Gián án Chapter 9 Chemical Bonding I:Lewis Theory

Tro, Chemistry: A Molecular Approach 13Lewis Symbols of Ions • Cations have Lewis symbols without valence electrons Lost in the cation formation • Anions have Lewis symbols with 8 vale

Bonding Theories

others are not

 why is water H 2 O, not HO or H 5 O

bonding

Lewis structures – that allow us to predict many

properties of molecules

 aka Electron Dot Structures

 such as molecular shape, size, polarity

Why Do Atoms Bond?

• processes are spontaneous if they result in a system

with lower potential energy

• chemical bonds form because they lower the potential energy between the charged particles that compose

atoms

• the potential energy between charged particles is

directly proportional to the product of the charges

• the potential energy between charged particles is

inversely proportional to the distance between the

charges

Chemistry, Julia Burdge, 2 nd e., McGraw Hill.

Potential Energy Between

Charged Particles

0 is a constant

 = 8.85 x 10 -12 C 2 /J∙m

• for charges with the same sign, Epotential is + and the

magnitude gets less positive as the particles get farther apart

• for charges with the opposite signs, Epotential is  and the magnitude gets more negative as the particles get



particles get closer together To bring them closer requires the addition of more energy.

The attraction between opposite-charged

particles increases as the particles get closer together Bringing them closer lowers the potential energy of the system.

Chemistry, Julia Burdge, 2 nd e., McGraw Hill.

• a chemical bond forms when the potential

energy of the bonded atoms is less than the

potential energy of the separate atoms

• have to consider following interactions:

nucleus-to-nucleus repulsion

electron-to-electron repulsion

nucleus-to-electron attraction

Tro, Chemistry: A Molecular Approach 7

electrons shared

metal to

electronspooled

Types of Bonding

Tro, Chemistry: A Molecular Approach 9

nonmetals have high electron affinities, relatively

good to add electrons to

Covalent Bonds

• nonmetals have relatively high ionization energies, so

it is difficult to remove electrons from them

• when nonmetals bond together, it is better in terms of potential energy for the atoms to share valence

Tro, Chemistry: A Molecular Approach 11

Determining the Number of Valence

Lewis Symbols of Atoms

• aka electron dot symbols

• use symbol of element to represent nucleus and

inner electrons

• use dots around the symbol to represent valence

electrons

pair first two electrons for the s orbital

put one electron on each open side for p electrons

then pair rest of the p electrons

Tro, Chemistry: A Molecular Approach 13

Lewis Symbols of Ions

• Cations have Lewis symbols without

valence electrons

Lost in the cation formation

• Anions have Lewis symbols with 8 valence

Stable Electron Arrangements

And Ion Charge

• Metals form cations by losing

enough electrons to get the

same electron configuration as

the previous noble gas

• Nonmetals form anions by

gaining enough electrons to get

the same electron configuration

as the next noble gas

• The noble gas electron

configuration must be very

stable

Electron Config

Electron Config

Tro, Chemistry: A Molecular Approach 16

 Li loses its one valence electron

 H shares or gains one electron

 though it commonly loses its one electron to become H +

 Be loses 2 electrons to become Be 2+

 though it commonly shares its two electrons in covalent bonds, resulting in 4 valence electrons

 B loses 3 electrons to become B 3+

 though it commonly shares its three electrons in covalent bonds, resulting in 6 valence electrons

 expanded octets for elements in Period 3 or below

 using empty valence d orbitals

Lewis Theory

• the basis of Lewis Theory is that there are

certain electron arrangements in the atom that are more stable

octet rule

• bonding occurs so atoms attain a more stable

electron configuration

more stable = lower potential energy

no attempt to quantify the energy as the calculation

is extremely complex

Tro, Chemistry: A Molecular Approach 18

Properties of Ionic Compounds

• hard and brittle crystalline solids

all are solids at room temperature

• melting points generally > 300C

• the liquid state conducts electricity

the solid state does not conduct electricity

• many are soluble in water

the solution conducts electricity well

Melting an Ionic Solid

Conductivity of NaCl

in NaCl(s), the

ions are stuck in

position and not

to the charged rods

Tro, Chemistry: A Molecular Approach 20

Lewis Theory and Ionic Bonding

• Lewis symbols can be used to represent the

transfer of electrons from metal atom to

nonmetal atom, resulting in ions that are

attracted to each other and therefore bond

Predicting Ionic Formulas Using Lewis Symbols

• electrons are transferred until the metal loses all its

valence electrons and the nonmetal has an octet

• numbers of atoms are adjusted so the electron transfer comes out even

Tro, Chemistry: A Molecular Approach 22

Energetics of Ionic Bond Formation

• the ionization energy of the metal is endothermic

Na(s) + ½Cl2(g) → NaCl(s) H°f = +147 kJ/mol

• However, the heat of formation of most ionic

compounds is exothermic and generally large; Why?

