AP Chemistry

Chemical Bonding
Chapter 4 & 5
Name ______________
Class Period ____

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 Lattice Energy
How strong are the ionic bonds in an ionic solid?

1. Consider the data presented in the above tables.

a. Which set of data could be analyzed to show the effect atomic size has on lattice energy? Explain your
choice.

b. Which set of data could be analyzed to show the effect ion charge has on lattice energy? Explain your
choice.

2. As the ions in the solid lattice get bigger, thus making the bond length longer, what happens to the lattice
energy of the solid? Note: Ignore the sign on the lattice energy. The sign indicates that energy is released.
You are interested in the magnitude of lattice energy. Describe how your answer relates to the law of
Coulombic attraction between charged particles.

3. When the ions in the solid lattice have higher charges, what happens to the lattice energy of the solid?
Describe how your answer relates to the law of Coulombic attraction between charged particles.

4. Circle the compound in each row would have the larger lattice energy and justify your answer.
MgO or MgCl2

MgCl2 or MgF2

MgO or CaO
AlCl3 or Al2O3

5. Match the ionic compounds below to their lattice energy.

Compound Lattice Energy kJ/mol
LiF -2800
Li2O -2240
KF -1030
KBr -820
K2O -680

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Ionic Nomenclature Practice
1. In the blanks above each column, write the charge of the ions formed by atoms in that column. Note:
group IA is done for you.

+1 __
__ __ __ __ __ __

2. Write the symbol and charges for the following ions:

Example: calcium = Ca2+

A) phosphide = ________ B) magnesium = ________ C) rubidium = ________

D) fluoride = ________ E) aluminum = ________ F) sulfide = _________

3. Write the formulas for the compound formed by the combination of the following pairs of atoms.

Example: nitrogen and magnesium = Mg3N2

A) sodium and sulfur = ______________ B) aluminum and chlorine = _______________

C) phosphorus and calcium = __________ D) barium and oxygen = __________________

E) nitrogen and strontium = __________ F) potassium and iodine = __________________

4. For each of the formulas you wrote in question 3 above, give the names for the compounds.

Example: Mg3N2 = magnesium nitride

A) ______________________________ B) _______________________________

C) ______________________________ D) _______________________________

E) ______________________________ F) _______________________________

5. Each of the following formulas is written incorrectly. Please rewrite them correctly.

Example: Ca2Cl = CaCl2

A) Ba2S = _____________ B) Rb2N = _____________ C) Li2Cl = ______________

D) Al3N3 = ____________ E) Mg3Br2 = ___________ F) O3Al2 = ______________

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Advanced Ionic Practice
Provide Formulas and charges for the following Polyatomic Ions:

1. Nitrate 2. Dihydrogen phosphate

3. Carbonate 4. Hypochlorite

5. Arsenite 6. Phosphate

7. dichromate 8. oxalate

Look up the charges for each cation and anion. Write the charges above the words. Balance the
charge and write the correct formula on the line.

12. Lithium acetate _____________________ 13. Sodium phosphate _________________

14. Magnesium hydroxide _______________ 15. Rubidium phosphide ________________

16. Ammonium sulfide __________________ 17. Potassium oxide ___________________

18. Aluminum phosphate ________________ 19. Sodium hydroxide _________________

20. Calcium chloride ___________________ 21. Strontium carbonate ___________________

Write the names for the following formulas.

22. CaO ____________________________ 23. BaCl2 ___________________________

24. K3PO4 ___________________________ 25. Mg(OH)2 ________________________

26. K2C2O4 ___________________________ 27. NaC2H3O2 ________________________

28. Li2SO4 ___________________________ 29. (NH4)2SO4 ________________________

30. Al(CN)3 __________________________ 31. Be(ClO3)2 _________________________

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Ionic and Transition Metals Practice
Write the formulas for each of the following compounds:

1. Manganese (IV) fluoride _______________ 2. Ammonium phosphate ________________

3. Nickel (II) nitrate ____________________ 4. Sodium nitride ______________________

5. Aluminum sulfate __________________ 6. Chromium (III) hydroxide ______________

7. Iron (II) phosphate _________________ 8. Copper (II) chloride __________________

Write the names for each of the following compounds:

9. CuO ____________________________ 10. FeSO4 ___________________________

11. V2O3 ____________________________ 12. Ni3(PO4)2 ________________________

13. Cr(OH)3 _______________________ 14. Ba(ClO3)2 ________________________

15. Mn3N4 ________________________ 16. Cu2CO3 __________________________

How many oxygen atoms are contained in each of the following compounds?

