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Pre-AP Chemistry: Chapters 7 & 8

Ionic, Metallic and Covalent Bonding

7.1 Valence Electrons

Valence electrons- electrons in the highest occupied energy level of an atom. Valence electrons are the only

electrons involved in the formation of chemical bonds.

Electron dot structures for atoms:

• each dot represents a valence electron

Examples: I N Xe

Al O Na Si

• One of the major “driving forces” in nature is the tendency to go to lower energy. Atoms lose, gain or

share electrons to become lower in energy and thus more stable.

Metals • lose electrons easily to become positively charged cations.

• lose their valence electrons to achieve a noble gas electron configuration.

Transition Metals • lose their highest energy level s and p electrons, d electrons still remain

• pseudo-noble gas electron configuration: not quite that of a noble gas but is still

stable.

o Ex: zinc loses its two electrons in 4s but keeps the ten electrons in 3d.

Nonmetals • tend to gain electrons to become stable

• form negatively charged anions.

• achieve a noble gas electron configuration.

7.2 Ionic Bonding

• Ionic compounds are created by the attraction of oppositely charged ions (cations and anions)

• When the electronegativity difference between two elements is large, the elements are likely to form a

compound by ionic bonding (transfer of electrons).

• The farther apart across the periodic table two Group A elements are, the more ionic their bonding will be.

We can use Lewis dot formulas to represent the formation of ionic compounds.

Na + Cl → NaCl 3Mg + N2 →Mg3N2

• Properties of Ionic Compounds:

1. They are usually crystalline solids with high melting points (>400oC)

2. Molten, or melted, ionic compounds and aqueous solutions conduct electricity well because they

contain mobile ions.

3. Ionic compound are brittle and when they are struck, like-charged ions collide, causing repulsion,

and the crystal shatters.

7.3 Bonding in Metals

• Metallic solids consist of positively charged metal cation nulcei in a “sea” of loosely held valence

electrons.

• This arrangement allows metals to have their unique properties.

o Ductility- metals can be pulled into a wire

o Malleability- metals can be hammered into a thin sheet because the valence electrons act as

“grease”, allowing the cations to slide past each other without colliding with each other and

shattering.

o Conductivity- metals can conduct heat and electricity easily. Electricity is a flow of electrons.

As electricity (electrons) enters one end of a piece of metal, an equal number of electrons exit the

other end.

Cubic Unit Cells

• Many solids, including metals and ionic solids, have one of three types of cubic unit cells. These three

unit cells are:

Name Net Spheres Coordination number Picture

Simple Cubic 1

(8 –1/8 spheres) 6

Body-Centered Cubic 2

(1 complete and 8-1/8 spheres) 8

Face-Centered Cubic (NaCl and other alkali halides are

face-centered cubic.)

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(6-1/2 spheres and 8 -1/8 spheres). 12

Alloys • Alloys are formed when metals are mixed with other metals. They are solutions of solids in solids.

o Substitutional alloy- atoms of one metal are substituted for atoms of a similar-sized metal in a

metallic crystal.

▪ Ex. brass, sterling silver, pewter

o Interstitial alloy- smaller metal atoms fit into holes in the crystal structure of a metal with larger

atoms

▪ Ex. steel (carbon in iron)

o Amalgam- alloy which contains mercury

8.1 Molecular Compounds

Hydrogen and nonmetals of Groups 4,5,6 & 7 often become stable and gain noble gas electron configurations

by sharing electrons to form covalent bonds. Atoms will usually share electrons to follow the octet rule (eight

electrons, like most noble gases) or the duet rule (2 electrons, like helium).

When atoms share one pair of electrons to form a covalent bond, it is called a single covalent bond. The

electrons shared between the atoms are a “shared pair”. A dash can be used instead of two dots to represent the

shared pair. Any other electrons on the atoms are “unshared pairs” or “lone pairs”.

H2 Cl2 HCl

Atoms must sometimes share more than one pair of electrons to become stable.

