Chapter 9

Models of Chemical Bonding

A general comparison of metals and nonmetals.Figure 9.1

Astatine may be a metalloid but it is radioactive.

BASIC CHEMICAL BONDING

Atoms form bonds in order to complete their outer shells of electrons

Main group elements try for 8. (Transition metals are more complex.)

Atoms can give up or gain electrons to form ions which then form ionic bonds, or they can share electrons to form covalent bonds

(Or in metals, share electrons throughout all atoms)

BASIC CHEMICAL BONDING

Atoms/ions seek lower energy states:

- ions in a bond have a lower energy than the separated ions

Data indicate lowest energy level is for atoms to share electrons at a certain distance = bond length = covalent radii

Types of Chemical Bonding

1. Metal with nonmetal:

electron transfer and ionic bonding

2. Nonmetal with nonmetal:

electron sharing and covalent bonding

3. Metal with metal:

electron pooling and metallic bonding

4. Cation with anion: ionic bonding (ammonium ion and nitrate ion)

Figure 9.2 The three models of chemical bonding.

“NEW” CONCEPT OF CHEMICAL BONDING: LEWIS THEORY

Valence electrons have fundamental role in chemical bonding

If e-(s) "transferred" ionic bond resultsIf e-(s) "shared" covalent bond resultsIf e-(s) “shared as a pool” metallic bond

resultsTrying to achieve noble gas configurationLewis symbol: Atomic symbol with valence

electrons for neutral atom, correct number of electrons for ion

Figure 2.14 from 4th ed.

The relationship between ions formed and the nearest noble gas. (from Chp 2)

Lewis Electron-Dot Symbols

For main group elements -

Example:

Nitrogen, N, is in Group 15 and therefore has 5 valence electrons.

N:.

..

:

N .. ..N :.

. :N ...

The group number digit gives the number of valence electrons.

Place one dot per valence electron on each of the four sides of the element symbol.

Pair the dots (electrons) until all of the valence electrons are used.

Figure 9.4

Lewis electron-dot symbols for elements in Periods 2 and 3.

How will these Lewis symbols look when these atoms form ions?

Lewis Symbols for Respresentative Elements

Lewis Structures For Ions:

Put individual ion structures together as group or use coefficient:

.. .. ..

CaF2 = [:F:]-[Ca]2+[:F:]- or [Ca]2+2[:F:]-

.. .. ..

Note that Ca: [Ca]2+ + 2 e-

each F + e- --> [F]-

SAMPLE PROBLEM 9.1 Depicting Ion Formation

SOLUTION:

PROBLEM: Use Lewis symbols to depict the formation of Na+ and O2- ions from the atoms, and determine the formula of the compound. Draw in the most correct way when I show you below. Ions should have brackets.

:Na

Na+ O

.

:

..

.

2 [Na]++[ O ]2-

::: :

Electron configurations

Li 1s22s1

Orbital diagrams

Lewis electron-dot symbols: focus on this way

+ F 1s22s22p5 [Li]+ 1s2 + [F]- 1s22s22p6

Three ways to represent the formation of Li+ and F- through electron transfer.

Figure 9.5

Li

1s 2s 2p

F

1s 2s 2p

+

[Li[+

1s 2s 2p

[F]-

1s 2s 2p+

.+ F: ::Li . [Li]+ + F -:

:

::

[They should both have brackets.]

Lewis Structures For Ions:

Practice drawing Lewis structure for ions and then use them to make ionic compounds of the two elements listed. Be prepared to show your work on the document camera for the class to see. Draw large structures in ink and use brackets where required!

Mg and O, Al and Cl, Na and S

LATTICE ENERGY

When ionic bonding occurs, energy is released – exothermic – because the positive and negative ions achieve their lowest energy state when surrounded by the opposite charge.

However, IE is endothermic and usually larger than EA, so the transfer of electrons is endothermic.

In chp 6 we will learn about Hess’ Law, which is how we actually determine the value of lattice energy.

LATTICE ENERGY

For now: the energy released when forming an ionic compound arranged in a crystal lattice from the elements is called enthalpy of formation. If we know the value of IE and EA, then we can determine the energy released when the ions are moved into the crystal lattice.

Lattice energy is affected by:

Ion size: the larger the radius, the lower the lattice energy

Ionic charge(s): the greater the charge, the higher the lattice energy

Figure 9.8

Electrostatic forces and the reason ionic compounds crack.

Figure 9.9 Electrical conductance and ion mobility.

