Transcript Patterns in the Periodic Table

Patterns in the Periodic Table
The periodic
table is made up of columns
(groups) and horizontal rows – (periods).
The elements in Groups have similar chemical
properties – due to the same number of outer
electrons .
Metals are on the left hand side. Non-metals are
on the right hand side.
Atomic Size
 As
we move from left to right along a period the atomic
size decreases.
 This is because we are adding in more electrons in the
same energy level.
 There will be a corresponding increase in the number of
protons in the nucleus.
 Therefore, more + and – charges which are attracted to
each other and so the atom will be smaller.
Atomic size
As
we go down a group the Atomic Size
increases.
This is because we are adding more energy
levels.
The electrons on the outside are “shielded” from
the nucleus.
Therefore, less attraction and the atom is bigger.
Ionisation Energy
 The
First Ionisation Energy is – the energy required to
remove 1 electron from each atom in 1 mole of an
element in the gaseous state!
 Example
 K(g) —>K+(g) + eΔH = + 425 kJ mol/l
 As we go down a group the First Ionisation Energy
decreases due to the “ shielding” effect of the extra
energy levels.( It’s easier to remove an electron! )
Patterns in the table
The First Ionisation Energy increases as
you go
left to right along a period in the periodic table.
This is due to increasing number of electrons
within the same energy level and so a
corresponding increase in protons in nucleus,
resulting in a greater electrostatic attraction.
Therefore it is more difficult to remove an
electron!
Descending a group
 As
you go down a group in the periodic table the First
Ionisation Energy decreases due to the increasing
energy levels – the outer electrons are “shielded” from
the nucleus by the extra energy levels and so are easier
to remove!
 Removing more than one electron!
 Example
 Mg+ (g) —>Mg 2+ (g) + e- ΔH = + 1460 kJ
 This is the Second Ionisation Energy
Electronegativity
 This
is a measure of the Attraction for Bonding Electrons
 Atoms of different elements will have a different
attraction for bonding electrons i.e. a different
electronegativity!
 Electronegavitvity increases as you go from left to right
along a period.
 Electronegativity deceases as you go down a group.
Bonding
Covalent Bonding
When non metal elements share electron clouds,
to achieve the stable electron arrangement of the
noble gases.
 Both + nuclei are attracted to the shared cloud
and so the atoms are held by an “ electrostatic”
attraction. A molecule is a group of atoms held
together by a covalent bond.

Ionic Bonds
Bond
formed between + metal ions and – non
metal ions. Metal atoms lose electrons to achieve
a stable electron arrangement, non metal atoms
gain electrons to achieve a stable electron
arrangement.
A large network of + and – ions form the crystal
lattice. Ionic bonds are very strong.
Metallic bonding
 The
atoms of metal elements are held together by
metallic bonding. The electrons are free moving i.e.
“ delocalised”. The move from one nucleus to the next,
allowing metals to conduct electricity. There are the
same number of protons in the nucleus as surrounding
electrons – therefore metals are neutral.
 Metallic bonds are very strong, since a lot of energy is
required to overcome the attraction between the
delocalised electrons and the nuclei. The greater the
number of delocalised electrons the greater charge on
the ions and the stronger the metallic bond.
 Metallic strength increases from group 1 to group 2 etc.
Intra molecular Forces
These are
the forces of attraction within
molecules or compounds.
Covalent, Ionic and metallic bonding are
examples of Intra molecular forces.
Intermolecular Forces
These are
forces of attraction between
DIFFERENT molecules.
Example
Van der Waal’s forces
Hydrogen Bonding
Polar Covalent Bonds
 When
H bonds with another H they form a covalent
bond.
 This is because they both have the same number of
protons in the nucleus and so both have an equal
attraction I.e. same electronegativity, for the shared
electron cloud.
 When H and Cl bond it is polar.
 The Cl has more protons in nucleus and so will have a
greater “pull” on the shared electron cloud. Cl has a
greater electronegativity than H.
Dipole
When there
is a difference in electronegativity
between atoms in a molecule a dipole can be
created.
This is when each end of the molecule becomes
“ slightly charged”. One end will be “slightly
negative” = δ- . The other end will be “slightly
positive” = δ+..