• If two pairs of electrons are shared, it is called a double bond.

• If three pairs of electrons are shared, it is a triple bond.

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O2 N2

8.2 The Nature of Covalent Bonding

Rules for Writing Lewis Structures (electron dot structures):

(Use pencil!)

1. Add up the valence electrons from all the atoms. Don’t worry about keeping track of which electrons come

from which atoms. If you are working with an ion, you must add or subtract electrons to equal the charge.

2. Use a pair of electrons to form a bond between each pair of bound atoms.

3. Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for everything else.

4. If necessary, change bonds to double or triple.

5. Remember, we cannot create or destroy electrons!

H2O # e- _____ NH3 # e- _____ NH4+ # e- _____

CO2 # e- _____ CCl4 # e- _____ CN- # e- _____

SO42- # e- _____ CO3

2- # e- _____

Coordinate Covalent Bond- Bond in which both shared electrons came from the same atom. This bond is not

really any different than any other single bond. Ex. Ammonium ion

Exceptions to the Octet Rule:

Less than an octet:

• They include beryllium or boron. These electron deficient compounds

are very reactive.

More than an octet:

• Elements in the third period and below can exceed the octet rule

• Extra electrons go in empty d orbitals.

• Elements in the second period cannot exceed the octet rule because there

are no 2d orbital for the extra electrons to go into.

• If it is necessary to exceed the octet rule, place the extra electrons on the central atom.

BF3

BeCl2

PCl5 # e- _____ SF6 # e- _____ I3- # e- _____

More practice: NF3 # e- _____ OF2

# e- _____ KrF4 # e- _____

BeH2 # e- _____ SO3

2- # e- _____ NO3- # e- _____ H2O2

# e- _____

Resonance occurs when more than one valid Lewis structure can be written for a molecule. The actual structure

is an average of all of the resonance structures.

Ex. NO3-

In nitrate, the experimental bond length is in-between that of a single bond and a double bond.

It acts like a 1 1/3 bond.

Ex. Benzene, C6H6

8.3 Bonding Theories

VSEPR (Valence Shell Electron Pair Repulsion)

• Lewis structures can be used to determine the shapes of molecules. Their shapes will tell us a lot about

their chemical behavior.

• The VSEPR theory tells us that valence electrons on the central atom repel each other.

• Valence electrons are arranged so there are minimum repulsions.

• When we are using VSEPR to determine molecular shape, we are really looking for regions of electron

density. Double and triple bonds count the same as single bonds in determining molecular shape.

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Practice Determining Molecular Shape: (remember to count electrons)

H2S CCl4 NH4+ BF3

NO2- PF6

- SbCl5 SO2

Some other VSPER Shapes: (you do not need to memorize these shapes)

Effective

Electron

Pairs

Bonding

Electron

Pairs

Nonbonding

Electron Pairs

VSEPR

Formula

Hybrid-

ization

Approx.

Bond Angles

3-D

Structure Examples

5 4 1 AX4E dsp3 <90o, <120o see-saw SF4, TeCl4

5 3 2 AX3E2 dsp3 <90o T-shaped ClF3

5 2 3 AX2E3 dsp3 180o linear ICl2-, I3

-, XeF2

6 5 1 AX5E d2sp3 <90o square pyramidal BrF5

6 4 2 AX4E2 d2sp3 90o square planar ICl4-, XeF4

Hybridization

When drawing Lewis structures to explain bonding, we have been using the Localized Electron Model of

bonding. This assumes that the electrons stay with (or close to) the atoms from which they originated and that

bonds are formed by the overlap of atomic orbitals. This model needs to be developed a little further to

explain experimental data.

Why hybridization? In methane, CH4, carbon has one 2s and three 2p orbitals available for bonding. Hydrogen

has one 1s orbital available. We could imagine that one of the bonds in methane would be

formed from the overlap of a hydrogen 1s orbital and a carbon 2s orbital and the other three

bonds would be formed from the overlap of a hydrogen 1s orbital and a carbon 2p orbital.

