chapter 8 bonding: general concepts. 8.1 types of chemical bonds l a bond is a force that holds...
TRANSCRIPT
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Chapter 8
Bonding: General Concepts
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8.1 Types of Chemical bonds
A bond is a force that holds atoms together and make them function together
Bond energy: the energy required to break a bond.
Why are compounds (atoms aggregate with each other) formed?– Because this situation gives the system
the lowest possible energy.
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Ionic Bonding An atom with a low ionization energy
(metal that looses its electron easily) reacts with an atom with high electron affinity (nonmetal).
The electron moves. Opposite charges hold the atoms
together. Closely packed oppositely charged ions
are held by strong electrostatic attraction forces
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The highly ordered solid collection of ions is called an ionic compound (crystal)
The net attractive electrostatic forces that hold the cations and anions together are ionic bonds
• Ionic crystals have great thermal stability and consequently acquire high melting points
When atoms lose or gain electrons, they acquire a noble gas configuration, but do not become noble gases
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Ionic Bond• When metals and nonmetals
combine, valence electrons usually are transferred from the metal to the nonmetal atoms giving rise to electrostatic attraction force called ionic bond
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Energy of interaction between a pair of ions: Coulomb's Law
E= 2.31 x 10-19 J · nm Q is the charge. r is the distance between the centers. If charges are opposite, E is negative Exothermic: the ion pair has less energy the ion pair has less energy
than separated ionsthan separated ions Same charge, positive E, requires energy to
bring them together.
)21
(r
)21
(r
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What about identical atoms?
The electrons in each atom are attracted to the nucleus of the other.
The electrons repel each other, The nuclei repel each other. They reach a distance with the lowest
possible energy. The distance between is the bond
length.
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0
En
erg
y
Internuclear Distance
Atoms are infinitely apart
Zero interacting
energy
How does a bonding force develop
between two identical atoms?
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0
En
erg
y
Internuclear Distance
Energy as a function of Internuclear distance
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0
En
erg
y
Internuclear Distance
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0
En
erg
y
Internuclear Distance
Most stable state
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0
En
erg
y
Internuclear Distance
Bond Length
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0
En
erg
y
Internuclear Distance
Bond Energy
The molecule is more stable than the two separate atoms
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Covalent Bonding Covalent bonding: electrons are
shared equally between two identical atoms.
on the other extreme (ionic bonding) electrons transfer completely to form oppositely charged ions
In between are polar covalent bondspolar covalent bonds. The electrons are not shared evenly. One end is slightly positive, the other
negative. Bond polarity is indicated using small
delta
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H - F+ -
Polar covalent bond Polar covalent bond
The density of electron cloud is shifted(some what) towards one of the two bonded atoms
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H - F+ -
H - F
+-H - F+
-
H - F
+-
H - F +-
H - F+-
H - F
+-
H - F
+-
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H - F+ -
H - F
+-H - F+
-
H - F
+-
H - F +-
H - F+-
H - F
+-
H - F
+-
+-Effe
ct of e
lec
tric field
on
HF
m
ole
cule
s
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H - F+ -
H - F+ -
H - F+ - H - F
+ -
H - F+ -
H - F+ -
H - F+ -
H - F+ -
- +
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8.2 Electronegativity The ability of an atom in a molecule to attract
shared electrons to itself. To measure the relative electronegativity imagine
an H-X molecule Pauling method: compare the measured
H-X bond energy with the expected H-X bond energy
= (H-X) actual - (H-X)expected
2
energy X-Xenergy H-Henergy X-H Expected
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Electronegativity
is known for almost every element Gives us relative electronegativities of all
elements. Tends to increase left to right. decreases as you go down a group. Noble gases aren’t discussed. Difference in electronegativity between
atoms tells about the polarity of the bond
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Electronegativity
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Two identical atoms have the same electronegativity and share a bonding electron pair equally. This is called a nonpolar covalent bond
Example: chlorine gas
EOS
All homonuclear diatomic molecules have nonpolar covalent bonds:
H2, N2, O2, F2, Cl2, Br2, I2
Electronegativity difference & bond type
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Electronegativity difference
Bond Type
Zero
Intermediate
Large
Covalent
Polar Covalent
Ionic
Co
valent C
haracter
decreases
Ion
ic Ch
aracter increases
Relationship between electronegativityand bond type
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Pauling’s Electronegativities
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Electronegativity Differences
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8.3 Bond polarity and dipole moments
A molecule with a center of negative charge and a center of positive charge is dipolar (two poles),
or has a dipole moment. Center of charge doesn’t have to be on
an atom. Dipoles will line up in the presence of an
electric field.
