Transcript Atomic Structure, Eelectronic Bonding, Periodicity, orbitals

Excited Atoms
& Atomic Structure
The Quantum Mechanical
Picture of the Atom
1.
2.
Basic Postulates of Quantum Theory
Atoms and molecules can exist only in certain energy states. In each energy
state, the atom or molecule has a definite energy. When an atom or
molecule changes its energy state, it must emit or absorb just enough energy
to bring it to the new energy state (the quantum condition).
Atoms or molecules emit or absorb radiation (light) as they change their
energies. The frequency of the light emitted or absorbed is related to the
energy change by a simple equation.
E  h 
hc

 mc 2
The Quantum Mechanical Picture of the Atom
The allowed energy states of atoms and molecules can be described by sets of
numbers called quantum numbers.
•
Quantum numbers are the solutions of the Schrodinger, Heisenberg &
Dirac equations.
Schroedinger 3-dimensional
time independent equation

(r , t )



 2  ( r , t )  V ( r , t )  ( r , t )  i
2m
t
 2
Heisenberg’s uncertainty
Equation
Dirac’s quantum
mechanical model
E. Schrodinger
1887-1961
• Four quantum numbers are necessary to describe
energy states of electrons in atoms
– n, , m, ms
W. Heisenberg
1901-1976
Quantum Numbers – n
1.
The Principal quantum number has the symbol – n.
n = 1, 2, 3, 4, ...... “shells”
n = K, L, M, N, ......
The electron’s energy depends principally on n and tells the average relative
distance of the electron from the nucleus.
– As n increases for a given atom, so does the average distance of the electrons from
the nucleus.
– Electrons with higher values of n are easier to remove from an atom.
n=1
n=2
n=3
n=4
n=5
n=6
n=7
Quantum Numbers – 
2. The azimuthal quantum number
has the symbol .
  describes the shape of the region of space occupied by the electron
 When linked with n defines the energy of the electron, All wave functions that have the
same value of both n and l form a subshell
= 0
=s
 = 0, 1, 2, 3, 4, 5, .......(n-1)
 = s, p, d, f, g, h, .......(n-1)
n=1
n=2
n=3
n=4
=2
=d
n=5
n=6
n=7
=3
=f
= 1
=p
Quantum Numbers – m
3. The symbol for the magnetic quantum number is m.
m = -  , (-  + 1), (-  +2), .....0, ......., ( -2), ( -1), 
• If  = 0 (or an s orbital), then m = 0.
– There is only 1 value of m. Thus there is one s orbital per n value. n  1
• If  = 1 (or a p orbital), then m = -1,0,+1.
– There are 3 values of m. Thus there are three p orbitals per n value. n  2
• If  = 2 (or a d orbital), then m = -2,-1,0,+1,+2.
– There are 5 values of m. Thus there are five d orbitals per n value. n  3
• If  = 3 (or an f orbital), then m = -3,-2,-1,0,+1,+2, +3.
– There are 7 values of m. Thus there are seven f orbitals per n value, n
– Theoretically, this series continues on to g,h,i, etc. orbitals.
• Practically speaking atoms that have been discovered or made up to this point
in time only have electrons in s, p, d, or f orbitals in their ground state
configurations.
• Each wave function with an allowed combination of n, l, and ml values describes
an atomic orbital, a particular spatial distribution for an electron
• For a given set of quantum numbers, each principal shell contains a fixed number
of subshells, and each subshell contains a fixed number of orbitals
Atomic Orbitals
•
•
Atomic orbitals are regions of space
where the probability of finding an
electron about an atom is highest.
s orbital properties:
• There is one s orbital per n level.
• =0
• 1 value of m
= 0
=s
n=1
n=2
n=3
n=4
n=5
n=6
n=7
© 2006 Brooks/Cole - Thomson
Atomic Orbitals
•
•
•
•
p orbitals are peanut or dumbbell shaped.
They are directed along the axes of a Cartesian coordinate system.
The first p orbitals appear in the n = 2 shell.
There are 3 p orbitals per n level.
– The three orbitals are named px, py, pz.
– They have an  = 1.
– m = -1,0,+1 3 values of m
Atomic Orbitals
• d orbital shapes
Atomic Orbitals
• f orbital shapes
Quantum Numbers
Quantum Numbers – ms
4.
–
The last quantum number is the spin quantum number which has the
symbol ms.
The spin quantum number only has two possible values.
» ms = +1/2 or -1/2
The Periodic Table and
Electron Configurations
Note that the 3d subshell is higher in energy than the
4s subshell so appears in the 4th period
Day 1
Building up the Periodic Table
•The Nucleus:
•The Aufbau Process
– Used to construct the periodic table
– First, Build by adding the required number of protons (the atomic number) and neutrons (the
mass of the atom)
– Second, Determine the number of electrons in the atoms then add electrons one at a time to
the lowest-energy orbitals available without violating the Pauli principle
Electrons:
Hund’s Rule states that each orbital will be filled singly before pairing begins.
The
singly filled orbitals will have a parallel spin.
– Each of the orbitals can hold two electrons, one with spin up , which is written first,
and one with spin down 
– A filled orbital is indicated by , in which the electron spins are paired
– The electron configuration is written in an abbreviated form, in which the occupied
orbitals are identified by their principal quantum n and their value of l (s, p, d, or f), with
the number of electrons in the subshell indicated by a superscript
Pauli’s Exclusion Principle states that paired electrons in an orbital will have
opposite spins.
Neon -
2p   
2s 
1s 
Building up the Periodic Table
• Valence electrons
– It is tedious to keep copying the configurations of the filled inner subshells
– The notation can be simplified by using a bracketed noble gas symbol to
represent the configuration of the noble gas from the preceding row
– Electrons in filled inner orbitals are closer and are more tightly bound to
the nucleus and are rarely involved in chemical reactions Now we can write a
complete set of quantum numbers for all of the electrons in these three elements as
examples.
• Na
• First for 11Na.
– When completed there must be one set of 4 quantum numbers for each of the 11
electrons in Na (remember Ne has 10 electrons)
3s
11 Na
Ne
3p
Configurat ion
Ne 3s1

