autumn lecture 3 (bonding) (1)
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Lecture 3: Structure and Bonding (Chemistry in Context 8.1–8.5 & 10.1–10.8 & ChemFactsheet 5)
The Chemical Bond
-classification of bond type
The Ionic Bond
-formation of positive and negative ions
-ionic bonding & lattices
The Covalent Bond
-dot and Cross Diagrams
-simple & Giant covalent structures
-dative Covalent Bonding
The Metallic Bond
-metallic lattices and electron delocalisation
-alloys
Intermediate Bond Types
-electronegativity and polarized bonds
-polarised ionic bonds and Fajan’s Rules
Periodicity (4.2-4.4)
-repeating chemical properties of the
chemical elements that occurs when
layed out in order of atomic number (Z)
is called periodicity
-elements in same valence electron
configuration similar chemistry
-blocks of periodic
table are named
after the orbitals
housing valence
electrons:
s-block f-block d-block p-block
Cl
Br
Sr
Mg
Ca
I
s2 s2p5
Cation and Anion formation
first
ionisation
second
ionisation
M M+ M2+
X first
electron
affinity
second
electron
affinity
X- X2-
(endothermic)
(exothermic)
(endothermic)
(endothermic)
Ionic Bonding (8.2)
when K atom and Cl atom
brought together it eventually
becomes favourable for electron
to jump from K to Cl:
(now both K+ and Cl- have stable
octet (noble gas configuration)
(Lewis formula of KCl)
K Cl K Cl+_
+DH
heat evolved
(lattice energy)
Ionic Bonding (8.2)
-both metal and
non-metal seek stable
noble gas electron
configuration
+
Cl
Cl -
[Ar]4s1
[Ar]
[Ne]2s22p5
[Ar]
chloride ion
chlorine
Sylvite
(KCl)
Ionic Bonding (8.2)
-both metal and
non-metal seek stable
noble gas electron
configuration
-MX ionic salts form
favourably when
M has a small ionisation
energy and X has a
large, exothermic
electron affinity
+
Cl
Cl -
[Ar]4s1 [Ne]2s22p5
[Ar]
chloride ion
chlorine
[Ar]
[Ar]
Dot-Cross Diagrams (8.2)
dot/cross diagrams show:
-compound’s electronic structure
-where the electrons originated
e.g. CaCl2:
[Ar]
[Ar]
Dot-Cross Diagrams (8.2)
dot/cross diagrams show:
-compound’s electronic structure
-where the electrons originated
e.g. CaCl2:
What salt is this?
[Ar]
[Ar]
Dot-Cross Diagrams (8.2)
dot/cross diagrams show:
-compound’s electronic structure
-where the electrons originated
e.g. CaCl2:
e.g. MgO:
[Ne] [Ne]
-salts are crystalline:
ions are arranged
regularly giving solids
with flat faces and
straight edges
Ionic Lattices (10.7)
small NaCl
crystals
-ionic bonds are very polar
dissolve well in polar solvents
solid salt (fixed ions)
electrical insulators
aqueous salt solution or molten salt
(ions free to move in liquid)
conduct electricity
-ionic bonds are strong!
high melting points
(750 °C for KCl, 3000 °C for MgO)
Physical Properties of Salts (10.8)
giant CaSO4 crystals, Mexico
What is a chemical bond? (8.1)
-an attractive force holding together atom(s) or ions(s) that makes
them function as a unit.
-bond forms if it makes the system is lower in energy than when
atoms are apart.
-energy input required to break bond = bond strength or bond energy
e.g. the ClCl bond has a strength of 243 kJ/mol
the O=O bond has a strength of 499 kJ/mol
the lattice energy of KCl is 715 kJ/mol
the lattice energy of MgO is 3930 kJ/mol
The way atoms are bonded together shapes physical and chemical
properties of a compound.
e.g. graphite is grey, soft, conductor of electricity
diamond is a transparent insulator, the hardest substance known
e.g. C and Si have similar chemistry but:
SiO2 is a brittle unreactive crystalline solid
CO2 is a gas
Types of Chemical Bond
-atoms can bond to each other in three main ways:
Ionic bonding (e.g. NaCl, CaCO3)
Covalent bonding (e.g. H2O, PCl5)
Metallic bonding (e.g. Fe, K, Hg)
Why Study Bonding?
