1. PDT 153 Materials Structure And Properties
Chapter 2: Atomic Structure and Interatomic Bonding
Prepared by: Dr. Tan Soo Jin

2. Basic Unit of an Element
Structure of Atoms
ATOM
Basic Unit of an Element
Diameter : 10 –10 m.
Neutrally Charged
Nucleus
Diameter : 10 –14 m
Accounts for almost all mass
Positive Charge
Electron Cloud
Mass : 9.11 x 10 –31 kg
Charge : x 10 –19 C
Accounts for all volume
Proton
Mass : 1.67 x 10 –27 kg
Charge : 1.60 x 10 –19 C
Neutron
Neutral Charge

3. Animation of an atom http://www.kscience.co.uk/animations/atom.htm
Atomic Structure
Each atom consists of a very small nucleus composed of protons and neutrons, which is encircled by moving electrons.
Animation of an atom

4. Atomic Numbers (Z)
The atomic number (Z) of an atom indicates the number of protons (positively charged particles) that are in the nucleus.
For electrically neutral atom, the atomic number equals to the number of electron.
The atomic number ranges in integral units from 1 for hydrogen to 92 for uranium, the highest naturally occurring elements.

5. Atomic Mass (A)
Sum of the masses of protons and neutrons within nucleus.
Although the number of protons is the same for all atoms of a given element, the number of neutrons may be variable (isotopes different atomic masses).
Example: Carbon has 6 Protons and 6 Neutrons. Atomic Mass = 12.
The relative atomic mass of an element is the mass in grams of x 1023 atoms (Avogadro’ number, NA) of that element.
One atomic mass unit (amu) is 1/12th of mass of a carbon atom.
One molar relative atomic mass of a carbon 12 has a mass of 12 g on this scale.
1 amu/atom (or molecule) = 1 g/mol
For example: 1 gram-mole of aluminium has a mass of g and contains x 1023 atoms.

7. Example Problem Number of gram moles of Cu =
A 100 gram alloy of nickel and copper consists of 75 wt% Cu and 25 wt% Ni. What are percentage of Cu and Ni Atoms in this alloy?
Given: 75g Cu Atomic Weight 63.54
25g Ni Atomic Weight 58.69
Number of gram moles of Cu =
Number of gram moles of Ni =
Atomic Percentage of Cu =
Atomic Percentage of Ni =

8. Electrons In Atoms (Atomic Models)
Bohr atomic model- in which electrons are assumed to revolve around the atomic nucleus in discrete orbitals, the position of any particular electron is more or less well defined in terms of its orbitals.

9. The Electronic Structure of Atoms
Electrons occupy discrete energy levels within the atom.
Each electron possesses a particular energy, with no more than two electrons in each atom having the same energy.
Electron may change energy by making a quantum jump either to an allowed higher energy (with absorption of energy) or to a lower energy (with emission of energy)

10. Bohr Model of Hydrogen Atom
Atomic orbital represent the energy state of an electron around the nucleus.
Electron is excited to a higher energy level, energy is absorbed;
Electron is dropped to a lower energy level, energy is emitted;
Energy change associated with a photondue to transition , ΔE =
Absorb Energy
(Photon)
Emit Energy
(Photon)
h=Planks Constant
= 6.63 x J.s
c= Velocity of light
= 3.00 x 108 m/s
λ = Wavelength of light
Energy in the form of electromagnetic radiation called photon
Energy levels

11. Principal quantum numbers
Bohr Model of Hydrogen Atom (Continue..)
Hydrogen atom has one proton and one electron
Energy of hydrogen atoms for different energy levels is given by
(n=1,2…..)
Example: If an electron undergoes transition from n=3 state to n=2 state, the energy of photon emitted is
Energy required to completely remove an electron from hydrogen atom is known as ionization energy.
Principal quantum numbers

12. Quantum Numbers
Every electron in an atom is characterized by four parameters called quantum numbers.
Four quantum numbers:
The Principal Quantum Number n
The Subsidiary Quantum Number l
The Magnetic Quantum Number ml
Electron Spin Quantum Number ms
The size, shape, and spatial orientation of an electron’s probability density are specified by these quantum numbers

