Topic 4
Bonding.
SL + HL
Ionic bond
Covalent bond
Intermolecular forces
Hydrogen bond
Dipole-dipole attraction
van der Waals’ forces
Metallic bond
4.1 Ionic bond
Ions = charged particles
Ionic bond= the electrostatic bond between
positively and negatively charged ions
Ionic compound= a compound built from ions,
i.e. a salt
• Atoms want to get a “Noble gas electron
configuration”. One way to get that structure
is to throw away the valence electron or steal
some electrons to get a full outer shell.
Na
2,8,1
+
Cl
2,8,7

Na+
2,8
+
Cl2,8,8
• Sodium gives one electron to chlorine and
both will have noble gas configuration. Which
noble gas?
• Sodium ion: positive electric charge. Cation.
Chloride ion: negative electric charge. Anion.
Ionic crystals
• + charge particles and – charge particle
attracts each other in three dimension and
builds up a lattice/crystal. Strong electrostatic
forces in three dimensions. Each cation is
surrounded by anions and vice versa.
Ions
• Group 1: H+, Li+, Na+, K+, Rb+, Cs+, Fr+
• Group 2: Be2+, Mg2+, Ca2+, Sr2+, Ba2+
• Group 3?/13: B3+, Al3+, Ga3+
• Group 6?/16: O2-, S2-,
• Group 7?/17: F-, Cl-, Br-, I-
Naming compounds
• Positive ions have the name of the atom:
Sodium ion
• Negative ions have the name of the atom (or
almost) + the ending –ide: Chloride ion
• Sodium + chlorine  Sodium chloride
• Lithium + oxygen  Lithium oxide
Formula of ionic compounds
Na+
Mg2+
Al3+
ClO2N3Write the chemical formula of the compounds formed between
the positive and negative ions above.
Write the name of the ionic compounds
Transition metals
• Transition metals can often form more than
one ion, e.g.:
– Fe2+ and Fe3+
– Cu+ and Cu2+
• HL: coloured because of d-orbitals
Some common polyatomic ions
•
•
•
•
•
Nitrate NO3Hydroxide OHCa(OH)2
Sulphate SO42Carbonate CO32Hydrogen carbonate HCO3(Bicarbonate)
• Phosphate PO43• Ammonium NH4+
Formula of ionic compounds-2
Na+
Mg2+
NH4+
OHSO42PO43Write the chemical formula of the compounds
formed between the positive and negative ions
above.
Write the name of the ionic compounds
When can we expect an Ionic bond?
• The quick rule:
Metal + non-metal => Ionic compounds
(Salts)
CuSO4 (s)
Electronegativity
• Electronegativity is a measure of an atoms
power to attract electrons.
• On the right side in the periodic table (group
7,6,5) the atoms have high values, they attract
electrons readily. Best is Fluorine, e-neg 4.
• On the left side the values are low. Low ability
to attract electrons.
Electronegativity values
FONClBrISCH
Ionic bond or not- calculate difference
in electronegativity
• If you want to do a more “precise” estimation
you can calculate it with the help of
electronegativity values.
• If the difference in electronegativity > 1,7 then
you can say it's an Ionic bond.
Ionic bond or not?
Use electronegativity values in the Chemistry Data booklet
•
•
•
•
•
NaCl
MgO
AlCl3
SiO2
Ca3N2
Typical properties of Salts
Hard, brittle,
• Conduct electricity in solution or melted
• High melting points => Strong bond
• Hydration of Ion in Water solution
4.2 Covalent bond
Electron pair bond, molecular bond
• If the De-neg < 1,7 the bond is considered to
be covalent
• Often between non-metals
• Polar or non-polar
• In a covalent bond the atoms share electrons with each
other to get a “Noble gas electron configuration”
• The bond has a direction (one atom to an other atom)
O
H H
• One bond consists of two electrons, an electron pair
• Single bond: the two atoms share two
electrons (1 pair)
• Double bond: the two atoms share four
electrons (2 pairs)
• Triple bond: the two atoms share six electrons
(3 pairs)
Lewis structures
• all valence electrons marked by dots or lines.