This would make 2 different kinds of bonds. Experimental data shows that all four bonds are

identical. Hybridization is an addition to the localized electron model that explains this

sp3 hybrid orbitals • Lets look at the hybridization of CH4.

• Carbon needs 4 identical orbitals to form 4 identical bonds.

• If we combine 1 s orbital and 3 p orbitals, the new orbitals have

properties of the s and of the p orbitals.

• Since they were made from one s and three p orbitals, we can

call them sp3 orbitals.

• We put four atomic orbitals in and got four hybrid orbitals out.

• The hybrid sp3 orbitals each overlap with the 1s orbitals from each hydrogen atom and make four identical

bonds having 109.5o bond angles.

• Atoms with four effective electron pairs have sp3 hybridization, even if they have unshared pairs as in H2O

and NH3. The electron pairs have a tetrahedral arrangement.

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sp2 hybrid orbitals • We can look at the hybridization in the bonding of ethene (C2H4).

• Each carbon atom has three effective pairs and needs three equal orbitals.

• To get three hybrid orbitals, hybridize one s and two p orbitals.

• We get three identical orbitals made from one s and two p orbitals.

• The three new orbitals are called sp2.

• This type of hybridization results in trigonal planar electron pair

arrangements and 120o bond angles.

• The overlap of the 1s orbitals from each hydrogen with the sp2 hybrid orbitals of carbon results in sigma

bonds.

o In a sigma bond, the orbitals overlap head-on and the electrons are in the region between the two

nuclei.

o All single bonds are sigma bonds.

o Each double bond and each triple bond contains one sigma bond.

• What happened to the p orbital that wasn’t hybridized?

o The remaining p orbitals from each carbon atom overlap above and below

the plane of the nuclei resulting in a sideways overlap that is called a pi bond.

o A pi bond can form only if there is also a sigma bond between the same two

atoms.

o Because the electrons in a pi bond are not directly between the nuclei, a pi bond is weaker than a

sigma bond.

o Double bonds are strong and shorter than single bonds because it has one sigma and one pi bond.

sp hybrid orbitals We can look at the hybridization and bonding in ethyne (acetylene), C2H2.

• Each carbon in ethyne has two effective electron pairs.

• To get three hybrid orbitals, hybridize one s and one p orbitals.

• The s orbital from each hydrogen atom overlaps with one of

the sp orbitals to form a sigma bond.

• The other sp orbital overlaps with one of the sp orbital on the

other carbon to forma sigma bond between the two C atoms.

• The remaining two p orbitals from each carbon atom overlap outside

the plane of the nuclei to form two pi bonds.

• A triple bond is one sigma and two pi bonds and is thus stronger and

shorter than either a single or a double bond.

• Linear molecules and 180o bond angles result from sp hybridization.

Other Hybridized orbitals • We learned that certain elements can exceed the octet rule by utilizing unfilled d orbitals. These d orbitals

may also be involved in hybridization and form dsp3 (or sp3d) orbitals. This form of hybridization results

in trigonal bipyramidal shapes and 90o and 120o bond angles.

• With SF6 we have six effective electron

pairs. We must hybridize one s, three p,

and two d orbitals to get six d2sp3 or sp3d2

orbitals. This results in the octahedral

shape with 90o bond angles.

Practice:

For each of the following, determine:

a) hybridization of each atom

b) bond angles

c) molecular shape

d) # of sigma bonds and # of pi bonds

SbCl5

A)

B)

C)

D)

NO3-

A)

B)

C)

D)

CHCl3

A)

B)

C)

D)

8.4 Polar Bonds and Molecules

• Covalent bonds within a molecule can be polar (electrons are shared unequally) or nonpolar (electrons are

shared equally).