Example – Hydrogen chloride – H δ+ Cl δ-
Van der Waal’s Forces
 This
is a very weak, temporary force of attraction
between different atoms or different molecules.
 Electrons are constantly moving round the nucleus.
Sometimes there may be an unequal distribution of the
electrons and this may end up in a temporary dipole
being created.
 The atoms with this temporary dipole attract other atoms
with an oppositely charged temporary dipole and so Van
der Waal’s force of attraction is created.
Permanent dipole – permanent
dipole attractions!
These are
additional intermolecular forces
between polar covalent molecules.
E.g. H2O , H Cl
These are stronger than the temporary Van der
Waal’s forces.
They result in an increase in the MP and BP of
polar covalent compounds.
Hydrogen Bonding
 Hydrogen
has a very low electronegativity.
 If it bonds to an element with a higher electronegativity
e.g. 0, N or F - the molecule will be highly polar.
 This strong intermolecular force is called a Hydrogen
bond.
 A hydrogen bond is the strongest intermolecular force
but it is not as strong as a covalent intra molecular bond.
 Hydrogen bonding increases the BP of the compounds.
Water
 The
density of water increases as the T falls to 4oC.
 But it starts to decrease between 4oC and 0oC.
 The water molecules start to move further apart and form
an open structure held together by Hydrogen bonds.
 As a result ice is less dense than water and so floats on
water.
 In pools the water freezes from the surface downwards,
insulating the water beneath it.
Polar Molecules
 A molecule
may have an unequal electronegativity within
atoms – creating a polar bond.
 However the polarity of the whole molecule is related to
symmetry.
 In polar molecules the bonds are not arranged in
symmetry, asymmetrical– e.g. HCl , NH3.,H2O
 Non polar molecules have bonds arranged in symmetry
– the permanent dipole – permanent dipole attraction
cancel each other out – e.g. CCl4, CO2
Structure and Properties of
Elements!
Element
Structure
M.P. Conductor?
Metals
Solids
High Yes
B, C, Si
Covalent network
solids
High No
P,S
Discrete covalent
molecules (solids)
Low No
H,N,O,F,Cl
Diatomic gases
Low No
Ne, Ar, He
Monatomic gases
Low No
Covalent Network Solids
Boron, Carbon
and Silicon form Covalent
Network Solids!
This is a large lattice of covalently bonded atoms.
They have High MP and BP.
Examples
Diamond – a large network of tetrahedrally
arranged carbon atoms. It has a rigid 3D
structure making it very hard!
Graphite
Carbon
atoms held in planar hexagonal rings ( 6
atoms joined together)
The rings slide over one another – powder
appearance.
Each carbon atom has 3 bonds – 1 electron is
delocalised allowing Graphite to conduct
electricity.
Other Covalent Networks
 Silicon
 Silicon
has a similar crystal structure as diamond.
Example - Silicon carbide( carborundum ), Silicon
dioxide – has a SiO4 tetrahedral structure
 Silicon is a semi conductor!
 Boron
 Forms interlocking B12 atoms.
 It can be as hard as diamond.
Covalent Molecular Solids Discrete
 Fullerines
 Buckminster
Fullerine – C5 and C6 rings form a dome
shaped solid ( football!) with a total of C60 atoms. Other
fullerines – C70 dome structures.S
 Fullerines can form addition reactions with halogens.
 Some are used as catalysts – (palladium containing
compounds.)S
 Phosphorous – P4 tetrahedral structure.
 Sulphur – S8 puckered rings.
Solubility
 Ionic
compounds dissolve in water because water is
polar!
 The H δ+ and O δ- pull apart the ionic lattice.
 The H δ+ part of the water molecule surrounds the
negative ion and the O δ- , of the water molecule,
surrounds the positive ion.
 Electrostatic attractions from between the ions and the
polar ends of the water molecule – this overcomes the
electrostatic attraction between the ions in the lattice.
Solubility patterns
Ionic or
polar substances dissolve in polar
solvents e.g. water.
Covalent, non polar substances dissolve in non
polar solvents e.g. ethanol, tetrachloromethane.
Ionisation in water
Polar compounds can ionise when dissolved in
water I.e. separate into separate ions – this
allows them to conduct electricity.