[Ne] = 1s22s22p6
Building up the Periodic Table
11Na
1s22s22p63s1
n

m
1
0
0
2 nd e - 1
0
0
3rd e -
2
0
0
4 th e -
2
0
0
5th e -
2
1
-1
6 th e -
2
1
0
7 th e -
2
1
1
8th e -
2
1
1
9 th e -
2
1
0
10 th e -
2
1
1
11th e -
3
0
0
1st e -
ms
 1/2 
1 s electrons
 1/2 
 1/2 
2 s electrons
 1/2 
 1/2 

 1/2 
 1/2 

2 p electrons
 1/2 
 1/2 

 1/2 

 1/23 s electron
Electron Configuration of the Elements
l=0 l=1 l=2 l=3
Periodic Trends in Atomic
Radii
In the periodic table, atomic radii
decrease from left to right across a row
because of the increase in effective
nuclear charge due to poor electron
screening by other electrons in the
same principal shell. Atomic radii
increase from top to bottom down a
column because the effective nuclear
charge remains constant as the
principal quantum number increases.
Ionization Energies
• There are two reasons for It takes more energy to remove the second
electron from an atom than the first, and so on. :
1. The second electron is being removed from a positively charged species
rather than a neutral one, so more energy is required.
2. Removing the first electron reduces the repulsive forces among the
remaining electrons, so the attraction of the remaining electrons to the
nucleus is stronger.
Electron Affinities
Electron affinity (EA) of an
element E is defined as the energy
change that occurs when an
electron is added to a gaseous atom:
E (g) + e--  E—(g)
energy change = EA.
Electron Affinities of Some Elements
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
Electron Affinity (kJ/mol)
• Electron affinities can be negative
(in which case energy is released
when an electron is added) or
positive (in which case energy must
be added to the system to produce
an anion) or zero (the process is
energetically neutral).
• Halogens have the most negative
electron affinities.
0
-50
-100
-150
-200
-250
-300
-350
-400
Atomic Number
Ionic Radii and
Isoelectronic Series
When one or more electrons is
removed from a neutral atom, two
things happen:
1. Repulsions between electrons
in the same principal shell
decrease because fewer
electrons are present.
2. The effective nuclear charge
felt by the remaining electrons
increases because there are
fewer electrons to shield one
another from the nucleus.
Electronegativity
Increases
Increases
The tendency of an element to gain or lose electrons is is called
electronegativity (), defined as the relative ability of an atom to
attract electrons to itself in a chemical compound.
An Overview of Chemical Bonding
• Chemical bond — the force that holds atoms together in a chemical
•
•
•
•
compound
Covalent bonding — electrons are shared between atoms in a molecule
or polyatomic ion
Ionic bonding — positively and negatively charged ions are held
together by electrostatic forces
Ionic compounds — dissolve in water to form aqueous solutions that
conduct electricity
Covalent compounds — dissolve to form solutions that do not conduct
electricity
• Chemical bonding
1. Ionic — one or more electrons are transferred completely from one atom to
another, and the resulting ions are held together by purely electrostatic
forces
2. Covalent — electrons are shared equally between two atoms
3. Polar covalent — electrons are shared unequally between the bonded atoms
4. Polar bond — bond between two atoms that possess a partial positive
charge (õ+) and a partial negative charge (õ-)
The Continuous Range of Bonding Types
•
Covalent and ionic bonding represent two extremes.
1.
2.
•
•
Most compounds fall somewhere between these two extremes.
All bonds have some ionic and some covalent character.
–
•
In pure covalent bonds electrons are equally shared by the atoms.
