Atomic Structure and Bonding between atoms

1. Module 1
We will cover the electronic structure of atoms
and the types of bonds that can form between
atoms.
Module 1
-Atomic number and Atomic mass
-Electrons
-Electronic configuration
-Periodic Table
-Bonding between atoms
1

2. Atom: Basic Element of Matter
Atomic Number
# of PROTONS in nucleus of atom
= # of ELECTRONS of neutral species
Atomic mass
Mass of PROTON + Mass of NEUTRON 2
Mass Charge
Electron 9.11x10-31 kg -1.69x10-19 C
Proton 1.67x 10-27 kg 1.69x10-19 C
Neutron 1.67x 10-27 kg 0

3. 3
Atomic mass unit (amu), 1 amu = 1/12 mass of 12C
=1.67x10-27kg 1 atom 12C=12 amu
Mole
The amount of substance that has as many atoms/molecules
as in 12 grams of Carbon.
Avogadro’s number
1 mole of a substance has 6.023 x 1023 number of particles
Atomic Mass Unit
This is defined in units of gram mole.
1 kg-mole will contain atoms6.023x10 26

4. 4
Isotopes
Variants of the same element having different
number of neutrons.
Atomic Weight
Weighted average of the atomic masses of the naturally
occurring isotopes of a particular element in atomic mass unit.
As elements have isotopes atomic
weight is not an integer for
example : C 12.001 or H 1.008
For example, atomic weight of copper is 63.54 amu/atom
or 63.54 g/mole

5. 5
 How many Fe atoms are there in one gram of Fe ?
Atomic mass of Fe = 55.85 g/mol
 How many grams in one amu of a material?

6. So, where are the electrons?
• Electrons (negatively charged particles) orbit
the nucleus.
• We will learn about two different models
describing the placement of the electrons
around the nucleus.
6

7. Structure of Atoms- Two Models
Quantum Mechanics- set of principles and laws
that govern systems of atomic and sub atomic
entities.
1) Bohr Model
Electrons revolve around nucleus in discrete orbitals.
Deals with position (electron orbitals) and energy (quantized energy).
2) Wave Mechanical Model
Electrons exhibit both wave-like and particle-like characteristics
Probability of electron being in a location
- Energy quantized in shells and subshells.
7

8. Bohr Model
• Protons and neutrons in the
nucleus and are held together by
the strong force (overcomes
electrostatic repulsion).
8
 Electrons revolve around a positively
charged nucleus in discrete orbits
(K, L, M or n=1, 2, 3 respectively) with
specific
levels of energy.
 Electrons positions are fixed as such,
however, an electron can jump to higher
or lower energy level by absorption or
emission of energy respectively

9. Wave Mechanical Model
 Electrons are in orbitals defined by a probability.
 Each orbital at discrete energy level determined by 4 quantum numbers.
Quantum numbers :
Four parameters that describe the distribution and position of the
electrons
Principal quantum number, n = principal (energy level-shell)
K, L, M, N, O (1, 2, 3, etc.)
• Determines the size of atom or distance of shell from nucleus
• Number of electrons in a shell is 2n2
9

10. 10
Angular quantum number, l = subsidiary (orbitals)
It signifies the subshells/electron orbital s, p, d, f
The values of l ranges from (0, 1,…, n -1)
For example , M shell n=3, so l is 0,1,2 s, p and d
Magnetic quantum number, ml
It signifies the number of energy states in each orbital.
1, 3, 5, 7 (-l to +l) determines number of energy states in an orbital
For example, s orbital l=0 , so ml = 0 One energy state
p orbital l=1, so ml = -1,0,+1 Three energy states

11. 11
Pauli’s exclusion Principle:
It states that no more than two electrons having opposite spin can
occupy the same energy state.
For example, s orbital has one energy state and can accommodate
only two electrons having opposite spin(ms = ±1/2)
Spin quantum number, ms
It signifies the spin moment associated with the electrons. It can
have only two values, +1/2 and -1/2

12. 12
n l
subshells
ml
Energy
states
ms Maximum
number of
electrons in shell
Electronic configuration
K shell 1 0 0 ±1/2 2
1s2
L shell 2 0 0 ±1/2 8 electrons 2s2 2p6
One 2s state and
Three 2p states
1 +1 ±1/2
0 ±1/2
-1 ±1/2
M shell 3 0 0 ±1/2 18 electrons 3s2 3p6 3d10
One 3s state
Three 3p state
Five 3d state
1 +1 ±1/2
0 ±1/2
-1 ±1/2
2 +2 ±1/2
+1 ±1/2
0 ±1/2
-1 ±1/2
-2 ±1/2
N shell 4 32 electrons 4s2 4p6 4d10 4f14
 It denotes the manner in which the electrons are distributed in the
orbital shells.
Electronic Configuration

13. Determining Electron Configuration: Fill Lowest
Energy Sites First
Aufbau principle: Electrons will
sit in the lowest energy positions
first.
13
Madelung’s rule:
 Orbitals fill in the order of
increasing (n+l) value.
 For elements with same values of
(n+l) the one with lower value of
n will be filled first.

