atomic structure & bonding lect

7/24/2019 Atomic Structure & Bonding Lect

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2

Chapter 2-

Atoms are made of protons, neutrons andelectrons

me=0.00091094x10-27= 9.1094x10-31kg = 0.511MeV

mp= 1.6726 x 10-27kg = 938.272 MeV

mn= 1.6749 x 10-27kg = 939.566 MeV =

mn= mp+ 1.293 MeV

proton & electron charge 1.6022 x 10-19C Howeverpare +ve and eareve

Atomic number (Z) describes the number ofprotons in the nucleus

Atomic mass (A) of an element isapproximately equal to the number of neutronsand protons the element has

Remember elements have isotopeselements canhave different numbers of neutrons (e.g. 12C, 13C,14C)

Atomic weight is the weighted average of theelement based on the relative amounts of itsisotopes (e.g. 1 mol/carbon = 12.0107 g/mol,NOT 12 g/mol!)

Basic concepts

Chapter 2-

Fundamental Concept

Atomic Weight

Weighted average of the atomic masses of an atom's

naturally occurring isotopes Atomic Mass Unit (amu)

Measure of atomic mass

1/12 the mass of C12atom

Mole

Quantity of a substance corresponding to 6.022X1023atoms

or molecules

1 amu/ atom (or molecule) = 1g/mol


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Chapter 2-

How many grams are there in one amu of a material?

The two major isotopes of carbon:

98.93% of 12C with an atomic weight of 12.00000 amu, and

1.07% of 13C with an atomic weight of 13.00335 amu.

Confirm that the average atomic weight of C is 12.011 amu.

Sum the product of the isotope atomic weight and the percent abundance.

(12 amu)*(.9893)+(13.00335 amu)*(.0107) = 12.011 amu

Examples

Chapter 2-

Electrons In Atoms

Bohr Atomic Model (old view)

Early outgrowth of

quantum mechanics

Electrons revolve aroundnucleus in discrete orbitals

Electrons closer to nucleus

travel faster then outer

orbitals

Principal quantum number

(n); 1stshell, n=1; 2ndshell,

n=2; 3rdshell, n=3


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Chapter 2-

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Quantum NumbersHydrogen atom

Chapter 2-

c02f03

Bohr AtomWave-mechanical atom


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Chapter 2-

Atomic Models

Wave-Mechanical

Model

Electron exhibits bothwave-like and particle-like

characteristics

Position is now considered

to be the probability of an

electron being at various

locations around the

nucleus, forming an

electron cloud

Chapter 2-

Atomic ModelsQuantum numbers

The size, shape, and spatial orientation of an

electrons probability density are specified by

three of these quantum numbers.

Principal quantum number n, represents a

shell

K, L, M, N, O correspond to n=1, 2, 3, 4,

5....

Quantum number l, signifies the subshell

Lowercase italicsletters, p, d, f; related to

the shape of the subshell

Quantum number ml, represents the

number of energy state

s, p, d, f have 1, 3, 5, 7 states respectively

Quantum number ms, is the spin moment

Each electron is a spin moment (either up

or down)

(+1/2) and (-1/2)

Each state can hold no more than 2

electrons which must have opposite spins


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Chapter 2-

Electron Configuration

Electron configuration

represents the manner in

which the states areoccupied

Valence electrons

Occupy the outermost

shell

Available for bonding

Tend to control chemical

properties

Ex. Silicon (Si)

Chapter 2-

Energy


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Chapter 2-

c02tf02

When some elements covalently

bond, they formsphybrid bonds,

e.g., C, Si, Ge

Chapter 2-

Examples

Give the electron configurations for the following:C

1s2 2s2 2p2

Br1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5

Mn+2

1s2 2s2 2p6 3s2 3p6 3d5

F-

1s2 2s2 2p6

Cr

1s2 2s2 2p6 3s2 3p6 4s1 3d5


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Chapter 2 - 15

Electronic Structure Electrons have wave-like and particle-like (old view)

properties.

We can better say that the wave-particle nature is the real

thing; individual wave and particle states are limiting cases;usually observed in measurements (collapse of the wave

function) To better understand electronic structure, we assume

Electrons reside in orbitals.

