chapter 9 chemical bonding. section 9.1: why does bonding occur in the first place? bonding lowers...

Chapter 9Chemical Bonding

Section 9.1: Why does bonding occur in the first place?

Bonding lowers the potential energy between positive and negative particles (p341).

What is potential energy?

1 type of Potential Energy: Gravitational P.E.

Ball On Top of a Hill

P.E. = mgh

h

m

Energy changes forms: P.E. Kinetic Energy (K.E.)



Energy changes forms

MechanicalEnergy

ElectricalEnergy

ChemicalEnergy

Light (Radiant)Energy

Heat (Thermal)Energy

Friction

EnginesGenerator

Motor

BatteryCharger

Battery

Photosynthesis

Chemiluminescence

SolarHeaterFIre



When chemical bonds form: Chemical P.E. changes to Heat Energy & Light Energy

MechanicalEnergy

ElectricalEnergy

ChemicalEnergy

Light (Radiant)Energy

Heat (Thermal)Energy


http://chemsite.lsrhs.net/chemKinetics/PotentialEnergy.html


Energy changes forms: Chemical P.E. Heat & Light Energy

Section 9.1: Three Type of Bonds

http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch10/non.php

Ionic bonding: Metal + Nonmetal (Valence e- transferred)Covalent bonding: Nonmetal + Nonmetal (Valence e- shared)Metallic bonding: Metal + Metal (“Sea” of e-)

http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch10/non.php

Review: Valence Electrons – e- involved in forming compounds (Ch 8, p315)

Boron (B) Magnesium (Mg) Hydrogen (H)

How many valence e-?

How many needed for full outer shell?

Total valance e-:

Concept Check

Section 9.1: Two Bond Types With Localized Electrons

Ionic & Covalent Bonding

Representative Elements

Transition Elements

For REPRESENTATIVE elements:• period (row) = shell # (n = 1, 2, 3, 4….n)• group (column) = # of e- in outer shell

Group # # of valence e-

IA 1

IIA 2IIIA 3

IVA 4

VA 5

VIA 6

VIIA 7

VIIA 8

Shellsof anatom

Section 9.1: Two Bond Types With Localized Electrons

Ionic & Covalent Bonding: Why do ionic bonds form instead of covalent bonds, and vice versa?

Covalent Bonds Ionic Bonds

“Bonding Continuum”

nonmetals + nonmetal metal + nonmetal

Polar Covalent BondNonpolar Covalent Bond Ionic Bond

Electrons are shared unequally. Electrons are transferred.

Extent of electron sharing in Covalent Bonds

e-’s shared between atomsof the same element:

Equal Sharing

e-’s shared between atomsof different elements:

Unequal Sharing

Unequal sharing – occurs because one of the atoms in a bond has a stronger attraction for the pair of e-’s than does the other atom

Why does one atom have a stronger attraction for e-?

Definition: electronegativity (E.N) – the ability of an atom to attract the shared electrons

Increasing E.N.D

ecre

asin

g E

.N.

Electronegativity

Rule for Bond FormationThe atom with the greater E.N. pulls the shared electrons closer to its nucleus resulting in (1) – charge on high E.N. atom (2) + charge on low E.N. atom

More later: Section 9.5



e- sharing2 nonmetals

e- transfermetal + nonmetal

E.N. difference > 1.70.4 < E.N. < 1.7 1.70.4 < E.N. Polar ColvalentNonpolar Colvalent

Why do ionic bonds form instead of covalent bonds, and vice versa?

Example: Oxygen (O) bonds withMagnesium (Mg): MgO

E.N. of O = 3.5E.N. of Mg = 1.2

E.N. difference = 2.3

Answer: Electronegativity Differences

Section 9.1: “The Other” Bond Type With Delocalized Electrons

Metallic Bonding

A messy “sea” of electrons

Metallic Bonding - Delocalized Covalent Bonding, Ionic Bonding - Delocalized

Electrons fit neatly into shells.