Solid ionic compound

Molten ionic compound

Ionic compound dissolved in water

Table 9.1 Melting and Boiling Points of Some Ionic Compounds

Compound mp (0C) bp (0C)

CsBr

661

1300

NaI

MgCl2

KBr

CaCl2

NaCl

LiF

KF

MgO

636

714

734

782

801

845

858

2852

1304

1412

1435

>1600

1413

1676

1505

3600

Covalent Bonding

A shared pair of e-s between two atoms

Other e-s in valence shell become Lone Pairs

Can have single, double or triple bonds (Bond Order)

Covalent Bonding

Energy is released when atoms join to form bonds

Energy must be absorbed to break bonds - called bond dissociation energy

Measured for gaseous species in kJ/molIncreases with multiple bondingIncreases with decreased bond length

See Bond Energy and Bond Length tables

Formation of a covalent bond between two H atoms. (from Chp 2)

Figure 2.13

Covalent bonds form when elements share electrons, which usually occurs between nonmetals.

BOND ENERGIES & Hrxn:

Hrxn approx same as difference between bond breaking energy and bond formation energy

REMEMBER THIS IS AN APPROXIMATION BECAUSE WE ARE USING AVERAGES!

Estimate Hrxn for N2 + 3 H2 2 NH3 First find kinds of bonds there areLook up average bond energies in table

Draw Lewis structure, look up bond energies, count number of bonds, Hrxn ~ BEprod – BEreact

SAMPLE PROBLEM 9.2 Comparing Bond Length and Bond Strength

SOLUTION:

PROBLEM: Using the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of decreasing bond length and bond strength: (NO CALCS, JUST THINK!)

(a) S - F, S - Br, S - Cl (b) C = O, C - O, C O

(a) Atomic size increases going down a group.

Bond length: S - Br > S - Cl > S - F

Bond strength: S - F > S - Cl > S - Br

(b) Using bond orders we get

Bond length: C - O > C = O > C O

Bond strength: C O > C = O > C - O

Section 9.4

We can’t really calculate bond energies until we can draw Lewis electron dot structures, which is in chp 10, and learn about enthalpy of reaction, which is in chp 6. We will return to this after chp 6.

 

Figure 9.13Strong covalent bonding forces within molecules

Weak intermolecular forces between molecules

Strong forces within molecules and weak forces between them.

Figure 9.14

Covalent bonds of network covalent solids.

CONCEPT OF ELECTRONEGATIVITY: from Linus Pauling

I define it as greediness for electrons in a covalent bond

Look at table – a scale with no unitsMetals are low and nonmetals are highDifference in their electronegativities:

If EN is < .5, bond is covalent (0 to 0.4)If > 2.0, it's ionic, and in between is polar covalent (0.5 to 1.9)

Determine something called "% ionic character" using the table in text

Figure 9.19

The Pauling electronegativity (EN) scale.

SAMPLE PROBLEM 9.4 Determining Bond Polarity from EN Values

SOLUTION:

PROBLEM: (a) Use a polar arrow to indicate the polarity of each bond: N-H, F-N, I-Cl.

(b) Rank the following bonds in order of increasing polarity: H-N, H-O, H-C.

(a) The EN of N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0

N - H F - N I - Cl

(b) The order of increasing EN is C < N < O; all have an EN larger than that of H.

H-C < H-N < H-O

Figure 9.21

EN

3.0

2.0

0.0

A. Boundary ranges for classifying ionic character of chemical bonds. B. Gradation in ionic character.

> 2.0 Ionic

0.5 to 1.9 Polar covalent

0 to 0.4 Nonpolar covalent

Use my simplified version!

Percent ionic character of electronegativity difference (EN).

Figure 9.21 continued

Figure 9.22

Properties of the Period 3 chlorides.

Metallic Bonding (not in text but YOU have to know it)Sea of electrons shared throughout the metal

piece, however large it isExplains both electrical and thermal conductivity,

since electrons can move throughout the pieceA mixture of metals, an alloy, also share electrons

throughoutMetallic bond strength varies, as you can see from

range of melting points, but most metals are solids (only Hg is liquid at room temperature)

Table 9.7 Melting and Boiling Points of Some Metals

Element mp(0C) bp(0C)

Lithium (Li) 180 1347

Tin (Sn) 232 2623

Aluminum (Al) 660 2467

Barium (Ba) 727 1850

Silver (Ag) 961 2155

Copper (Cu) 1083 2570

Uranium (U) 1130 3930

Figure 9.26

Melting points of the Group 1A(1) and Group 2A(2) elements.

Figure 9.27

metal is deformed

The reason metals deform.