Each orbital at discrete energy level is determined by

quantum numbers.c

Quantum # Designation

n= principal (energy level-shell) K, L, M, N, O (1, 2, 3, etc.)

l= angular (orbitals) s,p, d, f (0, 1, 2, 3,, n-1)

ml= magnetic 1, 3, 5, 7 (-lto +l)

ms= spin , -Chapter 2 - 16

Electron Configurations

Valence electronsthose in unfilled shells

Filled shells more stable

Valence electrons are most available for

bonding and tend to control the chemicalproperties

example: C (atomic number = 6)

1s2 2s22p2

valence electrons


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Chapter 2 - 17

Electronic Configurationsex: Fe - atomic # = 26

valence

electrons

Adapted from Fig. 2.4,Callister & Rethwisch 3e.

1s

2s2p

K-shell n = 1

L-shell n = 2

3s3p M-shell n = 3

3d

4s

4p4d

Energy

N-shell n = 4

1s2 2s2 2p6 3s23p6 3d6 4s2

Chapter 2-

Periodic Table

Elements classified according to electron configuration

Elements in a given column or group have similar valence electron

structures as well as chemical and physical properties

Group 0inert gases, filled shells and stable

Group VIIA halogen

Group IA and IIA - alkali and alkaline earth metals

Groups IIIB and IIB transition metals

Groups IIIA, IVA and VAcharacteristics between the metals and

nonmetals


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Chapter 2-19

The Periodic Table Columns: Similar ValenceStructure

Adapted fromFig. 2.6,Callister &

Rethwisch 8e.

Electropositive elements:Readily give up electronsto become + ions.

Electronegative elements:Readily acquire electronsto become - ions.

give

up

1e-

gi

ve

up

2e-

give

up

3e-

inertgases

ac

cept1e-

ac

cept2e-

O

Se

Te

Po At

I

Br

He

Ne

Ar

Kr

Xe

Rn

F

ClS

Li Be

H

Na Mg

BaCs

RaFr

CaK Sc

SrRb Y

Chapter 2 - 20

Atomic Bonding

Valence electrons determine all of the

following properties

1) Chemical2) Electrical

3) Thermal

4) Optical

5) Deteriorative

6) etc.


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Chapter 2-

tomic Bonding in Solids

Chapter 2-

Bonding in Solids

Bonding forces and energies

Far apart: atoms dont know about each other

As they approach one another, exert force on one another

Forces are Attractive (FA)slowly changing with distance

Repulsive (FR)typically short-range

Net force is the sum of these

FN= FA+ FR

At some point the net force is zero; at that position a state of

equilibrium exists


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Chapter 2-

Bonding Forces and

Energies

FN= FA+

FR

EN= EA+

ER

When 0 = FA+ FR,

equilibrium exists.

The centers of theatoms will remain

separated by the

equilibrium spacing

ro.

This spacing also

corresponds to the

minimum of the

potential energycurve. The energy

that would be

required to

separate two

atoms to an infiniteseparation is Eo

Figure 2.8

Chapter 2-

Bonding in Solids

Bonding forces and energies We are more accustomed to thinking in terms of potential energy

instead of forces in that case

RAN

r

R

r

AN

EEE

drFdrFE

The point where the forces are zero also corresponds to the minimumpotential energy for the two atoms (i.e. the trough in Figure 2.8), whichmakes sense because dE/dr = F =0 at a minimum.

The interatomic separation at that point (ro) corresponds to the potentialenergy at that minimum (Eo,it is also the bonding energy) The physical interpretation is that it is the energy needed to separate the atoms

infinitely far apart

FdrE

Setting our ZERO ENERGY reference at infinite

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Chapter 2-

Bonding Forces and Energies

A number of material properties depend on Eo,

the curve shape, and bonding type

Material with large Eotypically have higher melting

points

Mechanical stiffness is dependent on the shape of its

force vs. interatomic separation curve

A materials linear coefficient of thermal expansion

is related to the shapeof its Eovs. rocurve

Chapter 2-

ExamplesCalculate the force of attraction between ions X+and an Y-, the

centers of which are separated by a distance of 2.01 nm.

&

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Chapter 2-

Types of primary bonds found in solids

Ionic

Covalent

Metallic

As you might imagine, the type of bonding influences

propertieswhy?

Bonding involves the valence electrons!!!

