atomic structure & interatomic bonding · atomic structure mass charge ... ionization...
TRANSCRIPT
Atomic structure & interatomic
bonding
1
Chapter two
Atomic StructureMass Charge
Proton 1.67 х 10-27 kg + 1.60 х 10 -19 C
Neutron 1.67 х 10-27 kg Neutral
Electron 9.11 х 10-31 kg - 1.60 х 10 -19 C
2
�Electron mass = 1/1836 that of a proton
�Radius of an atom= 0.1 nm = 0.1 x10 -9 m (1Angstrom)
�50,000,000 atoms lined up measure 10mm !!!
�Nucleus takes up 10 -14 of the total volume of atom
and has diameter of 4 -15 fm (femtometer = 10 -15 m)
Precision: How closely measurements of the same quantity come to each other.Accuracy: How close an experimental observation lies to the true value.
� Atomic mass (A) = mass of protons + mass of neutrons
�# of protons gives chemical identification of the element
�# of protons = atomic number (Z)
�# of neutrons (N) defines isotope number
General Notes:
3
�The atomic mass unit (amu) is often used to express
atomic weight. 1 amu is defined as 1/12 of the atomic
mass of the most common isotope of carbon atom that
has 6 protons (Z=6) and six neutrons (N=6).
mproton ≈ mneutron = 1.67 х 10-27 kg = 1 amu.
The atomic mass of the 12C atom is 12 amu.
� Atomic mass (A) ≈ atomic number (Z) + # of neutrons (N)
�The atomic weight is often specified 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).
�Atomic weight = weighted average of the atomic
masses of the atoms naturally occurring isotopes.
Atomic weight of carbon is 12.011 amu.
4
�The number of atoms in a mole is called the Avogadro
number, Nav = 6.023 × 10 23.Nav = 1 gram/1 amu.
mole of carbon has a mass of 12 grams).
� Atomic weight of Fe = 55.85 amu/atom = 55.85 g/mol
�The number of atoms per cm3, n, for material of
density ρ (g/cm3) and atomic mass A (g/mol):
n = Nav × ρ / A
Graphite (carbon): ρ = 2.3 g/cm3, A = 12 g/mol
n = 6.023 × 1023 atoms/mol × 2.3 g/cm3 / 12 g/mol =
11.5 × 10 22 atoms/cm3
Diamond (carbon): ρ = 3.5 g/cm3, A= 12 g/mol
n = 6.023 × 1023 atoms/mol × 3.5 g/cm3 / 12 g/mol =
17.5 × 10 22 atoms/cm3
Examples:
5
17.5 × 10 22 atoms/cm3
Water (H2O) ρ = 1 g/cm3, Aw= 18 g/mol
n = 6.023 × 1023 molecules/mol × 1 g/cm3 / 18 g/mol =
3.3 × 10 22 molecules/cm3
For material with n = 6 × 10 22 atoms/cm3 calculate
mean distance between atoms L .
Atomic Models:
But quantum mechanics
tells us that this analogy
is not correct !!
The electrons form a cloud around the nucleus. This
picture looks like a mini planetary system.
6
is not correct !!
Electrons move not in circular
orbits, but in 'fuzzy‘ orbits.
Actually, we cannot tell how it
moves, but only can say what is
the probability of finding it at some distance from the
nucleus. Only certain “orbits” or shells of electron
probability densities are allowed.
�The shells & electrons are identified by four
quantum number, n, l, ml and ms
�The quantum numbers arise from solution of
Schrodinger’s equation.
�Pauli Exclusion Principle: only one electron
7
Pauli Exclusion Principle: only one electron
can have a given set of the four quantum
numbers.
�Now we can give a short description for the
quantum numbers.
Primary Quantum Number
� n
� Can have values from 1 to infinity, but they
can only be integers
8
� K, L, M, N
� Represents the energy of the orbital, which is
also related to the size of the orbital
� An orbital is the region of space where you
are likely to find the electron
Angular Momentum Quantum
Number� l
� Shape of the orbital
� Can have values from 0 to n-1
9
� Can have values from 0 to n-1
� s, p, d, f, g, h…..
� If there is more than one electron present, the
angular momentum quantum number also
affects the orbital energy (also called the
azimuthal quantum number)
Magnetic Quantum Number
� ml
� Can have integer values from –l to +l
� Thus, if n=1, l =0, and ml must equal 0
10
� Thus, if n=1, l =0, and ml must equal 0
� In other words, it can only have one value
� If n=2, then l can equal either 0 or 1
� If it equals 1, then mlcan equal –1, 0 or +1
� It can have three values
Remember
� s orbitals correspond to l = 0
� p orbitals correspond to l = 1
� d orbitals correspond to l = 2
1
3
5
11
� d orbitals correspond to l = 2
� f orbitals correspond to l = 3
� How many orbitals are possible for each of these types?
7
5
Spin Quantum Number
� ms
� +1/2
� -1/2
12
� -1/2
� Two electrons of opposite spin fill each orbital
� The first three quantum numbers define an orbital
� You need all four to define an electron
Shorthand Notation
13
� Germanium has 32 protons and 32 electrons
� 1s22s22p63s23p63d104s24p2
Shorthand Notation
Electron Shells
� Bonding occurs only with the electrons in the outer most shells – called the valence electrons
Inner electrons are called the core
14
� Inner electrons are called the core electrons
� The valence electrons are those in the outer s and p orbitals, and any unfilledd and p orbitals.
