The Ties That Bind

Chemical Bonding and Interactions

Chemical Bonding and Interactions1. Stable Electron Configurations2. Electron-Dot (Lewis) Structures

1. Drawing, Rules for Drawing2. The Octet Rule3. Some Exceptions to the Rule

3. Ionic Bonding1. Naming ionic compounds2. Drawing

4. Covalent Bonding1. Naming covalent compounds2. Drawing3. Electronegativity and Polar Covalent Compounds

5. Molecular Shapes and the VSEPR Theory6. Intermolecular Forces of Attraction

1. H-bonds, Dipole-Dipole, Ion-Dipole, London Dispersion Forces

It’s all about stability…

1. Fact: Noble gases are inert (undergo few, if any, chemical reactions)

2. Theory: The inertness of these gases is a result of their electronic structure

3. If elements could alter their electron structures like those of the noble gases, they would be less reactive.

Electron Configurations of Cations and Anions

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

Atoms lose electrons so that cation has a noble-gas outer electron configuration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Atoms gain electrons so that anion has a noble-gas outer electron configuration.

Of Representative Elements

+1

+2

+3 -1-2-3

Cations and Anions Of Representative Elements

Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne]

O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne]

Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

What neutral atom is isoelectronic with H- ?

H-: 1s2 same electron configuration as He

9.1

Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons thatparticpate in chemical bonding.

1A 1ns1

2A 2ns2

3A 3ns2np1

4A 4ns2np2

5A 5ns2np3

6A 6ns2np4

7A 7ns2np5

Group # of valence e-e- configuration

(Electron Dot Structures)

*We generally place the electrons one four sides of a square around the element symbol.*Octet rule: we know that s2p6 is a noble gas configuration. We assume that an atom is stable when surrounded by 8 electrons (4 electron pairs).

A violent bond: Sodium and Chlorine Sodium is a soft, reactive metal. Chlorine is a greenish-yellow gas which is

toxic (WWI poison) Sodium chloride is used for our food.

Ionic BondingIonic Bonding

Consider the reaction between sodium and chlorine:Na(s) + ½Cl2(g) NaCl(s)

Chemistry In Action:

Sodium Chloride

Mining Salt Solar Evaporation for Salt

9.2

Li + F Li+ F -

The Ionic Bond

1s22s11s22s22p5 1s21s22s22p6[He][Ne]

Li Li+ + e-

e- + F F -

F -Li+ + Li+ F -

IONIC BONDING-bonding due to

electrostatic attraction arising from an exchange of electrons.

ionic compounds consist of a combination of cations and an anions

• the formula is always the same as the empirical formula

• the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero

The ionic compound NaCl

2.6

Ionic BondingIonic Bonding

*NaCl forms a very regular structure in which each Na+ ion is surrounded by

6 Cl ions.

Similarly, each Cl ion is surrounded by six Na+ ions.

There is a regular arrangement of Na+ and Cl in 3D.

Note that the ions are packed as closely as possible.

Note that it is not easy to find a molecular formula to describe the ionic

lattice.

9.3

Lattice energy (E) increases as Q increases and/or

as r decreases.

cmpd lattice energyMgF2

MgO

LiF

LiCl

2957

3938

1036

853

Q= +2,-1

Q= +2,-2

r F < r Cl

SIDEBAR: Electrostatic (Lattice) Energy

E = kQ+Q-r

Q+ is the charge on the cation

Q- is the charge on the anionr is the distance between the ions

Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions.

9.3

An ion is an atom, or group of atoms, that has a net positive or negative charge.

cation – ion with a positive chargeIf a neutral atom loses one or more electronsit becomes a cation.

anion – ion with a negative chargeIf a neutral atom gains one or more electronsit becomes an anion.

Na 11 protons11 electrons Na+ 11 protons

10 electrons

Cl 17 protons17 electrons Cl-

17 protons18 electrons

2.5

A monatomic ion contains only one atom

A polyatomic ion contains more than one atom

2.5

Na+, Cl-, Ca2+, O2-, Al3+, N3-

OH-, CN-, NH4+, NO3

-

Test Yourself:

