Transcript Chapter 2 - MSE 235- Materials Science for Electronics Engineers

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
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 Matter
Physical 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
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
Chapter 2-
1869…
Chapter 2-
Today: Periodic Table of the
Elements
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?
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-
1909 – E. Rutherford
• Tested the Plum Pudding Model.
• Results:
– Majority of a particles transmitted (pass through) or
deflected through small angles
– Tiny fraction deflected through large angles
Chapter 2-
1909 – E. Rutherford
• 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 radiation while
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.
Chapter 2-
1912 – N. Bohr
• Many phenomena involving
electrons in solids could not be
explained in terms of
CLASSICAL MECHANICS.
• We need QUANTUM MECHANICS…
Chapter 2-
Bohr Postulates for the Hydrogen Atom
1.
2.
3.
4.
5.
Rutherford atom is correct
Classical EM theory not applicable to orbiting eNewtonian mechanics applicable to orbiting eEelectron = Ekinetic + Epotential
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-
BOHR ATOM
orbital electrons:
n = principal
quantum number
1
2
n=3
Adapted from Fig. 2.1,
Callister 6e.
Nucleus: Z = # protons
N = # neutrons
Atomic mass A ≈ Z + N
Chapter 2- 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 nucleus
l subshell, describes the shape of the subshell
m number of energy states in a subshell
s spin moment
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 particlelike 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 1924 de Broglie : dual character of electrons
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.
Schrodinger, math expression for the behavior of an electron around an
atom
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
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)
l
ml
ms = ±½
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-
SURVEY OF ELEMENTS
• Most elements: Electron configuration not stable.
Electron configuration
1s1
1s2
(stable)
1s22s1
1s22s2
Adapted from Table 2.2,
1s22s22p1
Callister 7e.
1s22s22p2
...
1s22s22p6
(stable)
1s22s22p63s1
1s22s22p63s2
1s22s22p63s23p1
...
1s22s22p63s23p6
(stable)
...
1s22s22p63s23p63d10 4s246
(stable)
• Why? Valence (outer) shell usually not filled completely.
Chapter 2- 5
STABLE ELECTRON CONFIGURATIONS
Stable electron configurations...
• have complete s and p subshells
• tend to be unreactive.
Adapted from Table 2.2,
Callister 6e.
Chapter 2- 4
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
Chapter 2-
THE PERIODIC TABLE
• Columns: Similar Valence Structure, Similar Properties
Electropositive elements:
Readily give up electrons
to become + ions.
Electronegative elements:
Readily acquire electrons
to become - ions.
Chapter 2- 6
ELECTRONEGATIVITY
• Ranges from 0.7 to 4.0,
• Large values: tendency to acquire electrons; reactivity
Metals like to give up, halogens like to acquire electrons !
Smaller electronegativity
Larger electronegativity
Chapter 2- 7
REVIEW OF ATOMIC STRUCTURE
(FRESHMAN CHEMISTRY)
ATOMS = (PROTONS+NEUTRONS) + ELECTRONS
NUCLEUS
BONDING
• 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 neutrons
Atomic 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
Chapter 2-
Atomic Structure
Valence electrons determine all of the following
properties:
1)
2)
3)
4)
Chemical
Electrical
Thermal
Optical
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
r
• 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
Chapter 2-
IONIC BONDING (I)
•
•
•
•
Occurs between + and – ions (anion and cation).
Requires electron transfer.
Large difference in electronegativity required.
Example: Na+ Cl-
Chapter 2- 8
Ionic bond – metal +
donates
electrons
nonmetal
accepts
electrons
Dissimilar electronegativities
ex: MgO
Mg
1s2 2s2 2p6 3s2
[Ne] 3s2
Mg2+ 1s2 2s2 2p6
[Ne]
O
1s2 2s2 2p4
O2- 1s2 2s2 2p6
[Ne]
Chapter 2-
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,ion
Hard materials, brittle, high melting temperature, electrically and thermally insulating
Chapter 2- 8
Ionic Bonding
• Energy – minimum energy most stable
– Energy balance of attractive and repulsive terms
EN = EA + ER =
-
A
r
-
B
rn
Repulsive energy ER
Interatomic separation r
Net energy EN
Adapted from Fig. 2.8(b),
Callister 7e.
Attractive energy EA
Chapter 2-
Examples: Ionic Bonding
• Predominant bonding in Ceramics
NaCl
MgO
CaF 2
CsCl
Give up electrons
Acquire electrons
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.
Chapter 2-
COVALENT BONDING (I)
• Requires shared electrons
• Example: CH4
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 6e.
Chapter 2- 10
COVALENT BONDING (II)
Diamond, sp3
Covalent bonds are formed by sharing of the valence
electrons
Covalent bonds are very directional
Covalent 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.
Chapter 2- 10
Primary Bonding
• Metallic Bond -- delocalized as electron cloud
• Ionic-Covalent Mixed Bonding
% ionic character =