Section 9.1: “The Other” Bond Type With Delocalized Electrons

Metallic Bonding


Metallic Bonding - Delocalized

Outer e-

Inner e-

Lewis Electron-Dot Symbols

(1) Element symbol – nucleus + inner electronsEx: The element lithium has an element symbol Li

(2) Surrounding dots – valence electrons (outer most shell)

Two parts:

Li

Different elements can have the same number of dots

Be

Mg

Same Group(Column)

Review: Ions

Ion – charged particles that form when an atom gains or loses one or more electrons(Ch2, p60)

Mg

Cl Cl-

Mg2+

Element Ion Ion Type

Cation

Anion

Review: Electron Configuration and Orbital Diagrams (Ch8, p304-317)

MgExample:

Concept Check

• End of Chapter Problems in-class (for now):

9.7, 9.9, 9.13, 9.15

Write the ion for the following elements: K, Br, Sr, Ar, O

For example, the ion for Mg is Mg2+.

• Suggested Optional Practice Problems (for outside of class):

9.6, 9.8, 9.10, 9.12, 9.14 (Answers in back of book or online)

Section 9.2: Ionic Bonding

Central idea: Electrons are transferred from metal atoms to nonmetal atoms to form ions that come together in a solid ionic compound.

HH

O O OC

Examples: Water (H2O) Carbon Dioxide (CO2)

Sodium chloride (NaCl)

Contrast with molecules formed during covalent bonding (more later).

Na – metal

Cl - nonmetal

Solid Ionic compound

Rule: The total number of e- lost by the metal atom equals the total number gained by the nonmetal atom.

Section 9.2: Ionic Bonding

Cl-Na+

lost gained

Why is the melting point of MgO higher than the melting point of KCl?

Behavior of Ionic Compounds

Lattice Energy(∆Hº

lattice)

Section 9.2: Lattice Energy

Definition – The enthalphy change that occurs when 1 mol of ionic solid separates into gaseous ions.

For Review of Enthalpy: Ch6, p243

Lattice Energy denoted as: ∆Hºlattice

∆Hºlattice cannot be measured directly, BUT it can be calculate using the:

Born-Haber cycle

Section 9.2: Born-Haber Cycle

Uses Hess’s Law: Total enthalpy of an overall reaction is the sum of the enthalpy changes of individual reactions. (∆Htotal = ∆Hrxn1 + ∆Hrxn2 +……….)

*Not actual steps.

Section 9.2: Trends in Lattice Energy

Coulomb’s Law (Ch2)

Section 9.2: Trends in Lattice Energy

So, why is the melting point of MgO higher than the melting point of KCl?

Behavior of Ionic Compounds

Concept Check

• End of Chapter Problems in-class (for now):

9.27, 9.30

• Suggested Optional Practice Problems (for outside of class):

9.26, 9.28 (Answers in back of book)

Problem 9.30

Section 9.3: Covalent Bonding

HH

O O OC

Examples: Water (H2O) Carbon Dioxide (CO2) Organic Compounds

Contrast with ionic solids formed during ionic bonding (discussed previously).

Sodium chloride (NaCl)

Na – metal

Cl - nonmetal

Nonmetal + Nonmetal

C C C H

H

HH

H

H

H

H

e- sharing – primary way that atoms interact


Why do covalent bonds form?

Lower P.E. = More stable


How are the electrons distributed?

ElectronDensity

• Bonding Pairs & Lone Pairs

• Bond Type – double, single, triple

In order for each atom to have a full outer shell (2 e- for H, He; 8 e- for others), the electrons arrange themselves in certain configurations:


Bond Energy (B.E.) – aka Bond Enthalpy or Bond Strength

Covalent Bond Strength – depends on strength of attraction between nuclei and shared electrons

Bond Energy – energy needed to overcome attraction and

break the bond


Bond Energy (B.E.)

Bond formation is exothermic: ∆Hº always +

Bond breakage is endothermic: ∆Hº always -

Absolute value of B.E. – Each bond has its own unique B.E. due to variations in:(1) e- density(2) charge(3) atomic radii


Strength of Bond different than E required to pull atoms apart (B.E.)