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Representation of the dipolar character
H - F+ -
• Any diatomic (two atoms) molecule with a polar bond will show a molecular dipole moment
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H F
FH
A covalent bond with greater electron density around one of the two atoms
electron richregion
electron poorregion e- riche- poor
+ -
Polar Covalent Bond
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Polarity of polyatomic molecules The effect of polar bonds on the polarity of
the entire molecule depends on the molecule shape– carbon dioxide has two polar bonds, and is
linear = nonpolar molecule
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Molecules with polar bonds but no resulting dipole moment
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Polar molecules The effect of polar bonds on the polarity of
the molecule depends on the molecular shape– water has two polar bonds and a bent
shape; the highly electronegative oxygen pulls the e- away from H = very polar!
Thus, H2O molecule has a dipole moment
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How to decide for molecular polarity?
Any diatomic molecule with a polar bond is a polar molecule
For a three or more atoms molecule there are two considerations:
– There must be a polar bond.
– Geometry can’t cancel it out.
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Geometry and polarity Three shapes will cancel them out.
Linear
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Planar triangles
120º
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Tetrahedral
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Others don’t cancel, e.g., Bent molecule
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8.4 Ions: electron configuration and size
Atoms tend to react to form noble gas configuration.
Nonmetals gain electrons from metals or share electrons with other nonmetals
Metals lose electrons to form cations Nonmetals can share electrons in covalent
bonds. Or they can gain electrons to form anions.
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Formation of Cations
Na 1s22s22p63s1 1 valence electron Na1+ 1s22s22p6 This is a noble gas
configuration with 8 electrons in the outer level.
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The noble gas configuration (Octet rule)
In forming compounds, atoms tend to achieve a noble gas configuration; 8 e- in the outer level is stable
Each noble gas (except He: 2 e-) has 8 electrons in the outer level
For H it is duet rule (Rule of 2)
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Li + F Li+ F -
Ionic bond and the octet rule
1s22s11s22s22p5 1s21s22s22p6[He][Ne]
Li
1s 2s 2p
F
1s 2s 2p
+
Li+
1s 2s 2p
F-
1s 2s 2p+
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Noble gas configuration (Octet) rule
and covalent bonding
Atoms are bonded with the electron configuration of a noble gas; that is, the atoms obey the octet rule
EOS
By double-counting the shared electrons in a Lewis structure, each atom appears to have a noble gas configuration
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Predicting formulas of Ionic Compounds
When the term “ionic compound” is used; it means solid state (crystalline) of that compound
Ions align themselves to maximize attractions between opposite charges,
and to minimize repulsion between like ions.
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Predicting formulas of ionic compounds
Stoichiometry is an important consideration …
MgO
Li2O
1s22s22p63s2 1s22s22p4
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Common ions with noble gas configurations in ionic compounds
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Predicting formulas
Predict the formula of the compound formed from Al and O
Al : [Ne]3s23p1 (should lose 3e- Ne) O : [He]2s22p4 ) should gain 2e- Ne) The compound to be electrically
neutral it has to be Al2O3
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Size of ions
Ion size increases down a group. Cations are smaller than the atoms they
came from. Anions are larger than atoms they came
from. across a row they get smaller, and then
suddenly larger. First half are cations. Second half are anions.
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Periodic Trends Across the period, the change is
complicated because of the change from predominantly metals on the left to nonmetals on the right.