1. MgO

2. MgBr2

3. The combination of the ion of magnesium and the ion of nitrogen

Formula of Ionic Compounds

Al2O3

2.6

2 x +3 = +6 3 x -2 = -6

Al3+ O2-

CaBr2

1 x +2 = +2 2 x -1 = -2

Ca2+ Br-

Na2CO3

1 x +2 = +2 1 x -2 = -2

Na+ CO32-

2.6

2.7

Chemical Nomenclature Ionic Compounds

often a metal + nonmetal anion (nonmetal), add “ide” to element name

BaCl2 barium chloride

K2O potassium oxide

Mg(OH)2 magnesium hydroxide

KNO3 potassium nitrate

2.7

Transition metal ionic compounds indicate charge on metal with Roman numerals

FeCl2 2 Cl- -2 so Fe is +2 iron(II) chloride

FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride

Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide

2.7

Test Yourself

1. Write the formula and draw the ionic structure1. Magnesium (II) bromide2. Iron (III) Oxide3. Sodium Azide4. Ammonium Chloride5. Magnesium Nitrate6. Ammonium Phosphate

2. Write the name1. FeO2. Fe2O3

3. LiBr4. KCrO4

5. K2Cr2O7

2.7

Covalent BondingCovalent Bonding

A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.

Why should two atoms share electrons?

F F+

7e- 7e-

F F

8e- 8e-

F F

F F

Lewis structure of F2

lone pairslone pairs

lone pairslone pairs

single covalent bond

single covalent bond

9.4

8e-

H HO+ + OH H O HHor

2e- 2e-

Lewis structure of water

Double bond – two atoms share two pairs of electrons

single covalent bonds

O C O or O C O

8e- 8e-8e-double bonds double bonds

Triple bond – two atoms share three pairs of electrons

N N8e-8e-

N N

triple bondtriple bond

or

9.4

Bond Type

Bond Length

(pm)

C-C 154

CC 133

CC 120

C-N 143

CN 138

CN 116

Lengths of Covalent Bonds

Bond Lengths

Triple bond < Double Bond < Single Bond 9.4

9.4

H F FH

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms

electron richregion

electron poorregion e- riche- poor

+ -

9.5

Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.

Electron Affinity - measurable, Cl is highest

Electronegativity - relative, F is highest

X (g) + e- X-(g)

9.5

9.5

Bond Polarity and ElectronegativityBond Polarity and Electronegativity

ElectronegativityElectronegativity

Drawing Lewis StructuresDrawing Lewis Structures1. Add the valence electrons.2. Identify the central atom (usually the one with the highest molecular mass, lowest electronegativity, or closest to the center of the periodic table).3. Place the central atom in the center of the molecule and add all other atoms around it.4.Place one bond (two electrons) between each pair of atoms.5.Complete the octet for the central atom.6.Complete the octets for all other atoms. Use double bonds if necessary.7. Place remaining electrons on the central atom.

Write the Lewis structure of nitrogen trifluoride (NF3).

Step 1 – N is less electronegative than F, put N in center

F N F

F

Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)

5 + (3 x 7) = 26 valence electrons

Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

9.6

Write the Lewis structure of the carbonate ion (CO32-).

Step 1 – C is less electronegative than O, put C in center

O C O

O

Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electrons

Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

9.6

Step 5 - Too many electrons, form double bond and re-check # of e-

2 single bonds (2x2) = 41 double bond = 4

8 lone pairs (8x2) = 16Total = 24

Trial

1. Write the Lewis stucture for1. H2O

2. CO2

3. NH3

4. CH4

5. N2

6. BH3

Exceptions to the Octet Rule

The Incomplete Octet

H HBeBe – 2e-

2H – 2x1e-

4e-

BeH2

BF3

B – 3e-

3F – 3x7e-

24e-

F B F

F

3 single bonds (3x2) = 69 lone pairs (9x2) = 18

Total = 24

9.9

Exceptions to the Octet Rule

Odd-Electron Molecules

N – 5e-

O – 6e-

11e-

NO N O

The Expanded Octet (central atom with principal quantum number n > 2)

SF6

S – 6e-

6F – 42e-

48e-

S

F

F

F

FF

F

6 single bonds (6x2) = 1218 lone pairs (18x2) = 36

Total = 48

9.9

Strengths of Covalent BondsStrengths of Covalent Bonds

Bond Enthalpies and Bond LengthBond Enthalpies and Bond Length

Covalent

share e-

Polar Covalent

partial transfer of e-

Ionic

transfer e-

Increasing difference in electronegativity

Classification of bonds by difference in electronegativity

Difference Bond Type

0 Covalent

2 Ionic

0 < and <2 Polar Covalent

9.5

Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; andthe NN bond in H2NNH2.

Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic

H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent

N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent

9.5