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Unit 1.4 - Bonding
The forces holding atoms together in compounds are known as chemical bonds
Bonds are broken and formed in chemical reactions
Bonding involves electrons in the outer shells of atoms
We will be studying ionic, covalent and metallic bonding
The octet rule
The first important thing to know about bonding is that atoms will generally react and form bonds in
such a way that they end up with an outer shell full of electrons. This principle is expressed in the
octet rule:
“When elements react they tend to do so in such a way that results in an _________ shell
containing __________ electrons”
Dot-cross diagrams
Dot-cross diagrams are an easy way of keeping track of what happens to ____________
during chemical reactions
Only the ___________ _________ electrons are shown as others are not involved in reacting
We will look at some examples drawing dot-cross diagrams for the formation of ionic bonds:
Sodium + chlorine
Lithium + fluorine
Magnesium + chlorine
Magnesium + oxygen
Ionic bonding
Ions are charged particles formed when atoms __________ or __________ electrons
Cations are ____________ ions
Anions are _____________ ions
Metals generally form ____________
Non-metals generally form ___________
Anions and cations are held together by
strong electrostatic forces of ____________
otherwise known as ionic bonds
Ionic radii
An atomic or ionic radius, measured in nanometers (nm, 10-9 m), is basically the size of an
atom or ion
The following table shows atomic and ionic radii for some common atoms and ions:
Cations Anions
Na Na+ F F-
Mg Mg2+ O O2-
K K+ Cl Cl2-
Ca Ca2+ N N3-
Trend
Cations are ______________ than the
corresponding atoms
Taking electrons away means that the
remaining electrons are held more tightly by
the __________________ force from the
______________ charged nucleus
Trend
Anions are _______________ than the
corresponding atoms
Adding more electrons means that the
__________________ force from the
_______________ is shared by more
electrons so they are held less tightly
Isoelectronic ions
The electronic configurations of common atoms and ions are shown below in pairs:
Remember… Think purrsitive!
Na+ 1s22s22p6
Ne 1s22s22p6Cl- 1s22s22p63s23p6
Ar 1s22s22p63s23p6
Because these pairs have the same electron configurations we say that they are ________________
Now try this…
1) Use a periodic table (p69) to determine the electron configuration of the following ions:
N3- 1s 2s 2p
O2- 1s 2s 2p
F- 1s 2s 2p
Na+ 1s 2s 2p
Mg2+ 1s 2s 2p
2) What do you notice about these different ions?
3) The circles in the diagram below represent these ions. Label the circles to show the order
of ionic radii:
0.072 nm
0.171 nm 0.140 nm 0.133 nm 0.102 nm
4) Can you explain why you chose to label the ions in this order?
Ionic compounds
Ionic compounds, when solid, form crystals with characteristic shapes and structures
We can use X-ray diffraction to study the structures of crystals
Hint: Think about the electrons and what holds them in place around the nucleus…
Lattice structures
We call the precise crystal structure of an ionic compound a ______________ structure
The lattice structure of a compound is the most _____________ way of arranging the ions
The forces of ______________ between anions and cations make lattices very strong and
rigid
unit cell . unit cell
Sodium chloride (NaCl) Caesium chloride (CsCl)
Face-centred cubic structure
a.k.a. rock salt structure
Body-centred cubic structure
Coordination number = 6 Coordination number = 8
What does all this mean?
Coordination number -
Unit cell -
Why are they different?
Hint: Think about the ionic radii (Na = 102 nm, Cs = 167 nm, Cl = 181 nm)
Properties of ionic compounds
The physical properties of all ionic compounds are very similar
Consider sodium chloride (NaCl):
Property Observation/ measurement
Explanation
Melting temperature801 C
Boiling temperature1413 C
Thermal conductivityPoor
Electrical conductivity
Solid -
Molten -
Solution -
Hardness
Malleability
NB Other ionic compounds behave similarly; you need to know their general properties and be able to explain how they arise
Evidence for the existence of ions
Ions are obviously far too small for us to see, so how do we know that they definitely exist?
