Chapter 9 Chemical Bonding I: Lewis Theory Chemistry: A Molecular Approach, 1 st Ed. Nivaldo Tro.

Chapter 9Chemical

Bonding I:Lewis Theory

Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro

Question

Complete the following sentence…

Properties of substances can be explained in terms of differences in chemical __________

e.g. -salt dissolves in water better than oil

-certain substances are electrolytes

-alcohol evaporates quicker than water

-wax melts at a lower temperature than salt

Tro, Chemistry: A Molecular Approach 3

Bonding Theories

• explain how and why atoms attach together• one of the simplest bonding theories is called Lewis Theory• Lewis Theory uses valence electrons to explain bonding• explains why some combinations of atoms are stable and others

are not• using Lewis Theory, we can draw models – called Lewis

structures – that allow us to predict many properties of moleculessuch as molecular shape, size, polarity


Types of Bonds

Types of Atoms Type of BondBond

Characteristic

metals to

nonmetalsIonic

electrons

transferred

nonmetals to

nonmetalsCovalent

electrons

shared

metal to

metalMetallic

electrons

pooled

10

Types of Bonding


Determining the Number of Valence Electrons in an Atom

• the column number on the Periodic Table tells us the no. valence e-

1A 2A 3A 4A 5A 6A 7A 8A

Li Be B C N O F Ne

1 e- 2 e- 3 e- 4 e- 5 e- 6 e- 7 e- 8 e-


Lewis Symbols of Atoms

• use symbol of element to represent nucleus and inner electrons

• use dots around the symbol to represent valence electronspair first two electrons for the s orbitalput one electron on each open side for p electrons then pair rest of the p electrons

Li Be

B

C

N

O

F

Ne


Lewis Symbols of Ions• Cations have Lewis symbols without valence e-

e.g. lithium

• Anions have Lewis symbols with 8 valence electrons

e.g. flourine

Li• Li+

F

F

e- loss

e- gain

Question

Draw Lewis dot structures of elemental magnesium and magneisum ion

Draw Lewis dot structures of elemental nitrogen and the nitride ion


Stable Electron ArrangementsAnd Ion Charge

• Metals form cations by losing e- to become isoelectric to the previous noble gas

• Nonmetals form anions by gaining enough e- to become isoelectric to the previous noble gas

[Ne] = 1s22s22p6

Atom Atom’s Electron Config

Ion Ion’s Electron Config

Na [Ne]3s1 Na+ [Ne]

Mg [Ne]3s2 Mg+2 [Ne]

Al [Ne]3s23p1 Al+3 [Ne]

O [He]2s22p4 O-2 [Ne]

F [He]2s22p5 F- [Ne]


Lewis Theory• the basis of Lewis Theory is that there are

certain electron arrangements in the atom that are more stableoctet rule

• bonding occurs so atoms attain a more stable electron configuration


Octet Rule• when atoms bond, they tend to gain, lose, or share e- to result in

8 valence e-

• ns2np6

noble gas configuration

• many exceptions H, Li, Be, B attain an electron configuration like He

He = 2 valence e- Li loses its one valence e-

H shares or gains one e-

though it commonly loses its one electron to become H+ Be loses 2 electrons to become Be2+

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

B loses 3 electrons to become B3+

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 and Ionic Bonding

• Transfer of e- from metal atom to nonmetal atom, resulting in ions that are attracted to each other and therefore bond, e.g. NaCl

ClNa +

ClNa + NaCl


Predicting Ionic FormulasUsing Lewis Symbols

• e- are transferred until the metal loses all its valence e- and the nonmetal obtains an octet

O

Li

Li

2

O2 Li + Li2O


Crystal Lattice

• Ionic substances exist as crystal lattices of repeating unit cells

Model of NaCl


Ionic BondingModel vs. Reality

• ionic compounds have high melting points and boiling pointsMP generally > 300°Call 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

Properties

• Describe the general properties of ionic compounds

• Metals react with non-metals:Crystalline solids3-D units extendedhigh mp/bp (all solids)brittleaqueous solutions conduct electricity


Ionic Bonding

• When ionic compounds are dissolved in water, they dissociate to form aqueous ions:

