IB DP1 Chemistry
Bonding
What makes atoms join together to make compounds?
Topic 4: Bonding (12.5 hours)
4.1 Ionic bonding
4.2.6 Predict the relative polarity of bonds from electronegativity
4.1.1 Describe the ionic bond as the electrostatic attraction between values
oppositely charged ions.
4.2.7 Predict the shape and bond angles for species with four, three
4.1.2 Describe how ions can be formed as a result of electron
and two negative charge centres on the central atom using the
transfer.
valence shell electron pair repulsion theory (VSEPR).
4.1.3 Deduce which ions will be formed when elements in groups 1, 4.2.8 Predict whether or not a molecule is polar from its molecular
2 and 3 lose electrons.
shape and bond polarities.
4.1.4 Deduce which ions will be formed when elements in groups 5, 4.2.9 Describe and compare the structure and bonding in the three
6 and 7 gain electrons.
allotropes of carbon (diamond, graphite and C60 fullerene).
4.1.5 State that transition elements can form more than one ion.
4.2.10 Describe the structure of and bonding in silicon and silicon
4.1.6 Predict whether a compound of two elements would be ionic dioxide.
from the position of the elements in the periodic table or from their 4.3 Intermolecular forces
electronegativity values.
4.3.1 Describe the types of intermolecular forces (attractions
4.1.7 State the formula of common polyatomic ions formed by non- between molecules that have temporary dipoles, permanent dipoles
metals in periods 2 and 3.
or hydrogen bonding) and explain how they
4.1.8 Describe the lattice structure of ionic compounds.
arise from the structural features of molecules.
4.2 Covalent bonding
4.3.2 Describe and explain how intermolecular forces affect the
4.2.1 Describe the covalent bond as the electrostatic attraction
boiling points of substances.
between a pair of electrons and positively charged nuclei.
4.4 Metallic bonding
4.2.2 Describe how the covalent bond is formed as a result of
4.4.1 Describe the metallic bond as the electrostatic attraction
electron sharing.
between a lattice of positive ions and delocalized electrons.
4.2.3 Deduce the Lewis (electron dot) structures of molecules and 4.4.2 Explain the electrical conductivity and malleability of metals.
ions for up to four electron pairs on each atom.
4.5 Physical properties
4.2.4 State and explain the relationship between the number of
4.5.1 Compare and explain the properties of substances resulting
bonds, bond length and bond strength.
from different types of bonding.
4.2.5 Predict whether a compound of two elements would be
covalent from the position of the elements in the periodic table or
from their electronegativity values.
Ionic Bonding
Crystals: 7 ‘perfect’ crystal shapes
Halite- rock salt- sodium chloride
Sodium chloride is an ionic compound
with ions arranged in a lattice
Ions
charged particles with electrostatic attraction between them
Na+
Cl-
Sodium and chloride ions formed when
electrons transfer
Na
2,8,1
+
Cl
2,8,7

Na+
2,8
+
Cl2,8,8
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-
Which is the smallest ion?
Na+
+3
Al
Cl3P
 Different sized atoms give
different mineral structures
as they pack in a different
way
Two or more electrons
can be transferred
Hexagonal Beryl
crystal; Image
Wikipedia
What is the formula of iron (III) oxide?
Fe2O
FeO
Fe3O2
Fe2O3
Polyatomic ions: charge distributed over
more than one atom
For example phosphate, PO4-3
can be found in products of
reactions of phosphoric acid
Some common polyatomic ions
 Nitrate NO3 Hydroxide OH Sulphate SO42 Carbonate CO32-
 Hydrogen carbonate HCO3(Bicarbonate)
 Phosphate PO43 Ammonium NH4+
Common Cations
Common Name Formula Alternative
name
Simple Cations
Aluminium
Al3+
Calcium
Ca2+
Copper(II)
Cu2+
cupric
Hydrogen
H+
Iron(II)
Fe2+
ferrous
Iron(III)
Fe3+
ferric
Magnesium
Mg2+
Mercury(II)
Hg2+
mercuric
Potassium
K+
kalic
Silver
Ag+
Sodium
Na+
natric
Polyatomic Cations
Ammonium
NH4+
Hydronium
H3O+
Common Anions
Common Name Formula
Simple Anions
Chloride
Cl−
Fluoride
F−
Bromide
Br−
Oxide
O2−
Polyatomic anions
Carbonate
CO32Hydrogen
HCO3−
carbonate
Hydroxide
OH−
Nitrate
NO32Phosphate
PO43Sulfate
SO42Anions from Organic Acids
Ethanoate
CH3COO−
Methanoate
HCOO−
Ethandioate
C2O4−2
Cyanide
CN-
Alternative
name
bicarbonate
acetate
formate
oxalate
Careful with...
