15.06.2011

1

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-

Atomic Models

15.06.2011

2

Chapter 2-

~ 400 BC - Democritus

• Ancient Greek philosopher

• Democritus coined the term átomos which means "uncuttable" or "the smallest indivisible particle of matter".

Structure of MatterPhysical world

“VOID + BEING”

Chapter 2-

1803 – John Dalton

• English instructor and natural philosopher

• “Each element consists of atoms of single unique type and can join to form chemical compounds.”

• Originator of the modern atomic theory

15.06.2011

3

Chapter 2-

1869 - Mendeleev• Building upon earlier

discoveries by scientists, Mendeleev published the first functional PERIODIC TABLE.

• Certain chemical properties of elements repeat periodically when arranged by atomic number.

Chapter 2-

Periodic Table

Draft of the first periodic table, Mendeleev, 1869

15.06.2011

4

Chapter 2-

1869…

Chapter 2-

Today: Periodic Table of the Elements

15.06.2011

5

Chapter 2-

The Structure of the Atom

Status report end of the 19th century

• Atom is electrically neutral

• Negative charge carried by electrons

• Electron has very small mass – bulk of the atom is positive,

– most mass resides in positive charge

Chapter 2-

The Structure of the Atom

particle symbol charge (C) mass (kg)

electron e– –1.6×10–19 9.11×10–31

proton p+ +1.6×10–19 1.673×10–27

neutron no 0 1.675×10–27

Question: what is the distribution of charge inside an atom?

15.06.2011

6

Chapter 2-

1897 – Sir J. J. Thomson

• Discovered the electron (1906 Nobel Prize in Physics).

• Plum Pudding (1904): “The atom as being made up of electrons swarming in a sea of positive charge.

Chapter 2-

• Results:– Majority of a particles transmitted (pass through) or

deflected through small angles

– Tiny fraction deflected through large angles

1909 – E. Rutherford

• Tested the Plum Pudding Model.

15.06.2011

7

Chapter 2-

• Conclusion:– Disproved the Plum-Pudding Model

– Large amount of the atom's charge and mass is concentrated into a small region

– Atom was mostly empty space

• Objections to Rutherford model– The laws of classical mechanics predict that the

electron will release electromagnetic radiationwhile orbiting a nucleus. Because the electron would lose energy, it would gradually spiral inwards, collapsing into the nucleus.

– This atom model is unsuccessful, because it predicts that all atoms are unstable.

1909 – E. Rutherford

Chapter 2-

1912 – N. Bohr

• Many phenomena involving electrons in solids could not be explained in terms of CLASSICAL MECHANICS.

• We need QUANTUM MECHANICS…

15.06.2011

8

Chapter 2-

Bohr Postulates for the Hydrogen Atom

1. Rutherford atom is correct

2. Classical EM theory not applicable to orbiting e-

3. Newtonian mechanics applicable to orbiting e-

4. Eelectron = Ekinetic + Epotential

5. e- energy quantized through its angular momentum: L = mvr = nh/2π, n = 1, 2, 3,…

6. Planck-Einstein relation applies to e-transitions:

ΔE = Ef-Ei= hν = hc/λ

c = νλ

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

15.06.2011

9

Chapter 2-

1853 - A. Ångström

Chapter 2-

1913 - Sommerfeld

• German theoretical physicist

• Modified the Bohr Model

• “suppose we have plurality of orbits” – a shell containing multiple orbits: ORBITALS

• How to capture these new ideas quantitatively?

• We need new quantum numbers: n, l, m, s n principal quantum number, distance of an electron from the nucleusl subshell, describes the shape of the subshellm number of energy states in a subshells spin moment

15.06.2011

10

Chapter 2-

Wave mechanics to arrive at same place: E=E(n,l,m,s)

• The Bohr model – significant limitations• Resolution: Wave-mechanical model

(electron is considered to exhibit both wave-like and particle-like characteristics).

– De Broglie: “If a photon which has no mass, can behave as a particle, does an electron which has mass can behave as a wave (1920)?” λ = h/p = h/mv

– Heisenberg: Uncertainty Principle

“I don’t know where any of one of electrons is, but I can tell you an average where any of one of them is likely to be”

– Schrodinger

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

15.06.2011

11

Chapter 2-

FUZZY ORBITS

Chapter 2-

What is the filling sequence of electrons in orbitals by n, l, m, s is not adequate?