• We can predict polar bonds by looking up the electronegativity values of each element in the bond and

subtracting the smaller value from the larger value to determine the electronegativity difference.

o If the electronegativity difference (EN) is zero, the bond is nonpolar covalent.

o If the EN is between 0 and 2.0, we can predict that the bond is polar covalent.

o If the EN is 2.0 or greater, the bond is usually considered to be ionic. There are no strict dividing

lines between covalent and ionic bonds.

• Physical properties are used to help determine whether something is covalent or ionic.

• Some covalent substances are so polar that they ionize partially or completely in water.

• Polarity of a bond or molecule can be represented by arrows or lowercase delta () symbols to show a

partial charge.

This shows that the hydrogen end of the molecule is more positive (less electronegative) and the fluorine end is

more negative (more electronegative). A charge difference such as this is called a dipole.

An entire molecule is polar if it has polar bonds that do not cancel. This happens in nonsymmetrical molecules.

• Molecules with all nonpolar bonds such as H2, O2, and Cl2, are always nonpolar. This means that there are

no positive and negative ends on the molecule.

• Heteronuclear diatomic molecules such as HCl, BrCl, or HF are always polar because the polar bonds

cannot cancel each other out.

Practice: Polar or Nonpolar?

NH3 SO2 H2O

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BF3 CH4

Bond Energies

• It requires energy to break chemical bonds.

• When bonds form, energy is released.

• The energy required to break a chemical bond is called the bond dissociation energy.

• Double and triple bonds have higher bond energy and shorter bond length than single bonds.

• The heat of reaction, H, can be determined using bond energies (on page 448) and the following formula:

H = Bond energy of bonds broken - Bond energy of bonds formed

Ex: H2 + Cl2 → 2HCl

Intermolecular Forces

Two terms describe the forces that hold substances together.

• Intramolecular bonding- the forces inside a molecule or formula unit due to shared electrons or charge

interactions.

• Intermolecular bonding- interactions between particles (atoms, molecules or ions)

o Changes in state are due to changes in intermolecular bonding, not intramolecular bonding.

There are three types of intermolecular bonding. These occur in molecular compounds

only!

1. Dipole-dipole attraction- attraction of polar molecules for each other. (negative-

positive)

• Approximately 1% as strong as covalent or ionic bonds.

• Molecules orient themselves to minimize repulsion and maximize attractions

2. Hydrogen bonding-

• Unusually strong dipole-dipole attractions involving hydrogen atoms which are

covalently bonded to a very electronegative element and a very electronegative

atom (F,O,N only) with unshared electrons.`

• Two reasons for strength of hydrogen bonds:

1. small size of H atom allows closeness

2. large polarity value across the molecule

• Substances with strong H-bonding have high boiling points compared to similar

substances. Ex. H2O, NH3, HF

3. London dispersion forces (LDFs)

• Relatively weak forces (usually) that exist between noble gas atoms and nonpolar molecules.

• LDFs also exist in compounds that have dipole-dipole and/or hydrogen bonding and may be the most

important force in large molecules of these types.

• LDFs occur because of momentary electron imbalance (temporary dipole) which can induce the same

to occur in adjacent molecules.

• This force is often very weak, thus the low freezing point of noble gases.

o The freezing point of noble gases increases going down the group because heavier atoms have

more electrons and an increased chance of temporary dipoles. o They also have a lower velocity and have more opportunity for attractions.

• This causes London dispersion forces to increase going down a group on the periodic table.

4. Covalent network solid

• Group 4 substances such as diamond, Si, SiC, and Ge form extensive covalent bonds and result in giant

molecules. They have an atom at each lattice point and are held together by covalent bonds. These

substances have the strongest attractions.

General trends in strength of attraction:

LDF<dipole-dipole<hydrogen-bonding<metallic bonding<ionic bonding<covalent network bonding

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1. __________________

2. __________________

3. __________________

4. __________________

5. __________________

6. __________________

7. __________________

8. __________________

9. __________________

10. __________________

11. __________________

12. __________________

13. __________________

14. __________________

Pre-AP Chemistry: Bonding Worksheet #1

Complete the following paragraph using terms from chapter 7.

