Transcript Slide 1

Chemical Bonding
Unit 5
Introduction
 Chemical bonding refers to the attractive
forces that hold atoms together in a compound
 Two major classes of bonding:
 Ionic
 Results from electrostatic interactions among ions
 Involves transfer of electrons
 Takes place b/w metals & non-metals
 Covalent
 Results from sharing one or more electron pairs b/w
atoms
 Takes place b/w non-metals
Ionic Compounds
1.
2.
3.
4.
5.
They are solids with high m.p.
(typically > 4000C)
Many are soluble in polar
solvents e.g. water
Most are insoluble in non-polar
solvents
Molten compounds conduct
electricity well because they
contain mobile ions
Aqueous solutions conduct
electricity well because they
contain mobile ions
Covalent Compounds
1.
2.
3.
4.
5.
They are gases, liquids or solids
with low m.p. (typically < 3000C)
Many are insoluble in polar
substances
Most are soluble in non-polar
solvents e.g. hexane, C6H12 &
carbon tetrachloride (CCl4)
Liquid and molten compounds do
not conduct electricity
Aqueous solutions are usually
poor conductors of electricity
because most do not contain
charged particles
Fig. 7-CO, p. 250
Lewis Dot Formulas of Atoms
 The number and arrangement of valence electrons
determine:
 Physical properties
 Chemical properties
 Chemical bonding
 Lewis dot formulas or Lewis structures show these
chemically important valence electrons
 Shows only electrons in outermost s and p orbitals as dots
 Electron pairs are represented as a pair of dots & an
unpaired paired electron as a single dot
 Not useful for representing transition and inner transition
metals
Table 7-1, p. 252
Formation of Ionic Compounds
 Ion  atom or group of atoms that carries an
electrical charge
 Cation- positively charged; more protons than
electrons
 Anion- negatively charged: more electrons than
protons
 Monoatomic ions – consists of only one atoms e.g.
Cl-, Mg2+
 Polyatomic ions - consists of a group of
covalently bonded atoms e.g. NH4+, SO42-
Formation of Ionic Compounds
 Ionic bonding can occur easily when elements
that low EN and low IE (i.e. metals) react with
elements that have high EN and very negative
electron affinities (i.e. non-metals)
 Many metals are easily oxidized i.e. they lose
electrons to form cations
 Many non-metals are easily reduced i.e. they gain
electrons to form anions
The farther apart across the periodic table two Group A elements are, the
more ionic their bonding will be
Formation of Ionic Compounds
 Lewis dot formula for ionic compounds e.g.
 Na+ is isoelectric with Ne
 In contrast, Cl- is isoelectric with Ar
p. 254
Formation of Ionic Compounds
 General representation of the reaction of 1A
metals with 7A elements (halogens):
p. 255
Formation of Ionic Compounds
 Like other simple ionic compounds, NaCl exists in a
regular, extended array of +ve and –ve ions.
 Distinct molecules of solid ionic substances do not
exist, therefore referred to as formula units
Fig. 7-1, p. 255
Formation of Ionic Compounds
 Reaction of 1A metals with 6A elements e.g.
 very small in size of Li+ gives it a much higher charge
density (ratio of charge to size) than that of Na+
 Similarly, the O2- is smaller than Cl- because of its
smaller size and double charge
 The more concentrated charges & smaller sizes bring
the Li+ and O2- closer together in Li2O than the ions in
NaCl  stronger ionic bond
 This is consistent with higher mp of Li2O (>1700oC) than
NaCl (801oC)
Formation of Ionic Compounds
 Reaction of 2A metals with 6A elements e.g.
 Ca2+ is about the same size as Na+ but carries
twice the charge, so its charge density is higher
 Attraction b/w two small, highly charged ions
is high  very strong ionic bond
 M.p. of CaO is 2580oC
p. 256
† group 1A and 2A can also form peroxides and superoxides
Table 7-2, p. 257
Energy relationships in Ionic Compounds

Why does ionic bonding occur?
 Why is solid NaCl more stable
than a mixture of individual Na
and Cl atoms?
 Consider a gaseous mixture of Na
and Cl atoms:
 Step 1: 1st IE of Na atoms is a
positive value  less stable than
original mixture of atoms
 Step 2: energy change for the gain
of 1 mole e-s by one mole of Cl
atoms is given by the electron
affinity of Cl
 This –ve value lowers the energy
of the mixture, but the mixture of
separated ions Na+ and Cl- ions
is still higher in energy ( less
stable) than original mixture of
atoms
Fig. 7-2, p. 258
Energy relationships in Ionic Compounds
 Thus, the formation of
ions does not explain
why the process occurs
 The strong attractive
forces b/w ions of
opposite charges draw
the ions together in a
regular array
 The energy associated
with this attraction (step
3) is the crystal lattice
energy
 This further lowers the
energy to (147789)kJ/mol = -642kJ/mol
Formation of Ionic Compounds
 d-transition metals
 Have s electrons in outermost shell and one d electrons one energy level
lower (e.g. 3d4s in 4th period transition elements)
 Outer s electrons are always lost before d electrons
 d- and f-transition elements form compounds that are essentially ionic
3d
21Sc Ar  
21
Sc
Ar 
4p
4s
Configurat ion
Ar  4s 2 3d1

