Unit 11 Reading Guide: Review of Bonding Part 1:

Type of

Cmpd

How is bond

formed?

Properties of

cmpd?

Usually composed of these

types of atoms:

Example (name and

formula)

Ionic

Metallic

Covalent

IONIC BONDING: READ IN YOUR TEXTBOOK AND COMPLETE EACH OF THE QUESTIONS BELOW.

EXTRA RESOURCES: HTTP://CHEMWIKI.UCDAVIS.EDU/THEORETICAL_CHEMISTRY/CHEMICAL_BONDING/COVALENT_BONDS_VS_IONIC_BONDS HTTPS://WWW.YOUTUBE.COM/WATCH?V=QXT4OVM4VXI

1. What is bond energy?

2. Why do sodium and chlorine combine to form sodium chloride?

3. How can we calculate the energy of attraction between a pair of ions?

4. What is the energy form of Coulomb’s law? What does each variable represent?

5. Why are ionic compounds solids at room temperature?

6. How can we predict the formula of an ionic compound?

7. Describe an ion’s size compared to its parent atom.

8. Describe the trend that summarizes ionic radius for an isoelectronic series.

9. What is used to indicate how strongly ions attract each other in the solid state?

10. Write an equation that represents lattice energy.

11. Consider the energy transition involved in the formation of lithium fluoride. Identify each equation

with the description of its energy term.

a) Li (s) Li (g) H = + 161 kJ/mol __________________

b) Li (g) Li + (g) H = + 520 kJ/mol _________________

c) ½ F2 (g) F (g) H = + ½ (154 kJ/mol) __________________

d) F (g) + e - F – (g) H = - 328 kJ/mol __________________

e) Li + (g) + F – (g) LiF (s) H = - 1047 kJ/mol __________________

12. Using Hess’s Law of Summation and the values from question #11, determine the enthalpy of

formation for lithium fluoride.

Li (s) + ½ F2 (g) LiF (s) H = ?

13. List the factors that determine the value of the lattice energy of an ionic compound?

14. Write the electron configuration for the following ions:

the most common ion of Ca the most common ion of S

14. Define isoelectronic. HINT: look at the results of question 14.

15. What is the octet rule?

Reading Guide Part 2: Draw the Lewis Structure for each atom, ion, or molecule below. Include all resonance structures

and check formal charge where needed to ensure proper charge.

1. Na

2. S

3. LiF

4. aluminum nitride

5. Cl –

6. NH4+

7. carbon disulfide

8. NO3 –

9. H2S

10. NH3

11. CF4

12. SO2

For questions above, write down the molecular geometry (shape) for each structure.

Also list the number of e- domains and type of hybridization the central atom would undergo.

Use one of the molecules above to explain resonance.

Lastly, attempt to draw all of the hybridized orbitals with overlap indicating bonds. See me if

you need help with this ’

AP CHEMISTRY: LEWIS STRUCTURE & MOLECULAR GEOMETRY PROBLEM SET

*DRAW THE LEWIS STRUCTURE FOR EACH MOLECULE OR ION BELOW. WRITE THE FORMAL CHARGES FOR EACH OF THE LEWIS STRUCTURES AND EXPLAIN

ANY INCONSISTENCIES.

* IDENTIFY ITS MOLECULAR SHAPE, BOND ANGLES, AND MOLECULAR

POLARITY

1) HCN 2) H2S 3) SnCl4 4) NCl3 5) SOF2 6) AsF2 +

7) NO3 – 8) PCl4

+ 9) ClO3 – 10) TeF4 11) AsF5 12) BrF5

13) ICl2 + 14) H3O

+ 15) AsF3 16) H2O 17) SiF2 18) SF6

19) I3-

BONDING GUIDED NOTES

To draw Lewis electron dot structures

1. Sum the ve- from all atoms. Add 1 for each negative charge and subtract

1 for each positive charge.

2. Choose a central atom (usually, singleton, or largest atom)

3. Draw a skeleton structure with atoms/ligands attached by single bonds.

4. Complete the octets of atoms bound to the central atom (outer atoms).

5. Place extra e- on the central atom. If the central atom still doesn’t have

an octet, try forming multiple bonds.

Draw correct Lewis Structures, shape, bond angles:

O2 N2 CF4 NH3 CH2O CO2 SO42- BeCl2

Formal Charge for an atom in a molecule =

(# ve- on free atom) – (# ve- assigned to the atom in the molecule)

Where the “number of ve- assigned to the atom in the molecule

“ is equal to the number of nonbonding electrons + 1/2 number

of bonding electrons

- if more ve- that normal ve-, then (-) charge | is less ve- than normal

ve-, then (+) charge

- A formal charge of zero for each atom in a molecule is a very

common result for a favorable Lewis structure.