Li+
Be2+
B3+
C4+
N3-
O2- F-
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Ato
mic an
d Io
nic
Rad
ii
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Size of Isoelectronic ions Iso - same Iso electronic ions have the same #
of electrons Al3+ , Mg2+, Na+, Ne, F-, O2- and N3-
All have 10 electrons. All have the configuration 1s22s22p6
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Isoelectronic Configurations
Elements that all have the same number of electrons
For isoelectronic species, the greater the nuclear charge, the smaller the species
Effective n
uclear ch
arge
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Size of Isoelectronic ions Positive ions have more protons
so they are smaller.
Al+3
Mg+2
Na+1 Ne F-1 O-2 N-3
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Ionic solid is formed because the aggregated oppositely charged ions have a lower energy than the original elements
How strongly the ions attract each other in the solid state is expressed by the lattice energy
Lattice energy:
The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid
M+(g) + X-(g) MX
8.5 Energy effects in binary ionic compounds
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Lattice energy It is usually defined as the energy
released released when ionic solid is formed from its gaseous ions (it has –ve sign)– M+(g) + X-(g) MX(s)
Lattice energy is a quantitative measure of the stability of ionic compound
The higher the lattice energy the more stable the compound
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Na(s) + ½F2(g) NaF(s)
First sublime Na Na(s) Na(g)H = 109 kJ/mol
Ionize Na(g) Na(g) Na+(g) + e-
H = 495 kJ/mol Break F-F Bond ½F2(g) F(g)
H = 77 kJ/mol Add electron to F F(g) + e- F-(g)
H = -328 kJ/mol Formation of NaF from Na+(g) & F-(g) (Lattice energy) Na+(g)+ F-(g) aF(s) H = -1281 kJ/mol
Energy changes associated with the formation of ionic solid
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Na(s) + ½F2(g) NaF(s) Lattice energy
Na(s) + ½F2(g) NaF(s)H = -928 kJ/mol
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Calculating Lattice Energy
k is a constant that depends on the structure of the crystal electron configuration of ions
r is internuclear distance (r = radius of cation+radius of anion)
Lattice energy is greater with more highly charged ions and distances between ions decrease
)r
QQk(energy Lattice 21
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8.6 Partial Ionic Character of covalent bonds
There are probably no totally ionic bonds between individual atoms.
Calculate % ionic character of a bond Compare measured dipole moment of X-
Y bonds to the calculated dipole moment of X+Y- the completely ionic case in the gaseous phase
%100)YX ofmoment dipole calculated
Y-X ofmoment dipole (bond a ofcharacter ionic
_X
measuredPercent
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% Io
nic
Ch
arac
ter
Electronegativity difference
25%
50%
75%
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The Relationship Between the Ionic Character of a Covalent Bond and the Electronegativity Difference of the Bounded Atoms
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What are the ionic compounds?
If bonds can’t be ionic, what are ionic compounds?
What about polyatomic ions?An ionic compound will be defined
as any substance that conducts electricity when melted
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8.7 The covalent chemical bond: A model
Bonds are the forces that cause a group of atoms to behave as a unit.
Why do chemical bonds occur?–Due to the tendency of atoms in a system
to achieve its lowest possible energy It takes 1652 kJ to dissociate a mole of CH4
into its separate atoms C & H. Thus, CH4 is stable relative to its separated
atoms
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To explain stability, chemical bond term was introduced
Since each hydrogen in CH4 is hooked to the carbon, we get the average energy = 413 kJ/mol = Bond EnergyBond Energy
CH3Cl has 3 C-H, and 1 C – Cl and a stabilization energy of 1578 kJ/mol
Thus, C-Cl bond can be calculated as 339 kJ/mol The bond is a human invention. It is a method of explaining the energy change
associated with forming molecules. Bonds don’t exist in nature, but are useful. We have a model of a bond.
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The bond Model
Explains how nature operates.Derived from observations. It simplifies the and categorizes the
information.A model must be sensible, but it
has limitations.
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Properties of a Model
Models are human inventions; it does not equal rality.
Models can be wrong, because they are based on speculations and oversimplification.
