1 chapter 16 covalent bonding. 2 section 16.1 the nature of covalent bonding l objectives: –use...
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Chapter 16Covalent Bonding
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Section 16.1The Nature of Covalent Bonding
OBJECTIVES:
–Use electron dot structures to show the formation of single, double, and triple covalent bonds.
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Section 16.1The Nature of Covalent Bonding
OBJECTIVES:
–Describe and give examples of coordinate covalent bonding, resonance structures, and exceptions to the octet rule.
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How does H2 form? The nuclei repel
++
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How does H2 form?
++
The nuclei repel But they are attracted to electrons They share the electrons
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Covalent bonds Nonmetals hold on to their valence
electrons. They can’t give away electrons to bond. Still want noble gas configuration. Get it by sharing valence electrons with
each other. By sharing, both atoms get to count the
electrons toward a noble gas configuration.
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Covalent bonding Fluorine has seven valence
electrons
F
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Covalent bonding Fluorine has seven valence
electrons A second atom also has seven
F F
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Covalent bonding Fluorine has seven valence
electrons A second atom also has seven By sharing electrons…
F F
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Covalent bonding Fluorine has seven valence
electrons A second atom also has seven By sharing electrons…
F F
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Covalent bonding Fluorine has seven valence
electrons A second atom also has seven By sharing electrons…
F F
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Covalent bonding Fluorine has seven valence
electrons A second atom also has seven By sharing electrons…
F F
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Covalent bonding Fluorine has seven valence
electrons A second atom also has seven By sharing electrons…
F F
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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 F
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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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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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A Single Covalent Bond is... A sharing of two valence electrons. Only nonmetals and Hydrogen. Different from an ionic bond because
they actually form molecules. Two specific atoms are joined. In an ionic solid, you can’t tell which
atom the electrons moved from or to.
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How to show how they formed It’s like a jigsaw puzzle. You put the pieces together to end
up with the right formula. Carbon is a special example - can it
really share 4 electrons?
–Electron promotion! Another example- show how water
is formed with covalent bonds.
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Water
H
O
Each hydrogen has 1 valence electron
Each hydrogen wants 1 more
The oxygen has 6 valence electrons
The oxygen wants 2 more
They share to make each other happy
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Water Put the pieces together The first hydrogen is happy The oxygen still wants one more
H O
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Water The second hydrogen attaches Every atom has full energy levels
H OHSample 16-1,
p.440
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Multiple Bonds Sometimes atoms share more than
one pair of valence electrons. A double bond is when atoms share
two pairs (4 total) of electrons A triple bond is when atoms share
three pairs (6 total) of electrons Table 16.1, p.443 - Know which
elements are diatomic (Oxygen?)
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Carbon dioxide CO2 - Carbon is central
atom ( more metallic ) Carbon has 4 valence
electrons Wants 4 more Oxygen has 6 valence
electrons Wants 2 more
O
C
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Carbon dioxide Attaching 1 oxygen leaves the
oxygen 1 short, and the carbon 3 short
OC
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Carbon dioxide Attaching the second oxygen
leaves both oxygen 1 short and the carbon 2 short
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more Requires two double bonds Each atom can count all the
electrons in the bond
OCO
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Carbon dioxide The only solution is to share more Requires two double bonds Each atom can count all the electrons in
the bond
OCO8 valence electrons
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Carbon dioxide The only solution is to share more Requires two double bonds Each atom can count all the electrons in
the bond
OCO8 valence electrons
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Carbon dioxide The only solution is to share more Requires two double bonds Each atom can count all the electrons in
the bond
OCO
8 valence electrons
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How to draw them? Add up all the valence electrons. Count up the total number of
electrons to make all atoms happy. Subtract; then Divide by 2 Tells you how many bonds - draw
them. Fill in the rest of the valence
electrons to fill atoms up.