13. Principal Quantum Number n Subsidiary Quantum Number l
Quantum Numbers of Electrons of Atoms
Principal Quantum Number n
Represents main energy levels for electron and shell.
Range of n = 1-7.
Larger the ‘n’ higher the energy.
Subsidiary Quantum Number l
Represents sub energy levels within the main energy level (subshell).
Range of l = 0, 1, 2,3….n-1.
Represented by letters s, p, d and f (orbitals).
s orbital (l=0)
n=2
n=1
p Orbital
(l=1)

14. Electron Spin Quantum Number ms.
Quantum Numbers of Electrons of Atoms
Magnetic Quantum Number ml.
Represents spatial orientation of single atomic orbital and has little effect on the energy of an electron.
Permissible values are –l to +l, including zero.
Example: if l=1,
ml = -1,0,+1.
i.e. 2l+1 allowed values.
maximum of 1s orbital, 3p orbitals, 5d orbitals, and 7f orbitals for each allowed s, p, d, and f subenergy level.
Electron Spin Quantum Number ms.
Specifies two allowed spin directions for an electron spinning on its own axis.
Directions are clockwise or anticlockwise.
Allowed values are +1/2 or –1/2.
Two electrons may on the same orbital but must have opposed spins.
Minor effect on energy.

15. Electronic Structure of Multielectron Atoms
Maximum number of electrons in each atomic shell is given by 2n2.
Atomic size (radius) increases with addition of shells.
Electron Configuration lists the arrangement of electrons in orbitals.
Example:
1s2 2s2 2p6 3s2
For Iron, (Z=26), Electronic configuration is
1s2 2s2 sp6 3s2 3p6 3d6 4s2
Orbital letters
Number of Electrons
Principal Quantum Numbers

16. Electronic Structure and Chemical Reactivity
Noble Gases
Chemical properties of the atoms of the elements depend on the reactivity of their outermost electrons (valence electron).
Except Helium, most noble gasses (Ne, Ar, Kr, Xe, Rn) are chemically very stable
All have s2 p6 configuration for outermost shell (high chemical stability).
Helium has 1s2 configuration
Electropositive and Electronegative Elements
Electropositive elements are metallic in nature and give up electrons to produce positive ions, cations (positive oxidation number).
Cations are indicated by positive oxidation numbers
Example: Fe : 1s2 2s2 2p6 3s2 3p6 3d6 4s2
Fe2+ : 1s2 2s2 2p6 3s2 3p6 3d6
Fe3+ : 1s2 2s2 2p6 3s2 3p6 3d5

17. Electronic Structure and Chemical Reactivity
Electropositive and Electronegative Elements
Electronegative elements accept electrons during chemical reaction to produce negative ions, anions.
Anions are indicated by negative oxidation number.
Some elements behave as both electronegative and electropositive (carbon, silicon, phosphorous…).
Electronegetivity
Electronegativity is the degree to which the atom attracts electrons to itself.
Measured on a scale of 0 to 4.1
Example : Electronegativity of Fluorine is 4.1
Electronegativity of Sodium is 1. 1 2 3 4 K Na N O Fl W Te Se H
Electro-
positive
negative

18. The Periodic Table Elements Group in Periodic Table
Inert Gases (He, Ne, Ar…)
Group VIIIA
Halogens (F, Cl, Br, I…)
Group VIIA
Alkali and alkaline earth metal (Li, Na, K, Be, Mg…)
Group IA and IIA
Transition metals (Fe, Zn, Cu…)
Group IIIB through IIB
Electronegativity increases in moving from left to right and from bottom to top.
Atoms are more likely to accept electron if their outer shells are almost full, and if they are less “shielded” from (closer to) the nucleus.
Periodic table gives useful information about formulas, atomic numbers, and atomic masses of elements.

