Bonding Bonding \u0026 \u0026 Properties Properties

1

Bonding Bonding & &

PropertiesProperties

TOPIC 2

Topic Contents(1) Fundamental Concepts(2) The Periodic Table(3) Bonding Forces and Energies(4) Primary Interatomic Bonds(5) Secondary Bonding or Van der Waals Bonding

2

At the end of the lecture, students will be able:1. To name the two atomic models cited, and note the

differences between them.2. To describe the important quantum-mechanical principle

that relates to electron energies.3. To name the four electron quantum numbers and for a

specific electron, note what each of its quantum numbers determines.

4. To write a definition of the Pauli exclusion principle.5. To cite the general characteristics of the elements that are

arrayed in each column of the periodic table.6. To write the equation that relates energy and force.

TOPIC OUTCOMES

7. To schematically plot attractive, repulsive, and net energies versus interatomic separation for two atoms or ions.

8. To note on this plot the equilibrium separation and the bonding energy.

9. To briefly describe ionic, covalent, metallic, hydrogen, and van der Waal's bonds and note what materials exhibit each of these bonding types.

10. To give the chemical formula for a material, be able to cite what bonding type(s) is (are) possible.

11. To give the electronegativities of two elements, compute the percent ionic character of the bond that forms between them.

TOPIC OUTCOMES

3

The structure of solid is our main interest.Solid is one of the 4 states of matter (solid, liquid, gas and plasma) that has been discovered so far.Technically we are studying solid state matter.

(1) FUNDAMENTAL CONCEPT

ISSUES TO ADDRESS...

• What promotes bonding?

• What types of bonds are there?

• What properties are inferred from bonding?

Some Terminologies

• Nucleus (Proton + Neutron):• Z (atomic number) = # protons

[1 for hydrogen to 94 for plutonium]

• N = # neutrons

Atomic mass (A) ≈ Z + N

orbital electrons: n = principal quantum number

n=3 2 1

ATOM

4

The mass and charge of Proton, Neutron, and Electron

-1.602 x 10-199.109 x 10-28Electron

0 1.675 x 10-24Neutron

+1.602 x 10-191.673 x 10-24Proton

Charge (C)Mass (g)

F 199

Example:

Determine the number of Proton, electron and neutron in a fluorine atom.

X AZ

Atomic Mass

Atomic Number

A = p + n = 19

Z = p = e = 9

n = A – Z = 19 – 9 = 10

• Proton = 9

• Electron = 9

• Neutron = 10

Answer:

5

Example: In nature

0.7%

99.3 %

Nuclear weapon

90% ↑U 23592

U 23892

Atomic mass (A) ≈ Z + N Variable Isotope

Isotope

Atomic Weight:

The weighted average of atomic masses of an atom’s naturally occurring isotopes.

Atomic Mass Unit (amu):

A measure of atomic mass; 1/12 of the mass of an atom C12

Mole:

1 mole of a substance = 6.023 x 1023

(Avogardo’s number) atoms or molecules.

1 amu/atom (or molecule) = 1 g/mole1 amu/atom (or molecule) = 1 g/mole

6

TWO ATOMIC MODELS

These models enable us to understand the behaviour of electrons in atoms.

Two Atomic Models:

Bohr Atomic Model

Wave-Mechanical Model

Bohr Atomic Model

Assumptions:• Electrons are revolved around the atomic nucleus in

discrete orbitals. • Position of any particular electron is more or less well

defined orbital.

To describe electrons in atoms, both in position(electron orbital) and in energy (quantized energy levels) – simplified & old model

7

Energies of electrons are quantized (electrons are permitted to have only specific values of energies).

Quantum-mechanical principle.

Change energy by quantum jump to an allowed:

Higher energy (with absorption of energy)

Lower energy (with emission of energy)

Allowed electron energies = energy levels or states.

Electrons tend to occupy lowest available energy state.

Wave-Mechanical Model

Electron:Wave-like characteristicsParticle-like characteristicsNot a particle moving in discrete orbitalElectron cloud (position is described by a PROBABILITY DISTRIBUTION).