NaCl(s) → Na+(aq) + Cl-(aq)

• The resulting solution conducts electricity and is called an electrolyte


Conductivity of NaCl

in NaCl(s), the ions are stuck in position and not allowed to move to the charged rods

in NaCl(aq), the ions are separated and allowed to move to the charged rods

Question

Use Lewis dot structures to represent the formation of aluminum bromide

Use Lewis dot structures to represent the formation of lithium hydride


Types of Bonds

Types of Atoms Type of Bond Bond Characteristic

metals to nonmetals Ionic e- transferred

nonmetals to nonmetals Covalent e- shared

metal to metal Metallic e- pooled

45

Types of Bonding


Single Covalent Bonds• two atoms share a pair of electrons

F••

•••• • F

•••••••

F••

••

•• ••••F••••

F F

e.g. fluorine


Single Covalent Bonds

HH O•• ••••••

H•H• O••

• •

••e.g. water

octet

duet duet

2 bonding pairs2 lone pairs


Double Covalent Bond• two atoms sharing two pairs of electrons

O••••O••

••••••

O••

• •

••O••

• •••

O O······ ··

e.g. oxygen


Triple Covalent Bond• two atoms sharing 3 pairs of electrons

N••

• •

•N••

• ••

N•••••••••• N

N N····

e.g. nitrogen


Covalent BondingPredictions 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 also shows that covalent bonds are highly directional the shared electrons are most stable between the bonding atoms resulting in molecules rather than an array


Ionic BondingModel 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


Bond Polarity• covalent bonding between unlike atoms results in unequal sharing

of the e-

one atom pulls the electrons in the bond closer to its sideone end of the bond has larger electron density than the other

• the result is a polar covalent bond bond polarity the end with the larger electron density gets a partial negative

charge the end that is electron deficient gets a partial positive charge


HF

FH

EN 2.1 EN 4.0EN 2.1


Electronegativity

• Ability of an atom to attract e- to itself in a chemical bond• increases across period (left to right) and• decreases down group (top to bottom)

59

Electronegativity and Bond Polarity• If ΔE.N. between bonded atoms is 0, the bond is pure covalent

equal sharing• If ΔE.N. between bonded atoms is 0.1 - 0.4, the bond is nonpolar covalent• If ΔE.N. between bonded atoms 0.5 - 1.9, the bond is polar covalent• If ΔE.N. between bonded atoms ≥ 2.0, the bond is ionic

“100%”

0 0.4 2.0 4.0

4% 51%

Percent Ionic Character

Electronegativity Difference

IONICPCNP


Bond Polarity

ENCl = 3.0ΔEN = 3.0 - 3.0 = 0

Pure Covalent

ENCl = 3.0ENH = 2.1

ΔEN = 3.0 – 2.1 = 0.9Polar Covalent

ENCl = 3.0ENNa = 1.0

ΔEN = 3.0 – 0.9 = 2.1Ionic


Lewis Structures of Molecules

• shows pattern of valence electron distribution in the molecule

• useful for understanding the bonding in many compounds

• allows us to predict shapes of molecules

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


Writing Lewis Structures of Molecules HNO3

1) 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 e-

sum the valence electrons for each atom

add 1 e- for each −ve charge subtract 1 e- for each +ve charge

ONOH

O

N = 5H = 1O3 = 3(6) = 18Total = 24 e-



3) Attach central atom to the surrounding atoms with pairs of e- and subtract from the total

ONOH

O

———

e-

Start 24Used 8Left 16



4) Complete octets, outside-in H is already complete with 2

1 bond

and re-count e-

:

::

——— ONOH

O

N = 5H = 1O3 = 3(6) = 18Total = 24 e-

e-

Start 24Used 8Left 16

e-

Start 16Used 16 (8 pairs)Left 0


Writing Lewis Structures of Molecules HNO35) 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

:

::

—— ONOH|

O


Practice - Lewis Structures

• CO2

• NO2-

• NH3

Draw Lewis structures for the following:

Tro, Chemistry: A Molecular ApproacH74

Writing Lewis Formulas of Molecules (cont’d)