AUFBAU PRINCIPLE

3 principles:1. Pauli Exclusion Principle:only one electron

can have a given set of four quantum numbers.

2. Electrons

Have discrete energy states

fill orbitals from lowest

en. to highest en.

3. Hund’s rule

Incr

easi

ng e

nerg

y

n=1

n=2

n=3

n=4

1s

2s

3s

2p

3p

4s4p

3d

15.06.2011

12

Chapter 2-

Niels Bohr, Werner Heisenberg, and Wolfgang Pauli talking in the Niels Bohr Institute lunchroom, possibly 1934 or 1936

Chapter 2-

Quantum Numbers (II)

mll ms = ±½

15.06.2011

13

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 1s 1

1s 2 (stable) 1s 22s 1 1s 22s 2 1s 22s 22p 1 1s 22s 22p 2 ...

1s 22s 22p 6 (stable) 1s 22s 22p 63s 1 1s 22s 22p 63s 2 1s 22s 22p 63s 23p 1 ...

1s 22s 22p 63s 23p 6 (stable) ...

1s 22s 22p 63s 23p 63d 10 4s 24 6 (stable)

SURVEY OF ELEMENTS

Adapted from Table 2.2, Callister 7e.

15.06.2011

14

Chapter 2- 4

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

Stable electron configurations...

Z Element Configuration

2 He 1s 2

10 Ne 1s 22s 22p 6

18 Ar 1s 2 2s 22p 63 s 23p 6

36 Kr 1s 2 2s 22p 63 s 23p 63d 10 4 s 24p 6

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

15.06.2011

15

Chapter 2-

He

Ne

Ar

Kr

Xe

Rn

iner

t gas

es

acce

pt 1

e

acce

pt 2

e

give

up

1e

gi

ve u

p 2

e

give

up

3e

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- 7

• 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; reactivityMetals like to give up, halogens like to acquire electrons !

ELECTRONEGATIVITY

15.06.2011

16

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 neutral

The atomic mass (A): total mass of protons + total mass of neutronsAtomic 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) + ELECTRONS

NUCLEUS BONDING

Chapter 2-

Atomic Structure

• Valence electrons determine all of the following properties

1) Chemical

2) Electrical

3) Thermal

4) Optical

15.06.2011

17

Chapter 2-

Atomic bonding in solids

Things are made of atoms—little particles that move around, attracting each other when they are a little distance apart, but repelling upon being squeezed into one another. In that one sentence ... there is an enormous amount of information about the world.

— Richard P. Feynman

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

15.06.2011

18

Chapter 2-

Na (metal) unstable

Cl (nonmetal) unstable

electron

+ - Coulombic Attraction

Na (cation) stable

Cl (anion) stable

8

• 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 acceptselectrons 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]

15.06.2011

19

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.

15.06.2011

20

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

NaCl

MgO

CaF2CsCl

Chapter 2- 10

• Requires shared electrons

• Example: CH4

C: has 4 valence e,needs 4 more

H: has 1 valence e,needs 1 more

Electronegativitiesare 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)

15.06.2011

21

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

1e

(XAXB)2

4

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

2

Ex: MgO XMg = 1.3XO = 3.5

15.06.2011

22

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 4.0

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

Si C

C(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…

15.06.2011

23

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 electronclouds

+ - + -secondary bonding

HH HH

H2 H2

secondary bonding

ex: liquid H2

H Cl H Clsecondary bonding

secondary bonding+ - + -

secondary bonding

Chapter 2-

Bonding Energies

15.06.2011

24

Chapter 2-

Type

Ionic

Covalent

Metallic

Secondary

Bond Energy

Large!

Variablelarge-Diamondsmall-Bismuth

Variablelarge-Tungstensmall-Mercury

smallest

Comments

Nondirectional (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 runstretched length

15.06.2011

25

Chapter 2- 16

• Elastic modulus, E

• E ~ curvature at ro

cross sectional area A o

L

length, Lo

F

undeformed

deformed

L F Ao

= E Lo

Elastic modulus

r

larger Elastic Modulus

smaller Elastic Modulus

Energy

r o 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

E

oE

o

15.06.2011

26

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 Elarge

Summary: Primary Bonds