Chapter 2-

ISSUES TO ADDRESS...• What promotes bonding?

• What types of bonds are there?

• What properties are inferred from bonding?

1

Today's Agenda

Chapter 2-

GECKO

WHY ?

Chapter 2-

Here is the answer

Chapter 2-

Chapter 2: Atomic Structure and Interatomic Bonding

• Atomic Structure• Electron Configuration• Periodic Table• Primary Bonding

– Ionic– Covalent– Metallic

• Secondary Bonding or van der Waals Bonding– Three types of Dipole Bonding

• Molecules

Chapter 2-

REVIEW OF ATOMIC STRUCTURE (FRESHMAN CHEMISTRY)

• Mass of an atom:– Proton and Neutron: ~ 1.67 x 10-27 kg – Electron: 9.11 x 10-31

kg • Charge:

– Electrons and protons: (±) 1.60 x 10-19 C

– Neutrons are neutralThe atomic mass (A): total mass of protons + total mass of neutrons Atomic weight ~ Atomic mass# of protons are used to identify elements (Z)# of neutron are used to identify isotopes ( e.g. 14C6 and 12C6 )

Isotopes are written as follows: AXZ , i.e. 1H1, 2H1, 3H1

ATOMS = (PROTONS+NEUTRONS) + ELECTRONSNUCLEUS BONDING

Chapter 2-

AMU and Mole ConceptThe atomic mass unit (amu) : 1 amu is defined as the 1/12 of the atomic mass of the most common isotope of carbon,

carbon 12 (12C6). Atomic mass of 12C6 is 12 amu : Carbon= 6 protons (Z=6) + 6 neutrons (N=6) Mproton ~ Mneutron = 1.67 x 10-24 g = 1 amu Atomic Mass = Z + N

The atomic weight is often expressed in mass per mole

A mole is the amount of matter that has a mass in grams equal to the atomic mass in amu of the atoms (A mole of carbon has a mass of 12 grams).

The number of atoms in a mole is called the Avogadro number, Nav = 6.023 × 1023.

Nav = 1 gram/1 amu.

• Example:Atomic weight of iron = 55.85 amu/atom = 55.85 g/mol

Chapter 2-

MOLE CONCEPT

ANY IDEAS WHY ?

Chapter 2-

Atomic Structure

• Valence electrons determine all of the following properties

1) Chemical2) Electrical 3) Thermal4) Optical

Chapter 2-

Atomic Models• Towards the end of 19th century physicists realized Newtonian physics

has serious difficulty in explaining many phenomena involving electrons => quantum mechanics– Bohr atomic model

• Electrons assume very well defined orbits around the nucleus (protons + neutrons)

• Electrons in each shell orbit assumes the same energy level– Severe issues when considering events involving electrons such as emission

spectra and photoelectrons…)– See figure 2.1 in Callister page 18 (page 13 for 6th ed.)

– Wave mechanical model• Electrons in an atom or molecule are permitted to have only specific values of

energy, energy is quantized…• Electrons do not move in circular orbits but in “fuzzy” orbits. At any given

time we can only talk about the probability of finding an electron at a radius from the orbit.

• Every electron is characterized by four quantum numbers. The size, shape, spatial orientation of an electrons probability density are specified by these numbers

Chapter 2-

Nucleus: Z = # protons

2

orbital electrons: n = principal quantum number

n=3 2 1

= 1 for hydrogen to 94 for plutoniumN = # neutrons

Atomic mass A ≈ Z + N

Adapted from Fig. 2.1, Callister 6e.

BOHR ATOM

Chapter 2-

Beyond Bohr’s Model

In 1927 Heisenberg : uncertainity, it is not possible to measure simultaneously both the momentum (or velocity) and the position of a microscopic particle with absolute accuracy.

In 1924 de Broglie : dual character of electrons

Schrodinger, math expression for the behavior of an electron around an atom

Chapter 2-

FUZZY ORBITS

Chapter 2-

• have discrete energy states• tend to occupy lowest available energy state.

3

Incr

easin

g en

ergy

n=1n=2

n=3

n=4

1s2s

3s2p3p

4s4p 3d

Electrons...

Adapted from Fig. 2.5, Callister 6e.

ELECTRON ENERGY STATES

Chapter 2-

Quantum Numbers (I)• A product of Schrodinger’s Equation

– n, l , ml , ms

• n principal quantum number, distance of an electron from the nucleus

• l subshell, describes the shape of the subshell• ml number of energy states in a subshell

• ms spin moment

• Pauli’s exclusion principle: only one electron can have a given set of four quantum numbers.

Chapter 2-

Quantum Numbers (II)

mll ms = ±½

Chapter 2-

Quantum Numbers (III)Electrons fill quantum levels in order of increasing energy ( only n and l make significant differences in energy configurations).

