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chemistry
organic chemistry

Questions and Answers of Organic Chemistry

Use Coulomb€™s law,to calculate the energy of interaction for the following two arrangements of charges, each having a magnitude equal to the electron charge.
Without using Fig. 13.3, predict the order of increasing electronegativity in each of the following groups of elements.a. C, N, Ob. S, Se, Clc. Si, Ge, Snd. Tl, S, Gee. Na, K, Rbf. B, O, GaFigure 13.3
Without using Fig. 13.3, predict which bond in each of the following groups is the most polar.a. COF, SiOF, GeOFb. POCl, SOClc. SOF, SOCl, SOBrd. TiOCl, SiOCl, GeOCle. COH, SiOH, SnOHf. AlOBr, GaOBr,
Repeat Exercises 15 and 16. This time use the values of the electronegativities of the elements given in Fig.. Are there any differences among your answers?Figure
Hydrogen has an electronegativity value between boron and carbon and identical to phosphorus. With this in mind, rank the following bonds in order of decreasing polarity: POH, OOH, NOH, FOH, COH.
Rank the following bonds in order of increasing ionic character: NOO, CaOO, COF, BrOBr, KOF.
List all the possible bonds that can occur between the elements P, Cs, O, and H. Predict the type of bond (ionic, covalent, or polar covalent) one would expect to form for each bond.
Some plant fertilizer compounds are (NH4)2SO4, Ca3(PO4)2, K2O, P2O5, and KCl. Which of these com-pounds contain both ionic and covalent bonds?
The following electrostatic potential diagrams represent CH4, NH3, or H2O. Label each, and explain your choices.
When an element forms an anion, what happens to the radius? When an element forms a cation, what happens to the radius? Why? Define the term isoelectronic. When comparing sizes of ions, which ion has
Consider the ions Sc3+, Cl-, K+, Ca2+, and S2-. Match these ions to the following pictures that represent the relative sizes of the ions.
For each of the following groups, place the atoms and ions in order of decreasing size. a. Cu, Cu+, Cu2+ b. Ni2+, Pd2+, Pt2+ c. O, O-, O2- d. La3+, Eu3+, Gd3+, Yb3+ e. Te2-, I-, Cs+, Ba2+, La3+
Write electron configurations for each of the following. a. the cations: Mg2+, Sn2+, K+, Al3+, Tl+, As3+ b. the anions: N3-, O2-, F-, Te2-
Write electron configurations for the most stable ion formed by each of the elements Rb, Ba, Se, and I (when in stable ionic compounds).
Give an example of an ionic compound where both the anion and the cation are isoelectronic with each of the following noble gases. a. Ne b. Ar c. Kr d. Xe
What noble gas has the same electron configuration as each of the ions in the following compounds? a. cesium sulfide b. strontium fluoride c. calcium nitride d. aluminum bromide
Which of the following ions have noble gas electron configurations? a. Fe2+, Fe3+, Sc3+, Co3+ c. Pu4+, Ce4+, Ti4+ b. Tl+, Te2-, Cr3+ d. Ba2+, Pt2+, Mn2+
Give three ions that are isoelectronic with krypton. Place these ions in order of increasing size.
Which compound in each of the following pairs of ionic substances has the most exothermic lattice energy? Jus-tify your answers. a. LiF, CsF b. NaBr, NaI c. BaCl2, BaO d. Na2SO4, CaSO4 e. KF, K2O f.
Predict the empirical formulas of the ionic compounds formed from the following pairs of elements. Name each compound. a. Al and S b. K and N c. Mg and Cl d. Cs and Br
Following are some important properties of ionic compounds:
Use the following data to estimate Î”Hof for potassium chloride.K(s) + 1/2 Cl2(g) †’ KCl(s)
Use the following data to estimate Î”Hof for magnesium fluoride.Mg(s) + F2(g) †’ MgF2(s)
Consider the following: Li(s) + 1/2 I2(g) → LiI(s) ΔH = –292 kJ. LiI(s) has a lattice energy of –753 kJ/ mol. The ionization energy of Li(g) is 520. kJ/ mol, the bond energy of I2(g) is 151
In general, the higher the charge on the ions in an ionic compound, the more favorable is the lattice energy. Why do some stable ionic compounds have 11 charged ions even though 14, 15, and 16
Consider the following energy changes:a. Magnesium oxide exists as Mg2+O2-, not as Mg+O-. Explain. b. What experiment could be done to confirm that magnesium oxide does not exist as Mg+O-?