1. Electron density maps
Sodium chloride (NaCl) 4-methoxybenzoic acid Electron density maps are produced using X-ray diffraction
X-rays are fired through crystals and the interference patterns produced tell us something
about where the ___________________ are
The big difference between electron density maps for ionic and covalent compounds is…
This shows that…
2. Electrolysis Watch the electrolysis of copper(II) chromate(VI), draw a diagram and use
it to explain how this experiment suggests the existence of ions
Lattice energy “The enthalpy of formation of one mole of an ionic compound from
gaseous ions under standard conditions”
Making bonds is an _____thermic process, in other words: energy is _______________ as
things become more stable
Forming an ionic lattice is a very _____thermic process as the ions take up their most stable
possible configuration
The energy released when forming an ionic lattice is known as the _______________ energy
Born-Häber cycles
We can determine lattice energies experimentally by measuring a number of other enthalpies and
combining them in a diagram called a Born-Häber cycle to give the lattice enthalpy:
Born-Häber cycle for the formation of NaCl:
energy
Definitions of terms:
∆Hθf = Standard enthalpy change of formation, measured using a bomb
calorimeter
∆Hθat = Standard enthalpy change of atomisation - changing a substance into a
monatomic gas
∆Hθi1 = First ionisation energy - energy required to remove 1 electron and form a cation,
from spectroscopic measurements
∆Hθe = First elecron affinity - energy required to add 1 electron and form and anion, also
from spectroscopic measurements
∆Hθlat = The lattice enthalpy of the substance - the enthalpy of formation of one mole of an
ionic compound from gaseous ions under standard conditions, calculated from the B-
H cycle
Steps involved in drawing a B-H cycle:
1) Start with the reactants in their standard forms
2) Draw in the enthalpy of formation (downwards)
3) Atomise the metal (upwards)
4) Ionise the metal (upwards)
5) Atomise the non-metal (upwards)
6) Add an electron to the non-metal (downwards)
7) Use the gap to calculate the lattice enthalpy:
Prep - Use the values in the table below to construct Born-Häber cycles for the formation of potassium iodide and potassium chloride and hence calculate the lattice energies:
∆Hθf[KI(s)] = -162.8 ∆Hθ
at[K(s)] = 89.0 ∆Hθat[1/2 Cls(g)] =
122.0∆Hθ
e[Cl(g)] = -349.0
∆Hθf[KCl(s)] = -436.7 ∆Hθ
i1[K(g)] = 418.8 ∆Hθat[1/2 Is(g)] =
107.0∆Hθ
e[I(g)] = -295.2
All values given in kJ mol-1
What factors affect lattice energy?
Compound
Radius of cation (nm)
Radius of anion (nm)
Lattice energy (kJ mol-1)
Compound Radius of cation (nm)
Radius of anion (nm)
Lattice energy (kJ mol-1)
NaF 0.102 0.133 -918 NaCl 0.102 0.180 -780
NaCl 0.102 0.180 -780 KCl 0.138 0.180 -711
NaBr 0.102 0.195 -742 MgF2 0.072 0.133 -2957
NaI 0.102 0.215 -705 CaF2 0.100 0.133 -2630
Trend in lattice enthalpies of sodium halides Trend in lattice enthalpies of group 1 & 2 halides
a) Size of cation…
b) Size of anion…
c) Charge on cation/anion…
This can all be explained using Coulomb’s law:
In other words:
Smaller ions can get __________ together so the force of
attraction between them is _____________
The greater the ____________ on an ion the greater its power
to ________________ another oppositely charged ion
Predicting stability - Why not NaCl 2?
Q ~ What are the formulae of the following compounds:
i. Sodium chloride
ii. Magnesium chloride
iii. Potassium chloride
iv. Calcium chloride
Why do these compounds always have the same formulae? Why is sodium chloride never NaCl2 for
example or magnesium fluoride never MgF3?
