‘’ATOMIC STRUCTURE AND

BONDING’’

IE-114 Materials Science and General Chemistry

Lecture-2

Outline

• Atomic Structure

(Fundamental concepts, Atomic models (Bohr and Wave-Mechanical Atomic Model), Electron

configurations)

• Periodic Table (Classification of elements, their characteristics)

• Atomic Bonding in Solid Materials(Primary Bonding:Ionic, covalent, metallic bonds)

(Secondary bonding: Fluactuating Induced Dipole, Polar Molecule-Induced

Dipole, Permanent Dipole Bonds)

Why atomic structure is important?

Some properties of solid materials depend on atomic nature and itsarrangement.

Type of atomic arrangements : crytalline or amorphous Type of bonding (interactions among the atoms) determines the melting temperature (Tm),

coefficient of thermal expansion(α), mechanical properties, i.e. Elastic modulus, E

Graphite: Soft and greasy feelDiamond: Hardest known material

Graphite and Diamond

different type of interatomic

bonding in graphite and diamond.

graphite diamond

Atomic Number (Z): Number of protons in the nucleus

Atomic Structure

Electrons are negatively(-) charged,

Protons are positively(+) charged

Neutrons are electrically neutral “particles.”

Atom = nucleus (protons+neutron) and electrons

Charge =1.60x10-19

C

Electrically neutral atom; # protons = #electrons

Atomic mass (A) for an atom = masses of protons(Z)+masses of neutrons(N)

(electrons are not considered, because ...?)Mass of proton = mass of neutrons=1.67x10-27 kg

Mass of electron= 9.11x10-31 kg

For all atoms of an element the number of protons are the same, but the number

of neutrons may vary, which vary the atomic mass,

Example: 12C, 13C, 14C

Atomic Weight:

Weighted average of the atomic masses of the atom’s naturally occuring

isotopes

Mass of an Atom

Atoms with two or more atomic masses (ISOTOPES)

The atomic mass unit (amu) is used for the computation of atomic

weight.

Scale: 1 amu=1/12 of the atomic mass of the Carbon (C)

(A=12.00000 for carbon 12 isotope)

1 amu/atom= 1 g/mol ( 1 mol of a substance=6.023x1023 atoms)

For example: Fe

A=55.85 amu/atom or 55.85 g/mol (this is most commonly used form)

Structure of Atom

BOHR ATOMIC MODEL:

(Used hydrogen atom)

1) Electrons are assumed to be positioned around the nucleus in discrete orbitals

2) Position of the electron is more or less well defined in its orbital.

Nucleus: Z = # protons

N = # neutrons

1) Bohr Atomic Model

2) Wave-Mechanical Model

An electron can change its energy level

To a higher level by absorbing energy, to a lower level by emitting energy

Example:

n=1 E1 = -13.6 eV

n=2 E2 = - 3.4 eV

Bohr Atomic model describe the electrons in terms of their positions

(orbitals) and energy (quantized energy levels by Rydberg equation).

Energy of Electrons

E= - (2π2me

4/n

2h

2) = - (13.6/n

2) eV

e: electron charge

m: electron mass

n: principal quantum number or principal energy levels(1,2,3,….)

If electron changes its energy level from 1 to 2 (from lower to higher energy level)

It must absorb energy, and the amount of energy absorbed;

E =E2 - E1 = -3.4 – (-13.4) = 10.0 eV

E2>E1

Allowed energy levels for hydrogen electron in Bohr Model

Ionization energy:

Energy required to remove the electron completely from the atom

Ionization energy for hydrogen electron is 13,6 eV

Figure.The first three electron energy states for the Bohr hydrogen aton

* Bohr’s model was not able to explain quantitatively the spectra of the atoms more complex than

hydrogen and the model could not have been modified.

WAVE-MECHANICAL MODEL:

Limitations of Bohr model was resolved by this model and electrons are

considered to behave both wave-like and particle-like.

