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8.E: Chemical Bonding and Molecular Geometry- Homework

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    Turn in your answers for the following questions - show your work

    Draw Lewis dot structures, determine the molecular shape, electronic geometry, bond length, determine the bond energy, determine the bond polarity, and determine the formal charge for each atom.

    1. a. H2
    2. b. CH4
    3. c. NH3
    4. d. H2O
    5. e. HBr
    6. f. PCl5
    7. g. CO2
    8. h. BF3
    9. i. CF2Cl2
    10. j. C4H8 (3 different structures)

     

    The Following Questions are for your practice - Do Not Turn In.  They include answers so you can check your work

    Ionic Bonding

    Q7.1.5

    Predict the charge on the monatomic ions formed from the following atoms in binary ionic compounds:

    1. Mg
    2. Al
    3. O
    4. Cl

    S7.1.5

    Mg2+; Al3+; O2–; Cl–

    Covalent Bonding

    Predict which of the following compounds are ionic and which are covalent, based on the location of their constituent atoms in the periodic table:

    1. Cl2CO
    2. MnO
    3. NCl3
    4. CoBr2

    ionic: (b), (d); covalent: (a), (c)

    From its position in the periodic table, determine which atom in each pair is more electronegative:

    1. Br or Cl
    2. N or O

    Cl; O

    Identify the more polar bond in each of the following pairs of bonds:

    1. HF or HCl
    2. NO or CO

    HF; CO

    Lewis Symbols and Structures

    Q7.4.1

    Write the Lewis symbols for each of the following ions:

    1. As3–
    2. I–

    S7.4.1

    eight electrons:

    CNX_Chem_07_03_Question1a_img.jpg

    eight electrons:

    CNX_Chem_07_03_Question1b_img.jpg

    Q7.4.3

    Write the Lewis symbols of the ions in each of the following ionic compounds and the Lewis symbols of the atom from which they are formed:

    1. MgS
    2. Al2O3

    (a)

    Two Lewis structures are shown. The left shows the symbol M g with a superscripted two positive sign while the right shows the symbol S surrounded by eight dots and a superscripted two negative sign. ;

    (b)

    Two Lewis structures are shown. The left shows the symbol A l with a superscripted three positive sign while the right shows the symbol O surrounded by eight dots and a superscripted two negative sign. ;

     

    Write Lewis structures for the following:

    1. O2
    2. H2CO
    3. AsF3

    (a)

    A Lewis structure shows two oxygen atoms double bonded together, and each has two lone pairs of electrons.

    (b)

    A Lewis structure shows a carbon atom that is single bonded to two hydrogen atoms and double bonded to an oxygen atom. The oxygen atom has two lone pairs of electrons. ;

    (c)

    A Lewis structure shows an arsenic atom single bonded to three fluorine atoms. Each fluorine atom has a lone pair of electrons. ;

     

     

    Write Lewis structures for the following:

    1. SeF6
    2. XeF4
    3. \(\ce{SeCl3+}\)

    SeF6:

    A Lewis structure shows a selenium atom single bonded to six fluorine atoms, each with three lone pairs of electrons. ;

    XeF4:

    A Lewis structure shows a xenon atom with two lone pairs of electrons. It is single bonded to four fluorine atoms each with three lone pairs of electrons. ;

    \(\ce{SeCl3+}\):

    A Lewis structure shows a selenium atom with one lone pair of electrons single bonded to three chlorine atoms each with three lone pairs of electrons. The whole structure is surrounded by brackets.

    Methanol, H3COH, is used as the fuel in some race cars. Ethanol, C2H5OH, is used extensively as motor fuel in Brazil. Both methanol and ethanol produce CO2 and H2O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.

    Two reactions are shown using Lewis structures. The top reaction shows a carbon atom, single bonded to three hydrogen atoms and single bonded to an oxygen atom with two lone pairs of electrons. The oxygen atom is also bonded to a hydrogen atom. This is followed by a plus sign and the number one point five, followed by two oxygen atoms bonded together with a double bond and each with two lone pairs of electrons. A right-facing arrow leads to a carbon atom that is double bonded to two oxygen atoms, each of which has two lone pairs of electrons. This structure is followed by a plus sign, a number two, and a structure made up of an oxygen with two lone pairs of electrons single bonded to two hydrogen atoms. The bottom reaction shows a carbon atom, single bonded to three hydrogen atoms and single bonded to another carbon atom. The second carbon atom is single bonded to two hydrogen atoms and one oxygen atom with two lone pairs of electrons. The oxygen atom is also bonded to a hydrogen atom. This is followed by a plus sign and the number three, followed by two oxygen atoms bonded together with a double bond. Each oxygen atom has two lone pairs of electrons. A right-facing arrow leads to a number two and a carbon atom that is double bonded to two oxygen atoms, each of which has two lone pairs of electrons. This structure is followed by a plus sign, a number three, and a structure made up of an oxygen with two lone pairs of electrons single bonded to two hydrogen atoms.

     

    The arrangement of atoms in several biologically important molecules is given here. Complete the Lewis structures of these molecules by adding multiple bonds and lone pairs. Do not add any more atoms.

    the amino acid serine:

    A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to a hydrogen atom and two other carbon atoms. One of these carbon atoms is single bonded to two hydrogen atoms and an oxygen atom. The oxygen atom is bonded to a hydrogen atom. The other carbon atom is single bonded to two oxygen atoms, one of which is bonded to a hydrogen atom.

    urea:

    A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to an oxygen atom and another nitrogen atom. That nitrogen atom is then single bonded to two hydrogen atoms.