 name of atom can change when ion is formed
chlorine atom (Cl)  chloride ion (Cl-)
 -ate is often a polyatomic ion with oxygen eg sulphate,
phosphate, etc.
 different ions often have similar names...
 nitrate NO3 nitrite NO2 nitride N-3
What is the formula of ammonium
sulphate?
 NH4SO4
 (NH4)2SO4
 NH4(SO4)2
 SO4(NH4)2
d-block (transition elements) can have
variable valencies
Mn2+
manganese(II)
Cr2+
chromium(II)/chromous
Mn3+
manganese(III)
Cr3+
chromium(III)/chromic
Mn4+
manganese(IV)
Cu1+
copper(I)/cuprous
Ni2+
nickel(II)/nickelous
Cu2+
copper(II)/cupric
Ni3+
nickel(III)/nickelic
Fe2+
iron(II)/ferrous
Pb2+
lead(II)/plumbous
Fe3+
iron(III)/ferric
Pb4+
lead(IV)/plumbic
Hg2+
mercury(I)/mercurous
Covalent bonding
Define electronegativity
Electronegativity is the tendency of an atom
to attract electrons towards itself. The atoms
with higher values attract electrons more
strongly.
Highest flourine (and rest of groups 7,6,5)
FONClBrISCH
Wikipedia table
How ionic is an ionic compound?
 bigger difference in electronegativity  more ionic
 (‘ionic’ usually De-neg> 1.8 difference)
 usually metal + non-metal
Which aluminium compounds will be ionic?
atom
Al
F
O
Cl
Br
electronegativity
1.5
4.0
3.5
3.0
2.8
Formula of aluminium
compound
De-neg
‘Ionic’ or ‘covalent’?
‘Sharing’ electrons De-neg < 1,7
covalent bonding forms molecules
Often between non-metals
Covalent bond formation- valence electrons
2, 4 or 6 electrons?
 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 (dot structures) show valence
electrons in pairs as dots, crosses or lines
skeletal formula for complex organic
molecules
Condensed formula
propanol CH3CH2CH2OH
Coordinate covalent bond (dative bond)
both electrons in the bond from the same atom
once formed, is the same as any other covalent bond
Bond lengths and Bond strengths
 As the number of shared electrons increases (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
Which bond has the highest bond polarity, δ
H-H
Cl-Cl
Al-F
Al-Br
Non-polar covalent bond
In, 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.
Bonds
100% Covalent bond  Polar covalent bond  Ionic bond
% ionic character of a bond: 0-90%
(there are no 100% ionic compounds)
Molecular shapes
What shape are molecules?
 VSEPR theory (Valence shell electron pair repulsion)
 pairs of electrons repel and sit as far away as possible from
each other
 double and triple bonds count as a pair
VSEPR: electron repulsion  molecular shape
 Structure of molecule given by pairs of electrons arranging
around an atom to be as far apart as possible
 non-bonded pairs repel more than bonded pairs
 double and triple bonds count as one
Build molecules from plasticine and
straws
 bond: 3cm length of straw
 atom: 1cm diameter plasticine ball
 unbonded pair of electrons 1cm straw length
Shapes of simple molecules
Number Name of shape
of charge
centres
Bond angles (s)
Example
2
linear
180
BeCl2
3
trigonal planar
120
BF3
4
tetrahedral
109.5
CH4
5
trigonal bipyramidal
90, 120, 180
6
octahedral
90, 180
http://en.wikipedia.org/wiki/Phosphorus_pentafluoride
http://en.wikipedia.org/wiki/Sulphur_hexafluoride
http://en.wikipedia.org/wiki/Boron_triflouride
Methane, Water and Ammonia
greater repulsion between
non-bonding pairs
smaller bond angles than
predicted
Intermolecular forces
Why do molecules stick together to form liquids and solids?