In pure ionic bonds electrons are completely lost or gained by one of the
atoms.
For example, HI is about 17% ionic
The greater the electronegativity differences the more polar the bond.
Day 2
OXIDATION NUMBERS
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1.
2.
3.
4.
5.
6.
Guidelines for assigning oxidation numbers.
The oxidation number of any free, uncombined element is zero.
The oxidation number of an element in a simple (monatomic) ion is the charge on the ion.
In the formula for any compound, the sum of the oxidation numbers of all elements in the
compound is zero.
In a polyatomic ion, the sum of the oxidation numbers of the constituent elements is equal to
the charge on the ion.
Fluorine has an oxidation number of –1 in its compounds.
Hydrogen, H, has an oxidation number of +1 unless it is combined with metals, where it has
the oxidation number -1.
–
7.
Oxygen usually has the oxidation number -2.
–
–
8.
Examples – LiH, BaH2
Exceptions:
In peroxides O has oxidation number of –1.
•
Examples - H2O2, CaO2, Na2O2
In OF2, O has oxidation number of +2. Use the periodic table to help with assigning oxidation
numbers of other elements.
1. IA metals have oxidation numbers of +1.
2. IIA metals have oxidation numbers of +2.
3. IIIA metals have oxidation numbers of +3.
4. VA elements have oxidation numbers of –3 in binary compounds with H, metals or NH4+.
5. VIA elements below O have oxidation numbers of –2 in binary compounds with H,
metals or NH4+.
OXIDATION NUMBERS
NH3
ClOH3PO4
MnO4Cr2O72C3H8
Formal Charges
― The most likely formula for a molecule or ion is usually the one in
which the formal charge on each atom is zero or as near zero as
possible
― Negative formal charges are more likely to occur on the more
electronegative elements
― Lewis dot formulas in which adjacent atoms have formula charges
of the same sign are usually not accurate
Cl=N-O
O=N-Cl
H
+
H N
H
H
∙∙
H O
H
Formal Charges
H
H B
-
H
Cl Al Cl
Cl
H
+
H
∙∙
H N
H
H
H
∙∙
S
∙∙
H
Covalent Bonding
•
•
•
•
Covalent bonds are formed when atoms share electrons.
If the atoms share 2 electrons a single covalent bond is formed.
If the atoms share 4 electrons a double covalent bond is formed.
If the atoms share 6 electrons a triple covalent bond is formed.
– The attraction between the electrons is electrostatic in nature
• The atoms have a lower potential energy when bound.
• Covalent bonds in which the electrons are shared equally are designated
as nonpolar covalent bonds.
– Nonpolar covalent bonds have a symmetrical charge distribution.
• To be nonpolar the two atoms involved in the bond must be the same
element to share equally.
Polar and Nonpolar Covalent Bonds
• Covalent bonds in which the electrons are not shared equally are designated
as polar covalent bonds
– Polar covalent bonds have an asymmetrical charge distribution
• To be a polar covalent bond the two atoms involved in the bond must have
different electronegativities.
• Some examples of polar covalent bonds.
• HF
H
F
Electroneg ativities 2.1
4.0



1.9
Difference  1.9
very polar bond
Polar and Nonpolar Covalent Bonds
• Compare HF to HI.
H
I
2.5
Electroneg ativities 2.1
2.1
2.5



0.4
Difference  0.4 slightly polar bond
Polar and Nonpolar Covalent Bonds
• Polar molecules can be attracted by magnetic and electric fields.
Dipole Moments
• Molecules whose centers of positive and negative charge do not coincide, have an
asymmetric charge distribution, and are polar.
– These molecules have a dipole moment.
• The dipole moment has the symbol .
•  is the product of the distance,d, separating charges of equal magnitude and opposite
sign, and the magnitude of the charge, q.


