Post on 09-Jul-2018
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Chapter 2-
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
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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 MatterPhysical 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
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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
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Chapter 2-
1869…
Chapter 2-
Today: Periodic Table of the Elements
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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?
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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-
• Results:– Majority of a particles transmitted (pass through) or
deflected through small angles
– Tiny fraction deflected through large angles
1909 – E. Rutherford
• Tested the Plum Pudding Model.
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Chapter 2-
• 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 radiationwhile 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.
1909 – E. Rutherford
Chapter 2-
1912 – N. Bohr
• Many phenomena involving electrons in solids could not be explained in terms of CLASSICAL MECHANICS.
• We need QUANTUM MECHANICS…
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Chapter 2-
Bohr Postulates for the Hydrogen Atom
1. Rutherford atom is correct
2. Classical EM theory not applicable to orbiting e-
3. Newtonian mechanics applicable to orbiting e-
4. Eelectron = Ekinetic + Epotential
5. 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-
Nucleus: Z = # protons
2
orbital electrons: n = principal quantum number
n=3 2 1
= 1 for hydrogen to 94 for plutoniumN = # neutrons
Atomic mass A ≈ Z + N
Adapted from Fig. 2.1, Callister 6e.
BOHR ATOM
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Chapter 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 nucleusl subshell, describes the shape of the subshellm number of energy states in a subshells spin moment
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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 particle-like 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 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.
In 1924 de Broglie : dual character of electrons
Schrodinger, math expression for the behavior of an electron around an atom
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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
Incr
easi
ng e
nerg
y
n=1
n=2
n=3
n=4
1s
2s
3s
2p
3p
4s4p
3d
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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)
mll ms = ±½
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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- 5
• Why? Valence (outer) shell usually not filled completely.
• Most elements: Electron configuration not stable.Element Hydrogen Helium Lithium Beryllium Boron Carbon ... Neon Sodium Magnesium Aluminum ... Argon ... Krypton
Atomic # 1 2 3 4 5 6
10 11 12 13
18 ... 36
Electron configuration 1s 1
1s 2 (stable) 1s 22s 1 1s 22s 2 1s 22s 22p 1 1s 22s 22p 2 ...
1s 22s 22p 6 (stable) 1s 22s 22p 63s 1 1s 22s 22p 63s 2 1s 22s 22p 63s 23p 1 ...
1s 22s 22p 63s 23p 6 (stable) ...
1s 22s 22p 63s 23p 63d 10 4s 24 6 (stable)
SURVEY OF ELEMENTS
Adapted from Table 2.2, Callister 7e.
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Chapter 2- 4
• have complete s and p subshells• tend to be unreactive.
Stable electron configurations...
Z Element Configuration
2 He 1s 2
10 Ne 1s 22s 22p 6
18 Ar 1s 2 2s 22p 63 s 23p 6
36 Kr 1s 2 2s 22p 63 s 23p 63d 10 4 s 24p 6
Adapted from Table 2.2, Callister 6e.
STABLE ELECTRON CONFIGURATIONS
Chapter 2-
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
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Chapter 2-
He
Ne
Ar
Kr
Xe
Rn
iner
t gas
es
acce
pt 1
e
acce
pt 2
e
give
up
1e
gi
ve u
p 2
e
give
up
3e
F Li Be
Metal
Nonmetal
Intermediate
H
Na Cl
Br
I
At
O
S Mg
Ca
Sr
Ba
Ra
K
Rb
Cs
Fr
Sc
Y
Se
Te
Po
6
• Columns: Similar Valence Structure, Similar Properties
Electropositive elements:Readily give up electronsto become + ions.
Electronegative elements:Readily acquire electronsto become - ions.
THE PERIODIC TABLE
Chapter 2- 7
• Ranges from 0.7 to 4.0,
Smaller electronegativity Larger electronegativity
He -
Ne -
Ar -
Kr -
Xe -
Rn -
F 4.0
Cl 3.0
Br 2.8
I 2.5
At 2.2
Li 1.0
Na 0.9
K 0.8
Rb 0.8
Cs 0.7
Fr 0.7
H 2.1
Be 1.5
Mg 1.2
Ca 1.0
Sr 1.0
Ba 0.9
Ra 0.9
Ti 1.5
Cr 1.6
Fe 1.8
Ni 1.8
Zn 1.8
As 2.0
• Large values: tendency to acquire electrons; reactivityMetals like to give up, halogens like to acquire electrons !
ELECTRONEGATIVITY
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Chapter 2-
REVIEW OF ATOMIC STRUCTURE (FRESHMAN CHEMISTRY)
• 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 neutronsAtomic 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
ATOMS = (PROTONS+NEUTRONS) + ELECTRONS
NUCLEUS BONDING
Chapter 2-
Atomic Structure
• Valence electrons determine all of the following properties
1) Chemical
2) Electrical
3) Thermal
4) Optical
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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
• 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
r
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Chapter 2-
Na (metal) unstable
Cl (nonmetal) unstable
electron
+ - Coulombic Attraction
Na (cation) stable
Cl (anion) stable
8
• Occurs between + and – ions (anion and cation).
• Requires electron transfer.
• Large difference in electronegativity required.