1s, 2s, 2p, 3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,….

When all electrons are at the lowest possible energy levels => ground state

Excited states do exist such as in glow discharges etc…

Valence electrons occupy the outermost filled shell.Valence electrons are responsible for all bonding !

Chapter 2-5

• Why? Valence (outer) shell usually not filled completely.

• 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)

SURVEY OF ELEMENTS

Adapted from Table 2.2, Callister 7e.

Chapter 2-4

• 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 CONFIGURATIONS

Chapter 2-

Electron Configurations• Valence electrons – those in unfilled shells• Filled shells more stable• Valence electrons are most available for bonding

and tend to control the chemical properties

– example: C (atomic number = 6)

1s2 2s2 2p2

valence electrons

Chapter 2-

He Ne Ar Kr Xe Rn

iner

t gas

es

acce

pt 1

e ac

cept

2e

give

up 1

e gi

ve u

p 2e

give

up 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

6

• Columns: Similar Valence Structure, Similar Properties

Electropositive elements:Readily give up electronsto become + ions.

Electronegative elements:Readily acquire electronsto become - ions.

THE PERIODIC TABLE

Chapter 2-

Periodic Table

Draft of the first periodic table, Mendeleev, 1869

Chapter 2-7

• Ranges from 0.7 to 4.0,

Smaller electronegativity Larger electronegativity

He -

Ne -

Ar -

Kr -

Xe -

Rn -

F 4.0Cl 3.0Br 2.8

I 2.5At 2.2

Li 1.0Na 0.9K

0.8Rb 0.8Cs 0.7Fr 0.7

H 2.1

Be 1.5Mg 1.2Ca 1.0Sr 1.0Ba 0.9Ra 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; reactivity Metals like to give up, halogens like to acquire electrons !

ELECTRONEGATIVITY

Chapter 2-

Concept ChecksQuestion: Why are the atomic weights of the elements generally not integers? Cite two reasons.

Answer: The atomic weights of the elements ordinarily are not integers because:(1) The atomic masses of the atoms normally are not integers (except for 12C), and (2) the atomic weight is taken as the weighted average of the atomic masses of an atom's naturally occurring isotopes.Question: Give electron configurations for the Fe3+and S2- ions.Answer: The Fe3+ ion is an iron atom that has lost three electrons. Since the electronconfiguration of the Fe atom is 1s22s22p63s23p63d64s2 (Table 2.2), the configuration for Fe3+ is 1s22s22p63s23p63d5.The S2- ion a sulfur atom that has gained two electrons. Since the electron configuration of the S atom is 1s22s22p63s23p4 (Table 2.2), the configuration for S2- is 1s22s22p63s23p6.

Chapter 2-

Atomic Bonding in Solids• Start with two atoms infinitely

separated• Attractive component is due to

nature of the bonding (minimize energy thru electronic configuration)

• Repulsive component is due to Pauli exclusion principle; electron clouds tend to overlap

• Essentially atoms either want to give up (transfer) or acquire (share) electrons to complete electron configurations; minimize their energy

– Transfer of electrons => ionic bond– Sharing of electrons => covalent– Metallic bond => sea of electons

r

Chapter 2-

Na (metal) unstable

Cl (nonmetal) unstable

electron

+ - Coulombic Attraction

Na (cation) stable

Cl (anion) stable

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• Occurs between + and – ions (anion and cation).• Requires electron transfer.• Large difference in electronegativity required.• Example: Na+ Cl-

IONIC BONDING (I)

Chapter 2-

Ionic bond – metal + nonmetal

donates accepts electrons electrons

 

Dissimilar electronegativities  

ex: MgO Mg 1s2 2s2 2p6 3s2 O 1s2 2s2 2p4

[Ne] 3s2 

Mg2+ 1s2 2s2 2p6 O2- 1s2 2s2 2p6 [Ne] [Ne]

Chapter 2-8

IONIC BONDING (II)

Oppositely charged ions attract, attractive force is coulombic. Ionic bond is non-directional, ions get attracted to one another in any direction.Bonding energies are high => 2 to 5 eV/atom,molecule,ionHard materials, brittle, high melting temperature, electrically and thermally insulating

Chapter 2-

Ionic Bonding• Energy – minimum energy most stable

– Energy balance of attractive and repulsive terms

Attractive energy EA

Net energy EN

Repulsive energy ER

Interatomic separation r

rA

nrBEN = EA + ER =

Adapted from Fig. 2.8(b), Callister 7e.

Chapter 2-

• Predominant bonding in Ceramics

Adapted from Fig. 2.7, Callister 7e. (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 Cornell University.