  H - F -
  H -I -
1.91 Debye units
•
•
0.38 Debye units
There are some nonpolar molecules that have polar bonds.
There are two conditions that must be true for a molecule to be polar.
1. There must be at least one polar bond present or one lone pair of electrons.
2. The polar bonds, if there are more than one, and lone pairs must be arranged
so that their dipole moments do not cancel one another.
Molecular Dipole Moments
• In complex molecules that contain polar covalent bonds, the
three-dimensional geometry and the compound’s symmetry
determine if there is a net dipole moment
• Mathematically, dipole moments are vectors; they possess both a
magnitude and a direction
• Dipole moment of a molecule is the vector sum of the dipole
moments of the individual bonds in the molecule
• If the individual bond dipole moments cancel one another,
there is no net dipole moment
• Molecular structures that are highly symmetrical (tetrahedral
and square planar AB4, trigonal bipyramidal AB5, and
octahedral AB6) have no net dipole moment; individual bond
dipole moments completely cancel out
• In molecules and ions that have V-shaped, trigonal pyramidal,
seesaw, T-shaped, and square pyramidal geometries, the bond
dipole moments cannot cancel one another and they have a
nonzero dipole moment
• Polar Molecules must meet two requirements:
1.
One polar bond or one lone pair of electrons on central atom.
2.
Neither bonds nor lone pairs can be symmetrically arranged that their polarities cancel.
Polar Covalent Bonds
© 2006 Brooks/Cole - Thomson
Polar Covalent Bonds
Ionic
Polar Covalent
Covalent
Determine Inductive effect
Count the number of electrons the element should have
Determine how equally electrons are shared (DEN) >1.7 consider it ionic
Oxidation number
Formal charge
Day 3
Writing Lewis Formulas: The Octet Rule
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•
•
•
•
•
•
•
The octet rule states that representative elements usually attain stable noble gas electron
configurations in most of their compounds.
Lewis dot formulas are based on the octet rule.
We need to distinguish between bonding (or shared) electrons and nonbonding (or
unshared or lone pairs) of electrons.
N - A = S rule
– Simple mathematical relationship to help us write Lewis dot formulas.
N = number of electrons needed to achieve a noble gas configuration.
– N usually has a value of 8 for representative elements.
– N has a value of 2 for H atoms.
A = number of electrons available in valence shells of the atoms.
– A is equal to the periodic group number for each element.
– A is equal to 8 for the noble gases.
S = number of electrons shared in bonds.
A-S = number of electrons in unshared, lone, pairs.
Writing Lewis Formulas: The Octet Rule
1.
2.
For ions we must adjust the number of electrons available, A.
a.
Add one e- to A for each negative charge.
b.
Subtract one e- from A for each positive charge.
The central atom in a molecule or polyatomic ion is determined by:
a.
The atom that requires the largest number of electrons to complete its octet goes
in the center.
b.
For two atoms in the same periodic group, the less electronegative element goes in
the center.
3. Select a reasonable skeleton
a. The least electronegative is the central atom
b. Carbon makes 2,3, or 4 bonds
c. Nitrogen makes 1(rarely), 2,3, or 4 bonds
d. Oxygen makes 1, 2(usually), or 3 bonds
e. Oxygen bonds to itself only as O2 or O3, peroxides, or superoxides
f. Ternary acids (those containing 3 elements) hydrogen bonds to the oxygen, not
the central atom, except phosphates
g. For ions or molecules with more than one central atom the most symmetrical
skeleton is used
4. Calculate N, S, and A
Writing Lewis Dot Formulas
N ever Have a Full Octet
Always Have a Full Octet
Sometimes Have a Full Octet
Sometimes Exceed a Full Octet
Writing Lewis Dot Formulas
1. Count the number of electrons brought to the party (element’s group number)
2. For ions we must adjust the number of electrons available.
a. Add one e- for each negative charge.
b. Subtract one e- for each positive charge.
3. Select a reasonable skeleton
a. The least electronegative is the central atom
b. See prior periodic table for number of electrons involved in bonding
i. Group I 2 electrons or 1 bond
ii. Group II 4 electrons or up to 2 bonds
iii.Group III Al and B, 6 or 8 electrons up to 3 or 4 bonds
iv.C,N,O,F must have 8 electrons (up to 4 bonds for C, 3 for N, 2 for O, and 1 bond
for F).
v. All others must have at least 8 electrons (up to 4 bonds), but may have more.
4. When a choice for the central atom in a molecule or polyatomic ion is unclear:
a. For ions or molecules with more than one central atom the most symmetrical skeleton
is used
b. The atom that requires the largest number of electrons to complete its octet goes in
the center.
c. For two atoms in the same periodic group, the less electronegative element goes in the
center.
5. Calculate Formal charges, adjust bonds for lowest numbers (zero preferred) and allow for
resonance structures
Writing Lewis Formulas:
• Write Lewis dot and dash formulas for hydrogen cyanide, HCN.
• Write Lewis dot and dash formulas for sulfur trioxide, SO3.
• Write Lewis dot and dash formulas for the sulfite ion, SO32-.
• There are three possible structures for SO32-.
– The double bond can be placed in one of three places.
When two or more Lewis formulas are necessary to show the bonding in a molecule, we must use equivalent
resonance structures to show the molecule’s structure. Double-headed arrows are used to indicate resonance formulas.
• Write dot and dash formulas for BBr3.
• Write dot and dash formulas for AsF5.
Stereochemistry
