Chemical Bonding

Chapter 12

GCC

CHM 130

12.1 Chemical Bonding Atoms want to be like noble gases (stable and happy) Goal = 8 outer valence electrons =Octet Rule (except H, He)

Metals lose electrons and become (+) cationsNonmetals gain electrons and become (-) anions

A metal / nonmetal compound is IONIC w/ IONIC bond (the + and – attract each other)

Who makes a good partner for Ca? Br? Li? Ionic Examples: KCl, CaBr2

Nonmetals can also share electrons with each other A nonmetal / nonmetal compounds is COVALENT

w/COVALENT bonds. Example: diatomic elements (H2, N2, F2…) Covalent Examples: H2O, CO2

Why is the formula CaBr2? Why one Ca per two Br? Any ideas?

Because Ca ion is +2 and Br ion is -1. So need two Br -1 ions to balance with one Ca +2 ion. +2-1-1=0 (Br-1Ca2+Br-1)

The answer is NOT because Br is diatomic

Note, Br is diatomic BY ITSELF (Br2) but when in a compound the Br ‘s break apart to bond with other atoms! The diatomic elements are NOT diatomic anymore once bonded with others.

Which compounds are Ionic?

• KBr

• SO3

• HCl

• Br2

• CO2

• MgCl2

Answer: The ones with a metal and a nonmetal.

KBr and MgCl2

12.2 Ionic BondingElectrons are completely transferred from metal to

nonmetal.

Draw electron dot structures for Mg and S atoms then Mg2+ and S2- ions in MgS. How many protons and electrons in Mg2+

and S2- ions?

Notice that Mg2+ and S2- are “like” noble gases. They are isoelectronic with Ne and Ar, and that is what makes them happy and stable. While the number of electrons changed, the number of protons did NOT. # Protons never change in chemical reactions.

Ionic Radius

• Cations have lost electrons, so there are more protons, so pull the electron orbits in closer to nucleus (smaller than the atom)

• Anions have gained electrons, so there are more electrons, they repel and push orbits farther from nucleus (larger than the atom)

True or False regarding an ionic bondbetween aluminum and iodine?

1. The aluminum atom loses electrons, and the

iodine atom gains electrons.

2. The aluminum atom is larger in radius than

the aluminum ion.

3. The iodine atom is smaller in radius than the

iodine ion.

4. The aluminum and iodine ions form a bond

by attraction.

True

True

True

True

…is when nonmetals share electrons.Single = 2, double = 4, triple = 6 e- shared

Note the bond length is less than r1+r2 due to orbital overlap

12.3 Covalent Bonding

Bond Energy

•=Energy required to break a bond.• •Breaking bonds always requires E.

• E is a reactant, it is absorbed.•Forming bonds always releases E.

• E is a product, it is produced.

H (g) + Cl (g) HCl (g) + heat

HCl (g) + heat H (g) + Cl (g)

Endothermic

Exothermic

True or False regarding H2S?

1.Electrons are shared in H2S.

2.The bond between H and S is

ionic.

3.The H-S bond length is less than

the sum of the two atomic radii.

4.Breaking the H-S bond releases

energy.

True

False

True

False

12.4 Draw electron dot structures for H2 and HCl

12.4 Electron Dot Structures

1. Add up the total number of valence electrons.

2. Surround the central atom with the other atoms and draw single bonds to them.

3. All atoms want octet 8e- except H wants 2 e-.

4. Final Check: Make SURE you use the total # of e-, no more or less.

bonding e-= shared e-

lone pairs = unshared e-

If single bonds don’t work, try double, then triple.

H2OTotal = 8e-

O::. .. .

H

H

• HCN

• CHCl3

• CO2

• NH3

Examples to put on

board:

The central atom is in bold.

12.5 Electron Dots of Polyatomic Ions

Add electrons for anions and subtract electrons for cations. Put brackets around the ion and charge in the right corner.

NH4+

+1 charge means one less e-

Total = 5 + 4(1) – 1 = 8 e-

Examples for the board:• BrO3

-

• SO42-

• CN-

The central atom is in bold.

12.10 Valence Shell Electron Pair Repulsion Theory

VSEPR

• Electron pairs (bonded and lone pairs)

repel each other and move as far away from

each other as possible.

• Molecular Shape or Geometry – the 3 D

arrangement of the atoms.

