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Topic 4. Chemical bonding

and structure

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There are three types of strong bonds:

 

– Ionic

 

– Covalent

 

– Metallic

 

 

 

 

 

?       
Some substances contain both covalent and ionic bonding or an
intermediate.

4.1 Ionic bonding

 

 

 

 

?       
Ionic bonding is an electrostatic attraction
between oppositely charged ions.

?       
One ore more electrons are transferred from the outer shell of one atom
to the outer shell of another atom.

 

?       
The charge of an ion depends on the number of electrons the atom needed
to loose or gain to achieve a full outer shell.

 

2 Na(s) + F2 (g) ?  2 NaF (s)

?     
The electrons are shown in
pairs, because each pair of electrons occupies an orbital.

 

 

 

?     
The successive energy
levels in the atoms and ions are shown getting closer together.

 

 

 

?     
The radius of a sodium atom
is approximately twice that of a chlorine atom.

 

 

 

?     
The radius of a sodium ion
is approximately half that of a sodium atom.

 

 

 

?     
The radius of a chlorine
ion is approximately twice that of a chlorine atom.

Cations

 

 

 

 

?       
If an
atom loses e-, it becomes a positively charged ion.

 

 

 

 

 

?       
Group 1:

 

?       
Group 2:

 

?       
Group 3

 

 

 

 

 

?       
Transition metals can form more than one ion, for example Cu+ and Cu2+, Fe2+ and Fe3+

Anions

 

 

 

 

?       
If an atom gains one or more e-, it becomes a negatively charged ion.

 

 

 

 

?       
Group
15:

 

?       
Group
16:

 

?       
Group
17:

Polyatomic ions

Ionic compounds

 

 

 

 

?       
Between metals (electropositive elements) and non-metals (elements with
high electronegativity).

 

 

 

 

?       
The difference in electronegativity values needs to be greater than
about 1.8.

Formulas of ionic compounds

 

 

 

 

?       
The
overall charge of the compound must be zero.

 

 

 

 

 

?       
Ex. CaF2

Lattice

 

 

 

 

 

 

 

 

?       
When an ionic compound is formed, the ions are packed in an organized
crystalline structure, a lattice.

 

 

 

 

?       
The sum of all the electrostatic attractions between the oppositely
charged ions is called the lattice energy.

?       
The lattice energy has a high value and this
energy is released when the ionic compound is formed.

 

 

 

 

?       
e.g. the formation of NaCl from Na(s) and Cl2(g) is an exothermic
reaction.

 

 

 

 

 

 

 

 

?       
The
value of lattice energy depends on:

 

?       
The
charge of the ions

 

?       
The size
of the ions

 

 

 

 

 

?       
The higher the value of lattice energy, the more stable is the ionic
compound.

Physical
properties

 

 

 

 

?       
Melting: The crystal structure is broken down, but there are still some
attractive forces between the particles.

 

 

 

 

 

?       
Boiling: The attractive forces between the particles are completely
broken.

 

 

 

 

?       
The stronger the bonds, the higher the boiling point.

Properties
of ionic compounds

 

 

 

?       High melting and boiling
points because of strong attractive forces between the ions in the
lattice (mp of Na 801º C)

 

 

 

 

?       Conducts electricity when
molten or dissolved in water.

 

 

 

 

?       When a salt dissolves, new
bonds are formed between the water molecules and the ions.

 

 

 

 

?       This process is called hydration
and the ions are said to be hydrated.

4.2 Covalent bonding

 

 

 

 

?       
Covalent bonding is the electrostatic attraction between a pair of
electrons and positively charged nuclei.

Multiple covalent bonds

 

 

 

 

 

 

 

 

 

 

 

?    Single bond: One shared electron pair with one electron from each atom.

 

?    Double bond: Two shared electron pairs with two electrons from each atom.

 

?    Triple bond: Three shared electron pairs with three electrons from each atom.

?       
The more pairs of electrons there are in a covalent bond:

 

–    the shorter the bond length

 

–    the stronger the bond

Polarity of molecules

?       
Molecules with polar bonds can be non-polar if they are
symmetrical, that is if the central atom is symmetrically surrounded by
identical atoms.

 

 

 

 

 

?       
In carbon dioxide the dipoles are exactly opposite in direction and
cancel each other.

 

 

 

 

 

O = C = O

Non-polar molecules

 

 

 

 

?       
In a chlorine molecule, the difference in electronegativities of the
atoms is 0.

 

 

 

 

?       
This means that the electronpair in the covalent bond is on average
shared EQUALLY between the 2 chlorine atoms.

 

 

 

 

 

?       
The bond is called a non-polar bond, thus making the molecule a non-polar
molecule.

Polar
molecules

 

 

 

 

?       
In hydrochloric acid, the difference in electronegativities is 1.0.

 

 

 

 

?       
The more electronegative chlorine atom draws the bonding pair of
electrons towards itself and becomes negatively charged.

 

 

 

 

?       
The
hydrogen atom then becomes positively charged.

 

 

 

 

 

?          
The bond
is polar and the molecule has a dipole moment.

4.3 Covalent structures

 

 

 

 

?       Lewis symbols show the
number of valence electrons of an element represented ass either dots or
crosses.

Drawing Lewis structures of molecules

 

 

 

 

 

Draw the Lewis structures for:

 

 

 

 

 

a) O2                                 b) N2                                 c) CO2                                     d) HCN

Shapes
of molecules and ions

 

 

 

 

 

 

 

 

?     The shape of a molecule or
ion can be predicted by the valence shell electron pair repulsion theory
(VSEPR).