F = k q1q2
r2
Where q1 and q2 are the charges on the ions, r is the distance between them and F is the force of attraction
Dot-cross diagram for NaCl Dot-cross diagram for NaCl2
Sodium loses _____ electron Sodium loses _____ electrons
Use the data-books (p20) to find
out the following information:
Sodium chloride
First ionisation energy of sodium
Second ionisation energy of sodium
Theoretical lattice energy of NaCl -2526 kJ mol-1
NaCl2 and MgCl3 will not form
because…
Magnesium chloride
First ionisation energy of magnesium
Second ionisation energy of magnesium
Third ionisation energy of magnesium
Theoretical lattice energy of MgCl3 -4500 kJ mol-1
Polarisation of bonds
Born-Häber cycles allow us to calculate lattice energies from ________________ results
Coulomb’s law allows us to calculate ________________ lattice energies
The table below compares experimental and theoretical lattice energies:
Q ~ What do you notice about the numbers in the table?
CompoundLattice energy / kJ mol-1
DifferenceBorn-Häber Theoretical
NaF -918 -912
NaCl -780 -770
NaBr -742 -735
NaI -705 -687
AgF -958 -920
AgCl -905 -833
AgBr -891 -816
AgI -889 -778
Hmmm….I think I see a pattern developing!
A polar bear!
But why is there a difference? We need to look closely at the bonding involved, specifically whether
there is polarisation:
(a) Totally ionic bond (b) Polarised ionic bond (c) Covalent bond
Ionic bonds can be distorted by the attraction of the positive ______________ for the outer
electrons of the negative _____________
This means that some ionic bonds show a degree of ________________ character
These polarised bonds are ______________ than totally ionic bonds and give rise to more
negative lattice energies than predicted theoretically; they are ____________ stable than
expected
What causes ionic bonds to become polarised?
Cation
Ionic radius - Small cations have
____________ charge density and so are
more polarising
Charge - Cations with a large positive charge
are __________ polarising because they
have stronger attraction for the outer
electrons
Anion
Ionic radius - Large anions hold their outer
electrons _________ tightly, so the larger the
anion, the more __________ it is polarised
Effect on physical properties
1. Because theoretical lattice enthalpies are calculated by assuming that the ions are spherical
and separate they cannot account for polarisation and so are generally lower than lattice
enthalpies calculated from experimental data using Born-Häber cycles.
2. The melting temperatures of silver halides are about 20% lower than those of the sodium
halides, which is in agreement with the idea that the silver compounds have a greater covalent
character to their bonding.
Covalent bonding
Covalent bonding, unlike ionic bonding which involves the _______________ of electrons,
involves the ______________ of electrons between atoms
Atoms share electrons in order to gain ________ ________ _________
Ionic or covalent?
Whether a compound involves ionic or covalent bonding is related to the potential lattice energy for
that compound:
Ionisation
energy
Lattice energy Ionisation
energy
Lattice energy
Ionic Covalent
Dot-cross diagrams
You’ve had lots of practice so you’ll find it easy to draw dot-cross diagrams for these:
Chlorine Water
Methane Oxygen
Electron density maps
Atoms consist of a nucleus containing __________
and ___________ which is surrounded by orbiting
_____________
In covalent bonds electrons are most likely to be
found in the region between the _____________
charged nuclei
The area of electron density between the nuclei
has a _____________ charge, which ____________ the nuclei inwards and keeps them
_______________ bonded together
How strong and long are covalent bonds?
The distance between two atoms involved in covalent bonding depends on two factors:
1. The attraction between the bonding _____________ and the nuclei
2. The repulsion between the like-charged _______________
How the two forces balance out can be shown in the following diagram:
Covalent bonds are of such a
length that attraction and repulsion
are balanced
Bond enthalpy is a measure of
bond strength
Bond Bond length / nm Bond enthalpy / kJmol-1
H-H 0.74 436
C-C 0.154 348
C=C 0.134 612
C-H 0.109 412
Dative covalent bonds
Sometimes both of the ______________ that make up a covalent bond come from the same atom,
this is called a dative covalent bond:
Ammonia
(NH4+)
Aluminium
chloride dimer
(Al2Cl6)
En
erg
y
Separation / nm
We represent dative bonds with an arrow:
Ammonia Aluminium chloride
Dative bonds are the same length and strength as normal covalent bonds
Another example of a compound with dative bonding is carbon monoxide (CO)
Shapes of covalent molecules
Different molecules have different shapes, remember methane, which is tetrahedral
Draw dot-cross diagrams for the following compounds:
Compoun
dDot-cross diagram Electron density diagram Shape
BeCl2
CH4
BF3
H2O
There is a strong force of _______________ between electron pairs
As a result of this the bonding and non-bonding pairs of electrons will spread out to be as far
apart from each other as possible
This gives rise to the molecular geometries shown above
Giant atomic structures
Many covalently bonded compounds exist as molecules e.g.