17. Ca(NO3)2 18. Al2O3 19. MgSO4 20. Ca3(PO4)2

21. Naming Ionic Compounds, Given the Ions complete the following chart.
Cation Anion Formula Name
Cu2+ OH
Ba2+ SO42
NH4+ Cr2O72
Ag+ C2H3O2
Fe3+ S2

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Bonding & Molecular Structure
LEWIS STRUCTURES
Indicate the # of VALENCE electrons for each species. Write the correct Lewis electron-dot structure for each.
F O K Al
# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

F O2 K+ Al3+
# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

F2 H2 HF NH3
# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

CH4 NF3 SiF4 C2H6

# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

MgH2 LiH AlH3 BH3

# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

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C2H4 C2F4 CO O2
# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

CO2 C2H2 (H C C H) N2 HCN

# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

CN SO42 PO43 ClO3

# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

CO32 NO3 SO2 O3 (O O O)

# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

SF6 XeF4 PCl5 SeF4

# of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____ # of valence e’s = ____

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I. Writing Formulas of Binary Molecular Compounds—Using Prefixes
Name Formula Name Formula

nitrogen trifluoride phosphorus trichloride

phosphorus
nitrogen monoxide
pentachloride
nitrogen dioxide sulfur hexafluoride
dinitrogen tetroxide disulfur decafluoride
dinitrogen monoxide xenon tetrafluoride

II. Naming Binary Molecular Compounds

Name Formula Name Formula

CCl4 HBr
P4O10 N2F4
ClF3 XeF3
BCl3 PI3
SF4 SCl2

III. Mixed Practice

Formula Name Formula Name

AgCl carbon dioxide

PCl5 ammonium carbonate

K2S sulfur dichloride

NiSO4 calcium iodide

ClF3 boron trifluoride

OF2 phosphorus triiodide

Al(OH)3 magnesium perchlorate

NCl3 potassium permanganate

(NH4)3PO4 aluminum phosphate

S2Cl2 dioxygen difluoride

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ACID NOMENCLATURE NAME______________________PER______

I. Write the correct name or formula for each of the following acids.

1. HNO3 _______________________________________________________

2. oxalic acid _______________________________________________________

3. H2SO4 ______________________________________________________

4. H3PO4 ______________________________________________________

5. HClO3 ______________________________________________________

6. HNO2 _______________________________________________________

7. hydroselenic acid _______________________________________________

8. HClO ______________________________________________________

9. hydrofluoric acid ________________________________________________

10. H2SO3 _____________________________________________________

11. hydrocyanic acid ________________________________________________

12. perchloric acid ____________________________________________________

13. NH4+ _____________________________________________________

14. CH3COOH _______________________________________________

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Additional Practice:
Name the following compounds:
1. Na3PO4
2. Fe2(Cr2O7)3
3. CH3COOH (two names)
4. NF3
5. PbO2
6. FeF2 • 8 H2O
7. HBr
8. B2O3
9. Li2SO4
10. MgSO4 • 2H2O
11. CdSO3
12. MnF6
13. PbCrO4
14. H3PO4
15. NH4C2H3O2
16. Ba(MnO4)2

Write the formulas for the following compounds.