Become more complicated with age. You must understand the assumptions in the
model, and look for weaknesses. We learn more when the model is wrong than
when it is right.
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8.8 Covalent bond energies and chemical reactions
Bond energies (BE) is H when 1 mole of bonds is broken in the gaseous state
H2(g) 2H(g) H = +436 kJ
Cl2(g) 2Cl(g) H = +243 kJ BE is always positive (endothermic) Energy is given off when bonds are formed
H(g) + F(g) HF(g) H = -565 kJ Bond energy increases with bond polarity C-F > N-F > O-F > F-F
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Covalent bond energies and chemical reactions
Consider stepwise decomposition of CH4 Each C-H bond has a different energy. CH4 CH3 + H H = 435 kJ/mol
CH3 CH2 + H H = 453 kJ/mol
CH2 CH + H H = 425 kJ/mol CH C + H H = 339 kJ/mol Each bond is sensitive to its environment. Average C-H bond energy = 1652/4=413kJ/mol
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Multiple bonds Single bond one pair of electrons is
shared. Double bond two pair of electrons are
shared. triple bond three pair of electrons are
shared. BE is larger for a multiple bond than
for a single bond between the same two atoms
More bonds, shorter bond length.
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Av
era
ge
Bo
nd
En
erg
ies
(kj/m
ol)
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Bond Type
Bond Length
)pm(
C-C 154
CC 133
CC 120
C-N 143
CN 138
CN 116
Lengths of Covalent Bonds
Bond Lengths
Triple bond < Double Bond < Single Bond
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Bond Energies and Enthalpy changes in reactions
H = D required– D released when BE bonds broken BE bonds formed
= BE (reactants) – SBE (products)
9.10
Enthalpy change for a reaction
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Use bond energies to calculate the enthalpy change for:H2 (g) + F2 (g) 2HF (g)
H = BE(reactants) – BE(products)
Type of bonds broken
Number of bonds broken
Bond energy (kJ/mol)
Energy change (kJ)
H H 1 436.4 436.4
F F 1 156.9 156.9
Type of bonds formed
Number of bonds formed
Bond energy (kJ/mol)
Energy change (kJ)
H F 2 568.2 1136.4
H = 436.4 + 156.9 – 2 x 568.2 = -543.1 kJ
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Find the energy for this
2 CH2 = CHCH3
+
2NH3 O2+
2 CH2 = CHC N
+
6 H2O
C-H 413 kJ/molC=C 614kJ/molN-H 391 kJ/mol
O-H 467 kJ/molO=O 495 kJ/molCN 891 kJ/mol
C-C 347 kJ/mol
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8.9 The localized electron bonding model
Simple model, easily applied to describe covalent bonds
A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
Electron pairs are assumed to be localized on a particular atom (Lone pairs) or in the space between two atoms (Bonding pairs)
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The Localized Electron Model
has three parts: 1. Description of valence electrons
arrangements in the molecule using Lewis structures
2. Prediction of geometry of the molecule using VSEPR (valence shell electron- pair repulsion model
3. Description of the types of orbitals used by the atoms to share electrons or hold lone pairs.
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8.10 Lewis Structure
Shows how the valence electrons are arranged around atoms in a molecule
One dot for each valence electron.A stable compound has all its atoms
with a noble gas configuration.Hydrogen follows the duet rule.The rest follow the octet rule.Bonding pair is the one between the
symbols.
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Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.
N:.
..
:
N .. ..N :.
. :N ...
Place one dot per valence electron on each of the four sides of the element symbol.
Pair the dots (electrons) until all of the valence electrons are used.
Lewis Dot Symbols
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1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.
2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.
3. Use one pair of electrons to form a bond (a single line) between each pair of atoms.
4. Arrange the remaining electrons to satisfy an octet for all atoms (duet for H), starting from outer atoms.
5. If a central atom does not have an octet, move in lone pairs to form double or triple bonds on the central atom as needed.
Rules for Writing Lewis Structures
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Examples Fluorine has seven valence electrons
9F: 1s22s22p5
A second atom also has seven By sharing electrons… …both end with full orbitals
F F8 Valence electrons
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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons… …both end with full orbitals
F F8 Valence electrons
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Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
F N F
F
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms.