20. Primary Interatomic Bonds: Ionic Bonding
Ionic bonds: relatively large interatomic forces are set up in this type of bonding by an electron transfer from one atom to another to produce ions that are bonded together by coulombic forces (attraction of positively and negatively charged ions).
Relatively strong nondirectional bond.
Can form between highly electropositive (metallic) elements and highly electronegative (nonmetallic) elements.
Electropositive
Element
Electronegative
Atom
Electron
Transfer
Cation
+ve charge
Anion
-ve charge
IONIC BOND
Electrostatic
Attraction

21. Example of Ionic Bonding
Sodium
Atom Na
Chlorine Cl
Sodium Ion
Na+
Chlorine Ion
Cl - I O N C B D

22. Ionic Force for Ion Pair
Nucleus of one ion attracts electron of another ion.
The electron clouds of ion repulse each other when they are sufficiently close.
Force versus separation
Distance for a pair of
oppositely charged ions

23. Ionic Force for Ion Pair (Continue..)
Z1,Z2 = Number of electrons removed or added during ion formation
e = Electron Charge
a = Interionic seperation distance
ε = Permeability of free space (8.85 x 10-12c2/Nm2)
(n and b are constants)

24. Example of Interionic Force
Force of attraction between Na+ and Cl- ions
Z1 = +1 for Na+, Z2 = -1 for Cl-
e = 1.60 x C , ε0 = 8.85 x C2/Nm2
a0 = Sum of Radii of Na+ and Cl- ions
= nm nm = 2.76 x m
Na+
Cl- a0

25. Primary Interatomic Bonds: Covalent Bonding
Relatively large interatomic forces are created by the sharing of electrons to form a bond with a localized direction.
Two atoms that are covalently bonded will each contribute at least one electron to the bond, and the shared electrons may be considered to belong to both atoms.
The covalent bond is directional and may be very strong.

26. Example of Covalent Bonding
In case of F2, O2 and N2, covalent bonding is formed by sharing p electrons
Fluorine gas (Outer orbital – 2s2 2p5) share one p electron to attain noble gas configuration.
Oxygen (Outer orbital - 2s2 2p4) atoms share two p electrons
Nitrogen (Outer orbital - 2s2 2p3) atoms share three p electrons

27. Primary Interatomic Bonds: Metallic Bonding
Occurs in solid metals.
Solid metals are visualized as consisting of positive-ion cores (atoms without their valence electrons) and of valence electrons dispersed in the form of an electron cloud that covers a large extent of space.
Metallic bonding is found in the periodic table for Group IA and IIA.

28. Secondary Bonding/ van der Waals Bonding
Physical bonds are wear in comparison to the primary or chemical bonding; bonding energies are typically on the order of only 10 kJ/mol (0.1 eV/atom).
Exists between virtually all atoms or molecules
Secondary bonding forces arise from atomic or molecule dipoles.
An electric dipole exists whenever there is some separation of positive and negative portions of an atom or molecule.
Dipole interactions occur between induced
dipoles and polar molecules
(which has permanent dipoles).
Hydrogen bonding, a special type
of secondary bonding.

29. Fluctuating Dipoles
Weak secondary bonds in noble gasses.
Dipoles are created due to asymmetrical distribution of electron
charges.
Electron cloud charge changes with time.
Symmetrical
distribution
of electron charge
Asymmetrical
Distribution
(Changes with time)

30. CH4 CH3Cl Permanent Dipoles
Dipoles that do not fluctuate with time are called Permanent dipoles.
Examples:
CH4
No Dipole
moment
CH3Cl
Asymmetrical
Tetrahedral
arrangement
Creates
Dipole
Symmetrical
Arrangement
Of 4 C-H bonds

31. Molecules
Composed of groups of atoms that bound together by strong covalent bonds; these include elemental diatomic molecules (F2, O2,H2, etc) as well as a host of compounds (H2O, CO2, HNO3, CH4, etc).
In the condensed liquid and solid states, bonds between molecules are weak secondary bonding.
Therefore, molecular materials have relatively low melting and boiling temperatures.
Hydrogen Bond