7) Assign formal charges to the atoms

a) formal charge = valence e- - lone pair e- - ½ bonding e-

b) follow the common bonding patterns

OSO

H

|

HOCCH

|||

OH

0 +1 -1

all 0sum of all the formal charges in a molecule = 0in an ion, total equals the charge


Practice - Assign Formal Charges

• CO2

• NO2-

• NH3

O N O ••

••

••

••

••••-


Practice - Assign Formal Charges

• CO2

• NO2-

• NH3

O N O ••

••

••

••

••••

all 0

-1

all 0

-


Resonance• when there is more than one Lewis structure for a molecule that

differ only in the position of the electrons, they are called resonance structures

• the actual molecule is a combination of the resonance forms – a resonance hybrid it does not resonate between the two forms, though we often

draw it that way

• look for multiple bonds or lone pairs

•••• •• ••••••••

•• ••O S O O S O•••••• ••••

••••

••••


Rules of Resonance Structures• Resonance structures must have the same connectivity

only electron positions can change• Resonance structures must have the same number of

electrons• Second row elements have a maximum of 8 electrons

bonding and nonbonding third row can have expanded octet

• Formal charges must total same• Better structures have fewer formal charges• Better structures have smaller formal charges• Better structures have − formal charge on more

electronegative atom


O N

O

O·· ··

········

··

··

Drawing Resonance Structures1. draw first Lewis structure that

maximizes octets2. assign formal charges3. move electron pairs from atoms

with (-) formal charge toward atoms with (+) formal charge

4. if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond

5. if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet.

-1

-1

+1

O N

O

O

·· ····

····

······

-1

-1 +1

0

-

-


Exceptions to the Octet Rule

• expanded octetselements with empty d orbitals can have more

than 8 electrons

• odd number electron species e.g., NOwill have 1 unpaired electronfree-radicalvery reactive

• incomplete octetsB, Al


Drawing Resonance Structures1. draw first Lewis structure that

maximizes octets2. assign formal charges3. move electron pairs from atoms

with (-) formal charge toward atoms with (+) formal charge

4. if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond

5. if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet.

O S

O

O

O

HH

·· ··

········

··

······

-1

-1

+2

O S

O

O

O

HH

··

······

··

······

0

0

0

Question

Draw Lewis structures with assigned formal charges of HCl, H2O2 and SF6


Metallic Bonds• low ionization energy of metals allows them to lose electrons

easily• the simplest theory of metallic bonding involves the metals

atoms releasing their valence electrons to be shared by all to atoms/ions in the metalan organization of metal cation islands in a sea of electronselectrons delocalized throughout the metal structure

• bonding results from attraction of cation for the delocalized electrons


Metallic Bonding


Metallic BondingModel vs. Reality

• metallic solids conduct electricity• because the free electrons are mobile, it allows the

electrons to move through the metallic crystal and conduct electricity

• as temperature increases, electrical conductivity decreases • heating causes the metal ions to vibrate faster, making it

harder for electrons to make their way through the crystal



• metallic solids conduct heat

• the movement of the small, light electrons through the solid can transfer kinetic energy quicker than larger particles

• metallic solids reflect light

• the mobile electrons on the surface absorb the outside light and then emit it at the same frequency



• metallic solids are malleable and ductile• because the free electrons are mobile, the direction of the

attractive force between the metal cation and free electrons is adjustable

• this allows the position of the metal cation islands to move around in the sea of electrons without breaking the attractions and the crystal structure



• metals generally have high melting points and boiling pointsall but Hg are solids at room temperature

• the attractions of the metal cations for the free electrons is strong and hard to overcome

• melting points generally increase to right across period• the charge on the metal cation increases across the period,

causing stronger attractions• melting points generally decrease down column• the cations get larger down the column, resulting in a larger

distance from the nucleus to the free electrons

Chapter 9 Chemical Bonding I: Lewis Theory Chemistry: A Molecular Approach, 1 st Ed. Nivaldo Tro.

Documents

Transcript of Chapter 9 Chemical Bonding I: Lewis Theory Chemistry: A Molecular Approach, 1 st Ed. Nivaldo Tro.