Step 4 – Arrange remaining 20 electrons to complete octets
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Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
O C O
O
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-
4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.
Step 4 - Arrange remaining 18 electrons to complete octets
Step 5 – The central C has only 6 electrons. Form a double bond.
2
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Writing Lewis Structures
If you run out of electrons before the central atom
has an octet…
…form multiple bonds until it does.
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Maximum number of bonds (or atoms) surrounding the central atom
Central atomMax # bondsH1H-O2 -O-
N3 -N-
C4 -C-
X-(F, Cl, Br, I)1X-
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Exceptions to the Octet Rule
The Incomplete Octet
H HBeBe – 2e-
2H – 2x1e-
4e-
BeH2
BF3
B – 3e-
3F – 3x7e-
24e-
F B F
F
3 single bonds (3x2) = 69 lone pairs (9x2) = 18
Total = 24
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The expanded octet
Some compounds have expanded valence shells, which means that the central atom has more than eight electrons around it
#Ve= 5 + 7(5)= 40 #Ve = 6 + 6(7)= 48
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Expanded octet
EOS
An expanded valence shell may also need to accommodate lone-pair electrons as well as bonding pairs
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Molecules with expanded octet Terminal atoms are most often halogens. In few
cases one or more O-atoms are at the end of the molecule
The central atom is a nonmetal in the 3rd, 4th or 5th period of the Periodic Table– 3rd P S Cl– 4th As Se Br Kr– 5th Sb Te I Xe
All atoms have d-orbitals available for bondig (3d, 4d, 5d) where the extra pairs of electrons are located
Elements of Period 2, NEVERPeriod 2, NEVER form compounds with expanded octet.
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Odd-Electron Molecules
N – 5e-
O – 6e-
11e-
NO N O
N O
Species with unpaired electrons show weak attraction to magnetic field
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A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.
O O O OOO
O C O
O
- -O C O
O
-
-
OCO
O
-
- 9.8
What are the resonance structures of the carbonate (CO3
2-) ion?
8.12 Resonance
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Resonance•In truth, the electrons that form the second C—O bond
in the double bonds below do not always sit between that C and that O, but rather can move among the two
oxygens and the carbon.•They are not localized, but rather are delocalized.
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Molecules that don’t follow the octet rule
EOS
Molecules with an odd number of valence electrons have at least one of them unpaired and are called free radicalsfree radicals
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Formal Charge
For molecules and polyatomic ions that exceed the octet there are several exceed the octet there are several different structuresdifferent structures.
Use charges on atoms to help decide which one is the real molecule.
Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.
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Formal chargeFormal charge is the difference between the number of valence electrons on the free atom and that assigned to the atom in the molecule.
CCff = N = Nvv – (N – (Nuu + ½ N + ½ Nbb))– Cf = formal charge
– Nv = #valence e- in the un-bonded atom
– Nu = # unshared e- owned by the atom
– Nb= # bonding e- shared by the atom
In molecules Cf is close to zero In ions should be equal to the charge on the ion
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Calculation of Formal Charge
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Using the assumption of formal charges to evaluate Lewis structure
Atoms in molecules try to achieve as low a formal charge (as close to zero) as possible
Negative formal charges are expected to be found on the most electronegative elements.
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Which is the most likely Lewis structure for CH2O?
9.7
H C O H
-1 +1 HC O
H
0 0
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8.13 Molecular Structure: VSEPR
Lewis structures tell us how the atoms are connected to each other.
They don’t tell us anything about shape. The shape of a molecule can greatly
affect its properties. VValencealence SShell hell EElectronlectron PPair air RRepulsionepulsion
Theory allows us to predict geometry
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VSEPRMolecules take a shape that puts
electron pairs as far awayfar away from each other as possible.
The electron-pairs surrounding an atom (valence electrons) repel one another and are oriented as far apart as possible
Structure around a given atom is determined pricipally by minimizing electron –pair repulsion
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To determine electron pairs Lewis structure should be drawn
Find bonding and nonbonding lone pairs
Lone pair take more space. Multiple bonds count as one pair.
