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Basic Chemical Bonding

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Title: Basic Chemical Bonding


1
Basic Chemical Bonding
Molecules are artwork just beautiful!
Cubane
Dodecahedrane
Side and top views of a single-wall
exohydrogenated carbon nanotube
2
Looking Back at Chemical Bonding
Bonding must be electric nature. 1852, E.
Frankland proposed the valence concept, using
for valence. 1857, F.A. Kekule figured out the
structure of benzene C6H6. 1874 J.H. van't Hoff
and le Bel postulated the tetrahedral arrangement
of 4 bonds around carbon. 1916 G.N. Lewis
propsed the dot symbol for valence electrons 1923
G.N. Lewis wrote Valence and the structure of
atoms and molecules. 1939 L. Pauling wrote The
nature of chemical bond 1940 N.V. Sidgwick and
H.E. Powell studied the lone pairs of valence
electrons.
3
Lewis Theory

G.N. Lewis (1875-1946) recognized valence
(outmost) electrons fundamental to
bondingelectron transfer resulting in ionic
bondssharing electrons resulting in covalent
bondsatoms tend to acquire a noble-gas
electronic configurations

The attraction between electrons of one atom to
the nucleus of another atom contribute to what is
known as chemical bonds.

4
Lewis Dot Structure
Lewis wrote in a memorandum dated March 28, 1902
5
Lewis Dot Structure 2
Lewis' Paper of 1916 In this paper, Lewis begins
by using cubes, but he moves away from them by
the end of the paper. Here is how he visualized
the elements lithium through fluorine
Please illustrate modern Lewis dot structures of
periods 2 and 3 elements. Chieh does that during
lecture.
6
Lewis Dot in Covalent Bond
Write the Lewis dot structures for these
molecules HH, HCl, HOH, NH3, H, He, Cl,
NeH3O, NH4, OH , (coordinate covalent)Cl2,
O2, (multiple bonds) N2, CO2 Explain the types
in each line and write the dot structures. Define
bond pair, lone pair, single bond, double bond,
triple bond
7
Polar Covalent Bond Electronegativity
Discuss the nature of these bonds HF, HCl
HOH (including lone pairs) Electronegativity
the ability of an element competing for bonding
electrons. The variation as a function of
atomic number and its trends on the Periodic
Table has been discussed previously, and the
Periodic Table showing electronegativity is shown
next.
8
Periodic Table of Electronegativity
9
Covalent and Ionic Bonds
The ionicity of a bond depends on the difference
in electronegativity. A difference of 1.7 is
given as 50 ionic, and usually considered ionic.
Analyze these
10
Electron Density of a Polar Bond LiH
Li ? Hdipole moment
11
Writing Lewis Dot Structures
Show all valence electrons. Each bond represents
two electrons. All electrons are paired, usually
(exceptions). Each atom acquires 8 valence
electrons, usually (exceptions). Multiple bonds
are needed sometimes.
Show class how to write Lewis structure for CF4,
(CX4, SiX4), NH3, H2O, HF C2H5OH, HC?N, H3PO4,
ONO
12
Formal Charge
The formal charge on any atom in a Lewis
structure is a number assigned to it according to
the number of valence electrons of the atom and
the number of electrons around it. The formal
charge of an atom is equal to the number of
valence electrons, Nv.e. subtract the number of
unshared electrons, Nus.e. and subtract half of
the bonding electrons, ½ Nb.e.. Formal charge