17. Diphosphorous tetrachloride
18. Magnesium arsenate
19. Copper II nitrate trihydrate
20. Manganese II sulfide
21. Sodium phosphate
22. Butane
23. Calcium chloride
24. Carbon dioxide
25. Cobalt III hydroxide
26. Tetraphophorous hexoxide
27. Copper II nitrate
28. Lead IV chromate
29. Iron III oxide
30. Aluminum hydrogen sulfate
31. Potassium citrate
32. Oxalic acid
33. Cobalt II carbonate
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Molecular Shapes

The shapes of molecules can be predicted from their Lewis Structures by using the VSEPR (Valence Shell Electron
Pair Repulsion) model, which states that electron pairs around a central atom will assume a geometry that keeps
them as far apart from each other as possible.

This theory is illustrated by the drawings below.

Six groups surrounding a central atom will form an Octahedron.

All of the groups in this structure are at 90o or 180o to each other. All positions
are equivalent.

Five groups will form a trigonal bipyramid. The two positions pointing up and
down are called the axial positions. They are at 180o to each other, and at 90o
to the other three, equatorial positions. The three equatorial positions are at
120o to each other.

Four groups will form a tetrahedron. All of the angles in a tetrahedron are
109.5o and all positions are equivalent.

Three groups will form a flat triangle (trigonal planar). Each of the angles is
120o and all positions are equivalent.

Two groups form a straight line (Linear) with 180o between them.

How does this apply to Chemistry?

 The groups occupying these geometric positions will be either, atoms bonded to the central atom, or lone
pair electrons on the central atom.
 Lone Pair electrons occupy more space than bonded electrons, so they will take the equatorial position in
the trigonal bipyramid.
 Lone pair electrons will also occupy positions that put them as far apart from each other as possible.

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1) Draw the Lewis structure for water, H2O.

a. How many "groups" (atoms and lone pairs) surround the central oxygen? ___________

b. What is the geometry of this molecule (look at atoms and lone pairs)? ______________________ Draw
this VSEPR structure next to the Lewis structure above.

c. What is the shape of this molecule (look only at the atoms)? _____________________

d. What is the H-O-H bond angle? ___________________

2) Draw the Lewis structure for NO2-.

a. How many "groups" (atoms and lone pairs) surround the central nitrogen? __________

b. What is the geometry of this molecule (look at atoms and lone pairs)? ______________________ Draw
this VSEPR structure next to the Lewis structure above.

c. What is the shape of this molecule (look only at the atoms)? _____________________

d. What is the O-N-O bond angle?

3) Draw the Lewis and VSEPR structures for the following 12 compounds and label them with their geometry.
Compound Lewis Structure VSEPR Geometry
Valence (Ball & Stick Model)
Electrons

SF6

ICl2-

ICl4-

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 SF4

CCl4

BrF5

BrF3

NH3

CO2

XeCl3-

SO3

PF5

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 Fill in the missing information in the chart using the structures you have drawn in problems 1 - 3.

compound atoms on central lone pairs on geometry shape

atom central atom

SF6 octahedral

5 1

4 octahedral

XeCl3-

5 0

4 1 seesaw

BrF3

trigonal bipyramidal linear

4 0

NH3

2 2 bent

trigonal planar

2 1

CO2

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 Molecular Polarity Lab
https://goo.gl/zQkLwf

Go to the website listed above. Click the play symbol on the simulation.

Part 1: Two Atom Investigation

❏ Select the Two Atom Investigation
❏ In the “View” Box on the right side of the screen, check bond dipole, partial charges and the
bond character. ( All three should be checked)
❏ Keep the electronegativity of A at low and increase and decrease the electronegativity of atom
B. Observe the arrow, partial charge and bond character. Fill in the following observation
As the electronegativity of atom B increases the...
arrow _________________________________________________________________________

partial charges ________________________________________________________________

bond character ________________________________________________________________

❏ Now put the electronegativity for A to the middle and vary the electronegativity of B and
observe the results.
Fill in the following information.
o The polarity arrow always points to the _________________________ electronegative atom.

o The partial positive charge is always on the ______________________ electronegative atom.

o The larger the electronegativity difference the more _________________ the bond character.