Electrons are no longer treated as a particle moving in discrite orbitals.

BOHR MODEL WAVE-MECHANICAL

MODEL

Position of electron is described by a probability distribution or electron cloud

Heisenberg’s uncertainty principle;

Position and momentum of a small particle such as an

electron can not be determined simultaneously.

Since the position of an electron can not be precisely determined,

an electron charge cloud density distribution is used

Motion of electron around its nucleus and its energy is

characterized by 4 QUANTUM NUMBERS (n, l, ml, m

s)

1) Principal quantum number,n:

Represents main energy levels for the electrons or shells (n=1 to 7 )

n=1 (first shell, K) n= 2 second shell(L), so forth... or

2) Secondary quantum number, l:

Specifies subenergy levels within the main energy levels(subshells) and related to shape of the electron subshell

Number designation of l: 0 1 2 3 4 5....Letter designation of l : s p d f g h

1 K s

2 L s,p

3 M s,p,d

4 N s,p,d,f

Principle quantum

number, n

Shell Designation Subshells

l=n-1

3) Magnetic quantum number, ml:

Number of orbitals or energy states for each subshell

Example: For a given l, ml can range from +l to –l

l=0 (s subshell) ml = 1 energy state (0)

l=1 (p subshell) ml = 3 energy states (+1, 0, -1)

l=2 (d subshell) ml = 5 energy states (+2,+1, 0, -1,-2)

ml=2l+1

4) Electron spin quantum number, ms:

Specifies two allowed spin directions for an electron.

ms = +1/2 and -1/2

Pauli’s Exclusion Principle:

No two electrons can have the identical values for all four of their quantum numbers

Comparison of electron energy states in

Bohr and Wave-mechanical Models

BOHR MODEL WAVE-MECHANICAL MODEL

***Electrons fill up the the lowest

possible energy states in the

electron shells and subshells

The maximum number of electrons in each shell in an atom is 2n2

Electron Configurations

Represents the manner in which the states are occupied

The number of electrons in each subshell is indicated by a superscript after the shell-subshell designation.

Example:

H 1s1

He 1s2

Na 1s22s22p63s1

Na 1s22s22p63s1

Principal quantum number,n

(SHELL K,L,M,..)

Secondary quantum number,l

(SUBSHELL; s,p,d,f)

s subshell has one orbital: ml=0

Orbital contains 2 electrons

p subshell has three orbitals:

ml=-1,0,+1, each of these orbitals

contains 2 electrons

• have complete s and p subshells

• tend to be unreactive.

Stable electron configurations

Valance Electrons:The electrons occupying the outermost shell

These electrons participate in bonding. Many of the physical and chemical

properties of solids are based on these valence electrons.

(inert, or noble, gases)

* Most of the elements are not stable.

Electron configuration 1s1

1s2 (stable) 1s22s1 1s22s2 1s22s22p1 1s22s22p2 ...

1s22s22p6 (stable) 1s22s22p63s1 1s22s22p63s2 1s22s22p63s23p1 ...

1s22s22p63s23p6 (stable) ...

1s22s22p63s23p63d104s246 (stable)

Survey of Elements

Periodic Table

Elements are classified according to electron configuration in this table.

Same column, or group, have similar chemical and physical properties due to

similar valence electron configurations

These properties change gradually and systematically across each period

moving horizontally.

Atomic Number(Z) +

*Group 0: inert gases (filled electron shells)

*Group IA and IIA are alkali (except H) and alkaline earth

metals

*Group IIIA, IVA and VA elements have characteristics between metal and nonmetals because of their valence electron configurations.

*Group VIIA (halogens) and VIA elements= one and two electrons deficient respectively from having stable configurations.

*Groups from IIIB to IIB are transition metals, with partially filled d electron states and in some cases one or two electrons in the next higher shell.

Groups are designated at the top by the numbers 0-7 and by the letters

A and B.