1s22s22p63s23p63d104s24p2
Core electrons Valence
electrons
1s22s22p63s23p63d24s2
15
1s22s22p63s23p63d24s2
Chemistry happens in the valence shell
Fundamental Concepts� First ionization energy (IE): it’s also called
ionization potential, it is the energy required
to remove the most weakly bound electron
from an isolated gaseous atom
16
Atom (g) + IE = positive ion (g) + e-
and can be calculated from the equation:
IE = 13.6 Z2 / n2
Fundamental Concepts
� Electron Affinity (EA): the reverse process to the ionization energy, it is the energy change associated with an isolated gaseous atom accepting one electron
17
accepting one electron
Atom (g) + e- = negative ion (g)
EA : positive if energy released.
: negative if energy required.
Fundamental Concepts
� Atomic and ionic radii : in general, positive ions are smaller than neutral atoms, while negative ions are larger.
18
atoms, while negative ions are larger.
Fundamental Concepts� Electronegativity (χ) : independent measure
for atom attraction to electrons from another atom in a bond forming. It can be calculated from:
(χ) = (IE + EA) / 2
19
(χ) = (IE + EA) / 2
or from:
(χ) = [ {0.31 (n + 1 ± c) } / r ] + 0.5
n: # of valence electrons.
c: any formal valence charge on the atom.
r: covalent radius.
Periodic TableS-block d-block p-block
20
f-block
Trends in The Periodic Table
(IE) (EA)
21
Atomic & ionic radii Electronegativity
Atomic Bonding:
� There are both attractive and repulsive
Binding Energy:
The bond is an electrostatic force that bind atoms
or molecules together.
22
� There are both attractive and repulsive forces acting on atoms
� When they are balanced a bond is formed
� When the total energy of a pair of atoms is minimized, a bond is formed
Interatomic separation Force
Attraction
Repulsion
Net force FN ( FA+FR)
ro
Attraction force FA
Repulsion force FR
23
Interatomic separation
Repulsion force FR
Potential energy; E
Attraction
Repulsion
Attraction energy
Repulsion energy
Net energy (E= min)Eo
Types of bonds:
Primary secondary
Strong weak
Eo : Bonding energy.
r0 : Bond length.
24
Strong weak
Chemical physical
Primary bond is created when there is direct interaction
of electrons between two or more atoms.
Secondary bond occur due to indirect interaction of
electrons in adjacent atoms or molecules.
Electronegativity controls how elements combine
(bond) with each other because it provides a
measure of the excess binding energy between atoms A and B, ∆A-B (in kJ/mol) :
Primary bonds:
25
The excess binding energy is related to the
energy required to separate two bonded atoms,
bond dissociation energy, DEAB :
∆A-B = 96.5 ( χA – χB ) 2
∆A-B = DEAB – [ (DEAA) (DEBB)] 1/2
Types of Primary bonds:Electronegativity difference > 2.0 Ionic bond
Electronegativity difference < 0.4 covalent bond
0.4 < Electronegativity difference < 2.0 Polar
covalent bond
26
covalent bond
Special types of primary bonds is metallic bond
Ionic Bonds
� Metal-Nonmetal
� Cation-anion
� Non-directional
27
� Non-directional
� Poor electrical conductivity
� Poor thermal conductivity
� Ceramics are formed from ionic bonds
� What is a molecule?
Covalent Bonds
� Nonmetal – nonmetal
� Directional bonds
� Poor electrical conductivity
28
� Poor electrical conductivity
� Poor thermal conductivity
� Polymers are covalently bonded
� Compounds
Partial Ionic character:
Polar covalent bond:
A bond neither truly ionic nor totally covalent.
(HF)
29
Partial Ionic character:
Ionic character % = 100 [ 1- exp { -0.25 ( χA – χB )2}]
Metallic Bonds
� Metal-Metal
� Non-directional
� Electrons are free to move around
30
� Electrons are free to move around
� Good electrical conductivity
� Good thermal conductivity
� What is a molecule?
Mixed bonding
� 2 or more metals may form an intermetallic
compound
� A mixture of ionic and metallic bonds
31
� Ceramics (nonmetal – metal) are usually a
mixture of ionic and covalent
� As the electronegativity difference increases,
the bond becomes more ionic
Secondary Bonding
� Metallic compounds and ionic compounds form crystals.
� But how do molecules of covalently
32
� But how do molecules of covalently bonded elements stick together?
� Secondary Bonds
� Van der Waal’s: Dipole-dipole forces.
� Dipole-induced dipole
� London dispersion forces
Types of Secondary Forces
33
� London dispersion forces
� Hydrogen bonding
What is dipole?
Atom or molecule that have some separation
of positive and negative portions
Van der Waal’s
� Interaction between permanent dipoles
� The interaction of permanent dipoles (analogous to magnets but having an
34
(analogous to magnets but having an electrostatic dipole moment).
CH3Cl CH3Cl
Dipole induced dipole
A permanent dipole moment can induce a
dipole in a neighboring molecule in
which the unperturbed centers of positive
and negative charge are otherwise
35
and negative charge are otherwise
coincident
CH4 CH3Cl
London dispersion forces:
•No permanent dipole
•It does have an instantaneous dipole moment
36
•These instantaneous dipoles orient
themselves with their neighbors to give an
overall force of attraction.
H-bonding
� Hydrogen bonding: can be viewed as the interaction between very strong dipoles in - OH or -NH2 groups. It is
37
dipoles in - OH or -NH2 groups. It is reflected in the very high boiling point of water compared with molecules of similar size.
H2O H2O
Bonding Energy
� How does bonding energy relate to melting point?
� Modulus of Elasticity?
38
� Modulus of Elasticity?
� Coefficient of Thermal Expansion?
� Hint: The higher the bonding energy the more tightly the atoms are held together.