Covalent Bonding
-a covalent bond forms when partially filled orbitals overlap:
-electrons are shared so that each contributing atom can experience a full outer shell of electrons (noble gas configuration)
-each H obtains full valence shell (n = 1) and the same electron
configuration as He H2 is stable
-other homonuclear diatomic molecules form by making similar
covalent bond
e.g. F has seven valence electrons so shares one electron:
dot-cross
structure of F2
each F obtains octet of electrons and electron configuration [Ne]
note each F has three non-bonded or lone pairs of electrons in the
valence shell (n = 2)
Octet rule: atoms proceed as far as possible toward completing
their octets by sharing electron pairs in covalent bonds
F
+ +
+
+
++
++ +
F
+ +
+ +
++
++
F +F
F is much more
electronegative than H
HF is a polar covalent bond
-in some diatomic molecules electrons are not distributed
evenly e.g. HF
H F
electron-poor electron-rich
dot-cross diagram
-covalent HF bond allows both
atoms to have full outer shell
(H obtains full n = 1 shell,
F obtains full n = 2 shell)
Electronegativity: the tendency of an atom to attract
electron density towards itself in a chemical bond
increasing electronegativity
What is Electronegativity? (8.2)
F is the most electronegative, Cs the least
The greater the difference in electronegativity between
two elements, the greater the polarity of the bond
If electronegativity difference >1.8
compound is mostly ionic:
P2O5 (difference = 1.4) completely covalent
LiI (difference = 1.5) partly ionic, partly covalent
Al2O3 (difference = 2.0) mostly ionic, slightly covalent
CsF (difference = 3.3) completely ionic
Polar Covalent Bonds
non-polar covalent: homonuclear diatomic
molecules e.g. H2, Cl2, both atoms identical
so electron density is arranged symmetrically
ionic bond – other extreme – very different
atoms so electrons completely transferred
very different electronegativity
H F
polar covalent: electrons
not transferred between
atoms but is unequally
shared due to slightly
different electronegativity
-polar bonds can give polarity to the overall molecule e.g.
-importantly for life, water is
a polar molecule:
polar molecules spontaneously align
themselves in an electric or magnetic field:
Cl
CClCl
Cl
H
CClCl
Cl
non polar
CCl4
polar
CHCl3
non-polar liquids unpeturbed
e.g. Hg, Br2, hydrocarbons
polar liquids deflected
e.g. H2O, alcohols, acetone
statically-charged
plastic rod
Dot-cross diagrams of polyatomic molecules (8.4)
-octet rule is obeyed for all atoms, even in larger molecules e.g.
ammonia (NH3)
N (1s22s22p3) is in group 5
must share 3 electrons to gain full octet
structure features three bonding pairs and one lone (non-bonded) pair
carbon dioxide (CO2)
O (1s22s22p4) is in group 6
must share 2 electrons to gain full octet
structure features four bonding pairs
(two double bonds) and
four lone (non-bonded) pairs
CO2 Lewis structure
Dative Bonding (8.5)
-sometimes both electron pairs from a covalent bond
are provided by one atom
e.g. formation of
ammonium ion:
-boron trifluoride is electron-deficient compound
(incomplete octet; only 6 electrons in outer shell)
forms dative bond with electron donors e.g. H2O, NH3
BF3
Metallic Bonding (10.5)
- close packing of metals has important
consequences:
-atoms’ outer (valence) electrons are
delocalised
(can move freely throughout lattice)
-metal structure is array of positive ions
immersed in a sea of electrons
-delocalised electrons:
-bind atoms strongly - strong high melting points (e.g. Fe 1530 ºC, W 3500 ºC)
-allow conduction of heat and electricity throughout solid (especially Ag, Cu)
delocalised ‘sea’
of electrons
Metallic Bonding (10.5)
-bonding between atoms is not
directional
so metals are
bendy and:
i) malleable - can be beaten into
thin sheets without fracturing
ii) ductile - can be drawn into
fine wires without fracturing
-close-packing of atoms explains metals’