.Print out Shape Table from the web page

Linear – AB and AB2

Examples: H2, HCl, CO2

A = Central Atom B = Outer Atom

E = Lone Pair on central atom

Bond Angle is 180

Trigonal Planar – AB3

Example: Formaldehyde, CH2O

C

O

H H

120°

Tetrahedral – AB4

Example: CH4, CF4, CH2F2

Bent – AB2E

Example: SO2

SO O

<120°

Trigonal Pyramidal – AB3E

Example: ammonia, NH3

Bent – AB2E2

Example: water, H2O

Summary

• Given any molecule or polyatomic ion you should be able to– Draw the Electron dot structure– Determine the shape and bond angles

Get used to the Table of Shapes online – you will get it on the exam over this chapter!!!

Practice: PH3 and ozone O3

12.6&7 Polar and Nonpolar Covalent Bonds

• A covalent bond where electrons are shared equally is a nonpolar bond. (no poles, no magnet)

• A covalent bond where electrons are shared unequally is a polar bond. (has poles like a magnet)

Symbols used to indicate polarity:Partially positive atom

Partially negative atom

points toward more EN atom

What does partial charge mean?• The atoms in ions are completely +1, +2,

-3, -2 and such

• Polar bonds make the atoms just a little bit + and -, like maybe +0.001 and -0.001

• So ions are WAY more + and – than polar covalent bonded atoms

• Ionic bond Na-Cl is completely +1 and -1

• Polar Covalent bond N-F is a little bit + and -

Electronegativity (EN) is the ability of a BONDED atom to attract electrons.

Noble Gases don’t have an ENWhy? Any ideas?

• EN = ability of an atom to pull BONDED electrons close

Well noble gases don’t BOND! So they can’t pull bonded electrons close. So what is the atom with the highest EN???

Yep, F. F pulls electrons closer than anything! F is an electron hog. Nothing holds electrons tighter than F.

•Note C and H are about the SAME in EN so also make nonpolar covalent bonds.

Example:C-H bond in CH4

Nonpolar covalent bonds•When an atom is bonded to itself, that bond is nonpolar because the electrons are shared equally between them. •Diatomic molecules have nonpolar covalent bonds.

Examples:H2, N2, F2, O2, I2, Cl2 , Br2

Polar covalent bonds

•In general, when two different nonmetal atoms are bonded, the bond is polar because the more EN atom pulls the electrons closer so they are shared unequally.

Examples of polar bonds:C-O, H-F, S-F, C-N

Examples:(a) Add the delta notation(b) Add the polarity arrow

C-F O-C C-Cl O-H

Ionic, polar covalent, or nonpolar covalent???

C=O bond

Cl-Cl bond

Na-O bond

C=C bond

Polar covalentNonpolar covIonicNonpolar cov

C-H

Polarity - Review

• This is really important and will come up later again and again and again.

• Think of this as a tug-of-war for bonded electrons. The more EN atom pulls them closer, and since e- are negative, that makes that atom a little bit -. By default the other atom is a little bit +. The bond is thus polar. If the atoms have the same EN, like C and H, then it is a tie (nonpolar).

Ionic, Polar, or Nonpolar Bonds?

• Na-Cl • H-Cl

• H-H

• Cl-C• C-H• O=O• K-O• P-F

IonicPolar CovalentNonpolar CovalentPolar CovalentNonpolar CovalentNonpolar CovalentIonicPolar Covalent

Metallic Bonding

• Pure metals have a freely moving “sea of electrons”.

• The electrons are shared among all the metal atoms.

• This is why they conduct heat and electricity so easily.

12.10 Polarity of Molecules

• All nonpolar bonds = nonpolar molecule.

• Polar bonds that don’t cancel out = polar molecule.

• Polar bonds that do cancel out = nonpolar molecule.

Polar Bonds BUT Nonpolar Molecule

Polar bonds cancel out = nonpolar molecule.

Summary

Draw the electron dot structure, determine shape, bond angle, determine if bonds are polar and if molecule is polar.

• Water

• Ammonia

• Carbon dioxide

Yes, we skipped sections 8 and 9

For Fun if time

• Why care about molecular shape? Well look at what cis-platin can do thanks to its shape! http://www.youtube.com/watch?v=Wq_up2uQRDo&feature=related

• Antioxidants and free radicals – what are they? Free radicals are when a molecule has an unpaired electron. They are considered bad for you. http://www.youtube.com/watch?v=KVyjmt10CH0&feature=related

Self Test

• Page 352

• Try 1-6, 10, 12, 14 (shape only), 16-17 (don’t worry about when it says electron geometry, worry about shape)

• Answers in Appendix J