ATOMIC ORBITAL THEORY: Describes covalent bonding as the mixing of native atomic orbitals to form a new

kind of orbital, a hybrid orbital.

O Hybrid orbitals are atomic orbitals (AOs) formed as a result of mixing

the AOs of similar energies of the atoms involved in the

covalent bond. The # of hybrid orbitals formed is the same as

the # of AOs mixed, and the type of hybrid orbital formed

depends on the types of AOs mixed.

AOs have particular shapes....

o When an atom is about to become part of a molecule, that atom’s AOs

must morph into a new set of orbitals with different shapes and

orientations than the original AOs. These new morphed orbitals are called

hybrid orbitals.

AX3 are sp2 hybridized = 3 effective pair around an atom will always require sp2

hybridization! *****Note that there is one unhybridized p orbital that is still

available for bonding (double bonds)

2 effective pair around an atom will always require sp hybridization!

5 effective pair = sp3d hybridization , blending of an s orbital, three p

orbitals, and one d orbital.

6 effective pair = sp3d2 hybridization, blend of one s, three p, and two d

orbitals

Sigma bonds (bond) - Bond in which the e- pair is shared in an area

centered on a line running between the atoms, internuclear axis, the e-

density is along line joining 2 nuclei

Pi bonds (bonds) - e- pair above and below the bond, form by parallel

orbitals, formed by sideways overlapping of orbitals, the e-density above

and below plane of nuclei; less overlap means weaker than sigma bonds

MOLECULAR ORBITAL THEORY:

When 2 AOs are added, 2 MOs are formed, 1 bonding and 1

antibonding. The bonding MO is of lower energy than the

antibonding MO.

Calculate Bond order and predict the stability of a particular

compound

BO = ½(# of e- in bonding orbitals - # of e- in antibonding

orbitals)

BO > 0 stable ; BO < or = 0 not stable

MOLECULAR GEOMETRY GUIDE

EFFECTIVE

ELECTRON PAIRS

ON

CENTRAL ATOM

SHARED

ELECTRO

N PAIRS

UNSHAR

ED

ELECTRO

N

PAIRS

Class SHAPE

(NAME & BOND

ANGLE)

NOTES

2

0

AX2

LINEAR

180

3

0

AX3

TRIGAONAL PLANAR

120

2

1

AX2E

ANGULAR <

120

4

0

AX4

TETRAHEDRAL

109.5

3

1

AX3E

PYRAMIDAL <

109.5

2 2 AX2E2

ANGULAR <<

109.5

1

3

AXE3

LINEAR

180

EFFECTIVE

ELECTRON PAIRS

ON

CENTRAL ATOM

SHARED

ELECTRO

N PAIRS

UNSHAR

ED

ELECTRO

N

PAIRS

Class SHAPE

(NAME & BOND

ANGLE)

NOTES

5

0

AX5

TRIGONAL

BIPYRMIDAL

120 AND 90

4

1

AX4E

SEE-SAW 120 &

90

3

2

AX3E2

T-SHAPED

90

2

3

AX2E3

LINEAR

180

6

0

AX6

OCTAHEDRAL

90

5

1

AX5E

SQUARE PYRAMIDAL

90

4

2

AX4E2

SQUARE PLANAR

90

Formula Lewis Structure Hybridization Geometry Polarity

CO2

SO2

AX2 H2O

XeF2

BCl3

AX3 NH3

ClF3

CH4

AX4

SCl4

XeF4

AX5

PCl5

ClF5

AX6

SF6

Choose 5 molecules to draw with hybridized orbitals. Draw them and staple to this chart.

AP Chemistry: Practice Lewis Structures, Hybridization, Geometry, and Polarity!

AP Chemistry: Bond Energy Problem Set Draw the structures and show work with your calculations of enthalpy of reaction for

each of the following: 1) H2 + Cl2 2HCl

2) N2 + 3H2 2NH3

3) CH2=CH2 + Br2 CH2BrCH2Br

4) C2H4 + H2O CH2OH – CH2OH

5) CH3NC CH3CN

6) C2H4 + O3 CH3CHO + O2

7) 5N2O4 + 4 N2H3CH3 12 H2O + 9N2 + 4CO2

Bonding Energies B O N D E N E R G I E S

Table 9.9 Some Average Single- and Multiple-Bond Energies (kJ/mol)

H C N O F Si P S Cl Br I

H 436 413 391 463 565 318 322 347 432 366 299

C 346 305 358 485 272 339 285 213

N 163 201 283 192

O 146 452 335 218 201 201

F 155 565 490 284 253 249 278

Si 222 293 381 310 234

P 201 326 184

S 226 255

Cl 242 216 208

Br 193 175

I 151

Multiple Bonds

N=N 418 C=C 602

NN 945 CC 835

C=N 615 C=O 732

CN 887 CO 1072

O=O (in O2) 498

Table 6.2 Standard Enthalpies of Formation (kJ/mol)

C2H6(g) ethane -84.7

H2O(g) water vapor -241.8

CO2(g) carbon dioxide -393.5

1. Write the balanced chemical equation for the complete combustion of ethane, C2H6(g).

2. Draw structural formulas (shortcut Lewis structures) for each of the species.

3. Calculate the energy needed to break the bonds in the reactants. ________

Calculate the energy released as the bonds in the products are formed. ________

4. What is the Hcombustion based on bond energies? ____________

5. On the back of this page, calculate the Hcombustion using Hess’s Law and the thermochemical

data from Chapter 6.