8

Wave-mechanical model

Bohr model

Quantum NumbersA set of 4 numbers, the values of which are used to label possible electron states.3 of the quantum numbers are integers (specify the size, shape, and spatial orientation of an electron’s probability density) + 1 quantum number of spin orientation.

n (Principle quantum number) size of electron shellℓ (2nd quantum number) shape of electron subshellmℓ (3rd quantum number) number of energy states of each subshell (magnetic field split of subshellstates)ms (spin moment, 4th quantum number) = spin orientation (±1/2).

9

Number of available electron states in some of the electron shells and subshells

32

261014

1357

spdf

N4

182610

135

spd

M3

826

13

sp

L2

221sK1

Per ShellPer Subshell

Number of ElectronsNumber of

StatesSubshellsShell Designation

Principle Quantum

Number, n

Allowed Values for the Quantum Numbers of Electrons

2+1/2-1/2

Spin quantum number

ms

2ℓ + 1integral values from -ℓ to +ℓ, including 0

Magnetic quantum number

mℓ

N allowed values of ℓ

ℓ = 0, 1, 2, 32nd quantum number or subsidiary quantum number

ℓ

All positive integersn = 1, 2, 3, …Principle quantum number

n

10

Incr

easi

ng e

nerg

yFree electron

Electrons tend to occupy lowest available energy state.

Electron Energy StatesBohr Model Wave-mechanical Model

11

• have complete s and p subshells• tend to be unreactive.

Stable electron configurations...

Z Element Configuration

2 He 1s2

10 Ne 1s22s22p6

18 Ar 1s22s22p63s23p6

36 Kr 1s22s22p63s23p63d104s24p6

Adapted from Table 2.2, Callister 6e.

STABLE ELECTRON CONFIGURATIONPauli exclusion principle:

Each electron state can hold no more than 2 electrons, which must have opposite spins.

• unfilled outer shell• gaining or losing electrons to form charged ions• sharing electrons with others atoms.

• The order by which electron fill up the orbitalsis as follow1s22s22p63s23p64s23d104p65s24d105p66s24f14

5d106p6…….

12

• The order for writing the orbitals for electron configurations (for this course) will be by increasing principal quantum number, as

1s22s22p63s23p63d104s24p64d104f145s25p65d105f146s2…….

13

Examples:

Write the electron configuration for the Following atoms by using conventional spdf notation

1. Fe atom (Z= 26) and the Fe 2+ and Fe 3+ ions

2. Cr atom (Z=24) and the Cr 2+ and Cr 6+ ions

Solution• Fe 1s22s22p63s23p64s23d6

1s22s22p63s23p63d64s2

• Fe2+ 1s22s22p63s23p63d6

• Fe3+ 1s22s22p63s23p63d5

Note: the outer 4s electrons are lost first since they have the highest energy & easier to remove

14

• Cr 1s22s22p63s23p64s13d5

1s22s22p63s23p63d54s1

• Cr2+ 1s22s22p63s23p63d4

• Cr6+ 1s22s22p63s23p6

SURVEY OF ELECTRON

• Why? Valence (outer) shell usually not filled completely.• Valence electron electrical, physical, chemical properties

• Most elements: Electron configuration not stable.Element Hydrogen Helium Lithium Beryllium Boron Carbon ... Neon Sodium Magnesium Aluminum ... Argon ... Krypton

Atomic # 1 2 3 4 5 6

10 11 12 13

18 ... 36

Electron configuration 1s1 1s2 (stable) 1s22s1 1s22s2 1s22s22p1 1s22s22p2 ... 1s22s22p6 (stable) 1s22s22p63s1 1s22s22p63s2 1s22s22p63s23p1 ... 1s22s22p63s23p6 (stable) ... 1s22s22p63s23p63d104s246 (stable)

Adapted from Table 2.2, Callister 6e.

15

(2) PERIODIC TABLE• Columns: Similar Valence Structure

Electropositive elements:Readily give up electronsto become + ions.