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Example NH3, which is ammonia
N - has 5 valence electrons, wants 8
H - has 1 (x3) valence electron, wants 2 (x3)
NH3 has 5+3 = 8
NH3 wants 8+6 = 14
(14-8)/2= 3 bonds 4 atoms with 3 bonds
N
H
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N HHH
Examples Draw in the bonds All 8 electrons are accounted for Everything is full
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Example HCN: C is central atom N - has 5 valence electrons, wants 8 C - has 4 valence electrons, wants 8 H - has 1 valence electron, wants 2 HCN has 5+4+1 = 10
HCN wants 8+8+2 = 18
(18-10)/2= 4 bonds 3 atoms with 4 bonds -will require multiple
bonds - not to H however
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HCN Put single bond between each atom Need to add 2 more bonds Must go between C and N
NH C
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HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add to
equal the 10 it has
NH C
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HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add Must go on N to fill octet
NH C
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Another way of indicating bonds
Often use a line to indicate a bond Called a structural formula Each line is 2 valence electrons
H HO = H HO
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Structural Examples
H C N
C OH
H
C has 8 e- because each line is 2 e-
same for N
same for C here same for O
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A Coordinate Covalent Bond... When one atom donates both
electrons in a covalent bond. Carbon monoxide CO
OC
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Coordinate Covalent Bond When one atom donates both
electrons in a covalent bond. Carbon monoxide CO
OC
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Coordinate Covalent Bond When one atom donates both
electrons in a covalent bond. Carbon monoxide CO
OCC O
Shown as:
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Coordinate covalent bond Most polyatomic cations and anions
contain covalent and coordinate covalent bonds
Table 16.2, p.445 Sample Problem 16-2, p.446
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Bond Dissociation Energies... The total energy required to break the
bond between 2 covalently bonded atoms
High dissociation energy usually means unreactive
Table 16.3, p448 Sample: Calculate the kJ to
dissociate the bonds in 0.5 mol CO2
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Resonance is... When more than one valid dot
diagram is possible. Consider the two ways to draw
ozone (O3) Which one is it? Does it go back and forth? It is a hybrid of both, like a mule;
shown by a double-headed arrow
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Exceptions to Octet rule For some molecules, it is impossible
to satisfy the octet rule–usually when there is an odd
number of valence electrons
–NO2 has 17 valence electrons, because the N has 5, and each O contributes 6
impossible to satisfy octet, yet the stable molecule does exist
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Exceptions to Octet rule Consider electrons as small,
spinning electrical charges creates a magnetic field when paired, they cancel each
other, because they are spinning in opposite directions
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Exceptions to Octet rule Substances in which all the
electrons are paired are called diamagnetic
–weakly repelled by external magnetic field
paramagnetic- substances that contain one or more unpaired e-
–attracted to external mag. field
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Exceptions to Octet rule Do not confuse with ferromagnetism
–attraction of Fe, Co, Ni to mag. fld. Oxygen: possible to write structure
with all electrons paired
–not true, because oxygen is paramagnetic
Another exception: Boron Top page 451 examples
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Section 16.2Bonding Theories
OBJECTIVES:
–Describe the molecular orbital theory of covalent bonding, including orbital hybridization.
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Section 16.2Bonding Theories
OBJECTIVES:
–Use VSEPR theory to predict the shapes of simple covalently bonded molecules.
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Molecular Orbitals are... Orbitals that apply to the overall
molecule, due to atomic orbital overlap. 2 types:
–1. Bonding orbital - energy is lower than the atomic orbitals from which it is formed
–2. Antibonding orbital - energy is higher than what formed them
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Molecular Orbitals Sigma bond- when two atomic
orbitals combine to form the molecular orbital that is symmetrical along the axis connecting the nuclei
Pi bond- the bonding electrons are likely to be found above and below the bond axis (weaker than sigma)
p.454 and 455
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VSEPR: stands for... Valence Shell Electron Pair Repulsion Predicts three dimensional geometry
of molecules. The name tells you the theory: Valence shell - outside electrons. Electron Pair repulsion - electron pairs
try to get as far away as possible. Can determine the angles of bonds.
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VSEPR Based on the number of pairs of
valence electrons both bonded and unbonded.