high density:
8 gcm-3 for Cu and Fe
as high as: 22 gcm-3 for Os, Au & Ir
Alloys (10.5)
-man-made mixture of metals
-made by mixing molten metals
in desired proportions
-metallic bonding but altered properties
e.g. bronze (10% Sn in Cu) – stronger than Cu
Cu-Be alloys are
strong and spark
resistant (used
on oil rigs)
Cr
Cr C C
weak (layers slide)
iron stainless steel
strengthens and
makes corrosion
resistant
solder (Pb-Sn)
very low melt point)
bronze
statue
Network Solids (Giant Covalent Lattices) (10.6)
-bonding is covalent and in an infinite lattice – ‘one giant molecule’
-e.g. diamond, quartz, ruby, sapphire – brittle but very hard
Quartz (SiO2) – a three-dimensional infinite
lattice of tetrahedral Si atoms linked by O atoms
-sand is mostly made of silica (10.10)
amethyst
Network Solids (Giant Covalent Lattices)
-bonding is covalent and in an infinite lattice – ‘one giant molecule’
-e.g. diamond, quartz, ruby, sapphire – brittle but very hard
diamond – a three-dimensional lattice of tetrahedral carbon atoms
-lattice very rigid: the hardest substance known
used in drill-bits and (powdered) as an abrasive
good thermal conductor (rigidity transfers atomic vibrations)
but electrical insulator (no delocalised electrons)
graphite - a two-dimensional array of trigonal carbon atoms
-hexagonal rings - strong bonding within the layers but weak
bonding between them
graphite is soft (used in pencils and as lubricant).
non-bonded electrons are delocalised throughout plane
graphite conducts electricity
examples bonding and
structure
conduct-
ivity
m.p./
b.p.
strength
metals Cu, Al,
Os, Hg,
brass,
solder
metallic lattice -
positive metal
ions held
together by
delocalised
electrons
good high malleable,
ductile,
bendy
ionic solids NaCl,
CaCO3
ionic lattice very low
(except in
solution)
high hard, rigid,
brittle
network
solids
graphite
diamond,
SiO2, BN
covalent lattice very low
(except
graphite)
very
high
very hard,
rigid,
brittle
molecular
solids
ice, sugar,
wax, I2,
PCl5, S8.
covalent
molecules held
by lattice of
weak forces
very low low soft, brittle
Summary: Bonding in Solids
-bonding in solids can be grouped into four categories :
metallic, ionic, giant covalent and simple covalent
Al CaCO3
S
C
What kind of bonding is
holding together the following?
candle
wax
SA GA
limestone
cliff
brass
ice
Hg
Classifying Bonding
‘bonding’ is general term referring to forces that hold together any
type of chemical species: molecules, groups of molecules, atoms or ions
Intramolecular bonding holds together atoms e.g.
-ionic bonding
-metallic bonding
-covalent bonding
-strong (150-500 kJ/mol)
Intermolecular bonding acts between molecules e.g.
-van der Waals’ forces
-dipole-dipole forces
-ion dipole forces
-hydrogen bonding
-weak (2-25 kJ/mol) but often responsible
for bulk, physical properties of matter
molecule’s functional groups
Functional Groups and Physical Properties
determine
strength of intermolecular forces present
determines compound’s
physical properties
e.g.
H3CO
CH3O
HH
H3CCH3
CH3H3C
H3CCl
CH3Cl
INTERMOLECULAR FORCES SUMMARY
(non-covalent interactions)
van der Waals’ force
(1-5 kJ/mol)
dipole-dipole interaction
(5-10 kJ/mol)
hydrogen bonding interaction
(10-30 kJ/mol)
d- d+
d+ d-
d+ d-
d- d+
weakest force
(lowest boiling
point)
stronger force
(higher boiling
point)
-what determines if one compound is soluble in another?
Functional Groups and Intermolecular Forces: Solubility
-polarity of a molecule’s bonds determine in which solvents
it is soluble
i.e. compound likely to be soluble in particular solvent if the
intermolecular forces are similar in compound itself and the solvent:
e.g.
water and
short-chain
alcohols are
miscible
methanol and
acetone are
miscible
“like dissolves like”
Functional Groups and Intermolecular Forces: Solubility
H3CO
H
OC
H3C
CH3
HO
H
HO
C2H5
H
OH5C2
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