Electronegative elements:Readily acquire electronsto become - ions.

He

Ne

Ar

Kr

Xe

Rn

ine

rt g

ase

s a

cc

ep

t 1

e a

cc

ep

t 2

e

giv

e u

p 1

e

giv

e u

p 2

e

giv

e u

p 3

e

F Li Be

Metal

Nonmetal

Intermediate

H

Na Cl

Br

I

At

O

S Mg

Ca

Sr

Ba

Ra

K

Rb

Cs

Fr

Sc

Y

Se

Te

Po

Adapted from Fig. 2.6, Callister 6e.

ELECTRONEGATIVE• Ranges from 0.7 to 4.0,

Smaller electronegativity Larger electronegativity

He -

Ne -

Ar -

Kr -

Xe -

Rn -

F 4.0

Cl 3.0

Br 2.8

I 2.5

At 2.2

Li 1.0

Na 0.9

K 0.8

Rb 0.8

Cs 0.7

Fr 0.7

H 2.1

Be 1.5

Mg 1.2

Ca 1.0

Sr 1.0

Ba 0.9

Ra 0.9

Ti 1.5

Cr 1.6

Fe 1.8

Ni 1.8

Zn 1.8

As 2.0

• Large values: tendency to acquire electrons.

Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by CornellUniversity.

16

The nature of various states of matter (except plasma) can be explained by using atomic forces (bonding forces) and potential energy of interatomic distance.

There are 2 forces (attractive and repulsive) that act on a collection of atoms, depending on the relative distances between various atoms.

Attractive forces acting on atoms pull them together (would they collide?).

Repulsive forces (short range force) acting between nuclei and electrons of individual atoms.

(3) Bonding Forces & Energies

No interactionInfinite separation

• Atoms approach• Each exerts forces

(attractive and repulsive) on the other

• Outer electron shells of the two atoms begin to overlap

• Repulsive force is significant

(1)

(2)

(3)

• Bond length, r, = interatomic distance

FF

r

FN = FA + FRNet Force Attractive

Force

Repulsive Force

Forces between 2 atoms

2 isolated atoms

17

r0

r

R1 R2

R1= cationR2 = anionr0 = R1 + R2

( )( )

120

221

1

20

221

20

21

4

44

+

+

−−=+=

−=

−=−=

nRAnet

nR

A

rnb

reZZFFF

rnbF

reZZ

reZeZF

πε

πεπε

Z1, Z2 = number of electrons removed or added from the atoms during ion formatione = electron charger interatomic separation distanceB and n = constantsε0 = permittiivity of free space = 8.85 x 10-12 C2/(Nm2)

18

r

r0

, r

nnet rb

reZZE

FdrE

++=

= ∫

0

221

4πε

19

attractive force, FA

repulsive force, FR

Interatomic distance, rForc

e, F Attr

activ

eR

epul

sion

0

+

-

Repulsive energy, ER

Attractive energy, EA

Interatomic distance, rPote

ntia

l Ene

rgy,

E

Rep

ulsi

onA

ttrac

tive

0

+

-

Net force, FN

Net energy, EN

r0

E0

∫= FdrE

EN = ER + EA

FA + FR = 0

Force-Potential energy relationship for 2 atoms:

From these relationships, 2 parameters can be perceived that will enable us to identify certain properties of solids.

(1) Equilibrium interatomic distance, r0 (or lattice constant, a0)

(2) Bonding energy, E0 [or Binding energy, Eb].

20

(1) Equilibrium interatomic distance, r0 (or lattice constant, a0)

Equilibrium distance at which these forces exactly balance each other.At this distance, attractive force (Coulombicforce) is exactly counterbalanced by repulsive force (short range force) between the two nuclei and between the two electron distributions.These forces tend to make the atoms move back to equilibrium position (stable position).Minimum potential energy.