Intermolecular forces hold molecules
together, affecting physical properties
 Melting and boiling points
 Strength
 Flexibility
 Viscosity
Intermolecular forces
Hydrogen bond
strong
Dipole-dipole
weaker
van der Waal’s forces
weakest
Why do molecules attract each other to
make liquids and gases?
Intermolecular forces: electrostatic attraction between
 permanent dipoles (polar molecules)
 permanent dipole and a temporary dipole (induced
polarity)
 temporary diploes (induced polarity)
Why do molecules attract each other?
electrostatic attraction between…
 permanent dipoles (in polar molecules)
 temporary diploes
A dipole is a overall charge imbalance in a molecule.
Which of the following molecules are polar?
Induced dipoles in all molecules (van der
Waal’s forces)
Movements in electron cloud  Temporary dipoles.
Temporary dipole in one molecule can induce a temporary dipole
in another.
Image:
http://www.uwec.edu/boulteje/Boulter103N
otes/11December.htm
van der Waals forces
 The strength increases with molar mass of the molecule.
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
Is a molecule polar?
A polar molecule
 Has 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.
Molecular polarity
HF
NH3
H2O
Images: http://en.wikipedia.org/wiki/Molecular_polarity
Molecular polarity
 http://phet.colorado.edu/en/simulation/molecule-polarity
Dipole-dipole
Electrostatic attraction between molecules with permanent
dipoles.
Stronger than vdW.
Hydrogen chloride
M= 36,5 g/mol b.p. 188 K
Fluorine
M= 38 g/mol
b.p.
85K
Induced dipole
Image:
http://www.uwec.edu/boulteje/Boulter103N
otes/11December.htm
Polar and non-polar liquids are immiscible
Image: http://en.wikipedia.org/wiki/Petroleum
Hydrogen bonding
 H bonded to a highly electronegative element eg F, O or N
 proton strongly attracts electronegative element in another
molecule
 important in water
Image: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force
Hydrogen bond
 In molecules that contain Hydrogen bonded to Oxygen,
Nitrogen or Fluorine (high electronegativity and non-
bonding electron pair).
 Interaction of the non-bonding electron pair in one molecule
and hydrogen (with high positive charge) in another
molecule.
Examples
 H2O b.p.=100oC
H2S b.p.= -61oC
 NH3 b.p.= -33 oC
PH3 b.p.= -88oC
 C3H8 b.p.
b.p. 20 oC
CH3CHO
42 oC
C2H5OH
78 oC
Examples
 H2O b.p.=100oC
H2S b.p.= -61oC
 NH3 b.p.= -33 oC
PH3 b.p.= -88oC
 C3H8 b.p.
b.p. 20 oC
CH3CHO
42 oC
C2H5OH
78 oC
Ice
Image:
http://en.wikipedia
.org/wiki/Ice
Trends in physical
properties
How strong are the forces between molecules?
Bond type
Dissociation energy (kJ/mol)
Covalent
1600
Hydrogen bonds
50–70
Permanent dipoles
2–8
Induced dipoles
<4
Data: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force
Trends in physical properties
Plot one graph showing melting point and boiling point (in Kelvin) against
molar mass for the halogens
Describe the pattern (2 sentences)
Explain the pattern (2 sentences)
melting point /C
boiling point /C
Flourine
-220
-188
Chlorine
-102
-34
Bromine
-7
59
Iodine
114
184
Astatine
302
337
Data: http://en.wikipedia.org/wiki/Halogen
How strong are the forces between molecules?