Ionic Bonds

• electrostatic attraction!!

no direct anion-cation pair

• ions arranged in a pattern called a crystal lattice

every cation surrounded by anions; and every anion surrounded by cations

maximizes attractions between + and - ions

Lattice Energy

• the lattice energy is the energy released when the

solid crystal forms from separate ions in the gas state

 always exothermic

 hard to measure directly, but can be calculated from

knowledge of other processes

• lattice energy depends directly on size of charges and inversely on distance between ions

For NaCl the energy

of forming a crystal lattice is -787 kJ/mol



Born-Haber Cycle

• method for determining the lattice energy of an

ionic substance by using other reactions

use Hess’s Law to add up heats of other processes

• H°f(salt) = H°f(metal atoms, g) + H°f(nonmetal atoms, g) + H°f(cations, g) + H°f(anions, g) + H°f(crystal lattice)

 H° f ( crystal lattice ) = Lattice Energy

 metal atoms (g)  cations (g), H°f = ionization energy

 don’t forget to add together all the ionization energies to get to the desired cation

 M 2+ = 1 st IE + 2 nd IE

 nonmetal atoms (g)  anions (g), H° = electron affinity

Born-Haber Cycle for NaCl

Na(s) → Na+(g) + 1 e ─ H° = +496 kJ/mol

Cl(g) + 1 e ─ → Cl ─(g) H° = ─ 349 kJ/mol

Na(s) + ½Cl2(g) → NaCl(s) H° f = +147 kJ/mol

NaCl the energy

of forming a crystal lattice is -787 kJ/mol

Practice - Given the Information Below,

Determine the Lattice Energy of MgCl2

Tro, Chemistry: A Molecular Approach 28

Practice - Given the Information Below, Determine the Lattice Energy of MgCl2

1450 (

kJ) 738 (

kJ) 121.3 (

2 kJ) 147.1 (

kJ) 3 641 (

-H

H 2 H

H H

2 H

H H

H H

2 H

H H

2 H

H

energy lattice

f

energy lattice

f

f 5 f

4 f

3 f

2 f

1 f

6 energy

lattice

f

energy lattice

f f

5 f

4 f

3 f

2 f

1 f

Trends in Lattice Energy

Ion Size

• the force of attraction between charged

particles is inversely proportional to the

distance between them

• larger ions mean the center of positive charge (nucleus of the cation) is farther away from

negative charge (electrons of the anion)

larger ion = weaker attraction = smaller lattice energy

Tro, Chemistry: A Molecular Approach 30

Trends in Lattice Energy

Ion Charge

• the force of attraction between

oppositely charged particles is

directly proportional to the product

of the charges

• larger charge means the ions are

more strongly attracted

larger charge = stronger attraction =

larger lattice energy

• of the two factors, ion charge

generally more important

Lattice Energy =

-910 kJ/mol

Lattice Energy =

-3414 kJ/mol

Tro, Chemistry: A Molecular Approach 32

Example 9.2 – Order the following ionic

compounds in order of increasing magnitude of

lattice energy.

CaO, KBr, KCl, SrO

First examine the ion charges and

order by product of the charges Ca

2+ & O 2- , K + & Br ─ , K +

& Cl ─ , Sr 2+ & O 2─

(KBr, KCl) < (CaO, SrO)

Then examine the ion sizes of

each group and order by radius;

Ionic Bonding Model vs Reality

• ionic compounds have high melting points and boiling points

 MP generally > 300°C

 all ionic compounds are solids at room temperature

• because the attractions between ions are strong,

breaking down the crystal requires a lot of energy

 the stronger the attraction (larger the lattice energy), the

higher the melting point

Tro, Chemistry: A Molecular Approach 34

Ionic Bonding Model vs Reality

• ionic solids are brittle and hard

• the position of the ion in the crystal is critical to establishing maximum attractive forces –

displacing the ions from their positions results

in like charges close to each other and the

repulsive forces take over

-Ionic Bonding Model vs Reality

• ionic compounds conduct electricity in the liquid state

or when dissolved in water, but not in the solid state

• to conduct electricity, a material must have charged

particles that are able to flow through the material

• in the ionic solid, the charged particles are locked in

position and cannot move around to conduct

• in the liquid state, or when dissolved in water, the ions have the ability to move through the structure and

therefore conduct electricity

Tro, Chemistry: A Molecular Approach 37

Single Covalent Bonds

• two atoms share a pair of electrons

Double Covalent Bond

• two atoms sharing two pairs of electrons

Tro, Chemistry: A Molecular Approach 39

Triple Covalent Bond

• two atoms sharing 3 pairs of electrons

•• •• •• ••

•• N

N N ··

··

Covalent Bonding Predictions from Lewis Theory

• Lewis theory allows us to predict the formulas of

molecules

• Lewis theory predicts that some combinations should be stable, while others should not

 because the stable combinations result in “octets”