Part 2: Three atom Investigation

❏ Keep the original “bent” arrangement of atoms and set the electronegativity’s of A and C to
low and B to middle. Click to show the bond
dipoles, molecular dipoles and partial charges.
❏ Draw the bond dipoles on the bonds in the
diagram
❏ In a different color show the overall
molecular dipole in the diagram
❏ Draw the partial charge symbols in another
color in the diagram.
❏ Click the electric field on. Describe what
happened to the molecule - be very
specific.

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❏ Click the reset button. Keep the “bent” arrangement of atoms and set the
electronegativity’s of A and C to high and B to
middle. Click to show the bond dipoles, molecular
dipoles and partial charges.
❏ Draw the bond dipoles on the bonds in the
diagram
❏ In a different color show the overall
molecular dipole in the diagram
❏ Draw the partial charge symbols in another
color in the diagram.
❏ Click the electric field on. Describe what
happened to the molecule - be very specific.

❏ Click the reset button. Change the arrangement of atoms to a “linear” arrangement (see
below) and set the electronegativity’s of A and C to high and keep B to middle. Click to show
the bond dipoles, molecular dipoles and partial
charges.
❏ Draw the bond dipoles on the bonds in the
diagram
❏ In a different color show the overall
molecular dipole in the diagram
❏ Draw the partial charge symbols in another color in the diagram.
❏ Click the electric field on. Describe what happened to the molecule - be very specific.

❏ Take off the electric field. Rotate the molecule to keep the “linear” arrangement (see below)
and set the electronegativity’s of A to high, keep B to middle and set C to low. Click to show
the bond dipoles, molecular dipoles and partial
charges.
❏ Draw the bond dipoles on the bonds in the
diagram
❏ In a different color show the overall
molecular dipole in the diagram
❏ Draw the partial charge symbols in another color in the diagram.
❏ Click the electric field on. Describe what happened to the molecule - be very specific.

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 ❏ Take off the electric field. Rotate the molecule to keep the “linear” arrangement (see below)
and set the electronegativity’s of A and C to low and keep B to middle and set C to low. Click to
show the bond dipoles, molecular dipoles and
partial charges.
❏ Draw the bond dipoles on the bonds in the
diagram
❏ In a different color show the overall
molecular dipole in the diagram
❏ Draw the partial charge symbols in another color in the diagram.
❏ Click the electric field on. Describe what happened to the molecule - be very specific.

❏ Take off the electric field. Rotate the molecule to keep the “linear” arrangement (see below)
and set the electronegativity’s of A, Band C to middle. Click to show the bond dipoles,
molecular dipoles and partial charges.
❏ Draw the bond dipoles on the bonds in the
diagram
❏ In a different color show the overall
molecular dipole in the diagram
❏ Draw the partial charge symbols in another
color in the diagram.
❏ Click the electric field on. Describe what happened to the molecule - be very specific.

Summary
Summarize what you learned by answering the following questions.
1. Which way do the bond arrows point?

2. Can a molecule have bond dipoles but not have a molecular dipole? Explain.

3. What happens when a molecule with a dipole is put in an electric field? Be specific.

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Part 3: Real Examples: Apply What You Learned
You will now apply what you learned to real molecules. The real molecules does not work on the
simulation so just apply what you know.
HF - The ball and stick structure for HF is shown. Answer the following and do what is asked
❏ Which atom is more electronegative? _______________
❏ Draw a bond polarity arrow (bond dipole)
❏ Draw the partial charges on the molecule
❏ Would you expect this to move in an electric field? Draw it in the field provided.

H2O - The ball and stick structure for H2O is given. Answer the following and do what is asked.
❏ Which atom is more electronegative? _______________
❏ Draw a bond polarity arrow (bond dipole)
❏ Place partial charges on the molecule
❏ In a different color, draw a molecular dipole arrow.
❏ Would you expect this to move in an electric field? Draw it in the field provided.

CO2 - The ball and stick structure for CO2 is given. Answer the following and do what is asked.
❏ Which atom is more electronegative? _______________
❏ Draw a bond polarity arrow (bond dipole)
❏ Place partial charges on the molecule
❏ In a different color, draw a molecular dipole arrow.
❏ Would you expect this to move in an electric field? Draw it in the field provided.

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