(2) Bonding energy, E0 [Binding energy, Eb].Energy per atom required to dissociate from the solid.A measure of the strength of the solid (in eV).This strength varies, depending on the nature of bonding.a ↑ E ↓ liquida ↑ ↑ E → ∞ gasTypes of bonding strength

1 eV = energy gained by accelerating 1 electron through 1 volts1 eV = 1.6x10-19 Joule.

1 eV = energy gained by accelerating 1 electron through 1 volts1 eV = 1.6x10-19 Joule.

21

PROPERTIES FROM BONDING: TM

• Melting Temperature, Tm

r

larger Tm

smaller Tm

Energy (r)

ro

Tm is larger if Eo is larger.

PROPERTIES FROM BONDING: E• Elastic modulus, E

• E ~ curvature at ro

cross sectional area Ao

ΔL

length, Lo

F

undeformed

deformed

r

larger Elastic Modulus

smaller Elastic Modulus

Energy

ro unstretched length

ΔL F Ao

= E Lo

Elastic modulus

22

PROPERTIES FROM BONDING: α• Coefficient of thermal expansion, α

• α ~ symmetry at ro

α is larger if Eo is smaller.

ΔL

length, Lo

unheated, T1

heated, T2

�

= α (T2-T1) ΔL Lo

coeff. thermal expansion

r

smaller α

larger α

Energy

ro

InteratomicInteratomicBondingBonding

Primary Bonding Secondary Bonding

• Ionic bonds• Covalent bonds• Metallic bonds

• Van der Waals bonds• Hydrogen bond

23

3 different types of primary or chemical bond are found in solids.

Ionic, covalent, and metallic.Involve valence electronsNature of bond depends on electron structures of the constituent atoms. Tendency of atoms to assume stable electron structure.

Secondary (or physical) forces and energies also found in many solid materials.

Weaker than primary onesInfluence physical properties of some material

Primary BondingPrimary Bonding> Ionic Bonding> Ionic Bonding> Covalent Bonding> Covalent Bonding> Metallic Bonding> Metallic Bonding

(4) Primary Interatomic Bonds

24

Formed between highly electropositive (metallic) elements and highly electronegative (nonmetallic) elements large difference in electronegativity.

Ionization: electrons are transferred from atoms of electropositive elements to atoms of electronegative elements, producing positively charged cations and negatively charge anions.

Ionic bonding: due to electrostatic or Coulombic force attraction of oppositely charged ions.

Nondirectional bonding magnitude of the bond is equal in all directions around an ion.

Binding energy large high melting temp.

Ionic material hard, brittle, electrically and thermally insulative.

Ionic Bonding

3s1 3p6

Sodium atom, NaAtomic radius = 0.192 nm

Chlorine atom, ClAtomic radius = 0.099 nm

+ -

Sodium ion, Na+

Ionic radius = 0.095 nm Chlorine ion, Cl-ionic radius = 0.181 nm

Coulombicbonding force

Unstable Unstable

Stable Stable

25

• Predominant bonding in Ceramics

Give up electrons Acquire electrons

He -

Ne -

Ar -

Kr -

Xe -

Rn -

F 4.0

Cl 3.0

Br 2.8

I 2.5

At 2.2

Li 1.0

Na 0.9

K 0.8

Rb 0.8

Cs 0.7

Fr 0.7

H 2.1

Be 1.5

Mg 1.2

Ca 1.0

Sr 1.0

Ba 0.9

Ra 0.9

Ti 1.5

Cr 1.6

Fe 1.8

Ni 1.8

Zn 1.8

As 2.0

CsCl

MgO

CaF2

NaCl

O 3.5

Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by CornellUniversity.

EXAMPLES: IONIC BONDING

26

Attractive energy between ion pair , EA

r

eZZ

r

eZeZAE

04

221

04

)2)(1(

πεπε−=−=

Z1, Z2 = number of electrons removed or added from the atoms during ion formatione = electron charger interatomic separation distanceε0 = permittiivity of free space = 8.85 x 10-12 C2/(Nm2)

Coulomb’s Law

UN

IVE

RS

ITI

SA

INS

MA

LAY

SIA Repulsive energy between ion pair , ER

nr

BRE =

r interatomic separation distancen and B = constants (n usually ranges from 7 to 9)

27

UN

IVE

RS

ITI

SA

INS

MA

LAY

SIA

Calculate the coulombic attractive energy between Na+ and Cl- ions.

