6. Answer the following questions that relate to chemical bonding. (a) Two Lewis structures can be drawn for the OPF3 molecule, as shown below. :Ö: :F-P-F: :F-P-F: :F: :F: Structure 1 Structure 2 (i) (ii) How many sigma and how many pi bonds are in Structure 1? 4 and 1 σ Does Structure 1 demonstrate resonance? Justify your decision. No it does not demonstrate resonance because the double bond is w/ the there oxygen mic and disn't another oxygen mild to transfer the bond to. (iii) Which one of the two structures best represents a molecule of OPF3? Justify your answer in terms of formal charge.

6. Answer the following questions that relate to chemical bonding. (a) Two Lewis structures can be drawn for the OPF3 molecule, as shown below. :Ö: :F-P-F: :F-P-F: :F: :F: Structure 1 Structure 2 (i) (ii) How many sigma and how many pi bonds are in Structure 1? 4 and 1 σ Does Structure 1 demonstrate resonance? Justify your decision. No it does not demonstrate resonance because the double bond is w/ the there oxygen mic and disn't another oxygen mild to transfer the bond to. (iii) Which one of the two structures best represents a molecule of OPF3? Justify your answer in terms of formal charge.

Organic Chemistry: A Guided Inquiry

2nd Edition

ISBN:9780618974122

Author:Andrei Straumanis

Publisher:Andrei Straumanis

Chapter1: Bond Angles And Shape

Section: Chapter Questions

Problem 12CTQ: Consider the following flat drawing of methane (CH4) . a. What is HCH bond angle implied by this...

See similar textbooks

Similar questions

Consider the following ion: BrO3−. What is the electronic geometry of BrO3−? What is its molecular shape? e) Does BrO3− have a dipole moment? Briefly justify your answer. f) On average, would you expect IO3− to have longer or shorter bonds than BrO3−? Briefly explain your answer.
“Inert” xenon actually forms several compounds, especially with the highly electronegative elements oxygen and fluorine.The simple fluorides XeF₂, XeF₄, and XeF₆ are all formed by di-rect reaction of the elements. As you might expect from the sizeof the xenon atom, the Xe-F bond is not a strong one. Calcu-late the Xe-F bond energy in XeF₆, given that the heat of for-mation is 402 kJ/mol.
Sulfur dioxide (SO2) and carbon dioxide (CO2) have somewhat different properties, in spite of of the fact that both are triatomic molecules with two terminal oxygen atoms. Explain why they are as different as they are? a. Draw the best Lewis Structure for both molecules in the space below: b. Write the correct formal charge value next to every single atom. c. Identify AND compare each of their following features in the Summary Table below: d. Draw all possible resonance structures (if any) and justify the best ones by citing the “Rule” for best Lewis Structure based on the formal charge distribution.
Draw Lewis structures for BrO4-, SiH4, PCl4, CH2Cl2 and BF4-.(a) Which species is/are likely to be paramagnetic?(b) What reaction(s) would occur for the paramagnetic species? (c) By VSEPR, suggest the molecular geometry of each species.
Sulfur tetrafluoride (SF4) reacts slowly with O2 to form sulfurtetrafluoride monoxide (OSF4) according to the followingunbalanced reaction: SF4(g) + O2(g)---->OSF4(g)The O atom and the four F atoms in OSF4 are bonded to acentral S atom. (a) Balance the equation. (b) Write a Lewisstructure of OSF4 in which the formal charges of all atomsare zero. (c) Use average bond enthalpies (Table 8.3) to estimatethe enthalpy of the reaction. Is it endothermic or exothermic?(d) Determine the electron-domain geometry ofOSF4, and write two possible molecular geometries for themolecule based on this electron-domain geometry. (e) Foreach of the molecules you drew in part (d), state how manyfluorines are equatorial and how many are axial.
Draw the Lewis structure for each of the following moleculesor ions, and predict their electron-domain and moleculargeometries: (a) AsF3, (b) CH3+, (c) BrF3, (d) ClO3-, (e) XeF2,(f) BrO2-.
Draw possible Lewis structures and using formal charge determine the best one for [AsSiSe]-. Show FC for ALL structures and then explain why you chose the one that is best.
The partial Lewis structure that follows is for a hydrocarbonmolecule. In the full Lewis structure, each carbon atomsatisfies the octet rule, and there are no unshared electronpairs in the molecule. The carbon—carbon bondsare labeled 1, 2, and 3. (a) How many hydrogen atomsare in the molecule? (b) Rank the carbon–carbonbonds in order of increasing bond length. (c) Whichcarbon—carbon bond is the strongest one?
5b) i) Discuss the bonding and extra.stability of benzene,ii) Using suitable examples, compare the reactivity of benzene and ethane,
Consider the following flat drawing of methane (CH4) . a. What is HCH bond angle implied by this drawing if you assume it is flat? b. Are the electron domains of this flat CH4 spread out as much as possible? c. Use model materials to make a model of CH4 (methane). If you assembled it correctly, thefour bonds (bonding electron domains) of your model will be 109.5° apart. d. In which representation, the drawing above or the model in your hand (circle one) are theH’s of CH4 more spread out around the central carbon? e. Confirm that your model looks like the following drawing. The wedgebond represents a bond coming out of the page, and the dash bondrepresents a bond going into the page f. You will often see methane drawn as if it were flat (like on the previous page). Why is thismisleading, and what is left to the viewer’s imagination when looking at such a drawing?
Define formal charge and explain how to calculate it. What is the purpose of the formal charge? Organic compounds are composed mostly of carbon and hydrogen but also may have oxygen, nitrogen, and/or halogens in the formula. Formal charge arguments work very well for organic compounds when drawing the best Lewis structure. How do C, H, N, O, and Cl satisfy the octet rule in organic compounds so as to have a formula charge of zero?
Two possible Lewis diagrams for sulfine (H2CSO) are (a) Compute the formal charges on all atoms. (b) Draw a Lewis diagram for which all the atoms in sulfine have formal charges of zero.