Weaker Bonds =Higher Energy

“Shallow Energy Well”

Stronger Bonds =Lower Energy

“Deeper Energy Well”

Less E needed to break.

Lower B.E.

More E needed to break.

Higher B.E.


Bond Energy (B.E.) and Bond Length

Bond Length – sum of the radii of the bonded atoms (analogous to distance in Coloumb’s Law)

At minimum E point.

R2 = 0.3155

0

50

100

150

200

250

300

150 250 350 450 550

Bond Energy

Bo

nd

Le

ng

th


Bond Energy (B.E.) and Bond Length

This relationshipholds, in general,ONLY for single

bonds.


Bond Type (Single, Double, Triple) also matters

Nuclei more attracted to 2 shared pairs of e- than one shared pair of e-.

Higher bond order = Shorter bond length = Higher Bond Energy

Same twoelements,

different B.E.


Periodic Table Trends Without Detailed Bond Lengths

The closer theatoms, the

stronger the bond.

Bond Energy:C—F > C—Cl > C—Br


Covalent Bonds are stronger than Ionic Bonds

So why, then, do covalent compounds have lower melting pointsthan ionic compounds?

Example: CCl4 m.p. = -23 ºC NaCl m.p. = 800 ºC

solid liquid gas

Strong covalent bonding forcesHold atoms together

Weak intermolecular forcesHold molecules together

(More in Chapter 12)

Chemical Reaction Phase Change

HH

O

-

++

Section 9.4: Bond Energy and Chemical Change

http://chemsite.lsrhs.net/chemKinetics/PotentialEnergy.html

Where does the heat that is released come from?


Total energy of a chemical system = K.E. + P.E.

Example of a chemical system

A container filled with molecules.

Kinetic Energy (K.E.)

Three types:(1) Vibrational(2) Rotational(3) Translational

• Does not change during chemical reaction (depends on T). Changes during a Phase Change (Chapter 12).

solid liquid gashttp://www.landfood.ubc.ca/courses/fnh/301/water/motion.gif


This leaves us with changes in P.E. during chemical reactions.

P.E. contributions can from electrostatic forces between:Separate Vibrating Atoms Nucleus & Electrons in AtomsProtons & Neutrons in NucleusNuclei and Shared Electron Pair in Each Bond = Bond Energy

Where does the heat that is released come from?

The energy released or absorbed during a chemical change is due to thedifferences between the reactant bond energies and the product bond energies.

B.E.reactants - B.E.products = Heat


∆Hºrxn = ∆Hºreactant bonds broken + ∆Hºproduct bonds formed

Heat of reaction, ∆Hºrxn

(∆Hºtotal = ∆Hºrxn1 + ∆Hºrxn2 +……+ ∆Hºlattice)

Lattice Energy, ∆Hºlattice Born-Haber cycle

Analogous to ionic compound formation:

Exothermic reaction: - ∆Hºrxn

Endothermic reaction: + ∆Hºrxn

∆Hºrxn = ∆BEreactant bonds broken – ∆BEproduct bonds formed


Example: H2 + F2 2 HF

Weaker BondsLess Stable, More Reactive

H2 and F2

Stronger BondMore Stable, Less Reactive

HF


Another way to looks at this reaction: H2 + F2 2 HF

∆Hºrxn = ∆Hºreactant bonds broken + ∆Hºproduct bonds formed

Heat of reaction, ∆Hºrxn

H2 + F2

2 H + 2 F

HF


Use bond energies to calculate ∆Hºrxn (Table 9.2)

H2 + F2 2 HF 9.39, 9.47, 9.49

Optional Homework Problems: 9.38, 9.46, 9.48, 9.50


Application: Energy Released From Combustion of Fuel

∆Hºrxn = ∆BEreactant bonds broken – ∆BEproduct bonds formed

Energy Released = B.E.(fuel + O2) – B.E.(CO2 + H2O)

Fuels with more weak bonds yield more energy than fuels with fewer weak bonds.