Draw Lewis structures for:
• F2
• NH3
• CO2
• N2
• C2H4
Number of bonds bond lengths and bond
strengths
• As the number of shared electrons increase (single to
triple) the bond lengths shortens and the bond
energy increase
Bond
Bond type
Lengths (pm)
Energy (kJ/mol)
CC
Single
154
347
CC
Double
134
614
CC
Triple
120
839
-COOH
Single
143
358
-COOH
Double
123
745
Non-polar covalent bond
• In, for example, H2 the two electrons in the
bond are shared equally between the two
hydrogen atoms
• H-H De-neg =0
• The electron distribution is symmetrical
Polar covalent bond
• If two different atoms form a covalent bond there
will be a difference in De-neg
• The atom with highest electronegativity will have
the electrons closer; they don’t share equally
• Unsymmetrical electron distribution
•
•
a)
b)
c)
d)
Which molecules contains at least one polar
covalent bond?
F2
HF
NF3
SiF4
Which bond in the pairs below have the
highest polar character?
C-O, N-O
H-O, S-O
H-O, H-S
Se-S, Se-F
Ionic, polar or non-polar covalent
bond?
Cl-Cl
H-Cl
Na+ Cl-
• % ionic character of a bond: 0-90%
(there are no 100% ionic compounds)
Dative covalent bond
• In a “normal” covalent bond both atoms
contribute with electrons to the bond.
• Sometimes only one atom contributes with
both electrons (the electron pair) to the bond
• Then the covalent bond is called a dative
covalent bond
Examples of dative bonds
• H2O + H+  H3O+
• NH3 + H+  NH4+
• C (4 ve-)+ O (6 ve-)  CO
VSEPR
• Valence Shell Electron Pair Repulsion
• Shape and bond angles
• Determine the molecules structure, the shape in
3 dimensions
• Structure around a given atom is determined
principally by minimising electron-pair repulsion
• Bonding or non-bonding pairs will be as far apart
as they can.
Linear
• Two negative centers
• 180o
• E.g. CO2
Trigonal planar (flat)
• Three negative centers
• 120o
• E.g. BF3
Tetrahedral arrangement
• Four negative centers
• 109.5 o
Method
1.
2.
3.
4.
Draw Lewis structure
Count electron pairs, minimise the repulsion
Positions of the atoms
Name of the structure
Ammonia, NH3
Tetrahedral, Trigonal planar 107o, one non-bonding
electron pair
• Non-bonding pair (lone pair) takes more space =>
reduce bond angels
Water, H2O
Tetrahedral, Non-linear (bent) 104o Two non-bonding
electron pairs
Non-bonding pair (lone pair) takes more space =>
reduces bond angles
Non-polar and polar
molecules (dipoles)
• Based on bond polarity and molecular shape
• May predict how a molecule will behave with
other compounds
• Polar molecule = Dipole
A polar molecule (a dipole)
• Must have polar covalent bonds.
– Look at the difference in electronegativity.
FONClBrISCH
AND
• Unsymmetrical shape according to charge
distribution.
• Otherwise it will be a non-polar molecule (NOT a
dipole)
Dipole or not?
•
•
•
•
•
H2O
CO2
NH3
CH4
CH3OH
Allotropes
• Some elements can be found in different
forms
• E.g:
– Carbon: Diamond, graphite, C60fullerene
– Oxygen, Ozone
See PPT: Carbon allotropes
Silicon
• Metalloid, Semiconductors, non-metallic
structure
• Similar structure as diamond
Silicon dioxide
• SiO2 Silica, giant structure similar to diamond, quarts
• Silicates, NaSiO4, tetrahedrical, silicon-oxygen single
bond
4.3 Intermolecular forces
• Holds molecules together in liquids or solids (No
Intermolecular forces between gaseous particles)
• Weaker than covalent and ionic bonds
• Hydrogen bond(Quite strong)
• Dipole-dipole
(Middle weak)
• van der Waal’s forces (~ London dispersion
forces)
(Very weak)
• Accounts for differences in aggregatio state and
solubility
van der Waal’s forces
• “Vibrations” in the electron cloud => Temporary
dipoles.