Primary Interatomic Bonds

Chapter 2 - 28

Occurs between + and - ions.

Requires electron transfer.

Large difference in electronegativity required.

Example: NaCl

Ionic Bonding

Na (metal)unstable

Cl (nonmetal)unstable

electron

+ -CoulombicAttraction

Na (cation)stable

Cl (anion)stable

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Chapter 2 - 29

Ionic bond: metal + nonmetal

donates accepts

electrons electrons

Dissimilar electronegativities

ex: MgO Mg 1s22s22p63s2 O 1s22s22p4

[Ne] 3s2

Mg2+ 1s22s22p6 O2- 1s22s22p6

[Ne] [Ne]

Chapter 2-


Ionic bonding Sodium chloride (NaCl)

Sodium gives up one its electrons to chlorine sodium becomespositively charged, chlorine becomes negatively charged

The attraction energy is electrostaticin nature in ionic solids

(opposite charges attract) The attractive component of the potential energy (for 2 point

charges) is given by

r

eZeZE

o

A

1

4

21

The repulsive term is given by

128~, nr

BE

nR

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Chapter 2-

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Chapter 2-

IONIC BONDING

Ionic bonding is non-directional magnitude of the bond is equal in

all directions around the ion

Many ceramics have an ionic bonding characteristic

Bonding energies typically in the range of 6001500 kJ/mol

Often hard, brittle materials, and generally insulators

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Chapter 2 - 33

Ionic Bonding

Energyminimum energy most stable

Energy balance of attractiveand repulsiveterms

Attractive energy EA

Net energy EN

Repulsive energy ER

Interatomic separation r

r

A

nr

BE

N= E

A+ E

R=

Adapted from Fig. 2.8(b),Callister & Rethwisch 3e.

Chapter 2 - 34

Predominant bonding in Ceramics

Adapted from Fig. 2.7, Callister & Rethwisch 3e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of theChemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.

Examples: Ionic Bonding

Give up electrons Acquire electrons

NaCl

MgO

CaF2CsCl

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Chapter 2-

Covalent bonding

Sharing of electrons between adjacent atoms

Most nonmetallic elements and molecules containing

dissimilar elements have covalent bonds

Polymers!

Bonding is highly directional! : between specific atomsand may exist only in the direction between one atom

and another that participates in electron sharing

Number of covalent bonds possible is guessed by the

number of valence electrons

Typically is 8N, where N is the number of valence

electrons

Carbon has 4 valence es 4 bonds (ok!)

Chapter 2 - 36

C: has 4 valence e-,

needs 4 more

H: has 1 valence e-,

needs 1 more

Electronegativities

are comparable.

Adapted from Fig. 2.10, Callister & Rethwisch 3e.

Covalent Bonding similar electronegativityshare electrons

bonds determined by valences&porbitals

dominate bonding

Example: CH4

shared electronsfrom carbon atom

shared electronsfrom hydrogenatoms

H

H

H

H

C

CH4

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Chapter 2- 11

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

He-

Ne-

Ar-

Kr-

Xe-

Rn-

F4.0

Cl3.0

Br2.8

I2.5

At2.2

Li1.0

Na0.9

K0.8

Rb0.8

Cs0.7

Fr0.7

H2.1

Be1.5

Mg1.2

Ca1.0

Sr1.0

Ba0.9

Ra0.9

Ti1.5

Cr1.6

Fe1.8

Ni1.8

Zn1.8

As2.0

SiC

C(diamond)

H2O

C2.5

H2

Cl2

F2

Si1.8

Ga1.6

GaAs

Ge1.8

O2.0

columnIVA

Sn1.8

Pb1.8

Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 isadapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.