•
•
•
Stereochemistry is the study of the three
dimensional shapes of molecules.
Valence Shell Electron Pair Repulsion Theory
• Commonly designated as VSEPR
• Principal originator
– R. J. Gillespie in the 1950’s
Valence Bond Theory
• Involves the use of hybridized atomic orbitals
• Principal originator
– L. Pauling in the 1930’s & 40’s
VSEPR Theory
• Regions of high electron density around the central atom are arranged as far apart as
possible to minimize repulsions.
• There are five basic molecular shapes based on the number of regions of high electron
density around the central atom.
• Lone pairs of electrons (unshared pairs) require more volume than shared pairs.
– Consequently, there is an ordering of repulsions of electrons around central atom.
• Criteria for the ordering of the repulsions:
1 Lone pair to lone pair is the strongest repulsion.
2 Lone pair to bonding pair is intermediate repulsion.
3 Bonding pair to bonding pair is weakest repulsion.
• Mnemonic for repulsion strengths
lp/lp > lp/bp > bp/bp
• Lone pair to lone pair repulsion is why bond angles in water are less than 109.5o.
• Valence-shell electron-pair repulsion (VSEPR) model predicts the shapes of
many molecules and polyatomic ions but provides no information about bond
lengths or the presence of multiple bonds
VSEPR Theory
• Frequently, we will describe two geometries for each molecule.
1. Electronic geometry is determined by the locations of regions of
high electron density around the central atom(s).
2. Molecular geometry determined by the arrangement of atoms around
the central atom(s).
Electron pairs are not used in the molecular geometry
determination just the positions of the atoms in the molecule are
used.
VSEPR Theory
• Two regions of high electron density around the central atom.
•
Three regions of high electron density around the central atom.
• Four regions of high electron density around the central atom.
VSEPR Theory
• Five regions of high electron density around the central atom.
• Six regions of high electron density around the central atom.
VSEPR Theory
• An example of a molecule that has different electronic and molecular geometries is water H2O.
• Electronic geometry is tetrahedral.
• Molecular geometry is bent or angular.
H
H C
H
H
• An example of a molecule that has the same electronic and molecular
geometries is methane - CH4.
• Electronic and molecular geometries are tetrahedral.
H
H C
H
H
The VSEPR Model
The VSEPR Model
•
The same basic approach will be used in every example of
molecular structure prediction:
The VSEPR Model
Day 4
Shorthand for organic molecules
Valence Bond (VB) Theory
• A more sophisticated treatment of bonding is a quantum mechanical
description of bonding, in which bonding electrons are viewed as being
localized between the nuclei of the bonded atoms
• The overlap of bonding orbitals is increased through a process called
hybridization, which results in the formation of stronger bonds
Regions of High Electron
Density
Electronic Geometry
Hybridization
2
Linear
sp
3
Trigonal planar
sp2
4
Tetrahedral
sp3
5
Trigonal bipyramidal
sp3d
6
Octahedral
sp3d2
Valence Bond Theory: A Localized Bonding
Approach
Hybridization of s and p Orbitals
• Localized bonding approach uses a process called hybridization, in which atomic
orbitals that are similar in energy but not equivalent are combined mathematically to
produce sets of equivalent orbitals that are properly oriented to form bonds.
• Spatial orientation of the hybrid atomic orbitals is consistent with the geometries
predicted using the VSEPR model.
• New combinations are called hybrid atomic orbitals because they are produced by
combining (hybridizing) two or more atomic orbitals from the same atom.
• Hybrid atomic orbitals are formed via promotion of an electron from a filled ns2
subshell to an empty np or (n – 1)d valence orbital, followed by hybridization.
Hybridization of s and p Orbitals
• The combination of an ns and an np orbital gives rise to two
equivalent sp hybrids oriented at 180º.
• Combination of an ns and two or three np orbitals produces
three equivalent sp2 hybrids or four equivalent sp3 hybrids.
Hybridization of s and p Orbitals
• Both promotion and hybridization require an input of energy; the
overall process of forming a compound with hybrid orbitals will be
energetically favorable only if the amount of energy released by the
formation of covalent bonds is greater than the amount of energy used
to form the hybrid orbitals.
© 2006 Brooks/Cole - Thomson
Hybridization Using d Orbitals
• Hybridization not restricted to ns and np atomic orbitals
• Bonding in compounds that have central atoms in the third
period and below can be described using hybrid atomic orbitals
• Central atom uses its valence (n – 1)d orbitals and its ns and np
orbitals to form hybrid atomic orbitals, which allows them to
accommodate five or more bonded atoms
• Using the ns orbital, all three np orbitals, and one (n – 1)d
orbital gives a set of five sp3d hybrid orbitals that point toward
the vertices of a trigonal bipyramid
• Combination of the ns orbital, all three np atomic orbitals, and
two (n – 1)d orbitals gives a set of six equivalent sp3d2 hybrid
orbitals oriented toward the vertices of an octahedron
Hybridization Using d Orbitals
Molecular Shapes and Bonding
• In the next sections we will use the following
terminology:
A = central atom
B = bonding pairs around central atom
U = lone pairs around central atom
• For example:
AB3U designates that there are 3 bonding pairs and 1
lone pair around the central atom.
Linear Electronic Geometry:AB2
Species (No Lone Pairs of Electrons
on A)
Be
1s
2s