Use the following data (in kJ/ mol) to estimate Î”H for the reaction S-(g) + e- †’ S2-(g). Include an estimate of uncertainty.S(s) †’ S(g) Î”H = 277 kJ/
Rationalize the following lattice energy values:
The lattice energies of FeCl3, FeCl2, and Fe2O3 are (in no particular order) –2631 kJ/ mol, –5339 kJ/ mol, and –14,774 kJ/ mol. Match the appropriate formula to each lattice energy.
Use bond energy values in Table to estimate Î”H for each of the following reactions in the gas phase.a. H2(g) + Cl2(g) †’ 2HCl(g)Table
Compare your answers from parts a and b of Exercise 43 with Î”H values calculated for each reaction using standard enthalpies of formation in Appendix 4. Do enthalpy changes calculated
Use data from Table to estimate Î”H for the combustion of methane (CH4), as shown below:Table
Use bond energies to estimate ΔH for the combustion of 1 mole of acetylene: C2H2(g) + 5/2 O2(g) → 2CO2(g) + H2O(g)
Consider the following reaction: A2 + B2 → 2AB ΔH = –285 kJ The bond energy for A2 is one- half the amount of the AB bond energy. The bond energy of B2 = 432 kJ/ mol. What is the bond energy of
The space shuttle orbiter uses the oxidation of methyl hydrazine by dinitrogen tetroxide for propulsion:5N2O4(g) + 4N2H3CH3(g) †’ 12H2O(g) + 9N2(g) + 4CO2(g)Use bond energies to estimate DH
Following are three processes that have been used for the industrial manufacture of acrylonitrile€” an important chemical used in the manufacture of plastics, synthetic rubber, and fibers.
Is the elevated temperature noted in parts b and c of Exercise 50 needed to provide energy to endothermic reactions?Exercise 50b.The nitrogen€“ oxygen bond energy in nitric oxide (NO) is
Acetic acid is responsible for the sour taste of vinegar. It can be manufactured using the following reaction:Use tabulated values of bond energies (Table 13.6) to estimate DH for this reaction.
Use bond energies (Table 6), values of electron affinities (Table 8), and the ionization energy of hydrogen (1312 kJ/ mol) to estimate DH for each of the following reactions.a. HF(g) †’
The standard enthalpies of formation of S(g), F(g), SF4(g), and SF6(g) are 1278.8 kJ/ mol, 179.0 kJ/ mol, €“775 kJ/ mol, and €“1209 kJ/ mol, respectively.a. Use these data to
Use the following standard enthalpies of formation to estimate the NOH bond energy in ammonia. Compare this with the value in Table. N(g) 472.7 kJ/ mol H(g) 216.0 kJ/ mol NH3(g) –46.1 kJ/ mol
Write Lewis structures that obey the octet rule for each of the following. Except for HCN and H2CO, the first atom listed is the central atom. For HCN and H2CO, carbon is the central atom. a. HCN b.
Draw a Lewis structure that obeys the octet rule for each of the following molecules and ions. In each case the first atom listed is the central atom. a. POCl3, SO42–, XeO4, PO43–, ClO4– b.
Draw Lewis structures for the following. Show all resonance structures, where applicable. Carbon is the central atom in OCN2 and SCN2. a. NO2-, NO3-, N2O4( N2O4 exists as O2NONO2.) b. OCN2, SCN2, N3-
Some of the pollutants in the atmosphere are ozone, sulfur dioxide, and sulfur trioxide. Draw Lewis structures for these three molecules. Show all resonance structures.
Peroxyacetyl nitrate, or PAN, is present in photochemical smog. Draw Lewis structures (including resonance forms) for PAN. The skeletal arrangement is
A toxic cloud covered Bhopal, India, in December 1984 when water leaked into a tank of methyl isocyanate, and the product escaped into the atmosphere. Methyl isocyanate is used in the production of
Explain the terms resonance and delocalized electrons. When a substance exhibits resonance, we say that none of the individual Lewis structures accurately portrays the bonding in the substance. Why
Benzene (C6H6) consists of a six-membered ring of car-bon atoms with one hydrogen bonded to each carbon. Draw Lewis structures for benzene, including resonance structures.