Bond type
Dissociation energy (kJ/mol)
Covalent
1600
Hydrogen bonds
50–70
Permanent dipoles
2–8
Induced dipoles
<4
Data: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force
Allotropes: different structural forms of
the same element
Oxygen
O2 diatomic
oxygen
O3 ozone
http://catalog.flatworldknowledge.com/bookhub/4309?e=averill_1.0-ch18_s04
Allotropes of Carbon
Diamond
 Hard, colourless, insulator
 Tetrahedral, giant structure
 Covalent bond => sp3 orbitals.
Graphite
 Slippery, black, conductor
 Layers of fused six-membered rings. Each carbon surrounded
by three others in a planar trigonal arrangement => sp2 + porbital
 The p-orbital is perpendicular to the layer and give close
packed p-orbitals
 stabilise the layers
 Delocalisation of electrons => electrical conductivity
Fullerene, C60
 Spherical molecule. Looks like a football. 12 pentagons and 20
hexagons.
 Bonds: C60 –hydration C60H60
(C2H4 + H2  C2H6 ; 1 H2 / double bond)
Each carbon has a double bond
Silicon
 Metalloid, Semiconductors, non-metallic structure
 Similar structure as diamond.
Silicon dioxide
 SO2 Silica, giant structure similar to diamond
 Silicates, SiO4, tetrahedrical, silicon-oxygen single bond
Physical properties
 Melting points (impurities lower the melting point)
 Boiling points
 Volatility (how easy a compound will convert to gas)
 Electrical conductivity
 Solubility
Properties
Structure type
Property
Hardness and
malleability
Melting and boiling
points
Giant
Metallic
Variable hard-ness,
malleable rather
than brittle
Giant
Ionic
Hard and brittle
Solubility
Insoluble, except as
alloys
Examples
Iron, copper
Molecular
Covalent
Hard and brittle
Usually soft and
malleable unless
hydrogen bonded
Very High
Low
Not as solids,
conduct in (aq) or
(l)
In Water mostly
No
No
Insoluble
Often more soluble
in other than water
except if H-bonded
NaCl, Na2SO4
Diamond,
SiO2 (Sand)
CO2, Cl2, ethanol,
sugar
Variable dep. On No High
of valence e-
Electrical and
Good in all states
thermal conductivity
Giant
Covalent
Ionic salts
 Typical properties
 Hard, brittle,
 Conduct electricity in solution or melted.
 High melting points => Strong bonds
 Hydration of Ion in Water solution
Metallic bond
 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 its location the electron will follow. The
bond is between the ion and the electrons not between the
ions.
Investigate a physical property of a
mixture related to intermolecular forces
Quantitative independent variable
(cause)
Quantitative dependent variable (effect)
viscosity, deflection by charged object, or
other physical property
Links
 Ionic bonding
http://www.teachersdomain.org/asset/lsps07_int_ionicbonding
/
 Covalent bonding
http://www.teachersdomain.org/asset/lsps07_int_covalentbon
d/
Polarity links
 http://phet.colorado.edu/en/simulation/molecule-polarity
 Viscosity http://www.youtube.com/watch?v=3KU_skfdZVQ
 States of matter http://phet.colorado.edu/en/simulation/states-
of-matter
Polarity links
 http://phet.colorado.edu/en/simulation/molecule-polarity
 http://antoine.frostburg.edu/chem/senese/101/liquids/faq/h-bonding-vs-london-
forces.shtml
 States of matter http://phet.colorado.edu/en/simulation/states-of-matter
 http://employees.oneonta.edu/viningwj/modules/CI_dipoleinduced_dipole_forces
_13_5a.html
 Notes: http://www.uwec.edu/boulteje/Boulter103Notes/11December.htm
 Snowflakes: http://www.its.caltech.edu/~atomic/snowcrystals/class/class.htm
 Ice crystals http://www.edinformatics.com/interactive_molecules/ice.htm
Links
 http://phet.colorado.edu/en/simulation/molecule-shapes
 http://en.wikipedia.org/wiki/Phosphorus_pentafluoride
 http://en.wikipedia.org/wiki/Sulphur_hexafluoride
 http://en.wikipedia.org/wiki/Boron_triflouride
Teaching notes