Write the Lewis structure for each molecule or ion. d. C₂H₄

Write the Lewis structure for each molecule or ion. a. H₃COCH₃

Write a Lewis structure that obeys the octet rule for each molecule or ion. Include resonance structures if necessary and assign formal charges to each atom. a. SeO₂

Write a Lewis structure that obeys the octet rule for each ion. Include resonance structures if necessary and assign formal charges to each atom. a. ClO₃^- b. ClO₄^- c. NO₃^- d. NH₄⁺

Use formal charges to identify the better Lewis structure.

How important is the resonance structure shown here to the overall structure of carbon dioxide? Explain.

In N₂O, nitrogen is the central atom and the oxygen atom is terminal. In OF₂, however, oxygen is the central atom. Use formal charges to explain why.

Draw the Lewis structure (including resonance structures) for the acetate ion (CH₃COO^-). For each resonance structure, assign formal charges to all atoms that have formal charge.

Draw the Lewis structure (including resonance structures) for methyl azide (CH₃N₃). For each resonance structure, assign formal charges to all atoms that have formal charge.

What are the formal charges of the atoms shown in red?

What are the formal charges of the atoms shown in red?

Write the Lewis structure for each molecule (octet rule not followed). b. NO₂

Write the Lewis structure for each molecule (octet rule not followed). a. BBr₃ b. NO c. ClO₂

Write Lewis structures for each molecule or ion. Use expanded octets as necessary. a. PF₅

Write Lewis structures for each molecule or ion. Use expanded octets as necessary. b. AsF₆^-

Order these compounds in order of increasing carbon–carbon bond strength: HCCH, H₂CCH₂, H₃CCH₃.

Order these compounds in order of decreasing carbon–carbon bond length: HCCH, H₂CCH₂, H₃CCH₃.

Which of the two compounds, H₂NNH₂ and HNNH, has the strongest nitrogen-nitrogen bond, and which has the shorter nitrogen-nitrogen bond.

Hydrogenation reactions are used to add hydrogen across double bonds in hydrocarbons and other organic compounds. Use average bond energies to calculate ΔH_rxn for the hydrogenation reaction. H₂C=CH₂(g) + H₂(g) → H₃C–CH₃(g)

Ethanol is a possible fuel. Use average bond energies to calculate ΔHrxn for the combustion of ethanol. CH₃CH₂OH(g) + 3 O2(g) → 2 CO₂(g) + 3 H₂O(g)

Ethane burns in air to form carbon dixode and water vapor.

2 H3C¬CH3( g) + 7 O2( g)¡4 CO2( g) + 6 H2O( g)

Use average bond energies to calculate ΔHrxn for the reaction.

In the Chemistry and the Environment box on free radicals in this chapter, we discussed the importance of the hydroxyl radical in reacting with and eliminating many atmospheric pollutants. However, the hydroxyl radical does not clean up everything. For example, chlorofluorocarbons—which destroy stratospheric ozone—are not attacked by the hydroxyl radical. Consider the hypothetical reaction by which the hydroxyl radical might react with a chlorofluorocarbon: OH(g) + CF₂Cl₂(g) → HOF(g) + CFCl₂(g) Use bond energies to explain why this reaction is improbable. (The C–F bond energy is 552 kJ/mol.)

Write an appropriate Lewis structure for each compound. Make certain to distinguish between ionic and molecular compounds. b. ClF₅

Each compound contains both ionic and covalent bonds. Write ionic Lewis structures for each, including the covalent structure for the ion in brackets. Write resonance structures if necessary. a. BaCO₃

Each compound contains both ionic and covalent bonds. Write ionic Lewis structures for each, including the covalent structure for the ion in brackets. Write resonance structures if necessary. b. Ca(OH)₂ c. KNO₃ d. LiIO

Carbon ring structures are common in organic chemistry. Draw a Lewis structure for each carbon ring structure, including any necessary resonance structures. a. C₄H₈ b. C₄H₄ c. C₆H₁₂ d. C₆H₆