Atoms are held together in __1__ by chemical bonds. These bonds result from the

sharing or transfer of __2__ between atoms. Atoms share or transfer electrons to

attain __3__ electron configuration. Ionic compounds are produced when atoms that

have lost electrons to form __4__ and those that have gained electrons to form __5__

are attracted to one another. Nearly all ionic compounds form repeating structures

called __6__ and are __7__ at room temperature. Ionic solids have __8__ melting

points and are only able to conduct electricity when __9__.

Metals consist of metal cations packed together and surrounded by a sea of their

__10__. This arrangement accounts for many of the properties of metals, like their

ability to be hammered into sheets, or __11__, their ability to be pulled into wires, or

__12__ or their tendency to easily conduct __13__ and __14__.

1. Determine the hybridization of each carbon in the following:

A. B. C.

2. Determine the hybridization of the central atom in each of the following:

A. B. C.

3.

How many sigma () bonds are in the molecule? ____

How many pi () bonds are in the molecule? ____

4. One of the first drugs to be approved for use in treatment

of AIDS is azidothymidine (AZT).

a. How many carbon atoms are sp3 hybridized? ____

b. How many carbon atoms are sp2 hybridized? ____

c. Which atom is sp hybridized? ____

d. How many sigma bonds are in the molecule? ____

e. How many pi bonds are in the molecule? ____

f. What is the N – N – N bond angle in the lowest group? ___

g. What is the H – O – C bond angle in the side group attached to the 5-

membered ring? ____

h. What is the hybridization of the oxygen atom in the –CH2OH group? ____

5. Draw Lewis structures for CO2, H2, SO3, and SO32- and predict the shape of each species.

6. NF3 and PF5 are stable molecules. Write the electron dot formulas for these molecules. On the basis of

structural and bonding considerations, account for the fact that NF3 and PF5 are stable molecules but NF5

does not exist.

7. Using principles of chemical bonding and/or intermolecular forces, explain each of the following.

(a) Xenon has a higher boiling point than neon has.

(b) Solid copper is an excellent conductor of electricity, but solid copper chloride is not.

(c) SiO2 melts at a very high temperature, while CO2 is a gas at room temperature, even though Si and C are

in the same chemical family.

(d) Molecules of NF3 are polar, but those of BF3 are not.

8. Draw the following molecules. Write their shape, bond angles, hybridization of the central atom and

polarity in the spaces provided.

CS2

Shape:__________

Bond angle ______

Hybridization ____

Polar or nonpolar?

NF3

Shape:__________

Bond angle ______

Hybridization ____

Polar or nonpolar?

H2Se

Shape:__________

Bond angle ______

Hybridization ____

Polar or nonpolar?

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PAP Bonding Worksheet #2

Multiple Choice: (2 pts each)

