Carbon dioxide can be represented using multiple resonance structures, which distribute electrons differently while maintaining overall charge neutrality. Two possible resonance structures, I and II, are shown above.
Calculate the formal charge on each oxygen atom in the two structures.
Determine which resonance structure is more stable. Justify your answer in terms of formal charge minimization and electronegativity.
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Boron trifluoride (BF3) and bromine trifluoride (BrF3) are two compounds that exhibit different bonding behaviors.
Use the concept of formal charge to explain why BF3 does not obey the octet rule and prefers an incomplete octet rather than forming a dative bond with fluorine.
Determine the electron domain geometry and molecular geometry of BrF3 using VSEPR theory.
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The sulfate ion (SO42-) can be represented by multiple valid Lewis diagrams, depending on whether sulfur follows or exceeds the octet rule.
Draw two valid Lewis structures for SO42-:
One in which all atoms obey the octet rule.
One in which sulfur has an expanded octet of 12 electrons.
Calculate the formal charge on sulfur in both structures and determine which structure is more stable based on formal charge minimization.
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A student is studying the bonding and reactivity of several phosphorus-containing molecules: PF3, PCl5, and POCl3. These molecules illustrate key concepts such as formal charge, expanded octets, molecular polarity, and bonding limitations.
Draw the complete Lewis structure for PF3.
Predict the molecular geometry of PF3. Justify your answer.
Draw the Lewis diagram for PCl5.
Explain why PCl5 is an exception to the octet rule.
Represent a complete Lewis diagram for POCl3. Minimize formal charges and show all bonding.
Determine the formal charge on each atom in the Lewis structure of POCl3 in your answer to part (e) and explain why this is the stable resonance structure.
Predict whether PF3 or POCl3 is more polar. Justify your answer.
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Nonmetal | C | N | O | Ne | Si | P | S | Ar |
|---|---|---|---|---|---|---|---|---|
Formula of Compound |
CF4 |
NF3 |
OF2 | No compound |
SiF4 |
PF3 |
SF2 | No compound |
Some binary compounds that form between fluorine and various nonmetals are listed in the table above. A student examines the data in the table and poses the following hypothesis: the number of F atoms that will bond to a nonmetal is always equal to 8 minus the number of valence electrons in the nonmetal atom.
Based on the student’s hypothesis, what should be the formula of the compound that forms between chlorine and fluorine?
In an attempt to verify the hypothesis, the student researches the fluoride compounds of the other halogens and finds the formula ClF3 . In the box below, draw a complete Lewis electron-dot diagram for a molecule of ClF3 .
Two possible geometric shapes for the ClF3 molecule are trigonal planar and T-shaped. The student does some research and learns that the molecule has a dipole moment.
Which of the two shapes is consistent with the fact that the ClF3 molecule has a dipole moment? Justify your answer in terms of bond polarity and molecular structure.
In an attempt to resolve the existence of the ClF3 molecule with the hypothesis stated above, the student researches the compounds that form between halogens and fluorine, and assembles the following list.
Halogen | Formula(s) |
|---|---|
F | F2 |
Cl |
|
Br | BrF, BrF3 , BrF5 |
I | IF, IF3 , IF5 , IF7 |
Based on concepts of atomic structure and periodicity, propose a modification to the student’s previous hypothesis to account for the compounds that form between halogens and fluorine.
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Draw the Lewis structure of the carbonate ion, CO32-.
Explain why the carbonate ion exhibits resonance and describe how this affects the distribution of electron density in the molecule.
Explain how resonance influences bond length in the CO32- ion.
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Ozone (O3) is an important molecule in Earth’s atmosphere that absorbs ultraviolet radiation. The bonding within ozone cannot be represented by a single Lewis structure.
Draw all valid resonance structures for ozone (O3), ensuring correct placement of lone pairs and formal charges.
Explain why both O–O bonds in ozone are the same length, despite one being represented as a double bond and the other as a single bond in individual resonance structures.
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