Unbonded pair are called lone pair. CH4 - draw the structural formula Has 4 + 4(1) = 8 wants 8 + 4(2) = 16 (16-8)/2 = 4 bonds
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VSEPR Single bonds fill
all atoms. There are 4
pairs of electrons pushing away.
The furthest they can get away is 109.5º
C HH
H
H
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4 atoms bonded Basic shape is
tetrahedral. A pyramid with a
triangular base. Same shape for
everything with 4 pairs. CH H
H
H 109.5º
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Other angles…p.456 Ammonia (NH3) = 107o
Water (H2O) = 105o
Carbon dioxide (CO2) = 180o
Note shapes in Fig. 16.16, p.457
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Linear 2 atoms around
central atom No lone pairs
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Tetrahedral
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4 atoms around central atom
No lone pairs
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Pyramidal
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3 atoms around central atom
One lone pair
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Bent
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2 atoms around central atom
2 lone pairs
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Trigonal planar
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3 atoms around central atom
No lone pairs
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Section 16.3Polar Bonds and Molecules
OBJECTIVES:
–Use electronegativity values to classify a bond as nonpolar covalent, polar covalent, or ionic.
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Section 16.3Polar Bonds and Molecules
OBJECTIVES:
–Name and describe the weak attractive forces that hold groups of molecules together.
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Bond Polarity Covalent bonding = shared electrons
–but, do they share equally? Electrons are pulled, as in a tug-of-
war, between the atoms nuclei
–In equal sharing (such as diatomic molecules), the bond that results is called a nonpolar covalent bond
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Bond Polarity When two different atoms bond
covalently, there is an unequal sharing
–the more electronegative atom will have a stronger attraction, and will acquire a slightly negative charge
–called a polar covalent bond, or simply polar bond.
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Bond Polarity Refer to Table 14.2, p.405 Consider HCl
H = electronegativity of 2.1
Cl = electronegativity of 3.0
–the bond is polar
–the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge
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Bond Polarity Only partial charges, much less
than a true 1+ or 1- as in ionic bond Written as:
HCl the positive and minus signs (with
the lower case delta ) denote partial charges.
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Bond Polarity Can also be shown:
the arrow points to the more electronegative atom.
Table 16.4, p.462 shows how the electronegativity can also indicate the type of bond that tends to form
H Cl
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Polar molecules Sample Problem 16-4, p.462 A polar bond tends to make the
entire molecule “polar”
–areas of “difference” HCl has polar bonds, thus is a polar
molecule.
–A molecule that has two poles is called dipole
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Polar molecules The effect of polar bonds on the
polarity of the entire molecule depends on the molecule shape
–carbon dioxide has two polar bonds, but is linear:
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Polar molecules The effect of polar bonds on the
polarity of the entire molecule depends on the molecule shape
–water also has two polar bonds, but the highly electronegative oxygen pulls the e- away from H:
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Attractions between molecules The weakest called van der Waal’s
forces - there are two kinds:
1. Dispersion forces
weakest of all, caused by motion of e-
increases as # e- increases
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2. Dipole interactions Occurs when polar molecules are
attracted to each other. Fig. 16.23, p.464 Dipole interaction happens in water
–positive region of one water molecule attracts the negative region of another water molecule.
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2. Dipole interactions Occur when polar molecules are
attracted to each other. Slightly stronger than dispersion
forces. Opposites attract, but not
completely hooked like in ionic solids.
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Dipole Interactions
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Hydrogen bonding Are the attractive force caused by
hydrogen bonded to F, O, or N. F, O, and N are very electronegative
so it is a very strong dipole. The strongest of the intermolecular
forces.
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Hydrogen bonding When a hydrogen is covalently
bonded to a highly electronegative atom, AND is also weakly bonded to an unshared electron pair of another electronegative atom.
–The hydrogen is left very electron deficient, thus it shares with something nearby
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Hydrogen Bonding
HH
O+ -
+
H HO+-
+
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Hydrogen bonding
HH
O H HO
HH
O
H
H
OH
HO
H
HO HH
O