(X A -X B )2 


4
1e

x (100%)




where XA & XB are Pauling electronegativities
Ex: MgO
XMg = 1.3
XO = 3.5

(3.5 -1.3)2

4
% ionic character  1 - e




 x (100%)  70.2% ionic


Chapter 2-
EXAMPLES: COVALENT BONDING
H2
H
2.1
Li
1.0
Na
0.9
K
0.8
Be
1.5
Mg
1.2
Ca
1.0
Rb
0.8
Cs
0.7
Sr
1.0
Fr
0.7
Ra
0.9
•
•
•
•
Ba
0.9
column IVA
H2O
C(diamond)
SiC
Ti
1.5
Cr
1.6
Fe
1.8
Ni
1.8
Zn
1.8
Ga
1.6
C
2.5
Si
1.8
Ge
1.8
F2
He
O
2.0
As
2.0
Sn
1.8
Pb
1.8
F
4.0
Ne
-
Cl
3.0
Ar
Kr
-
Br
2.8
I
2.5
At
2.2
Cl2
Xe
-
Rn
-
GaAs
Adapted from Fig. 2.7, Callister 6e. (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.
Molecules with nonmetals
Molecules with metals and nonmetals
Elemental solids (RHS of Periodic Table)
Compound solids (about column IVA)
Chapter 2- 11
METALLIC BONDING
• Arises from a sea of donated valence electrons
(1, 2, or 3 from each atom).
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…
Adapted from Fig. 2.11, Callister 6e.
• Primary bond for metals and their alloys
Chapter 2- 12
SECONDARY BONDING
Arises from interaction between dipoles
• Fluctuating dipoles
asymmetric electron
clouds
+
-
+
secondary
bonding
-
ex: liquid H 2
H2
H2
H H
H H
secondary
bonding
Adapted from Fig. 2.13, Callister 7e.
• Permanent dipoles-molecule induced
-general case:
-ex: liquid HCl
-ex: polymer
+
-
H Cl
secondary
bonding
+
secondary
bonding
H Cl
Adapted from Fig. 2.14,
Callister 7e.
secondary bonding
Chapter 2-
Bonding Energies
Chapter 2-
Summary: Bonding
Comments
Type
Bond Energy
Ionic
Large!
Nondirectional (ceramics)
Covalent
Variable
large-Diamond
small-Bismuth
Directional
(semiconductors, ceramics
polymer chains)
Metallic
Variable
large-Tungsten
small-Mercury
Nondirectional (metals)
Secondary
smallest
Directional
inter-chain (polymer)
inter-molecular
Chapter 2-
Properties From Bonding: Tm
• Bond length, r
• Melting Temperature, Tm
Energy
r
• Bond energy, Eo
ro
Energy
r
smaller Tm
unstretched length
ro
r
Eo =
“bond energy”
larger Tm
Tm is larger if Eo is larger.
Chapter 2-
PROPERTIES FROM BONDING: E
• Elastic modulus, E
Elastic modulus
F
L
=E
Ao
Lo
• E ~ curvature at ro
Energy
unstretched length
ro
r
E is larger if Eo is larger.
smaller Elastic Modulus
larger Elastic Modulus
Chapter 2- 16
Summary: Primary Bonds
Ceramics
(Ionic & covalent bonding):
Metals
(Metallic bonding):
Polymers
(Covalent & Secondary):
Large bond energy
large Tm
large E
small a
Variable bond energy
moderate Tm
moderate E
moderate a
Directional Properties
Secondary bonding dominates
small Tm
small E
large a
Chapter 2-