AP Chemistry: Unit Review Bonding

1. What are the two factors that affect lattice energy? Which is more important?

2. For each of the following pairs of ionic compounds, box the compound that has the highest

exothermic lattice energy.

a. LiF, LiCl b. Mg(OH)2 , MgO c. NaCl, Na2O

d. MgO, BaS e. Fe(OH)2, Fe(OH)3 d. NaCl, KCl

3. Using the periodic trend in electronegativity, place each group of elements in order of increasing

electronegativity.

a. C, N, O b. S, Se, Cl c. Si, Ge, Sn d. Tl, S, Ge

4. Using Pauling’s electronegativty values, predict bond type for the following compounds.

a. C-F b. Si-F c. Si-Cl d.Ge-F

5. Place the bonds from #4 in order of decreasing polarity.

6. List three ions that are isoelectronic with the krypton atom.

7. Use the following data to estimate Hf for sodium chloride.

a. Write the equation for the formation of sodium chloride.

8. Use the table of bond energies to calculate the H for the following reactions:

a.. CH3NC(g) CH3CN (g)

b. C2H4(g) + O3(g) CH3CHO(g) + O2(g)

9. Write Lewis structures that obey the octet rule for each of the following. Indicate the polarity

of the bond, the structure (including bond angles) of the molecule and if a dipole moment exists.

Indicate the hybridization.

a. HCN

b. CHCl3

c. BF4-

d. PH3

e. NH4+

f. SeF2

g. COCl2

h. C2N2

Now for ones that do not follow the octet rule:

a. XeOF4

b. BCl3

c. XeF6=

10. Order the following species with respect to C-O bond length (longest to shortest) and order from

weakest to strongest: CO, CO2, CO32-, CH3OH

Lattice energy -786kJ/mol

Ionization energy for Na 495kJ/mol

Electron affinity of Cl -348kJ/mol

Bond Energy for Cl2 239 kJ/mol

Enthalpy of sublimation for Na 108kJ/mol

11. Complete all questions for the following compounds

a. lycopene:

i. How many carbons are sp3 hybridized? sp2 hybridized? sp hybridized?

b. The acetate ion has both oxygen atoms bonded to the same carbon:

i. Draw the Lewis structure and all resonance structures

ii. Label the hybridization around each carbon

iii. Pick one resonance structure and label the hybridization of each oxygen

iv. How many sigma and pi bonds are present?

v. Which atom carries the formal negative charge?

c. Acylonitrile

i. ii. What are the hybridizations & bond angles around each of the numbered atoms? iii. How many sigma and pi bonds are there?

d. Caffeine

i. ii. What are the hybridizations & bond angles around each of the numbered atoms? iii. How many sigma and pi bonds are there?

12. Metallic bonds are different from ionic and covalent bonds because the electrons in metallic

bonds are delocalized. The electrons are not assigned to a specific atom. This is important

because it gives metals the ability to conduct heat and electricity, makes them malleable (easily

shaped) and ductile (very thin). When two atoms are smushed together as they would be if the

metal was being hammered into a new shape, the electrons repel but are able to move apart easily.

In an ionic bond, because electrons are localized (assigned a specific atom) if two like charges

come into contact, they repel and cause the material to crack. This explains the brittleness of

ionic compounds. The localized electron model also explains why ionic compounds can’t conduct

electricity – the ions can’t move past each other unless they are either liquefied or dissolved in

water. Although metals and salts have similar lattice structures, metals make good materials for

electrical wiring. Why aren’t salts used instead? (more than one reason)

13. Which has greater potential energy, a noble gas or a

metal? Explain.

14. Define all terms in box to left:

15. Make a table on a separate sheet and draw what these

intermolecular forces look like on an atomic level. Give

examples if it helps you draw realistic forces (see the

list below in #16). Here is a helpful doc

http://www.smallscalechemistry.colostate.edu/Powerful

Pictures/ChemicalBonding.pdf or pg 495 in book

16. Rank the following in order of increasing strength: dispersion forces, hydrogen bonding, dipole-

dipole forces (aka permanent dipole-dipole), ion-permanent dipole force, ion-induced dipole force,

permanent dipole –induced dipole force.