Others, like diamond (diagram right) exist as atomic
crystals or molecular crystals
These giant atomic structures form crystal lattices
held together by ___________ bonds
Other examples of giant atomic structures are silicon
carbide (SiC) and silicon(IV) oxide (SiO2)
Compounds with giant atomic structures all have
similar physical properties:
Metallic bonding
Properties of metals
How does the theory of metallic bonding explain the following properties of metals? (p92)
Electrical conductivity:
Most simple model of bonding in metals
Metal crystal containing metal cations
surrounded by a delocalised ‘sea’ of the
electrons that they have lost
High thermal conductivity:
High melting/boiling temperatures:
Malleability and ductility:
Predicting the properties of metals
The theory of _______________ bonding can be used to predict the properties of metals
The strength of metallic bonding depends on the charge density of the cloud of delocalised
_________________ and the metal ________
Higher charge density means _______________ electrostatic attraction
Lower charge density means ____________ electrostatic attraction
Q ~ Use the theory of metallic bonding to explain the differences shown in the following table:
a) Melting temperature…
Metal Melting temp / ºC Thermal conductivity / W/cm/K
Sodium 98 1.35
Magnesium 649 1.5
Copper 1083 3.85
b) Thermal conductivity…
Summary of chapter 1.4 - Bonding
Ionic bonding
Covalent bonding
Practice questions:
Ionic bonding1. Magnesium reacts spontaneously and extremely vigorously with fluorine gas
a) Write an equation, including state symbols, for the reaction of magnesium with fluorine.
b) Draw a dot-cross diagram to show the bonding in magnesium fluoride.
2. Compare and contrast the structures of sodium chloride and caesium chloride. Why do they differ?
3. Sodium chloride is, under standard conditions, an ionic solid and is used as a flavour enhancera) Under what conditions does sodium chloride conduct electricity? Explain your answer.
b) For which compound would you expect the lattice enthalpy to be more negative: sodium chloride or magnesium chloride? Explain your choice.
Metallic bonding
c) Why does sodium chloride always exist as NaCl and never NaCl2?
4. Why is the theoretical lattice enthalpy for silver iodide so different to the experimentally determined value?
Covalent bonding1.Define the term covalent bond.
2. Nitrogen is the most abundant gas in the Earth’s atmosphere. It occurs as a covalently bonded molecule, N2
a) Draw an electron density diagram for nitrogen
c) Draw a dot-cross diagram to show the bonding in nitrogen
3. Beryllium chloride, BeCl2, is a chemical compound that can be said to be ‘electron-deficient’. It forms pairs of linked molecules known as dimers when solid.
a) Draw a dot-cross diagram for a single BeCl2 molecule
b) Draw a dot-cross diagram for a Be2Cl4 dimer
c) What type of bonding is present in Be2Cl4?
4. Covalent bonds vary in length depending on what atoms are involved. Describe the factors which influence the length of a covalent bond; you can include a diagram as part of your answer.
Metallic bonding1. Sodium is shiny, grey metal with a relatively low melting temperature. It is also relatively soft and can be cut with a sharp knife at room temperature.
a) Describe the type of bonding that is present in solid sodium metal
b) Explain why iron is much harder than sodium
c) Explain why copper conducts thermal energy more readily than sodium
2. Most metals are said to be malleable. What does the word malleable mean and how does the theory of metallic bonding account for this property?
Now go away and REVISE! There is a test on this chapter just around the corner and that’s going to be closely followed by the school exam, which will also include questions on… guess what… BONDING!