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Chapter 9 Chemical Bonding. 1 type of Potential Energy : Gravitational P.E. Ball On Top of a Hill. m. P.E. = mgh. h. 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?.

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slide2

1 type of Potential Energy: Gravitational P.E.

Ball On Top of a Hill

m

P.E. = mgh

h

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?

Energy changes forms:

P.E.  Kinetic Energy (K.E.)

slide3

Energy changes forms

Mechanical

Energy

Friction

Motor

Generator

Engines

Electrical

Energy

Heat (Thermal)

Energy

Solar

Heater

Battery

FIre

Battery

Charger

Light (Radiant)

Energy

Chemiluminescence

Chemical

Energy

Photosynthesis

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

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

slide4

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

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

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

Mechanical

Energy

Electrical

Energy

Heat (Thermal)

Energy

Light (Radiant)

Energy

Chemical

Energy

slide5

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

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

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

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

slide6

Section 9.1: Three Type of Bonds

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

slide7

Concept Check

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-:

slide8

Representative Elements

For REPRESENTATIVE elements:

• period (row) = shell # (n = 1, 2, 3, 4….n)

• group (column) = # of e- in outer shell

Group ## of valence e-

Transition Elements

IA 1

IIA 2

IIIA 3

IVA 4

VA 5

VIA 6

Shells

of an

atom

VIIA 7

VIIA 8

Section 9.1: Two Bond Types With Localized Electrons

Ionic & Covalent Bonding

slide9

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

Nonpolar Covalent Bond

Polar Covalent Bond

Ionic Bond

Electrons are shared unequally.

Electrons are transferred.

slide10

Extent of electron sharing in Covalent Bonds

e-’s shared between atoms

of the same element:

Equal Sharing

e-’s shared between atoms

of 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-?

slide11

Electronegativity

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

Increasing E.N.

Decreasing E.N.

Rule for Bond Formation

The 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

slide12

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

e- sharing

2 nonmetals

e- transfer

metal + nonmetal

“Bonding Continuum”

Covalent Bonds

Ionic Bonds

Polar Colvalent

Nonpolar Colvalent

0.4 < E.N.

0.4 < E.N. < 1.7

E.N. difference > 1.7

1.7

Answer: Electronegativity Differences

Example:

Oxygen (O) bonds with

Magnesium (Mg): MgO

E.N. of O = 3.5

E.N. of Mg = 1.2

E.N. difference = 2.3

slide13

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

Metallic Bonding

Metallic Bonding - Delocalized

Covalent Bonding, Ionic Bonding

- Delocalized

A messy “sea” of electrons

Electrons fit neatly into shells.

slide14

Outer e-

Inner e-

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

Metallic Bonding

Metallic Bonding - Delocalized

A messy “sea” of electrons

slide15

Li

Different elements can have the same number of dots

Be

Mg

Same Group

(Column)

Lewis Electron-Dot Symbols

Two parts:

(1) Element symbol – nucleus + inner electrons

Ex: The element lithium has an element symbol Li

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

slide16

Mg2+

Cl-

Review: Ions

Ion – charged particles that form when an atom gains or loses one or more electrons

(Ch2, p60)

Element

Ion

Ion Type

Mg

Cation

Cl

Anion

slide17

Mg

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

Example:

slide18

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)

slide19

O

Examples: Water (H2O) Carbon Dioxide (CO2)

H

H

O

O

C

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.

Solid Ionic compound

Na – metal

Cl - nonmetal

Sodium chloride (NaCl)

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

slide20

Cl-

Section 9.2: Ionic Bonding

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

gained by the nonmetal atom.

Na+

lost

gained

slide21

Behavior of Ionic Compounds

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

slide22

Lattice Energy

(∆Hºlattice)

slide23

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

slide24

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.

slide27

Behavior of Ionic Compounds

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

slide28

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)

slide30

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

Na – metal

Cl - nonmetal

Sodium chloride (NaCl)

Section 9.3: Covalent Bonding

e- sharing – primary way that atoms interact

Nonmetal + Nonmetal

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

O

O

O

H

H

H

C

C

C

C

H

H

H

H

H

H

H

slide31

Section 9.3: Covalent Bonding

Why do covalent bonds form?

Lower P.E. = More stable

slide32

Section 9.3: Covalent Bonding

How are the electrons distributed?

Electron

Density

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:

• Bonding Pairs & Lone Pairs

• Bond Type – double, single, triple

slide33

Bond Energy – energy needed

to overcome attraction and

break the bond

Section 9.3: Covalent Bonding

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

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

shared electrons

slide34

Section 9.3: Covalent Bonding

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

slide35

Section 9.3: Covalent Bonding

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

Less E needed

to break.

Lower B.E.

Weaker Bonds =

Higher Energy

“Shallow Energy Well”

Stronger Bonds =

Lower Energy

“Deeper Energy Well”

More E needed

to break.

Higher B.E.

slide36

Section 9.3: Covalent Bonding

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.

slide37

Section 9.3: Covalent Bonding

Bond Energy (B.E.) and Bond Length

This relationship

holds, in general,

ONLY for single

bonds.

slide38

Section 9.3: Covalent Bonding

Bond Type (Single, Double, Triple) also matters

Same two

elements,

different B.E.

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

Higher bond order = Shorter bond length = Higher Bond Energy

slide39

Section 9.3: Covalent Bonding

Periodic Table Trends Without Detailed Bond Lengths

The closer the

atoms, the

stronger the bond.