At equilibrium, r → r0

rNa+ = 0.095 nm; rCl- = 0.181 nmr0 = rNa+ + rCl- = 2.76x10-10 m

JEmNmC

CE

r

eZZE

A

A

A

19

102212

219

1034.8)1076.2)](/(1085.8[4

)1060.1)(1)(1(04

221

−

−−

−

×+=

×××−+

−=

−=

π

πε

Repulsive energy = -8.34 x 10-19 J

Q & A

Bonding Energies and Melting Temperatures for Various Substances

8012800

3.35.2

6401000

NaClMgO

Ionic

1410>3550

4.77.4

450713

SiC (diamond)

Covelent

-39660

15383410

0.73.44.28.8

68324406849

HgAlFeW

Metallic

-780

0.360.52

3551

NH3

H2OHydrogen

-189-101

0.080.32

7.731

ArCl2

Van derWaals

eV/atom, ion, moleculekJ/mol

Melting Temperature

(oC)

Bonding Energy

SubstanceBonding Type

28

In covalent bonding stable electron configurations are assumed by sharing of electrons between adjacent atoms.

Two atoms that are covalently bonded will each contribute at least one electron to the bond, and the shared electrons may be considered to belong to both atoms.

H• + H• → H:H (1s1 electron from hydrogen atom)

Covalent Bonding

shared electrons from carbon atom

shared electrons from hydrogen atoms

H

H

H

H

C

CH4

• Example: CH4

C: has 4 valence e,needs 4 more

H: has 1 valence e,needs 1 more

Electronegativitiesare comparable.

Many nonmetallic elemental molecules (H2, Cl2, F2, etc)

Molecules containing dissimilar atoms (CH4, H2O, HNO3, HF, etc)

Other elemental solids:diamond (carbon), silicon, germanium

Other solid compounds composed of elements that are located on the RHS of the periodic table (GaAs, indium antimonide, InSb, SiC)

29

• Molecules with nonmetals• Molecules with metals and nonmetals• Elemental solids (RHS of Periodic Table)• Compound solids (about column IVA)

He -

Ne -

Ar -

Kr -

Xe -

Rn -

F 4.0

Cl 3.0

Br 2.8

I 2.5

At 2.2

Li 1.0

Na 0.9

K 0.8

Rb 0.8

Cs 0.7

Fr 0.7

H 2.1

Be 1.5

Mg 1.2

Ca 1.0

Sr 1.0

Ba 0.9

Ra 0.9

Ti 1.5

Cr 1.6

Fe 1.8

Ni 1.8

Zn 1.8

As 2.0

SiC

C(diamond)

H2O

C 2.5

H2

Cl2

F2

Si 1.8

Ga 1.6

GaAs

Ge 1.8

O 2.0

co

lum

n IV

A

Sn 1.8Pb 1.8

Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 isadapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.

EXAMPLES: COVALENT BONDING

Number of covalent bonds in an atom is determined by number of valence electron.

N’ valence electron an atom can covalently bond with at most 8-N’ other atoms

N’ = 7 (chlorine), and 8-N’ = 1 (one Cl atom can bond to only one other atom, as in Cl2)

Highly directional type of bonding.

Binding energy & melting temp for covalently bonded materials very high (diamond) to very weak (bismuth, polymeric material)

Very few compounds exhibit pure covalent bonding (or ionic bonding).

Possible of having interatomic bonds (partially ionic and partially covalent).

30

Degree of ionic (covalent) relative positions of constituent atoms in periodic table or the difference in their electronegativities.

% ionic character = {1-exp[-(0.25)(XA – XB)2]} x 100

XA and XB = electronegativities for respective elements.