14. What is the electronic configuration
for Carbon and Iron?
14

15. What is the electronic configuration
for Carbon and Iron?
• Carbon
– Symbol ‘C’ on the periodic chart
– Atomic number 6
– Had 6 electrons
• Iron
– Symbol ‘Fe’ on the periodic chart
– Atomic number 26
– Had 26 electrons
15

16. Example: Carbon
16

17. Determining Electron Configuration for Carbon
17
1s
2s
2p
K-shell n = 1
L-shell n = 2
3s
3p M-shell n = 3
3d
4s
4p
4d
Energy
N-shell n = 4
L-shell with two
subshells
K-shell with one
subshells

18. Example: Iron
18
Lowest energy orbitals are filled first .

19. Adapted from Fig. 2.4,
Callister 7e.
1s
2s
2p
K-shell n = 1
L-shell n = 2
3s
3p M-shell n = 3
3d
4s
4p
4d
Energy
N-shell n = 4
Determining Electron Configuration for Iron
19

20. 20
Electron configuration
(stable)
...
...
1s22s22p 63s23p 6 (stable)
...
1s22s22p 63s23p 63d 10 4s24p6 (stable)
Atomic #
18
...
36
Element
1s11Hydrogen
1s22Helium
1s22s13Lithium
1s22s24Beryllium
1s22s22p 15Boron
1s22s22p 26Carbon
...
1s22s22p 6 (stable)10Neon
1s22s22p 63s111Sodium
1s22s22p 63s212Magnesium
1s22s22p 63s23p 113Aluminum
...
Argon
...
Krypton
Valence electrons are most available for bonding and tend to
control the chemical properties
 Filled shells more stable
Adapted from Table 2.2,
Callister & Rethwisch 3e.
Stable configurations
are lowest energy and
do not take part in any
chemical reaction.

21. Periodic Table
21
18 numbered groups (columns) and 7 periods (horizontal rows)

22. 22
Electropositive elements:
Readily give up electrons
to become + ions.
Electronegative elements:
Readily acquire electrons
to become - ions.
giveup1e-
giveup2e-
giveup3e-
inertgases
accept1e-
accept2e-
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
Periodic Table
Determine properties from valence electrons

23. Electronegativity is a measure of how readily an
atom accepts an electron to form a negatively
charged ion.
Electronegativity
23
Smaller electronegativity Larger electronegativity

24. 24
• Atomic radii describes the distance of the
outermost energy level from the nucleus.
• Atomic radii decrease as move from left to right
across of period and increase down a group.
• For ions , For example Fe, Fe2+ and Fe3+ , Fe3+ <
Fe2+ < Fe
Atomic Radii
n is constant
‘n’
increases

25. Bonding Between Atoms
Source: R.E. Smallman, Ch1, F1.3
25
• Type of bond between atoms will depend on
each atom’s electronic structure.
• The valence electrons take part in bonding.
• The atoms – ‘acquire’, ‘loose’or ‘share’
valence electrons to achieve lowest energy
configuration.

26. 26
Primary Classification
Primary Bonds-
High amount of energy needed to break bond
Secondary Bonds-
Lower amount of energy needed to break bonds

27. Na (metal)
unstable
Cl (nonmetal)
unstable
electron
+ -
Coulombic
Attraction
Na (cation)
stable
Cl (anion)
stable
27
• Requires electron transfer.
• Large difference in electronegativity required.
Example: NaCl

28. Examples: Ionic Bonding
Give up electrons Acquire electrons
NaCl
MgO
CaF2
CsCl
28

29. Example: MgO
Mg 1s2 2s2 2p6 3s2 O 1s2 2s2 2p4
Mg2+ 1s2 2s2 2p6 O2- 1s2 2s2 2p6
29
Donates
electrons
Accepts electrons
Ionic bonding found between atoms with large difference in
electronegativity ---in compounds composed of a metal and non-metal
and require the transfer of valence electrons
Bonds are non-directional in
nature.

30. Covalent Bonding
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 7e.
 Share valence electrons
 Formed between atoms of similar electronegativity
 s & p orbitals dominate bonding
shared electrons
from carbon atom
shared electrons
from hydrogen
atoms
H
H
H
H
C
CH4
30
Example: CH4
These bonds are directional in nature.

31. 31
Covalent bonding form through sharing of valence electrons
between atoms of similar electronegativity.