• A temporary dipole in one
molecule can induce a
temporary dipole in another
molecule
• Exist between all molecules
van der Waal’s forces, cont.
• The strength increases with molar mass of the
molecule/atom
E.g. He b.p 4 K
Xe b.p. 165 K
• Only effective over short range so the molecule
“area” is also important.
E.g: Pentane, C5H12, b.p. 309 K
Dimethylpropane, (CH3)4C b.p. 283 K
Dipole-dipole bond
• Electrostatic attraction between molecules
with permanent dipoles
• Stronger than van der Waals bond
Hydrogen chloride
M= 36,5 g/mol b.p. 188 K
Fluorine
M= 38 g/mol
b.p. 85K
Hydrogen bond
• In molecules that contain Hydrogen bonded to
Oxygen, Nitrogen or Fluorine (high electronegativity
and non-bonding electron pair)
• Stronger than dipole-dipole bonds
• Interaction of the non-bonding electron pair in
one molecule and hydrogen (with high positive
charge) in another molecule.
Which intermolecular bond?
• H2O b.p.= 100oC
H2S b.p.= -61oC
• NH3 b.p.= -33 oC
PH3 b.p.= -88oC
Which intermolecular bond?
• C2H6 b.p.
b.p. -89 oC
CH3CHO
20 oC
C2H5OH
78 oC
Boiling points of hydrogen compounds
What kind of intermolecular bonds can be
expected to dominate between the
molecules/atoms?
a) NH3(l)
b) C5H12(l)
c) Br2(l)
d) HI(I)
e) Ar(l)
f) PCl3(l)
g) CH3OH(l)
h) CHCl3(l)
4.4 Metallic bonding
• Metals have low electronegativity
• The atoms are packed close together in a
lattice
• The valence electrons are delocalised among
all atoms.
– The valence electron have no “home”
– The atoms can be seen as positive ions in a see of
electrons that keep them together
This can explain the metallic properties
• Electrical conductivity: electrons float around.
If you put in one, one will fall out.
• Malleability (smidbarhet) and Ductility
(sträckbarhet): if the atom is pushed from it’s
location the electron will follow. The bond is
between the ion and the electrons not
between the ions.
Summary
• The different kind of bonds is very important
for the behaviour of a compound, solution or
a mixture
4.5 The strength/type of the bond
affects:
• Melting points (impurities lower the melting
point)
• Boiling points
• Volatility (how easy a compound will convert
to gas)
• Electrical conductivity
• Solubility and miscibility
Changes in state (1)
Fe (s)
1538 ºC
Fe (l)
2861 ºC
801 ºC
NaCl (s)
Fe (g)
1413 ºC
NaCl (l)
Which bonds are broken?
Which bond is strongest?
NaCl (g)
Changes in state (2)
H2O (s)
0 ºC
H2O (l)
100 ºC
-85 ºC
HCl (s)
-61 ºC
HCl (l)
HCl (g)
-162 ºC
-184 ºC
CH4 (s)
H2O (g)
CH4 (l)
Which bonds are broken?
Which bond is strongest?
CH4(g)
Solubility
• Unpolar compounds
• Polar compounds
• Molecules with van
der Waals bonds
• Ionic bonds
• Molecules with
hydrogen bonds
• Molecules with
dipol-dipol bonds
Intermolecular bonds in solutions
• Ion-hydrogen bonds or ion-dipole bonds
Giant
Metallic
Structure type
Property
Hardness and Variable hardmalleability
ness,
malleable
rather than
brittle
Melting and
Variable dep.
boiling points On No of
valence eElectrical and Good in all
thermal
states
conductivity
Solubility
Insoluble,
except as
alloys
Examples
Iron, copper
Giant
Ionic
Giant
Covalent
Hard and
brittle
Hard and
brittle
High
Very High
Not as solids, No
conduct in (aq)
or (l)
In Water
Insoluble
mostly
NaCl, Na2SO4 Diamond,
Molecular
Covalent
Usually soft
and malleable
unless
hydrogen
bonded
Low
No
Often more
soluble in other
than water
except if Hbonded
CO2, Cl2,