarity of Bonds:

  • C=O: Polar bonds due to the electronegativity difference.

4. HCN (Hydrogen Cyanide)

a. Lewis Structure:

  • Hydrogen (H) bonded to carbon (C), which is triple-bonded to nitrogen (N).
H-C≡N

b. Electronic Geometry: Linear

  • Two regions of electron density (C-H and C≡N).

c. Molecular Shape: Linear

  • The molecule is linear due to the two regions of electron density.

d. Bond Angle: 180°

  • The ideal bond angle for a linear molecule.

e. Polarity of Bonds:

  • C-H: Slightly polar.
  • C≡N: Polar due to the electronegativity difference.

5. SO3 (Sulfur Trioxide)

a. Lewis Structure:

  • Central sulfur (S) atom with double bonds to three oxygen (O) atoms.
  O=S=O
  |   |
  O   O

b. Electronic Geometry: Trigonal Planar

  • Three regions of electron density (three O atoms).

c. Molecular Shape: Trigonal Planar

  • The distribution of electrons results in a trigonal planar shape.

d. Bond Angle: 120°

  • The ideal bond angle for a trigonal planar molecule.

e. Polarity of Bonds:

  • S=O: Polar bonds due to the electronegativity difference.

6. NO2^- (Nitrite Ion)

a. Lewis Structure:

  • Nitrogen (N) is bonded to two oxygen (O) atoms, one of which has a negative charge.
 O=N-O^-

b. Electronic Geometry: Trigonal Planar

  • Three regions of electron density (two bonds with oxygen and one lone pair on nitrogen).

c. Molecular Shape: Bent

  • The lone pair on nitrogen causes the molecule to adopt a bent shape.

d. Bond Angle: Approximately 120°

  • The bond angle in a bent structure is slightly less than 120° due to the lone pair.

e. Polarity of Bonds:

  • N-O: Polar bonds due to the electronegativity difference.

7. SF6 (Sulfur Hexafluoride)

a. Lewis Structure:

  • Central sulfur (S) atom bonded to six fluorine (F) atoms.
      F
      |
 F - S - F
      |
      F

b. Electronic Geometry: Octahedral

  • Six regions of electron density (six F atoms).

c. Molecular Shape: Octahedral

  • The symmetrical distribution of electrons gives it an octahedral shape.

d. Bond Angle: 90°

  • The ideal bond angle for an octahedral molecule.

e. Polarity of Bonds:

  • S-F: Polar bonds due to the electronegativity difference.

8. XeF4 (Xenon Tetrafluoride)

a. Lewis Structure:

  • Central xenon (Xe) atom bonded to four fluorine (F) atoms.
    F
    |
 F - Xe - F
    |
    F

b. Electronic Geometry: Octahedral

  • Six regions of electron density (four F atoms and two lone pairs on Xe).

c. Molecular Shape: Square Planar

  • The lone pairs on xenon cause the molecular shape to be square planar.

d. Bond Angle: 90°

  • The ideal bond angle for a square planar shape.

e. Polarity of Bonds:

  • Xe-F: Polar bonds due to the electronegativity difference.

300-Word Explanation:

The Lewis structures of these molecules/ions represent how atoms are bonded and how lone pairs are distributed. For each molecule, we can determine the electronic geometry, which describes the spatial arrangement of electron regions around the central atom. Molecular shape follows from the electronic geometry but takes into account only the positions of atoms, not lone pairs of electrons. The bond angles reflect the ideal angles between bonds based on the geometry.

For example, in CHCl3, the tetrahedral electronic geometry results in bond angles of 109.5° and a tetrahedral molecular shape, as the three chlorine atoms and one hydrogen atom form bonds with carbon. However, SO4^2- has a tetrahedral geometry with symmetrical bonding, making its molecular shape also tetrahedral with bond angles of 109.5°.

In CO2, the central carbon atom is double-bonded to two oxygen atoms in a linear arrangement, leading to a bond angle of 180°. The NO2^- ion is bent because of the lone pair on nitrogen, which pushes the bonds closer together and results in a bond angle slightly less than 120°.

Understanding the polarity of bonds is essential in recognizing how the differences in electronegativity between atoms in a molecule cause electron density shifts, leading to dipoles that affect molecular interactions. The SF6 and XeF4 molecules demonstrate how lone pairs can influence molecular shape, causing distortions from idealized bond angles.

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