1. How many valence electrons are there in the phosphate ion?

A. 26 B. 32 C. 31 D. something else

2. BCl3 is a trigonal planar molecule. This means that the Cl-B-Cl bond angle is:

A. 90o B. 109.5o C. 120o D. 180o

3. Which is the best statement concerning the bond lengths in the carbonate ion, CO32-?

A. Two bonds are longer single bonds, the third is a shorter single bond.

B. Two bonds are shorter double bonds, the third is a longer single bond.

C. All three bonds are equal length, between that of a double and a single bond.

D. All three bonds are of equal length, that of a single bond.

4. The molecular geometry of BeCl2 is:

A.linear B.bent C.trigonal planar D.tetrahedral

5. What is the hybridization of the nitrogen atom in NH3?

A. sp B. sp2 C. sp3 D. dsp3

6. The fact that BCl3 is trigonal planar implies that the B-Cl bonds involve ____ hybrid orbitals.

A. sp B. sp2 C. sp3 D. no

7. In a double bond, there are ___ sigma and ___ pi bonds, respectively.

A. 1,1 B. 3,0 C. 2,1 D. 1,2 E. 0,3

8. Which of the following elements is most likely to form compounds that do not obey the octet rule?

A. Na B. P C. O D. C

9. Which of the following molecules possesses two pi and one sigma bond?

A. nitrogen B. oxygen C. chlorine D. iodine

10. An example of a covalent network solid is:

A. aluminum B. diamond C. iodine D. sodium nitrate

11. How many sigma bonds are present in a molecule of methyl ether, C2H6O?

A. 6 B. 8 C. 10 D. 7

12. The hybridization of the carbon atoms in ethylene (C2H4) is:

A. sp B. sp2 C. s2p D. sp3 E. sp4

13. What is the shape of the water molecule?

A. linear B. tetrahedral C. trigonal planar D. bent

14. Which of these forces of intermolecular attraction is generally the strongest?

A. dipole-dipole attraction B. London dispersion forces C. hydrogen bonds

15. In which compound is the bonding primarily ionic?

A. KCl B. OCl2 C. BCl3 D. SiCl4

16. Which pair of elements is most likely to react to form a covalently bonded species?

A. P and O B. Ca and O C. K and S D. Zn and Cl

17. Hydrogen bonding is shown by molecules containing hydrogen attached to:

A. highly electronegative atoms C. any atoms

B. highly electropositive atoms D. only oxygen

18. Which of the following molecules or ions is polar?

A. NH3 B. BF3 C. CO32- D. CH4

19. How many pi bonds are present in a molecule of C2H6?

A.6 B. 1 C. 0 D. 7

20. Which of these molecules violates the octet rule?

A. H2SO4 B. PCl3 C. PCl5 D. ClO2

21. The melting temperature of potassium chloride is relatively:

A. high C. low B. variable D. Potassium chloride does not melt.

22. How many electrons does boron have to give up in order to achieve a noble-gas electron configuration?

A. 1 B. 2 C. 3 D. 4 E. 5

23. What is the basis of a metallic bond?

A. the attraction of metal ions for mobile electrons

B. the attraction between neutral metal atoms

C. the neutralization of protons by electrons

D. the attraction of oppositely-charged ions

E. the sharing of two valence electrons between

two atoms

24. The octet rule states that, in chemical reactions, atoms change their numbers of electrons in order to:

A. obtain the electron configuration of a noble gas.

B. become neutral rather than charged.

C. have eight electrons in their principle energy level.

D. give off electrons.

25. Intermolecular attractions between nonpolar molecules are called:

A. hydrogen bonds B. metallic bonds C. dipole-dipole attractions D. London dispersion forces

26. Write electron configurations for the following: (3 pts)

A. Fe3+

B. Cl-

C. Cu+1

27. Write electron dot structures for the following: (3 pts)

S

S2-

Ge

28. Describe how metals conduct an electrical current. (3 pts)

29. Draw the electron dot structures for the following: (4 pts)

A. SF4 B. XeCl4

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30. Draw the electron dot structure and predict the shape and type of molecule (polar or nonpolar) for each of

the following: (16 pts)

A. NF3

Shape:

Type:

B. SiH4

Shape:

Type:

C. PCl3

Shape:

Type:

D. SCl2

Shape:

Type:

31. Using a table of bond energies, calculate the total bond energy of ethanol, CH3CH2OH.

Show your electron dot structure, also. (5 pts)

32. Draw all resonance structures for NO3-. (3 pts)

AP Essay Bonding Questions for Pre-AP Chemistry

1974: The boiling points of the following compounds increase in the order in which they are listed below:

CH4< H2S<NH3

Discuss the theoretical considerations involved and use them to account for this order.

1990: Use simple structure and bonding models to account for each of the following.

A) the bond length between the two carbon atoms is shorter in C2H4 than in C2H6.

B) Th H-N-H bond angle is 107.5o, in NH3

C) The bond lengths in SO3 are all identical and are shorter than a sulfur-oxygen single bond.

1992: Nitrogen is the central atom in each of the given species: NO2 NO2- NO2

+

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.