Very similar to covalent bondingValence electrons

are shared (1, 2, or 3 from each atom) by a great many atoms.are essentially free electrons and move (drift) through out the metal not directional high conductivityare not restricted to strong bond positionsea of electron or electron cloud

Remaining nonvalence electrons and atomic nuclei ion coresGroup IA and IIA elementsAll elemental metalsHighly conductiveDuctile, binding energy & melting temp (wide range)

Metallic Bonding

31

+ + + +

+ + + +

+ + + +

+ + + +

Ion cores

Sea of valence electrons

Schematic illustration of metallic bonding

Van der Waals bond or physical bond

Binding energy (typical) in the order of 10 kJ/mol (0.1 eV/atom)

Exist between virtually all atoms or molecules.

The presence of any of the 3 primary bonding types may obscure it.

This force arises from atomic or molecular dipoles electric dipoles.

Electric dipoles

separation of positive and negative portions of an atom or molecule

coulombic attraction between +ve end of one dipole and –ve end dipole.

Schematic illustration of van der Waals bonding between two dipolesSchematic illustration of van der Waals bonding between two dipoles

Atomic or molecular dipoles

+ +

(5) Secondary BondsVan der Waals Bonds

32

Dipole interactions occur

Dipole interactions occur

Fluctuating Induced Dipole BondsFluctuating Induced Dipole Bonds

Polar Molecule-Induced Dipole BondsPolar Molecule-Induced Dipole Bonds

Permanent Dipole BondsPermanent Dipole Bonds

an electrically symmetric atom

Schematic representations of:Schematic representations of:

an induced atomic dipole

≡

Atomic nucleusElectron cloud

+

Atom/molecule (electrically symmetric) induced dipole.

overall spatial distribution of

electrons is symmetric with

respect to the positively

charged nucleus.

Atoms experiences constant vibration motion

instantaneous or short-live distortions of the symmetric

One of these dipoles can in turn produce a displacement of electron distribution of an adjacent molecule or atom, which induces the second one also to become a dipole.

Weak force

(a) Fluctuating Induce Dipole Bonding

33

It exists in some molecules due to asymmetrical arrangement of positively and negatively charged regions

POLAR MOLECULES.

Schematic representation of a polar hydrogen chloride (HCl) molecule.

H Cl +

Polar moleculePolar molecule Non-polar moleculeNon-polar moleculeInduce dipole

Magnitude of strength > fluctuating induced dipolesMagnitude of strength > fluctuating induced dipoles

(b) Polar Molecule-Induced Dipole Bonding

UN

IVE

RS

ITI

SA

INS

MA

LAY

SIA

Van der Waals forces exist between adjacent polar molecules

Binding energy > other induced dipole bondings

Hydrogen is covalently bonded to Fluorine (as in HF)

Hydrogen end of the bond is a positively charged bare proton that is unscreened by any electrons.

This highly positive charged end of molecule is capable of a strong attractive force with the negative end of an adjacent molecule.

Hydrogen bond

Strongest secondary bonding

H F H F

Hydrogen bond

(c) Permanent Dipole Bonding

secondary bonding

polymer

34

Summary

Electrons in Atoms2 models

Electron energy statesQuantum numbersElectron configuration of an atom & Pauli exclusion principle

Periodic Table of elements & arrangement of elements according to valence electron configurationElectronegative & electropositive

Summary

Type

Ionic

Covalent

Metallic

Secondary

Bond Energy

Large!

Variablelarge-Diamondsmall-Bismuth

Variablelarge-Tungstensmall-Mercury

smallest

Comments

Nondirectional (ceramics)

Directionalsemiconductors, ceramics

polymer chains)

Nondirectional (metals)

Directionalinter-chain (polymer)

inter-molecular

35

Summary

Ceramics(Ionic & covalent bonding):

Metals(Metallic bonding):

Polymers(Covalent & Secondary):

secondary bonding

Large bond energylarge Tm

large Esmall α

Variable bond energymoderate Tm

moderate Emoderate α

Directional PropertiesSecondary bonding dominates

small Tsmall Elarge α