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VSEPR The number of pairs determines
–bond angles–underlying structure
The number of atoms determinesThe number of atoms determines –actual shape
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Valence shell electron pair repulsion (VSEPR) model
Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs.
AB2 2 0
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
10.1
linear linear
B B
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Cl ClBe
2 atoms bonded to central atom0 lone pairs on central atom 10.1
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AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
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F
F F
B
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AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
10.1
AB4 4 0 tetrahedral tetrahedral
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H
H
H
H
C
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AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
AB4 4 0 tetrahedral tetrahedral
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
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P
Cl
Cl
Cl
Cl
Cl
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AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
AB4 4 0 tetrahedral tetrahedral
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB6 6 0 octahedraloctahedral
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S
F
F
F
F
F
F
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Effect of lone pairs on Geometry
• Unshared pair of electrons (under the influence of one nucleus) spreads out over a volume larger than a bonding pair (under the influence of two nuclei).
• The electron pair geometry is approximately same as that observed when only single bonds are involved
• The bond angles are either equal to the ideal values or little less
• The molecular geometry is quite different when line pairs are involved.
• Molecular geometry refers only to the positions of the bonded atoms
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Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
AB2E 2 1trigonal planar
bent
< 120o
SO2
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Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3E 3 1
AB4 4 0 tetrahedral tetrahedral
tetrahedraltrigonal
pyramidal
< 109.5o
107o
NH3
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Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB4 4 0 tetrahedral tetrahedral
AB3E 3 1 tetrahedraltrigonal
pyramidal
AB2E2 2 2 tetrahedral bent
H
O
H
< 109.5o
104.5o
H2O
ABE31 3 LinearH-B
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Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB4E 4 1trigonal
bipyramidaldistorted
tetrahedron
SF4
Seesaw90o, 120o, 180o
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Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB4E 4 1trigonal
bipyramidaldistorted
tetrahedron
AB3E2 3 2trigonal
bipyramidalT-shaped
ClF
F
FClF3
90o, 180o
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Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB4E 4 1trigonal
bipyramidaldistorted
tetrahedron
AB3E2 3 2trigonal
bipyramidalT-shaped
AB2E3 2 3trigonal
bipyramidallinear
I
I
II3 180o
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Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB6 6 0 octahedraloctahedral
AB5E 5 1 octahedral square pyramidal
Br
F F
FF
F
BrF5
90o, 180o
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Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB6 6 0 octahedraloctahedral
AB5E 5 1 octahedral square pyramidal
AB4E2 4 2 octahedral square planar
Xe
F F
FFXeF4
90o, 180o
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10.1
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The VSEPR model and multiple bond
• For the geometry purposes:
A multiple bond behaves exactly as if it were a single electron-pair
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Molecules containing no single central atom
• The central atoms of the molecule should be labeled first.
• Geometry can be predicted by focusing on each central atom by counting the electron pairs around each central atom.
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END
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Predicting Molecular Geometry1. Draw Lewis structure for molecule.
2. Count number of lone pairs on the central atom and number of atoms bonded to the central atom.
3. Use VSEPR to predict the geometry of the molecule.
What are the molecular geometries of SO2 and SF4?
SO O
AB2E
bent
S
F
F
F F
AB4E
distortedtetrahedron
10.1
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Dipole Moments and Polar Molecules
10.2
H F
electron richregion
electron poorregion
= Q x rQ is the charge
r is the distance between charges
1 D = 3.36 x 10-30 C m
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10.2
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10.2
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10.2
Which of the following molecules have a dipole moment?H2O, CO2, SO2, and CH4
O HH
dipole momentpolar molecule
SO
O
CO O
no dipole momentnonpolar molecule
dipole momentpolar molecule
C
H
H
HH
no dipole momentnonpolar molecule
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Does CH2Cl2 have a dipole moment?
10.2
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10.2
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How well does it work? Does an outstanding job for such a
simple model. Predictions are almost always
accurate. Like all simple models, it has
exceptions.