EXAMPLES: COVALENT BONDING

Chapter 2-

Bonding in Solids

Many materials have bonding that is both ionic andcovalent in nature (very few materials actually exhibit pure

ionic or covalent bonding)

Easy (empirical) way to estimate % of ionic bondingcharacter:

XA, XBare the electronegativities of atomsA and Binvolved

100x))(25.0(exp1characterionic% 2BA

XX

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Chapter 2-

Primary Interatomic Bonds Metallic Bonding

Found in metals and their alloys

1 to 3 valence electrons that form asea of electrons or an electroncloud because they are more orless free to drift through the entiremetal

Nonvalence electrons and atomicnuclei form ion cores

Bonding energies range from weakto strong

Good conductor of both electricityand heat

Most metals and their alloys fail ina ductile manner

Ion

Cores

Sea of Valence

Electrons

+

+

+

+

+

+

+

+

+

- -

- -

Chapter 2-

Bonding in Solids

Metallic bonding

Most metals have one, two, or at most three valence electrons

These electrons are highly delocalized from a specific atomhavea sea of valence electrons

Free electrons shield positive core ofions from one another (reduce ER)

Metallic bonding is also non-

directional

Free electrons also act to holdstructure together

Wide range of bonding energies,typically good conductors (why?)

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Chapter 2 -

Secondary Bonding or van der

Walls Bonding Also known as physical bonds

Weak in comparison to primary or chemical

bonds Exist between virtually all atoms and molecules

Arise from atomic or molecular dipoles

bonding that results from the coulombic attraction

between the positive end of one dipole and the

negative region of an adjacent one

a dipole may be created or induced in an atom or

molecule that is normally electrically symmetricChapter 2 -


Waals Bonding Fluctuating Induced Dipole Bonds

A dipole (whether induced or instantaneous)

produces a displacement of the electron distribution

of an adjacent molecule or atom and continues as a

chain effect

Liquefaction and solidification of inert gases

Weakest Bonds

Extremely low boiling and melting pointAtomic nucleus

Atomic nucleus

Electron

cloud

Electron

cloud

Instantaneous

Fluctuation

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Chapter 2 -

Secondary Bonding or van der Waals Bonding

Polar Molecule-Induced Dipole Bonds

Permanent dipole moments exist by virtue of an

asymmetrical arrangement of positively and negatively

charged regions

Polar molecules can induce dipoles in adjacent nonpolar

molecules

Magnitude of bond greater than for fluctuating induced

dipoles

+ -

Polar

Molecule

Induced

Dipole

Atomic nucleus

Electron Cloud

Chapter 2 -


Waals Bonding Permanent Dipole Bonds

Stronger than any secondary bonding with induced

dipoles

A special case of this is hydrogen bonding: existsbetween molecules that have hydrogen as one of the

constituents

H Cl H Cl

Hydrogen Bond

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Chapter 2-

c02tf03

Chapter 2-

c02f16

Many molecules do not have asymmetric distribution/arrangementof positive and negative charges(e.g. H2O, HCl)

MATERIAL OF IMPORTANCE

Water

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Chapter 2-

c02uf01

Chapter 2 - 48

Bond length, r

Bond energy, Eo

Melting Temperature, Tm

Tmis larger if Eois larger.

Properties From Bonding: Tm

ror

Energy

r

larger Tm

smaller Tm

Eo =

bond energy

Energy

ro r

unstretched length

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Chapter 2 - 49

Coefficient of thermal expansion, a

a~ symmetric at ro

ais larger if Eois smaller.

Properties From Bonding : a

= a (T2-T1)DL

Lo

coeff. thermal expansion

DL

length, Lo

unheated, T1

heated, T2

ror

smaller a

larger a

Energy

unstretched length

Eo

Eo

Chapter 2- 16

Elastic modulus, E

DLFAo

= ELo

Elastic modulus

PROPERTIES FROM BONDING: E

E ~ dF/dr|ro elastic modulus

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http://www.wiley.com/college/callister/0471470147/gallery/ch07/pages/Fig07_08.html


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Chapter 2 - 51

Ceramics

(Ionic & covalent bonding):

Large bond energylarge Tmlarge E

small a

Metals(Metallic bonding): Variable bond energymoderate Tmmoderate E

moderate a

Summary: Primary Bonds

Polymers(Covalent & Secondary):

Directional PropertiesSecondary bonding dominates

small Tmsmall Elarge a

Chapter 2 - 52

Type

Ionic

Covalent

Metallic

Secondary

Bond Energy

Large!

Variablelarge-Diamondsmall-Bismuth

Variablelarge-Tungstensmall-Mercury

smallest

Comments

Nondirectional (ceramics)

Directional(semiconductors, ceramicspolymer chains)

Nondirectional (metals)

Directional

inter-chain (polymer)inter-molecular

Summary: Bonding

atomic structure & bonding lect

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