17. Give an example of two organic isomers with different functional groups. Draw these isomers and

explain how they fit the definition of “isomer.” Which of them has stronger intermolecular

forces?

Complete the PES Questions below:

18. According to the shell model, why do first ionization energies increase across a row (period) of

the periodic table?

19. According to the shell model, why do first ionization energies decrease down a column (group) of

the periodic table?

20. Do all of the electrons in a given shell have the same energy? Why or why not?

21. Why are the number of peaks in the PES for H and He the same?

22. Roughly sketch the photoelectron specta for Al and S. Give the relative intensities for each peak.

23. What element do you think should give rise to the PES shown below (Fig 3P1)? Explain your

reasoning…

a. London Dispersion Forces

b. Dipole/Dipole Moment

c. Temporary/Induced Dipole

d. Permanent Dipole

e. Dipole-Dipole Forces

f. Hydrogen Bonding

g. Van der Waals Forces

Ch 8, 9, 10 MC Self Assessment Questions (highly recommended, but not required)

Chemistry in the Environment: Required

~ Page 407 - read and answer these questions about free radicals

~ Page 414 - read and answer these questions about ozone

AP Chemistry – Bonding FRQs

FRQ 2005 (Required)

FRQ 2003 (Required)

FRQ 1997 D (Required)

Consider the molecules PF3 and PF5.

(a) Draw the Lewis electron-dot structures for PF3 and PF5 and predict the molecular geometry of each.

(b) Is the PF3 molecule polar, or is it nonpolar? Explain.

(c) On the basis of bonding principles, predict whether each of the following compounds exists. In each case,

explain your prediction.

(i) NF5

(ii) AsF5

FRQ 1978 D (Required)

Dehydration of 3-hexanol yields a mixture of four isomers each with the molecular formula C6H12. Draw

structures of the four isomers and name each of them.

1970 (Required)

What is meant by the lattice energy of an ionic compound? What quantities need to be determined and how are

they used to calculate the lattice energy of an ionic compound.

FRQ 2011 (Required)

FRQ(Required)

FRQ 1992 D (required)

NO2 NO2- NO2

+

Nitrogen is the central atom in each of the species given above.

(a) Draw the Lewis electron-dot structure for each of the three species.

(b) List the species in order of increasing bond angle. Justify your answer.

(c) Select one of the species and give the hybridization of the nitrogen atom in it.

(d) Identify the only one of the species that dimerizes and explain what causes it to do so.

1996 D (Required)

Explain each of the following observations in terms of the electronic structure and/or bonding of the compounds

involved.

(a) At ordinary conditions, HF (normal boiling point = 20ºC) is a liquid, whereas HCl (normal boiling point = -114ºC)

is a gas.

(b) Molecules of AsF3 are polar, whereas molecules of AsF5 are nonpolar.

(c) The N-O bonds in the NO2- ion are equal in length, whereas they are unequal in HNO2.

(d) For sulfur, the fluorides SF2, SF4, and SF6 are known to exist, whereas for oxygen only OF2 is known to exist.

AP Chemistry: Lab Molecular Geometry

The physical properties of molecular substances are related to the polarity of the molecules. This polarity is

determined in part by the arrangement of atoms around the central atom in a molecule. This molecular geometry

can be predicted by the VSEPR theory.

According to the Valance Shell Electron Pair Repulsion theory, the electron pairs around the central atom will take

positions as far apart as possible. This gives rise to the five possible electron arrangements shown below. In this

lab, you will draw Lewis Structures for several molecules and build the best model that fits according the VSEPR.

LINEAR TRIGONAL PLANAR TETRAHEDRAL

TRIGONAL BIPYRAMIDAL OCTAHEDRAL

A> OCTET

Consider the following molecules that obey the octet rule:

CO2 SO3 SO2 CH4 NF3 OF2 HCl

1. Draw the Lewis Structure for each molecule.

2. Build a model from the kits provided with the electrons as far apart as possible (linear, trigonal planar, or

tetrahedral).

3. Have you instructor check the 7 models that you built.

4. Draw a 3-D representation of your model.

5. Indicate the shape and bond angles. Choose a name for the shape from the following: linear, angular, trigonal

planar, tetrahedral, pyramidal.

B> EXPANDED OCTET

Consider the following molecules with an expanded octet:

PCl5 SF4 ClF3 XeF2 SF6 BrF5 XeF4

1. Draw the Lewis Structure for each molecule.

2. Build as many different models as possible using an electron arrangement of either trigonal bipyramidal or

octahedral.

3. Calculate the total repulsive units (RU) for each model by considering the electron pair interactions (90 or

less) as:

lone pair to lone pair 3.5 RU

lone pair to bond pair 2.0 RU

bond pair to bond pair 1.0 RU

4. Draw a 3-D representation of each model indicating the total repulsive units for that model.

5. Choose the arrangement that gives the lowest total repulsive units and name the shape of this model.