3d
3
4s
4p
Configurat ion
Ar 
p. 256
Formation of Ionic Compounds
 d-transition metals
3d
4s
4p
30 Zn Ar        
3d
4s
Configurat ion
Ar  4s 2 3d10
4p
2
Ar      
Zn
30
Many transition metals ions are
highly coloured
Configurat ion
Ar  4s 0 3d10
Covalent Bonding
 Covalent bonding occurs when the
electronegativity difference b/w elements
(atoms) is zero or relatively small.
 In covalent compounds the bonds b/w atoms
within a molecule (intramolecular bonding)
are relatively strong BUT the attractive forces
between molecules (intermolecular forces) are
relatively weak.
 Hence covalent compounds have lower mp and bp
than ionic compounds
Covalent Bonding e.g. b/w hydrogen atoms
a)
Two H atoms are separated by a
large distance
b)
As the atoms approach, the e- cloud
of each atom is attracted by the
+vely charged nucleus of the other
atom (blue arrow). At the same time
the e- clouds repel one another, as do
the two nuclei
c)
The 2 e-s can both be in the region
where the two 1s orbitals overlap;
the e- density is highest b/w the
nuclei of the two atoms
Fig. 7-3, p. 259
The potential energy of the H2 molecule as a function of the
distance b/w the to nuclei
• The bonded atoms are lower in energy (more stable) than the separate atoms
• The result of sharing is that each atoms gains an electron configuration of the
nearest noble gas
Fig. 7-4, p. 259
Polar and Non-Polar Covalent Bonds
 Nonpolar bond- the electrons are shared
equally b/w atoms e.g. H2 or H- H
 Both H atoms have the same electronegativity 
the shared e-s are equally attracted to both H
nuclei and spend equal amts of time near each
nucleus
The covalent bond in all homonuclear diatomic
molecules must be nonpolar
Polar and Non-Polar Covalent Bonds
 Polar bonds – the e- pairs are shared unequally
e.g. H-F
 There is a large difference in electronegativity b/w
H (2.1) and F (4.0)
 e-s spend more time close to the F nucleus
p. 276
The separation of charge in a polar
covalent bond creates an electric
dipole
p. 276
Polar and Non-Polar Covalent Bonds
 Each halogen can form a single bond to another
halogen to form an interhalogen
 Bond polarities decreases as the electronegativity differences
b/w atoms decreases:
p. 276
Dipole Moments
 The polarity of a molecule is
indicated by its dipole moment,
which is given by:
 μ = d x q where
 d = distance separating opposite
charges of equal magnitude
 q = magnitude of charge
Table 7-5, p. 277
Dipole Moments
 Measured by placing a
sample of the molecule in an
electric field
 Polar molecules e.g. H-F tend to line up slightly in a direction
opposite to the field.
 Non-polar molecules are not oriented in by an electric field
Fig. 7-5, p. 278
Dipole Moments
NOTE:
 Dipole moments of individual bonds can only
be measured for diatomic molecules
 Dipole moments reflect overall polarities of
molecules
 For polyatomic molecules overall dipole is
affected by molecular geometry and the
presence of lone pairs of e-s.
Bond lengths & Bond Energies
 The internuclear distance at which the attractive &
repulsive forces balance and the bond is most stable is
called the bond length
 Bond dissociation energy is the energy needed to
separate the atoms, breaking the covalent bond
p. 260
Comparison of Carbon-carbon bond lengths and energies
 Two nuclei are more strongly attracted to 2 shared e- pairs
than to 1 pair
 Atoms are pulled closer together & more difficult to pull apart
 Multiple bonds are shorter & are stronger than single bonds.
p. 260
 The greater the difference in electronegativity
between bonded atoms, the stronger the polar bond,
and the greater the "extra" bond energy.
Table 7-3, p. 261
Trends
 Stronger bonds tend to be
shorter
 Bonds b/w H and 2nd row
elements are very strong
 Bonds become longer &
weaker as atomic number
increases e.g. C-halogen
bonds
 Bonds b/w C and 2nd row
elements are reasonably
strong
 N-O & O-O bonds weaker
than C-C bonds because
of repulsion b/w unshared
e-s on N and O
Lewis Formulas
 A covalent bond is represented by writing each
shared electron pair as either:
 A pair of two dots b/w atoms
 As a dash connecting atoms
OR
p. 261
Lewis Formulas
 Showing double bonds…
p. 262
Lewis Formulas
 Polyatomic ions may also be represented in this way
 For the ammonium ion, only 8 valence e-s are shown even
though N has 5 valence e-s and each H has 1 (a total of 5
+4 = 9)
 The charge of +1 implies the species has one less e- than
the original atoms
• Lewis formula is an e- bookkeeping
method that is useful as a 1st guide to
bonding schemes
• They only show:
• # of valence e-s
• The # and kinds of bonds
• Order in which atoms are connected
• They are not intended to show 3D shape
of the molecule
p. 262
Guide to Writing Lewis Formulas
See Handout
Resonance
 A molecule or polyatomic ion for which two or more
Lewis formulas with the same arrangements of atoms
can be drawn is said to exhibit resonance.
 Resonance does not mean:
 The molecules changes from one structure to the next
 Experiments show that the C-O bond is neither a
double nor a single bond but has intermediate bond
length and strength.
p. 274
Resonance
 Another way to represent this situation is by
delocalization of bonding electrons:
 The dashed lines indicate that some of the e-s shared
b/w C and O are delocalized (spread across) all four
atoms
p. 275
Resonance
 Example: Draw two resonance structures for
SO2.
 Solution:
S
O
N = 1(8) + 2(8)
= 24eA = 1(6) + 2(6) = 18eS=N–A
= 6e- shared
 The resonance structures are:
 Or showing the delocalization of the electrons as follows:
p. 275
In Summary…
 Chemical bonding can be described as a continuum
that may be represented as:
∆E for the
bonding atoms
Bonding type
zero
non-polar covalent

intermediate
 large

polar covalent
 ionic
p. 279