• Example: Na+ Cl-
IONIC BONDING (I)
Chapter 2-
Ionic bond – metal + nonmetal
donates acceptselectrons electrons
Dissimilar electronegativities
ex: MgO Mg 1s2 2s2 2p6 3s2 O 1s2 2s2 2p4
[Ne] 3s2
Mg2+ 1s2 2s2 2p6 O2- 1s2 2s2 2p6
[Ne] [Ne]
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Chapter 2- 8
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,ionHard materials, brittle, high melting temperature, electrically and thermally insulating
Chapter 2-
Ionic Bonding
• Energy – minimum energy most stable– Energy balance of attractive and repulsive terms
Attractive energy EA
Net energy EN
Repulsive energy ER
Interatomic separation r
rA
nrBEN = EA + ER =
Adapted from Fig. 2.8(b), Callister 7e.
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Chapter 2-
• Predominant bonding in Ceramics
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.
Examples: Ionic Bonding
Give up electrons Acquire electrons
NaCl
MgO
CaF2CsCl
Chapter 2- 10
• Requires shared electrons
• Example: CH4
C: has 4 valence e,needs 4 more
H: has 1 valence e,needs 1 more
Electronegativitiesare comparable.
shared electrons from carbon atom
shared electrons from hydrogen atoms
H
H
H
H
C
CH4
Adapted from Fig. 2.10, Callister 6e.
COVALENT BONDING (I)
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Chapter 2- 10
COVALENT BONDING (II)
Covalent bonds are formed by sharing of the valence electronsCovalent bonds are very directionalCovalent 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.
Diamond, sp3
Chapter 2-
Primary Bonding
• Metallic Bond -- delocalized as electron cloud
• Ionic-Covalent Mixed Bonding
% ionic character =
where XA & XB are Pauling electronegativities
%)100(x
1e
(XAXB)2
4
ionic 70.2% (100%) x e1 characterionic % 4)3.15.3(
2
Ex: MgO XMg = 1.3XO = 3.5
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Chapter 2- 11
• Molecules with nonmetals• Molecules with metals and nonmetals• Elemental solids (RHS of Periodic Table)• Compound solids (about column IVA)
He -
Ne -
Ar -
Kr -
Xe -
Rn -
F 4.0
Cl 3.0
Br 2.8
I 2.5
At 2.2
Li 1.0
Na 0.9
K 0.8
Rb 0.8
Cs 0.7
Fr 0.7
H 2.1
Be 1.5
Mg 1.2
Ca 1.0
Sr 1.0
Ba 0.9
Ra 0.9
Ti 1.5
Cr 1.6
Fe 1.8
Ni 1.8
Zn 1.8
As 2.0
Si C
C(diamond)
H2O
C 2.5
H2
Cl2
F2
Si 1.8
Ga 1.6
GaAs
Ge 1.8
O 2.0
colu
mn
IVA
Sn 1.8Pb 1.8
Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 isadapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.
EXAMPLES: COVALENT BONDING
Chapter 2- 12
• Arises from a sea of donated valence electrons(1, 2, or 3 from each atom).
• Primary bond for metals and their alloys
+ + +
+ + +
+ + +Adapted from Fig. 2.11, Callister 6e.
METALLIC BONDING
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…
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Chapter 2-
Arises from interaction between dipoles
• Permanent dipoles-molecule induced
• Fluctuating dipoles
-general case:
-ex: liquid HCl
-ex: polymer
Adapted from Fig. 2.13, Callister 7e.
Adapted from Fig. 2.14,Callister 7e.
SECONDARY BONDING
asymmetric electronclouds
+ - + -secondary bonding
HH HH
H2 H2
secondary bonding
ex: liquid H2
H Cl H Clsecondary bonding
secondary bonding+ - + -
secondary bonding
Chapter 2-
Bonding Energies
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Chapter 2-
Type
Ionic
Covalent
Metallic
Secondary
Bond Energy
Large!
Variablelarge-Diamondsmall-Bismuth
Variablelarge-Tungstensmall-Mercury
smallest
Comments
Nondirectional (ceramics)
Directional(semiconductors, ceramicspolymer chains)
Nondirectional (metals)
Directionalinter-chain (polymer)inter-molecular
Summary: Bonding
Chapter 2-
• Bond length, r
• Bond energy, Eo
• Melting Temperature, Tm
Tm is larger if Eo is larger.
Properties From Bonding: Tm
r o r
Energyr
larger Tm
smaller Tm
Eo =
“bond energy”
Energy
r o runstretched length
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Chapter 2- 16
• Elastic modulus, E
• E ~ curvature at ro
cross sectional area A o
L
length, Lo
F
undeformed
deformed
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.
PROPERTIES FROM BONDING: E
Chapter 2-
• Coefficient of thermal expansion,
• ~ symmetry at ro
is larger if Eo is smaller.
Properties From Bonding :
= (T2-T1)LLo
coeff. thermal expansion
L
length, Lo
unheated, T1
heated, T2
r or
larger
smaller
Energy
unstretched length
E
oE
o
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Chapter 2-
Ceramics(Ionic & covalent bonding):
Metals(Metallic bonding):
Polymers(Covalent & Secondary):
Large bond energylarge Tm
large Esmall
Variable bond energymoderate Tm
moderate Emoderate
Directional PropertiesSecondary bonding dominates
small Tm
small Elarge
Summary: Primary Bonds