2p
1s
 
sp hybrid 2p


Trigonal Planar Electronic Geometry:
AB3 Species
(No Lone Pairs of Electrons on A)
B
1s 2s 2p
  
1s
 
sp2 hybrid
  
Tetrahedral Electronic Geometry: AB4
Species
(No Lone Pairs of Electrons on A)
2s
C [He] 
2p

.
Tetrahedral Electronic Geometry: AB4 Species
Valence Bond Theory (Hybridization)
C [He]
2s

2p
.
four sp3 hybrids

.
Tetrahedral Electronic Geometry: AB3U Species
2s
N [He]

2p four sp3 hybrids

Tetrahedral Electronic Geometry: AB2U2 Species
O [He]
2s
 
3
2p four sp hybrids
 
Tetrahedral Electronic Geometry:
ABU3 Species
(Three Lone Pairs of Electrons on A)
Valence Bond Theory (Hybridization)
F [He]
four sp3 hybrids
   
2s
2p
  
··
H
F
··
··
tetrahedral
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
4s
4p
   
4d
As [Ar] 3d10
_______________

five sp3 d hybrids
4d
    
____________
Day 5
Compounds Containing
Double Bonds
Valence Bond Theory (Hybridization)
C atom has four electrons.
Three electrons from each C atom are in sp2 hybrids.
One electron in each C atom remains in an unhybridized p orbital
2s 2p
three sp2 hybrids 2p
C    


•
An sp2 hybridized C atom has this shape.
Remember there will be one electron in each of the three lobes.
Top view of
an sp2 hybrid
Compounds Containing
Double Bonds
• The single 2p orbital is perpendicular to the trigonal planar sp2 lobes.
The fourth electron is in the p orbital.
Side view of sp2 hybrid
with p orbital included.
Compounds Containing
Double Bonds
•
Two sp2 hybridized C atoms plus p orbitals in proper orientation to form C=C
double bond.
• The portion of the double
bond formed from the
head-on overlap of the sp2
hybrids is designated as a
s bond.
•
The other portion of the
double bond, resulting from the
side-on overlap of the p
orbitals, is designated as a p
bond.
Compounds Containing Triple Bonds
A s bond results from the head-on overlap of two sp hybrid
orbitals.
The unhybridized p orbitals form two p bonds.
Note that a triple bond consists of one s and two p bonds.
Summary of Electronic & Molecular
Geometries
Day 6