Nv.e. - Nus.e. - ½ Nb.e.
Stability rulesFormulas with the lowest
magnitude of formal charges are more stable.More
electonegative atoms should have negative formal
charges.Adjacent atoms should have opposite
formal charges.
Explain workout formal charge judge stability
of a formula
13
Find Formal Charge
SO42 Find FC in these structures
Confirm these FCs
14
Resonance
When several structures with different electron
distributions among the bonds are possible, all
structures contribute to the electronic structure
of the molecule. These structures are called
resonance structures. When two or more plausible
Lewis structures can be written but the correct
structure cannot be written is called resonance.
For example . . . . . . O O O O O
O O O O Please
complete the dot structure and find the formal
charge for the above structures.
15
Draw Resonance Structures
Draw resonance structures for these CO2
OCO (plus two more dots for each of
O) NO2 .NO2 (bent molecule due to the odd
electron) NO2- NO2- (same number of valence
electron as O3 SO2) HCO2- H-CO2 ( ditto) O3
ozone SO3 consider O-SO2, and the resonance
structures NO3 flat same number of valence
electron as CO32-
Draw all resonance structures of all these
16
Exceptions to the Octet Rule
Molecules with odd number of valence electrons,
NO (compare to CO), CH3, OH, H, NO2
etc. Molecules with incomplete octets, BeCl2
AlCl3, (gas and polymeric for both), BF3,
compare with NH3BF3, BF4, Expanded valence
shells, PCl3, PCl5, SF6, H2SO4, H3PO4
Cl Cl Cl / \
/ \ / \M M
M M \ / \ / \
/ Cl Cl Cl M Al or
Be
Draw Lewis dot structures of all these molecules
to see the exceptions
17
Bond Properties
Bond length distance between the nuclei of
bonded atoms bond angle angles for any two bonds
around an atom bond energy energy required to
break the bond bond-dissociation energy
length energy Compound Bond
(pm) (kJ/mol) H2 H H 74 436 HF F H 92
565 H2O O H 96 464 NH3 N H 101 389 CH4 C
H 109 414
Bond Length Energy C C 154 348C C 134 614C
? C 120 839 O O 148 145O O 121 498
Discuss the variations of bond length and bond
energy
18
VSEPR Theory
Valence-Shell Electron Repulsion Theory The
VSEPR model counts both bonding and nonbonding
(lone) electron pairs (E), and call the total
number of pairs number of electron groups (Neg).
If the element A has m atoms bonded to it and n
nonbonding pairs (E), then Neg m n
Discuss the electronic and molecular structures
of CH4, ENH3, E2OH2. All have Neg 4. Bond
angles in these structure indicates that E E
repulsion is stronger than that of bonding
electrons. CH4 ENH3 H2OE2 HFE3
19
Shape of Molecules
During the lecture, we will discuss structures of
the following AX2 linearBeCl2 AX3, AX2E
triangular planar, bentBF3, SO2E AX4, AX3E AX2E2
tetrahedral, pyramidal, bent AX5, AX4E, AX3E2,
AX2E3 triangular pyramidal, butterflyPCl5, SF4E,
ClF3E2, XeF2E3 T-shape, linear AX6, AX5E,
AX4E2 octahedral, square pyramidal, square
planarSF6, BrF5E, XeF4E2, ICl4E2
Make sure you can draw and name the geometrical
shape of these structures.
AX4E, whats my shape?
20
Chemistry and Molecular Shapes
Neg Example Descriptor 2 BeCl2, CO2 Linear
3 BF3, SO3 Trigonal planar SO2E, OO2E
Bent 4 CH4 Tetrahedral NH3E pyramidal H2OE2 Ben
t 5 PF5 Trigonal bypyramidal SF4E Seesaw,
butterfly ClF3E2 T-shape