Bond Energy:

C—F > C—Cl > C—Br

slide40

O

solid  liquid  gas

H

H

Strong covalent bonding forces

Hold atoms together

Weak intermolecular forces

Hold molecules together

(More in Chapter 12)

-

+

+

Chemical Reaction

Phase Change

Section 9.3: Covalent Bonding

Covalent Bonds are stronger than Ionic Bonds

So why, then, do covalent compounds have lower melting points

than ionic compounds?

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

slide41

Section 9.4: Bond Energy and Chemical Change

Where does the heat that is released come from?

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

slide42

solid  liquid  gas

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).

Section 9.4: Bond Energy and Chemical Change

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

Example of a chemical system

A container filled with molecules.

http://www.landfood.ubc.ca/courses/fnh/301/water/motion.gif

slide43

= Bond Energy

Where does the heat that is released come from?

The energy released or absorbed during a chemical change is due to the

differences between the reactant bond energies and the product bond energies.

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

Section 9.4: Bond Energy and Chemical Change

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 Atoms

Protons & Neutrons in Nucleus

Nuclei and Shared Electron Pair in Each Bond

slide44

Analogous to ionic compound formation:

Lattice Energy, ∆Hºlattice Born-Haber cycle

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

Section 9.4: Bond Energy and Chemical Change

Heat of reaction, ∆Hºrxn

Exothermic reaction: - ∆Hºrxn

Endothermic reaction: + ∆Hºrxn

∆Hºrxn = ∆Hºreactantbonds broken + ∆Hºproductbonds formed

∆Hºrxn = ∆BEreactantbonds broken – ∆BEproductbonds formed

slide45

Section 9.4: Bond Energy and Chemical Change

Example: H2 + F2 2 HF

Weaker Bonds

Less Stable, More Reactive

H2 and F2

Stronger Bond

More Stable, Less Reactive

HF

slide46

Section 9.4: Bond Energy and Chemical Change

Another way to looks at this reaction:

H2 + F2 2 HF

Heat of reaction, ∆Hºrxn

2 H + 2 F

H2 + F2

HF

∆Hºrxn = ∆Hºreactantbonds broken + ∆Hºproductbonds formed

slide47

Section 9.4: Bond Energy and Chemical Change

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

slide48

∆Hºrxn = ∆BEreactantbonds broken – ∆BEproductbonds 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.

Fats:

More

C-H

C-C

Carbs:

More

O-H

C-O

Food fuels the body:

Section 9.4: Bond Energy and Chemical Change

Application: Energy Released From Combustion of Fuel

slide49

e- sharing

2 nonmetals

e- transfer

metal + nonmetal

“Bonding Continuum”

Covalent Bonds

Ionic Bonds

Polar Colvalent

Nonpolar Colvalent

0.4 < E.N.

0.4 < E.N. < 1.7

E.N. difference > 1.7

1.7

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-

slide50

atomic size

E.N.

Section 9.5: Between the Extremes

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

slide51

Section 9.5: Between the Extremes

Nonmetals are more electronegative than metals.

slide52

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

Section 9.5: Between the Extremes

Electronegativity and Oxidation Number (O.N.)

(Review of O.N.: Section 4.5)

Oxidation-reduction (redox) reactions: The net movement of electrons from one

reactant 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

slide53

Oxidation Number and Electronegativity

When dead organisms (such as plankton) fall to the bottom of the sea, their

dead 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 NO32- are present in the sediments.

Which might they use?

slide54

Section 9.5: Between the Extremes

  • Electronegativity and Oxidation Number (O.N.)
  • E.N. is used to determine an atom’s O.N. in a given bond.
  • The more E.N. atom in a bond is assigned ALL the SHARED e-; The less
  • E.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
slide55

Section 9.5: Between the Extremes

e- sharing

2 nonmetals

e- transfer

metal + nonmetal

O

H

H

“Bonding Continuum”

-

Covalent Bonds

Ionic Bonds

+

+

Polar Colvalent

Nonpolar Colvalent

0.4 < E.N.

0.4 < E.N. < 1.7

E.N. difference > 1.7

1.7

Polar Covalent Bonds

  • This bond type is indicated by:
  • polar arrow ( ) pointing toward negative pole H–F
  • delta symbol ()
slide56

Section 9.5: Between the Extremes

e- sharing

2 nonmetals

e- transfer

metal + nonmetal

“Bonding Continuum”

Covalent Bonds

Ionic Bonds

Polar Colvalent

Nonpolar Colvalent

0.4 < E.N.

0.4 < E.N. < 1.7

E.N. difference > 1.7

1.7

Polar Covalent vs. Nonpolar Covalent

slide57

Section 9.5: Between the Extremes

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 mostpartial ionic character.

slide58

Section 9.5: Between the Extremes

e- sharing

2 nonmetals

e- transfer

metal + nonmetal

“Bonding Continuum”

Covalent Bonds

Ionic Bonds

Polar Colvalent

Nonpolar Colvalent

0.4 < E.N.

0.4 < E.N. < 1.7

E.N. difference > 1.7

1.7

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

#1: Arbitrary cutoffs used in bonding continuum.

slide59

Section 9.5: Between the Extremes

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 the

behavior 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- sharing

occurs to some extent in every bond.

slide60

Section 9.5: Between the Extremes

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

Why nonmetals that bond with other nonmetals form covalent bonds.

slide61

Section 9.5: Between the Extremes

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

slide62

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

A messy “sea” of electrons

Electrons fit neatly into shells.

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

slide63

Section 9.6: Metallic Bonding (More in Chap 12)

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-:

slide64

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/mol

F–F = 159 kJ/mol

?