     

    (a)

    A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to a hydrogen atom and two other carbon atoms. One of these carbon atoms is single bonded to two hydrogen atoms and an oxygen atom. The oxygen atom is bonded to a hydrogen atom. The other carbon is single bonded to two oxygen atoms, one of which is bonded to a hydrogen atom. The oxygen atoms have two lone pairs of electron dots, and the nitrogen atom has one lone pair of electron dots. ;

    (b)

    A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to an oxygen atom and one nitrogen atom. That nitrogen atom is then single bonded to two hydrogen atoms. The oxygen atom has two lone pairs of electron dots, and the nitrogen atoms have one lone pair of electron dots each. ;

     

    Formal Charges and Resonance

    Draw all possible resonance structures for each of these compounds. Determine the formal charge on each atom in each of the resonance structures:

    1. SO2
    2. \(\ce{NO2-}\)

    (a)

    Two Lewis structures are shown, with a double-headed arrow in between. The left structure shows a sulfur atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. The sulfur atom also double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read, “( negative 1 ), ( positive 1 ), ( 0 ).” The right structure appears as a mirror image of the left and the symbols and numbers below this structure read, “( 0 ), ( positive 1 ), ( negative 1 ).” ;

    (b)

    [Two Lewis structures are shown, with brackets surrounding each with a superscripted negative sign and a double ended arrow in between. The left structure shows a nitrogen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons and double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read “open parenthesis, 0, close parenthesis, open parenthesis, 0, close parenthesis, open parenthesis, negative 1, close parenthesis. The right structure appears as a mirror image of the left and the symbols and numbers below this structure read “open parenthesis, negative 1, close parenthesis, open parenthesis, 0, close parenthesis, open parenthesis, 0, close parenthesis.] 

    Strengths of Ionic and Covalent Bonds

    Using the bond energies in Table, determine the approximate enthalpy change for each of the following reactions:

    1. \(\ce{H2}(g)+\ce{Br2}(g)⟶\ce{2HBr}(g)\)
    1. −114 kJ;

     

    Complete the following Lewis structure by adding bonds (not atoms), and then indicate the longest bond:

    A Lewis structure is shown that is missing its bonds. It shows a horizontal row of six carbon atoms, equally spaced. Three hydrogen atoms are drawn around the first carbon, two around the second, one above the fifth, and two by the sixth.

    A Lewis structure is shown. A carbon atom that is single bonded to three hydrogen atoms is bonded to a second carbon atom. The second carbon atom is single bonded to two hydrogen atoms. The second carbon atom is single bonded to a third carbon atom that is triple bonded to a fourth carbon atom and single bonded to a fifth carbon atom. The fifth carbon atom is single bonded to a hydrogen atom and double bonded to a sixth carbon atom that is single bonded to two hydrogen atoms.

    The C–C single bonds are longest.

    Molecular Structure and Polarity

    Predict the electron pair geometry and the molecular structure of each of the following molecules or ions:

    1. SF6
    2. PCl5
    3. (c) BeH2
    4. \(\ce{CH3+}\)
    1. Both the electron geometry and the molecular structure are octahedral.
    2. Both the electron geometry and the molecular structure are trigonal bipyramid.
    3. (c) Both the electron geometry and the molecular structure are linear.
    4. Both the electron geometry and the molecular structure are trigonal planar.

     

    Identify the electron pair geometry and the molecular structure of each of the following molecules:

    1. ClNO (N is the central atom)
    2. CS2
    3. (c) Cl2CO (C is the central atom)
    4. Cl2SO (S is the central atom)

    (a) electron-pair geometry: trigonal planar, molecular structure: bent (120°); (b) electron-pair geometry: linear, molecular structure: linear; (c) electron-pair geometry: trigonal planar, molecular structure: (d) trigonal planar; electron-pair geometry: tetrahedral, molecular structure: trigonal pyramidal

     

    Which of the following molecules have dipole moments?

    1. CS2
    2. SeS2
    3. (c) CCl2F2
    4. PCl3 (P is the central atom)
    5. ClNO (N is the central atom)

    SeS2, CCl2F2, PCl3, and ClNO all have dipole moments.

     

    Describe the molecular structure around the indicated atom or atoms:

    1. the sulfur atom in sulfuric acid, H2SO4 [(HO)2SO2]
    2. the chlorine atom in chloric acid, HClO3 [HOClO2]
    3. (c) the oxygen atom in hydrogen peroxide, HOOH
    4. the nitrogen atom in nitric acid, HNO3 [HONO2]

    (a) tetrahedral; (b) trigonal pyramidal; (c) bent (109°); (d) trigonal planar; bent (109°); 

    A molecule with the formula AB2, in which A and B represent different atoms, could have one of three different shapes. Sketch and name the three different shapes that this molecule might have. Give an example of a molecule or ion for each shape.

    Three Lewis diagrams are shown. The first diagram shows the letter A single bonded to the left and right to the letter B. An example, “C O subscript 2,” and the term, “linear,” are written beside this diagram. The second diagram shows the letter A with two lone pairs of electrons, single bonded to the left and lower right to the letter B. An example, “H subscript 2 O,” and the term, “bent with an approximately 109 degree angle,” are written beside this diagram. The third diagram shows the letter A with one lone electron pair, single bonded to the left and lower right to the letter B. An example, “S O subscript 2,” and the term, “bent with an approximately 120 degree angle,” are written beside this diagram.

    Use the Molecule Shape simulator to explore real molecules. On the Real Molecules tab, select H2O. Switch between the “real” and “model” modes. Explain the difference observed.

    The structures are very similar. In the model mode, each electron group occupies the same amount of space, so the bond angle is shown as 109.5°. In the “real” mode, the lone pairs are larger, causing the hydrogens to be compressed. This leads to the smaller angle of 104.5°.

     

    This page titled 8.E: Chemical Bonding and Molecular Geometry- Homework is shared under a CC BY license and was authored, remixed, and/or curated by Scott Van Bramer.