Transcript Chapter 2

Today's Agenda
ISSUES TO ADDRESS...
• What promotes bonding?
• What types of bonds are there?
• What properties are inferred from bonding?
Chapter 2- 1
GECKO
WHY ?
Chapter 2-
Here is the answer
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-
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-
AMU and Mole Concept
The atomic mass unit (amu) :
1 amu is defined as the 1/12 of the atomic mass of the most common isotope of carbon,
carbon 12 (12C6).
Atomic mass of 12C6 is 12 amu : Carbon= 6 protons (Z=6) + 6 neutrons (N=6)
Mproton ~ Mneutron = 1.67 x 10-24 g = 1 amu
Atomic Mass = Z + N
The atomic weight is often expressed 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).
The number of atoms in a mole is called the Avogadro number, Nav = 6.023 × 1023.
Nav = 1 gram/1 amu.
• Example:
Atomic weight of iron = 55.85 amu/atom = 55.85 g/mol
Chapter 2-
MOLE CONCEPT
ANY IDEAS WHY ?
Chapter 2-
Atomic Structure
• Valence electrons determine all of the
following properties
1)
2)
3)
4)
Chemical
Electrical
Thermal
Optical
Chapter 2-
Atomic Models
• Towards the end of 19th century physicists realized Newtonian physics
has serious difficulty in explaining many phenomena involving
electrons => quantum mechanics
– Bohr atomic model
• Electrons assume very well defined orbits around the nucleus (protons +
neutrons)
• Electrons in each shell orbit assumes the same energy level
– Severe issues when considering events involving electrons such as emission spectra
and photoelectrons…)
– See figure 2.1 in Callister page 18 (page 13 for 6th ed.)
– Wave mechanical model
• Electrons in an atom or molecule are permitted to have only specific values of
energy, energy is quantized…
• Electrons do not move in circular orbits but in “fuzzy” orbits. At any given
time we can only talk about the probability of finding an electron at a radius
from the orbit.
• Every electron is characterized by four quantum numbers. The size, shape,
spatial orientation of an electrons probability density are specified by these
numbers
Chapter 2-
BOHR ATOM
orbital electrons:
n = principal
quantum number
1
2
n=3
Adapted from Fig. 2.1,
Callister 6e.
Nucleus: Z = # protons
= 1 for hydrogen to 94 for plutonium
N = # neutrons
Atomic mass A ≈ Z + N
Chapter 2- 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-
ELECTRON ENERGY STATES
Electrons...
• have discrete energy states
• tend to occupy lowest available energy state.
Adapted from Fig. 2.5,
Callister 6e.
Chapter 2- 3
Quantum Numbers (I)
• A product of Schrodinger’s Equation
– n, l , ml , ms
• n principal quantum number, distance of an electron
from the nucleus
• l subshell, describes the shape of the subshell
• ml number of energy states in a subshell
• ms spin moment
• Pauli’s exclusion principle: only one
electron can have a given set of four
quantum numbers.
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
Periodic Table
Draft of the first periodic table, Mendeleev, 1869
Chapter 2-
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
Concept Checks
Question: Why are the atomic weights of the elements generally not integers?
Cite two reasons.
Answer: The atomic weights of the elements ordinarily are not integers because:
(1) The atomic masses of the atoms normally are not integers (except for 12C),
and
(2) the atomic weight is taken as the weighted average of the atomic masses of
an atom's naturally occurring isotopes.
Question: Give electron configurations for the Fe3+and S2- ions.
Answer: The Fe3+ ion is an iron atom that has lost three electrons. Since the
electron
configuration of the Fe atom is 1s22s22p63s23p63d64s2 (Table 2.2), the
configuration for Fe3+ is 1s22s22p63s23p63d5.
The S2- ion a sulfur atom that has gained two electrons. Since the electron
configuration of the S atom is 1s22s22p63s23p4 (Table 2.2), the configuration for
S2- is 1s22s22p63s23p6.
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-
COVALENT BONDING (III)
Very few materials have pure ionic or covalent bonding; electronegativity
inpart defines how much time electrons spend between ion cores…
Chapter 2- 10
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
Properties From Bonding : a
• Coefficient of thermal expansion, a
length, L o
coeff. thermal expansion
unheated, T1
L
= a(T2 -T1)
Lo
L
heated, T 2
• a ~ symmetry at ro
Energy
unstretched length
ro
E
o
E
o
r
a is larger if Eo is smaller.
smaller a
larger a
Chapter 2-
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-
ANNOUNCEMENTS
Read Chapter 3!!!
Chapter 2- 0