A group elements- Representative or main group elements

B group elements- Transition elements

Electropositive elements:

Readily give up electrons

to become + ions.

Electronegative elements:

Readily acquire electrons

to become - ions.

Electropositive elements: Elements capable of giving up their electrons to

become positively charged ions (located on the left of the table.)

Electronegative elements: Elements ready to accept electrons to form

negatively charged ions or to share their electrons.

Electropositive and Electronegative Elements

• Ranges from 0.7 to 4.0

• Large values: tendency to acquire electrons.

Electronegativity

The degree to which an atom attracts electrons to itself

Smaller electronegativity Larger electronegativity

Atomic Bonding in Solids

Physical properties are related

to interatomic forces that bind the

atoms together

Atomic bonding can be

explained by interaction of two

isolated atoms

Net force is zero (equilibrium state)

r0=equilibrium spacing

Energy required to seperate these two atoms to an infinite seperation

Bonding energy (Eo); Solids >Liquids > Gases

1) Bonding arises from the tendency of the atoms to assume stable

electron structures

2) Valance electrons are involved

3) Nature of bond depends on the electron structure

Bonding energy, Eo

Bond energy, Eo

1) Melting Temperature, Tm

Tm is larger if Eo is larger.

Properties from bonding

• α ~ symmetry at ro

is larger if Eo is smaller.

= (T2-T1) L

Lo

coeff. thermal expansion

2) Coefficient of thermal expansion, α

• E ~ curvature at ro

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.

3) Elastic Modulus, E

1)Primary Bonding:

Ionic bonding

Covalent bonding

Metallic bonding

2)Secondary bonding:

Fluactuating Induced Dipole

Polar Molecule-Induced Dipole

Permanent Dipole Bonds

Types of Bondings

Found in compounds formed by metallic and nonmetallic elements

(occurs between + and – ions)

Requires electron transfer.

Large difference in electronegativity is required.

• Example: NaCl

Ionic Bonding

The ionic bonding is nondirectional, that is the magnitude of the bond is

equal in all directions.

The predominant bonding in ceramics is ionic.

Enet= Eatt. + Erep.

Enet= - (A/r) + B/rn

B, and n are constants. n is

approximately 8.

(A = (Z1Z2e2/4π o) + B/rn)

Give up electrons Acquire electrons

Covalent Bonding

Stable electron configurations are assumed by sharing of electrons

between adjacent atoms.

• Example: CH4(methane)

C: has 4 valence e, needs 4 more

H: has 1 valence e, needs 1 moreElectronegativities are comparable.

H feels like helium electron configuration, while C feels like neon electron configuration.

Covalent bonding is directional

It forms between two specific atoms and may exist only in the direction

between one atom and another.

• Compound solids (about column IVA)

• Molecules with nonmetals

• Molecules with metals and nonmetals

• Elemental solids (RHS of Periodic Table)

N’= number of valence electrons

Example: Cl atom7 valence electrons, an atom can have maximum 1 more bond (completing the valence orbital electron number to eight)

Covalent bonds may be extremely strong (like in diamonds) or may be weak (like in Bismuth).Polymeric materials are covalently bonded materials.

Some bonds are partially ionic and partially covalent. The degree of either bond is controlled by the electronegativities of the composing atoms.

%ionic character = (1-e-(0.25)(XA-XB)2)x100

(XA and XB are the electronegativities of the respective elements)

As the electronegativity difference gets higher, the bonding becomes more ionic.

Number of covalent bonds

# = 8-N’

%Ionic character

Valence electrons are not bound to any particular atom in the solid and they

are more or less free to move throughout the entire metal.

Metallic bond is nondirectional.

Metallic Bonding

Primary bond for metals and their alloys

Metalling bonding explains the heat and electric conductivity of the

metallic materials as well as their ductility.

Arises from interaction between dipoles

• Permanent dipoles-molecule induced

• Fluctuating dipoles

-general case:

-ex: liquid HCl

-ex: polymer

Secondary Bonding