An important observation supporting the need for resonance in the LE model is that there are only three different structures of dichlorobenzene (C6H4Cl2). How does this fact support the need for the
Borazine (B3N3H6) has often been called “inorganic” benzene. Draw Lewis structures for borazine. Borazine is a six-membered ring of alternating boron and nitrogen atoms.
Draw all the possible Lewis structures for dimethylborazine [(CH3)2B3N3H4]. (See Exercise 67.) Would there be a different number of structures if there was no resonance?
Which of the following statements is(are) true? Correct the false statements. a. It is impossible to satisfy the octet rule for all atoms in XeF2. b. Because SF4 exists, OF4 should also exist because
Lewis structures can be used to understand why some molecules react in certain ways. Write the Lewis structures for the reactants and products in the reactions de-scribed below. a. Nitrogen dioxide
The most common type of exception to the octet rule are compounds or ions with central atoms having more than eight electrons around them. PF5, SF4, ClF3, and Br3- are examples of this type of
SF6, ClF5, and XeF4 are three compounds whose central atoms do not follow the octet rule. Draw Lewis structures for these compounds.
Consider the following bond lengths:In the CO32- ion, all three COO bonds have identical bond lengths of 1.36 Ã…. Why?
Order the following species with respect to the carbon– oxygen bond length (longest to shortest): CO, CO2, CO32-, CH3OH What is the order from the weakest to the strongest carbon–oxygen bond?
Place the species below in order of the shortest to the longest nitrogen– oxygen bond. H2NOH, N2O, NO+, NO2-, NO3- (H2NOH exists as H2NOOH.)
Use the formal charge arguments to rationalize why BF3 would not follow the octet rule.
Use formal charge arguments to explain why CO has a much smaller dipole moment than would be expected on the basis of electronegativity.
Nitrous oxide (N2O) has three possible Lewis structures:Given the following bond lengths, Rationalize the observations that the NON bond length in N2O is 112 pm and that the NOO bond length is 119
Draw Lewis structures that obey the octet rule for the following species. Assign the formal charge to each central atom. a. POCl3 b. SO42- c. ClO4- d. PO43- e. SO2Cl2 f. XeO4 g. ClO3- h. NO43-
Draw the Lewis structures that involve minimum formal charges for the species in Exercise 79. Species in Exercise 79 a. POCl3 b. SO42- c. ClO4- d. PO43- e. SO2Cl2 f. XeO4 g. ClO3- h. NO43-
When molten sulfur reacts with chlorine gas, a vile-smelling orange liquid forms that has an empirical formula of SCl. The structure of this compound has a for-mal charge of zero on all elements in
Oxidation of the cyanide ion produces the stable cyanate ion (OCN2). The fulminate ion (CNO2), on the other hand, is very unstable. Fulminate salts explode when struck; Hg(CNO)2 is used in blasting
Write the Lewis Structure for O2F2 (O2F2 exists as F–O–O–F). Assign oxidation states and formal charges to the atoms in O2F2. This compound is a vigorous and potent oxidizing and fluorinating
Benzoic acid is a food preservative. The space-filling model for benzoic acid is shown below.Draw the Lewis structure for benzoic acid, including all resonance structures in which all atoms have a
Write the name of each of the following molecular structures.
State whether or not each of the following has a permanent dipole moment.
Predict the molecular structure and the bond angles for each molecule or ion in Exercises 57, 58, and 60.
