Step 4 C CICC C Clear A C (Refer C My ArC Create C Home C Tor-aC EXERCCI ChecO8 https://east.cengagenow.com/ilrn/takeAssignment/takeCovalentActivity.do?locator=assignment-takeClearAvailable Electrons: 16||Cl- c-clCreate the Lewis structure of COCl2 using the following five steps.NOTE:• Work only on the currently bolded step below don't work ahead to solve the final structure.• Scroll down to see steps 3-5.Step 1.Count valence electrons in the molecule or ion. Do this by adding the periodic group numbers for each atom in the struccharge. Enter and check your answer.Step 2.Connect each pair of bonded atoms with a single bond by dragging bonds onto the molecule. Each bond decreases the nuelectrons by 2. Check your answer.Step 3.Drag lone pairs of electrons to each terminal atom unt each has an octet of electrons, or until all electrons are used. You mpair by clicking on it. Check your answer.Step 4.Place remaining electrons on the central atom. If the structure is already correct, or no electrons remain, check yourStep 5.Examine the structure. If the central atom has less than eight valence electrons, create multiple bonds by changing a terminalto a second or third bond pair. Only C, N, O, P, and S commonly make multiple bonds. If the structure is already correct, checkeWEBdDISPLAYOne Touch Web DeO CC

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Principles of Modern Chemistry

8th Edition

ISBN:9781305079113

Author:David W. Oxtoby, H. Pat Gillis, Laurie J. Butler

Publisher:David W. Oxtoby, H. Pat Gillis, Laurie J. Butler

Chapter3: Atomic Shells And Classical Models Of Chemical Bonding

Section: Chapter Questions

Problem 99AP: A stable triatomic molecule can be formed that contains one atom each of nitrogen, sulfur, and...

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A stable triatomic molecule can be formed that contains one atom each of nitrogen, sulfur, and fluorine. Three bonding structures are possible, depending on which is the central atom: NSF, SNF, and SFN. (a) Write a Lewis diagram for each of these molecules, indicating the formal charge on each atom. (b) Often, the structure with the least separation of formal charge is the most stable. Is this statement consistent with the observed structure for this molecule—namely, NSF, which has a central sulfur atom? (c) Does consideration of the electronegativities of N, S, and F from Figure 3.18 help rationalize this observed structure? Explain.
The percent ionic character of a bond can be approximated by the formula 16+3.52 , where is the magnitude of the difference in the electronegativities of the atoms (see Fig. 3.18). Calculate the percent ionic character of HF, HCl, HBr, HI, and CsF, and compare the results with those in Table 3.7.
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.
The molecular ion S3N3 has the cyclic structure All SN bonds are equivalent. (a) Give six equivalent resonance hybrid Lewis diagrams for this molecular ion. (b) Compute the formal charges on all atoms in the molecular ion in each of the six Lewis diagrams. (c) Determine the charge on each atom in the polyatomic ion, assuming that the true distribution of electrons is the average of the six Lewis diagrams arrived at in parts (a) and (b). (d) An advanced calculation suggests that the actual charge resident on each N atom is 0.375 and on each S atom is +0.041 . Show that this result is consistent with the overall +1 charge on the molecular ion.
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?
How can we drsw the Lewis structures for the 4 boxes? And how many charges are in Oxime?
Represent the bonding in SF2 (F-S-F) with Lewis diagrams. Include the formal charges on all atoms. The dimerof this compound has the formula S2F4. It was isolated in1980 and shown to have the structure F3S-SF. Draw apossible Lewis diagram to represent the bonding in thedimer, indicating the formal charges on all atoms. Is it possible to draw a Lewis diagram for S2F4 in which all atomshave valence octets? Explain why or why not
Draw the Lewis structures for the following species: (a)CN-; (b) [N3]-; (c) [SCN]- For (b) and (c), there are more than one canonical structure that satisfy the Octet rule, and resonance is likely applied. Which canonical structures do you think would be the major contributor to the overall structure (i.e., resonance hybrid)?
Use Lewis symbols and Lewis structures to diagram the formationof PF3 from P and F atoms, showing valence-shellelectrons. (a) How many valence electrons does P have initially?(b) How many valence electrons does each F haveinitially? (c) How many valence electrons surround the P inthe PF3 molecule? (d) How many valence electrons surroundeach F in the PF3 molecule? (e) How many bonding pairs ofelectrons are in the PF3 molecule?
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?
The hypochlorite ion, ClO-, is the active ingredient inbleach. The perchlorate ion, ClO4-, is a main componentof rocket propellants. Draw Lewis structures for both ions. (a) What is the formal charge of Cl in the hypochlorite ion?(b) What is the formal charge of Cl in the perchlorate ion, assumingthe Cl—O bonds are all single bonds? (c) What is theoxidation number of Cl in the hypochlorite ion? (d) Whatis the oxidation number of Cl in the perchlorate ion, assumingthe Cl—O bonds are all single bonds? (e) In a redox reaction,which ion would you expect to be more easily reduced?
Im taking organic chemistry and in class we have to determine whether the moelcules are stable, unstable or impossible from looking at the lewis structure diagrams What steps do i take to determine their stability?