32. Metallic Bonding
• Valence electrons sea/ cloud
of electrons which are shared
by all atoms.
• Rest of electrons + Nuclei
ion core.
• Sea of electrons shields ion
cores.
32
Metallic bonding arises from interaction between sea of electrons
and ion Core

33. 33
Secondary Bonding – Van der Waals
bonding
 Dipole-Dipole
 Dipole induced Dipole
 Hydrogen Bonding
 Fluctuating dipole (weakest)
H Cl H Clsecondary
bonding
secondary
bonding
+ - + -
Example: Liquid HCl
Dipole-Dipole interaction:
Between molecules with permanent dipole moments.
Adapted from Fig. 2.13, Callister 7e.
~3.3kJ/mole

34. 34
Hydrogen Bonding
Bond strength: 10-50 kJ/mole
 A form of dipole-dipole interaction.
 Interaction between hydrogen atom and
another atom of high electronegativity
such as Fluorine or Oxygen or Nitrogen.

35. 35
Water- Anomalous expansion on cooling ---
decrease in density

36. 36
Polarity vs. Dipole
Asymmetric distribution of electrical charge over
the atoms joined by a bond or a molecule
http://sscchemistry.weebly.com/polaritydipole-and-hybridization.html
Non-polar
Polar
Bond between two
non-metal atoms with
greater difference in
electronegativity –
Polar covalent bond.
H 𝛿 +
−𝐶l 𝛿 −
Oxygen – 1s2 2s2 2p4

37. 37
Dipole Induced Dipole:
Polar molecules (with asymmetric arrangement of positively and
negatively charged regions) can induce dipoles in adjacent nonpolar
molecules.
Example:
Dipole in water molecule
(polar) induces dipole in
oxygen molecule (non-
polar)

38. 38
asymmetric electron
clouds
+ - + -
secondary
bonding
HH HH
H 2 H 2
secondary
bonding
ex: liquid H2
Fluctuating Dipole
Adapted from Fig. 2.14, Callister 7e.
 Weak attraction forces between dipoles due to natural
oscillation of atoms leading to momentary breakdown of
charge symmetry and temporary dipoles.

39. Bonding of Atoms
Types
Ionic
Covalent
Metallic
Secondary
Bond Energy
Large!
large-Diamond
small-Bismuth
Variable
large-Tungsten 850 kJ/mole
small-Mercury- 68 kJ/mole
smallest
39
 Large Bond energy
 Lack of free electrons
--Poor conductors of
heat and electricity.
 Brittle in nature.
 Moderate Bond Energy
 Valence electrons are free to
move
-- Good thermal and electrical
conductivities.
-- Bond strength increases with
atomic number—increase in
melting point, hardness and
strength.
-- Agile electrons- Ductile and
malleable
4-50 kJ/mole
600-1500 kJ/mole
100-1200kJ/mole

40. • Energy – minimum energy most stable
Energy balance of attractive and repulsive terms
Attractive energy EA
Net energy EN
Repulsive energy ER
Interatomic separation r
r
A
n
r
B
EN = EA + ER = +-
Adapted from Fig. 2.8(b),
Callister 7e.
40

41. Interactions Between Atoms
Relationship between interaction energy E and force F
Force
Total energy:
41
r
A
n
r
B
E = EA + ER = +-
 rF(r)dE
dr
dE(r)
F 

42. 42
How do we find forces
and bond length?

43. Find Attractive Force:
43
Find Repulsive Force!
Find Bond Length!

44. Secondary Bonding Example
Consider the van der Waals bonding in solid Argon. The potential energy as a
function of interatomic separation can be modeled as
Calculate the equilibrium bond length and bond energy.
A=10.37x10-78J-m6 B=16.16x10-135J-m12
E r( ) = -Ar-6
+ Br-12
44

45. 45
The Potential we just looked at
is called
Lennard-Lones potential














-






612
4
rr
ULG


ε is the depth of the potential well
σ is the finite distance at which the inter-particle potential is zero
Find equilibrium bond length

46. 46
Indicate the Equilibrium bond distance and Bond
Energy in the following Energy-interatomic
distance plots:
Eo
“bond energy”
Energy
r
o
r

47. • Bond length, r
• Bond energy, Eo
• Melting Temperature, Tm
Tm is larger if Eo is larger.
r
o r
Energy
r
larger Tm
smaller Tm
Eo =
“bond energy”
Energy
r
o r
unstretched length
47

48. 48
A steep slope at the equilibrium bond distance indicates
high modulus of elasticity of stiffness.
Force vs. Interatomic separation distance plots
of Material A and Material B.
r0
r0

49. 49
Ceramics
(Ionic & covalent bonding):
Metals
(Metallic bonding):
Polymers
(Covalent & Secondary):
Large bond energy
large Tm
large E
small α
Variable bond energy
moderate Tm
moderate E
• moderate α
Directional Properties
Secondary bonding dominates
small Tm
small E
• large α

50. Example
50
Calculate the force of attraction between a Ca2+ and an O2- ion the centers of
which are separated by a distance of 1.25 nm.
Charge on electron =1.67 x10-19 Coulomb

51. 51
• Bonding energies between 600 and 1500 kJ/mol (or 3 to 8 eV/atom) are
considered to be relatively large with high Tm