Predict the molecular structure and the bond angles for each of the following. a. SeO3 b. SeO2 c. PCl3 d. SCl2 e. SiF4
There are several molecular structures based on the trigonal bipyramid geometry. Three such structures areWhich of the compounds or ions in Exercises 71 and 72 have these molecular
Predict the molecular structure and the bond angles for each of the following. (See Exercises 89 and 90.) a. XeCl2 b. ICl3 c. TeF4 d. PCl5
Predict the molecular structure and the bond angles for each of the following. a. ICl5 b. XeCl4 c. SeCl6
Which of the molecules in Exercise 88 have net dipole moments (are polar)? SeO3 SeO2 PCl3 SCl2 SiF4
Which of the molecules in Exercises 91 and 92 have net dipole moments (are polar)? Exercise 91. a. XeCl2 b. ICl3 c. TeF4 d. PCl5 Exercise 92. a. ICl5 b. XeCl4 c. SeCl6
Give two requirements that should be satisfied for a molecule to be polar. Explain why CF4 and XeF4 are nonpolar compounds (have no net dipole moments), whereas SF4 is polar (has a net dipole
What do each of the following sets of compounds/ ions have in common with each other? Reference your Lewis structures for Exercises 88, 91, and 92. a. XeCl4, XeCl2 b. ICl5, TeF4, ICl3, PCl3, SCl2,
Which of the following statements is(are) true? Correct the false statements. a. The molecules SeS3, SeS2, PCl5, TeCl4, ICl3, and XeCl2 all exhibit at least one bond angle which is approximately 120
Consider the following Lewis structure, where E is an unknown element:What are some possible identities for element E? Predict the molecular structure (including bond angles) for this ion.
Consider the following Lewis structure, where E is an unknown element:What are some possible identities for element E? Predict the molecular structure (including bond angles) for this ion. (See
Although the VSEPR model is correct in predicting that CH4 is tetrahedral, NH3 is pyramidal, and H2O is bent, the model in its simplest form does not account for the fact that these molecules do not
Draw Lewis structures and predict the molecular structures of the following. (See Exercises 89 and 90.) a. OCl2, KrF2, BeH2, SO2 c. CF4, SeF4, KrF4 b. SO3, NF3, IF3 d. IF5, AsF5 Which of the above
Which of the following molecules have net dipole moments? For the molecules that are polar, indicate the polarity of each bond and the direction of the net dipole moment of the molecule. a. CH2Cl2,
The molecules BF3, CF4, CO2, PF5, and SF6 are all non-polar, even though they contain polar bonds. Why?
Although both the Br3- and I3- ions are known, the F3- ion does not exist. Explain.
The structure of TeF5 2 isDraw a complete Lewis structure for TeF5-, and explain the distortion from the ideal square pyramidal structure. (See Exercise 90.)
There are two possible structures of XeF2Cl2, where Xe is the central atom. Draw them, and describe how measurements of dipole moments might be used to distinguish among them.
Which member of the following pairs would you expect to be more energetically stable? Justify each choice. a. NaBr or NaBr2 c. SO4 or XeO4 b. ClO4 or ClO4- d. OF4 or SeF4
Many times, extra stability is characteristic of a molecule or ion in which resonance is possible. How could this feature be used to explain the acidities of the following compounds? (The acidic
Arrange the following in order of increasing radius and increasing ionization energy. a. N+, N, N- b. Se, Se-, Cl, Cl+ c. Br-, Rb+, Sr2+
Draw a Lewis structure for the N, N-dimethyl formamide molecule. The skeletal structure isVarious types of evidence lead to the conclusion that there is some double- bond character to one of the
A compound, XF5, is 42.81% fluorine by mass. Identify the element X. What is the molecular structure of XF5?
The study of carbon- containing compounds and their properties is called organic chemistry. Besides carbon atoms, organic compounds also can contain hydrogen, oxygen, and nitrogen atoms (as well as
Predict the molecular structure for each of the following. (See Exercises 89 and 90.) a. BrFI2 b. XeO2F2 c. TeF2Cl3- For each formula, there are at least two different structures that can be drawn
Predict the molecular structure of KrF2. Using hypercon-jugation, draw the Lewis structures for KrF2 that obey the octet rule. Show all resonance forms.
Consider the following computer- generated model of caffeine.Draw a Lewis structure for caffeine in which all atoms have a formal charge of zero.
Given the following information: Heat of sublimation of Li(s) = 166 kJ/ mol Bond energy of HCl = 427 kJ/ mol Ionization energy of Li(g) = 520. kJ/ mol Electron affinity of Cl(g) = –349 kJ/ mol
Use data in this chapter and Chapter 12 to discuss why MgO is an ionic compound but CO is not an ionic compound.

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