Chemistry UNIT 3
PERIODICITY
3.1 The periodic table
•3.1.1 Describe the arrangement of elements in the periodic
table in order of increasing atomic number
Metalloids
Metalloids are the elements found along the stairstep line that distinguishes metals from nonmetals.
This line is drawn from between Boron and
Aluminum to the border between Polonium and
Astatine.
The only exception to this is Aluminum, which is
classified under "Other Metals".
3.1.2 Distinguish between the term group
and period
• Group
• These are the numbers represented by Roman Numerals
at the top of the periodic table.
• This number tells us the number of valence electrons in
an atom.
• Valence electrons are important because they determine
the chemical reactivity of elements.
• Period: Represent the number of energy levels /Shells
3.1.1 Apply the relationship between the electron arrangement of
elements and their position in the periodic table up to Z = 20
The electron configuration of the first 20 elements
Hydrogen
Helium
Lithium
H
He
Li
1s
1s2
1s2 2s1
Beryllium
Boron
Carbon
Nitrogen
Be
B
C
N
1s2 2s2
1s2 2s2 2p1
1s2 2s2 2p2
1s2 2s2 2p3
Oxygen
Neon
Sodium
O
Ne
Na
1s2 2s2 2p4
1s2 2s2 2p6
1s2 2s2 2p6 3s1
Sulphur
S
1s2 2s2 2p6 3s2 3p4
Argon
Potassium
Ar
K
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p6 4s1
Calcium
Ca
1s2 2s2 2p6 3s2 3p6 4s2
3.1.4 Apply the relationship
between the number of
electrons in the highest
occupied energy level for an
element and its position in
the periodic table.
• Do the following questions: 1-3 p. 75
students text book
• Question from work book. Exercise 3.1
P. 97-99
• 3.2 Physical Properties
• Elements in the same group therefore have
different physical properties
• Physical properties include:
• Effective nuclear charge
• Atomic radius
• Ionic radius
• Ionization energy
• Electronegativity
• Melting point
The trend in the physical and chemical properties
are governed by the effective nuclear charge.
• The nuclear charge of the atom is given by the atomic
number. As you go across the periodic table the
atomic number increases by one, as one proton is
added to the nucleus.
• The outer electron (that determine many of the
physical and the chemical properties) do not
experience the full attraction of this charge (protons)
as they are shielded from the nucleus and repelled by
the inner electrons.
• The presence of the inner electrons reduces the attraction of the nucleus for
the outer electrons.
• The effective charge experienced by the outer electron is less than th efull
nuclear charge.
elements
Nuclear
Charge
Electron
arrangement
Na
11
Mg
12
2,8,1 2,8,2
Al
13
Si
14
2,8,3
2,8,4
• As you go across a period, one proton is added to the nucleus
and one electron is added to the outer electron shell. The
effective charge increases with the nuclear charge as there is no
change in the number of the inner electrons
• Chemical property
• Group I Metals (Alkali Metals)
• These metals react with water to form alkaline
solutions
• They have low melting points
• Low boiling points
• Soft
• Low density
• Very reactive
3.2 Physical properties: Trends in Melting and boiling
points
You will see that both the melting points and boiling points fall as you go
down the Group.
Group II Metals
The alkaline earth elements
are metallic elements found in the second group of the periodic
table. All alkaline earth elements have an oxidation number of +2,
making them very reactive. Because of their reactivity, the alkaline
metals are not found free in nature.
The Alkaline Earth Metals are:
Beryllium
Magnesium
Calcium
Strontium
Barium
Radium
• Extracted from oxides in the earth’s crust
• Less reactive than the alkali metals
Metalloids have properties of both metals and non-metals. Some of
the metalloids, such as silicon and germanium, are semiconductors.
This means that they can carry an electrical charge under special
conditions.
This property makes metalloids useful in computers and calculators
The Metalloids are:
Boron
Silicon
Germanium
Arsenic
Antimony
Tellurium
Polonium
• Properties of Metals
- They have high melting and boiling
points
• Properties of Non-metals
– Have low densities
– Brittle
Group VII THE HALOGENS
The halogens are five non-metallic elements found in group 7 of
the periodic table. The term "halogen" means "salt-former" and
compounds containing halogens are called "salts".