 

?     The theory states that
electron pairs (= electron domains) repel each other, and are
therefore located as far away from
each other as possible.

 

 

?     The order of repulsion strength is:

 

 

 

lone pair-lone-pair > lone pair-bond pair
> bond pair-bond pair

?    If one or more of the
negative charge centres is a non-bonding pair, this will influence the final
shape of the molecule.

 

 

?    e.g NH3 and H2O

Resonance structures

 

 

 

?       
For some molecules it is possible to write more than one correct Lewis
structure.

 

 

 

 

?       
These structures are called resonance structures and true
structure is an intermediate form known as a resonance hybrid.

 

?       
Ex. All of the C-C bonds in
benzene have the same bond length:

Coordinate covalent bonds

 

 

 

 

?       
In coordinate covalent bonds (dative covalent bonds) the shared pair of
electrons comes from the same atom.

Covalent network solids

 

 

 

?       
Pure
carbon has several different structural forms:

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

?       
These forms have different physical properties and they are called allotropes.

 

 

 

 

?       
Allotropes are crystalline forms of the same element, in which the atoms
are bonded differently.

Silicon

 

 

 

 

?       
Tetrahedral
arrangement

Silicon
dioxide, SiO2 (quartz)

 

 

 

 

 

?       
Strong

 

?       
Insoluble
in water

 

?       
High
melting point

 

?       
Non-conductor
of electricity

?    
A common impure form of silicon dioxide is sand, which is colored
yellow by the presence of iron (III) oxide.

Metallic bonding

 

 

 

?       
Delocalized valence electrons move freely through the metal.

 

 

 

 

?       
The attraction between these electrons and the cations holds the piece
of metal intact.

Electrical
conductivity

 

 

 

 

 

 

 

 

 

?       
The delocalization electrons enables free movement in response to electric
fields.

Thermal conductivity

 

 

 

?       
Tight packing of cations and delocalized electrons transmit kinetic
energy rapidly.

Malleability

 

 

 

 

?       
Individual atoms are not held to any other specific atoms, hence atoms
slip easily past one another.

4.4 Intermolecular
forces

 

 

 

 

 

12

 

 

 

10

 

 

 

8

 

 Column 1

 

6  Column
2

 

 Column 3

 

4

 

 

 

2

 

 

 

0

 

Row 1                                                    Row 2                                                    Row 3                                                    Row 4

 

 

 

 

 

 

 

 

 

 

 

 

 

 

?       
Intramolecular forces:

 

–    holds the atoms together within a molecule

 

–   affects molecular geometry and reactivity

 

 

 

 

 

?       
Intermolecular forces:

 

–   between the molecules within a compound

 

–   affects melting and boiling points

London forces (dispersion forces)

 

 

 

?     Attractive forces that
exist between ALL atoms and molecules.

 

?     These forces are only temporary and very weak.

 

 

 

?     Compounds that only have
London forces have very low boiling points (they are gaseous at room
temperature)

Factors that affect the
magnitude of the London forces

 

 

 

 

1. Number of electrons in an atom

 

 

– The more electrons, the stronger the London forces.

 

 

 

The more
electrons, the further they are from the nucleus = less attraction ? the
electron cloud is more easily polarized

2. Size of the electron cloud

 

 

–    The longer the carbon
chain, the larger the electron cloud ? the stronger the London forces and the
higher the boiling point

 

 

3. Shapes of molecules

 

–   The more contact area for
the molecules, the stronger the forces.

 

 

 

?    Van der Waal´s forces are
due to the motions of electrons, which causes temporary dipoles.

 

 

?     These forces generally
increase in strength as the number of electrons in a molecule increases or if
the surface area between the molecules increases.

 

 

?     These forces are so weak
that non-polar molecules have low boiling-points (many of them are gases at
room temperature).

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Dipole- dipole bonding

 

 

 

 

?    Between permanent dipoles

 

 

?    The negative pole of one
polar molecule is attracted to the positive pole of another polar molecule.

Hydrogen bonding

 

 

 

 

?        
In molecules where hydrogen is directly bonded to a small highly
electronegative element such as oxygen, nitrogen or fluorine.

?    Small molecules can have
surprisingly high boiling points due to hydrogen bonds.

The lattice structure of ice

14.1 Further aspects of covalent bonding and
structure (HL)

 

?    The octet is the most
common electron arrangement because of its stability.

 

?    Exceptions:

a)    Fewer electrons (incomplete
octet) if the

central atom is a small atoms, e.g. Be and B

 

 

 

 

 

 

b)    More than eight electrons (expanded
octet) if the central atom is a 3rd row element or below,

 

e.g. P and S

Species
with five negative charge centres

 

 

 

 

 

 

 

?        
If a molecule has five charge centres and they all are bonding
electrons, the shape is triangular bipyramidal.

?       
If one or more of these five negative charge centres is a non-bonding
pair, this will influence the final shape of the molecule.

 

 

 

 

 

?       
One: Tetrahedron

 

 

 

 

?

Two:
T-shaped

ClF3

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Three:
Linear

?

I3

 

 

Species
with six negative charge centres

 

 

 

 

?       
Molecules with six charged centres that are all bonding have an octahedral shape, e.g. SF6.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

?       
One non-bonding pair: square pyramidal BrF5

 

?       
Two
non-bonding pairs: square planar XeF4

Formal Charge

 

 

 

?        
Formal charges are assigned to atoms that have an “abnormal” number of
bonds.

 

 

 

 

Formal charge

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

?        
Ex. For the nitrogen in ammonium: formal charge = 5- 8/2 – 0 = +1

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