Neg Example Descriptor 6 SF6,
OIF5 Octahedral BrF5E Pyramidal XeF4E2 Square
planar
21
Structures with Multiple Covalent Bonds
We will talk about pi (p) bonding later. At this
stage, you may consider all electrons in a
multiple bond are confined around the lines
connecting the two atoms. Thus the number of
electron groups Neg for a multiple bond is 1. For
example, Neg 3 for
H \ C O /H
. .S/ \\ O O
What is the Neg forSO42, COS, N2O?
22
Molecules with more than one central atom
Describe the structure of methyl isocyanate,
CH3NCO. Draw the skeleton and add all valence
electrons H3C N C O Draw the Lewis dot
structure that satisfy the octet rule.
N C O
180o
109o
120o
H - C
H H
What are the formal charges of all atoms in both
structures? Describe the structures of C2H5OH,
CH3CO2H, and H2NCH2CH2(OH)COOH.
23
Dipole Moment
The product of magnitude of charge on a molecule
and the distance between two charges of equal
magnitude with opposite sign is equal to dipole
moment D (unit is debye, 1 D 3.34E30 C m
(coulumb.metre) representation Cl?H, a vector
) Dipole moment charge x distanceSymbol µ
e x d dq dbond
For Cl?H, µ 1.03 D, dHCl 127.4 pmTwo
ways of lookint at H?Cl, dq 1.033.34e30 C
m / 1.274e-12 m 2.70e-20 C (charge
separation by HCl )Ionic character dq / e
0.17 17 d 3.44e-30 C m / 1.60e 19 C (e
charge) 2.15E11 m 0.215 pm (?e by
0.215 pm)
mHCl 1.03 DmHF 1.9 D, find d and ionic
character for them.
24
Dipole moment of H2O
Verify please The dipole moment of individual
water molecules measured by Shostak, Ebenstein,
and Muenter (1991) is 6.187?1030 C m (or 1.855
D). This quantity is a vector resultant of two
dipole moments of due to OH bonds. The bond
angle HOH of water is 104.5o. Thus, the dipole
moment of a OH bond is 5.053?1030 C m. The bond
length between H and O is 0.10 nm, and the
partial charge at the O and the H is therefore q
5.053?1020 C, 32 of the charge of an
electron (1.6022?1019 C). Of course, the dipole
moment may also be considered as separation of
the electron and positive charge by a distance
0.031 nm. For the water molecule, a dipole moment
of 6.187?1030 C m many be considered as
separation of charge of electron by 0.039 nm.
25
Dipole moment and Molecular Shape
Dipole moments are vectors. The net dipole moment
of a molecule is the resultant (vector sum) of
all bond-dipole-moment. Answer explain these
mHH ____mOCO _____ mCH4 _____
mCCl4 _____ mBF3 _____ mH2O 1.84 D
mO3 0.534 D (implication of long pair) Which
are polar and non-polar, SF6, H2O2, C2H4,
Cl3CCH3, PCl5, I-Cl, NO, SO2, CH2Cl2, NH3, (put
your skill to tell molecular shape at work)
26
Review 1
Predict the molecular geometry of the polyatomic
anion ICl4 Hint Draw the Lewis dot structure
for Cl and I (figure out the valence es) Drew
the Lewis dot structure for ICl4 What is the
number of unshared e of the above? Drew the ion,
and describe this shape in proper term.
Do the same for NCl3, POCl3, COS, H2CO,
27
Review 2
Apply bond energy for thermochemistry
calculation In a chemical reaction, add (ve)
energy released from bonds formed and (ve)
energy required to break the bonds is the energy
of the reaction DHrxno. What is the heat of
reaction for 2 H2 (g) O2 (g) ? 2 H2O (g),
DHrxno 2 D (HH) D(OO) 4 D(HO)
2 436 489 4464
495 kJ compare to DHfo 248 kJ mol1 of H2O
Data D(OO) 498 kJ mol1D(HH) 436
D(HO) 464 2 H (g) ? H2 (g), DH
D (HH) H2 (g) ? H (g), DH D (HH)
Work on example 11-14 on page 423
28
Review 3
Data D(C-H) 414 kJ mol1D(ClCl) 243
D(C-Cl) 339 D(HCl) 431
What is the energy of reaction for CH4
(g) Cl2 (g) ? CH3Cl (g) HCl
(g)? Solution H3C H Cl Cl ? H3C
Cl H Cl 414 243
339 431 kJ DHrxno 414 243
339 431 kJ 113 kJ
Answer 113 kJ is released in this reaction.
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