Examples: Ionic Bonding

Give up electrons Acquire electrons

NaClMgOCaF2CsCl

Chapter 2-10

• Requires shared electrons• Example: CH4

C: has 4 valence e, needs 4 moreH: has 1 valence e, needs 1 moreElectronegativities are comparable.

shared electrons from carbon atom

shared electrons from hydrogen atoms

H

H

H

H

C

CH4

Adapted from Fig. 2.10, Callister 6e.

COVALENT BONDING (I)

Chapter 2-10

COVALENT BONDING (II)

Covalent bonds are formed by sharing of the valence electronsCovalent bonds are very directionalCovalent bond model: an atom can have at most 8-N’ covalent bonds, where N’ = number of valence electrons

Covalent bonds can be very strong, eg diamond, SiC, Si, etc, also can be very weak, eg Bismuth Polymeric materials do exhibit covalent type bonding.

Diamond, sp3

Chapter 2-

Primary Bonding• Metallic Bond -- delocalized as electron cloud  

• Ionic-Covalent Mixed Bonding

% ionic character =  where XA & XB are Pauling electronegativities

%)100( x

1 e (XA XB )2

4

ionic 70.2% (100%) x e1 characterionic % 4)3.15.3( 2

Ex: MgO XMg = 1.3XO = 3.5

Chapter 2-10

COVALENT BONDING (III)Very few materials have pure ionic or covalent bonding; electronegativity inpart defines how much time electrons spend between ion cores…

Chapter 2-11

• 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 4.0Cl 3.0Br 2.8

I 2.5At 2.2

Li 1.0Na 0.9K

0.8Rb 0.8Cs 0.7Fr 0.7

H 2.1

Be 1.5Mg 1.2Ca 1.0Sr 1.0Ba 0.9Ra 0.9

Ti 1.5

Cr 1.6

Fe 1.8

Ni 1.8

Zn 1.8

As 2.0

SiCC(diamond)

H2O

C 2.5

H2

Cl2

F2

Si 1.8

Ga 1.6

GaAs

Ge 1.8

O 2.0

colu

mn

IVA

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

Chapter 2-12

• Arises from a sea of donated valence electrons (1, 2, or 3 from each atom).

• Primary bond for metals and their alloys

+ + +

+ + +

+ + +Adapted from Fig. 2.11, Callister 6e.

METALLIC BONDING

Ion cores in the “sea of electrons”.

Valance electrons belong no one particular atom but drift throughout the entire metal.

“Free electrons” shield +’ly charged ions from repelling each other…

Chapter 2-

Arises from interaction between dipoles

• Permanent dipoles-molecule induced

• Fluctuating dipoles

-general case:

-ex: liquid HCl

-ex: polymer

Adapted from Fig. 2.13, Callister 7e.

Adapted from Fig. 2.14, Callister 7e.

SECONDARY BONDING

asymmetric electron clouds

+ - + -secondary

bonding

HH HH

H2 H2

secondary bonding

ex: liquid H2

H Cl H Clsecondary bonding

secondary bonding+ - + -

secondary bondingsecondary bonding

Chapter 2-

Bonding Energies

Chapter 2-

TypeIonic

Covalent

Metallic

Secondary

Bond EnergyLarge!

Variablelarge-Diamondsmall-Bismuth

Variablelarge-Tungstensmall-Mercury

smallest

CommentsNondirectional (ceramics)

Directional(semiconductors, ceramicspolymer chains)

Nondirectional (metals)

Directionalinter-chain (polymer)inter-molecular

Summary: Bonding

Chapter 2-

• Bond length, r

• Bond energy, Eo

• Melting Temperature, Tm

Tm is larger if Eo is larger.

Properties From Bonding: Tm

r o r

Energyr

larger Tm

smaller Tm

Eo = “bond energy”

Energy

r o r

unstretched length

Chapter 2-16

• Elastic modulus, E

• E ~ curvature at ro

cross sectional area Ao

L

length, Lo

F

undeformed

deformed

L F Ao = E Lo

Elastic modulus

r

larger Elastic Modulus

smaller Elastic Modulus

Energy

ro unstretched length

E is larger if Eo is larger.

PROPERTIES FROM BONDING: E

Chapter 2-

• Coefficient of thermal expansion,

• ~ symmetry at ro

is larger if Eo is smaller.

Properties From Bonding :

= (T2 -T1)LLo

coeff. thermal expansion L

length, Lo

unheated, T1

heated, T2

r or

larger

smaller

Energy

unstretched length

Eo

Eo

Chapter 2-

Ceramics(Ionic & covalent bonding):

Metals(Metallic bonding):

Polymers(Covalent & Secondary):

Large bond energylarge Tm

large Esmall

Variable bond energymoderate Tm

moderate Emoderate

Directional PropertiesSecondary bonding dominates

small Tm

small E large

Summary: Primary Bonds

secondary bonding

Chapter 2-

Read Chapter 3!!!

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