• Lewis theory predicts in covalent bonding that the

attractions between atoms are directional

 the shared electrons are most stable between the bonding atoms

resulting in molecules rather than an array

Tro, Chemistry: A Molecular Approach 41

Covalent Bonding Model vs Reality

• molecular compounds have low melting points and

boiling points

 MP generally < 300°C

 molecular compounds are found in all 3 states at room

temperature

• melting and boiling involve breaking the attractions

between the molecules, but not the bonds between

the atoms

 the covalent bonds are strong

 the attractions between the molecules are generally weak

 the polarity of the covalent bonds influences the strength of the intermolecular attractions

Intermolecular Attractions vs Bonding

Tro, Chemistry: A Molecular Approach 43

Ionic Bonding Model vs Reality

• some molecular solids are brittle and hard, but many are soft and waxy

• the kind and strength of the intermolecular

attractions varies based on many factors

• the covalent bonds are not broken, however, the polarity of the bonds has influence on these

attractive forces

Ionic Bonding Model vs Reality

• molecular compounds do not conduct electricity in the liquid state

• molecular acids conduct electricity when dissolved in water, but not in the solid state

• in molecular solids, there are no charged particles

around to allow the material to conduct

• when dissolved in water, molecular acids are ionized, and have the ability to move through the structure and therefore conduct electricity

Tro, Chemistry: A Molecular Approach 45

Bond Polarity

• covalent bonding between unlike atoms results in unequal sharing of the electrons

Tro, Chemistry: A Molecular Approach 47

Electronegativity

• measure of the pull an atom has on bonding electrons

• increases across period (left to right) and

• decreases down group (top to bottom)

fluorine is the most electronegative element

francium is the least electronegative element

• the larger the difference in

electronegativity, the more polar the bond

negative end toward more electronegative atom

Electronegativity Scale

Electronegativity and Bond Polarity

• If difference in electronegativity between bonded atoms is 0, the bond is pure covalent

 equal sharing

• If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent

• If difference in electronegativity between bonded atoms 0.5 to

1.9, the bond is polar covalent

• If difference in electronegativity between bonded atoms larger

than or equal to 2.0, the bond is ionic

“100%”

4% Percent Ionic Character51%

Electronegativity Difference

ENCl = 3.0

ENNa = 1.0 3.0 – 0.9 = 2.1

Ionic

Tro, Chemistry: A Molecular Approach 51

Bond Dipole Moments

polarity of a bond

 a dipole is a material with positively and negatively charged ends

 measured

• dipole moment, , is a measure of bond polarity

 it is directly proportional to the size of the partial charges and

directly proportional to the distance between them

Tro, Chemistry: A Molecular Approach 53

Dipole Moments

Water – a Polar Molecule

attracted

to a charged glass rod

Tro, Chemistry: A Molecular Approach 55

Example 9.3(c) - Determine whether an N-O bond is ionic, covalent, or polar covalent

• Determine the electronegativity of each element

N = 3.0; O = 3.5

• Subtract the electronegativities, large minus small

(3.5) - (3.0) = 0.5

• If the difference is 2.0 or larger, then the bond is

ionic; otherwise it’s covalent

difference (0.5) is less than 2.0, therefore covalent

• If the difference is 0.5 to 1.9, then the bond is

polar covalent; otherwise it’s covalent

difference (0.5) is 0.5 to 1.9, therefore polar covalent

• allows us to predict shapes of molecules

• allows us to predict properties of molecules and how they will interact together

Tro, Chemistry: A Molecular Approach 57

Lewis Structures

• use common bonding patterns

 C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be

= 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs

 often Lewis structures with line bonds have the lone

pairs left off

 their presence is assumed from common bonding patterns

• structures which result in bonding patterns

different from common have formal charges

Writing Lewis Structures of Molecules

HNO31) Write skeletal structure

 H always terminal

 in oxyacid, H outside attached to O’s

 make least electronegative atom central

 N is central

2) Count valence electrons

 sum the valence electrons for each

atom

 add 1 electron for each − charge

 subtract 1 electron for each + charge

O N

O H

O

N = 5

H = 1

O3 = 3∙6 = 18Total = 24 e-

Tro, Chemistry: A Molecular Approach 59

Writing Lewis Structures of Molecules

HNO33) Attach central atom to the surrounding atoms with

pairs of electrons and subtract from the total

O N

O H

Used 8Left 16

Writing Lewis Structures of Molecules

HNO3

4) Complete octets, outside-in

 H is already complete with 2

ElectronsStart 24

Used 8Left 16

ElectronsStart 16

Used 16Left 0

Tro, Chemistry: A Molecular Approach 61

Writing Lewis Structures of Molecules

HNO3

5) If all octets complete, give extra electrons to central

atom

 elements with d orbitals can have more than 8 electrons

 Period 3 and below

6) If central atom does not have octet, bring in

electrons from outside atoms to share

 follow common bonding patterns if possible