Food fuels the body:

Fats:MoreC-HC-C

Carbs:MoreO-HC-O

Section 9.5: Between the Extremes

Scientific models are idealized descriptions of reality.

Electronegativity – the relative ability of a bonded atom to attract the shared e-






Electronegativity – inversely related to atomic size (radius) WHY?


atomic size E.N.


Nonmetals are more electronegative than metals.


Electronegativity and Oxidation Number (O.N.) (Review of O.N.: Section 4.5)

Oxidation-reduction (redox) reactions: The net movement of electrons from onereactant to the other.

Oxidation – the loss of e- (LEO), Reduction – the gain of e- (GER) “LEO the lions says GER!”

Oxidizing agent – becomes reduced; Reducing agent – becomes oxidized

Which element is oxidized? Reduced? Which is the oxidizing agent? Reducing agent?

Oxidation Number and Electronegativity

When dead organisms (such as plankton) fall to the bottom of the sea, theirdead bodies are eaten (respiration) by bacteria living in the ocean sediments:

CH2O + O2 CO2 + H2O

What might be a problem for bacteria trying to eat CH2O deep in sediments?

In addition to O2: SO42- and NO3

2- are present in the sediments.

Which might they use?


Electronegativity and Oxidation Number (O.N.)

E.N. is used to determine an atom’s O.N. in a given bond.

(1)The more E.N. atom in a bond is assigned ALL the SHARED e-; The lessE.N. atoms is assigned NONE

Example: HCl Cl: 8 H: 0

(2) O.N. = # valence e- - # shared e-

Example: O.N.Cl = 7 – 8 = -1 O.N.H = 1 – 0 = +1


Polar Covalent Bonds






This bond type is indicated by:

(1)polar arrow ( ) pointing toward negative pole H–F

(2)delta symbol ()

HH

O

-

++







Polar Covalent vs. Nonpolar Covalent


Partial Ionic Character – related directly to the electronegativity difference (∆EN)

Why?

A greater ∆EN results in larger partial charges () and a higher partial ionic character.

Example: HCl, LiCl, Cl2

Arrange these compounds in order of least to most partial ionic character.


Two approaches for getting a sense of a compound’s ionic character:

#1: Arbitrary cutoffs used in bonding continuum.







Two approaches for getting a sense of a compound’s ionic character:

#2: Calculate the percent ionic character (increases with ∆EN)

Compare actual behavior of a polar molecule in an electric field with thebehavior it would show if the e- were completely transferred (pure ionic).

50 % is dividing line.

Notice: Cl2 is 0% ionic, but no molecule has 100 % ionic character (e- sharingoccurs to some extent in every bond.


Notice, now: Why metal that bond with nonmetals form ionic bonds. Why nonmetals that bond with other nonmetals form covalent bonds.


Properties of substances are indicative of their ionic or covalent character.

Section 9.6: Metallic Bonding (More in Chap 12)

Electron Sea Model

In reactions with nonmetals, metals (Na) transfer their outer e- to form ionic solids (NaCl).

What holds together bonded metals (Na)? All metal atoms contribute their valence e-,which are shared among all the atoms in a sample.


Metallic Bonding - Delocalized Covalent Bonding, Ionic Bonding - Localized

Electrons fit neatly into shells.

Alloys - more than one metal element involved in a metallic “sea”

Properties of metal substances are explained by the electron sea model.

Most metals are solids.

High m.p. = attractions b/w cations and anions need not be broken

Much higher b.p. = attractions b/w cations and anions broken

m.p. depends on # of valence e-:

Section 9.6: Metallic Bonding (More in Chap 12)

Problems for today

9.62, 9.64, 9.66

What would you expect the B.E. of a H–F bond to be given that:H–H = 432 kJ/molF–F = 159 kJ/mol ?

chapter 9 chemical bonding. section 9.1: why does bonding occur in the first place? bonding lowers...

Documents

kinetic energy

chemical bonding slide

heat light energy slide

type of potential energy

ionic bonding

atom slide

valence electrons e

covalent bonds e s