The Halogens are:
Fluorine
Chlorine
Bromine
Iodine
Astatine
• The halogens exist, at room temperature, in all
three states of matter:
• Solid- Iodine, Astatine
• Liquid- Bromine
• Gas- Fluorine, Chlorine
The Halogens are:
Fluorine
Chlorine
Astatine
Iodine
Bromine
Group VIII The Noble Gases
The six noble gases are found in group VIII (8) of the periodic table.
These elements are very stable and do not react with anything else.
They are sometimes referred to as inert gases
Helium
Neon
Argon
Krypton
Xenon
Radon
•These gases have low melting points and boiling points.
PERIODS
• Period Number indicates the number of
occupied shells.
• Elements in the same period share a gradual
change in their physical and chemical
properties.
Trends in the Periodic Table
• The trends in the periodic table are
- atomic radius
- first ionisation energy
- electronegativity
- melting and boiling points
- density.
3.2.2 Trends in Atomic Radius
• ATOMIC SIZE
The radius of an atom is found from the distance
between the nuclei in a molecule of two touching
atoms, and then halving that distance.
Trends in atomic radius in Periods 2 and 3
Trends in atomic radius down a group
• It is fairly obvious that the atoms get bigger as
you go down groups. The reason is equally
obvious - you are adding extra layers of
electrons.
Trends in atomic radius across periods
• You have to ignore the noble gas at the end of
each period. Because neon and argon don't
form bonds.
• Leaving the noble gases out, atoms get
smaller as you go across a period.
You can see that the atomic radius increases as you go down the
Group.
Explaining the increase in atomic radius
The radius of an atom is governed by
• the number of layers of electrons around the
nucleus
• the pull the outer electrons feel from the
nucleus.
• Compare lithium and sodium:
– Li
– Na
In each case, the outer electron feels a net pull of 1+ from the
nucleus.
The positive charge on the nucleus is cut down by the
negativeness of the inner electrons.
• The only factor which is going to affect the size of the
atom is therefore the number of layers of inner
electrons which have to be fitted in around the atom.
• The more layers of electrons you have, the more space
they will take up - electrons repel each other.
• This means that the atoms are bound to get bigger as
you go down the Group.
Trends in First Ionisation Energy
• First ionisation energy is the energy needed
to remove the most loosely held electron from
one mole of gaseous atoms.
• The state symbols - (g) - are essential. When you are
talking about ionisation energies, everything must be
present in the gaseous state.
• Ionisation energies are measured in kJ mol-1
(kilojoules per mole). They vary in size from 381 KJ
(which you would consider very low) up to 2370 KJ
(which is very high).
• All elements have a first ionisation energy.
Example: Helium (1st I.E. = 2370 kJ mol-1) Large
amount of energy that is needed to remove one of its
electrons. (Break a complete shell)
Patterns of first ionisation energies in the Periodic Table
The first 20 elements
• Factors affecting the size of ionisation energy
- Ionisation energy is a measure of the energy needed to pull a
particular electron away from the attraction of the nucleus. A
high value of ionisation energy shows a high attraction
between the electron and the nucleus.
The size of that attraction will be governed by:
• The charge on the nucleus.
- The more protons there are in the nucleus, the more
positively charged the nucleus is, and the more strongly
electrons are attracted to it.
• The distance of the electron from the nucleus.
- Attraction falls off very rapidly with distance. An electron
close to the nucleus will be much more strongly attracted than
one further away.
The patterns in periods 2 and 3
• Explaining the general trend for ionisation energies
across periods 2 and 3
• The general trend is for ionisation energies to increase across a
period.
• The major difference is the increasing number of protons in
the nucleus as you go from lithium to neon.
• This causes greater attraction between the nucleus and the
electrons and so increases the ionisation energies.
• The increasing nuclear charge also drags the outer electrons in
closer to the nucleus.
• That increases ionisation energies still more as you go across
the period.
Removal of successive electrons
• Second ionisation energy is defined by the equation:
• It is the energy needed
to remove a second electron
:
from each ion in 1 mole of gaseous 1+ ions to give
gaseous 2+ ions.
More ionisation energies
• You can then have as many successive ionisation
energies as there are electrons in the original atom.
• The first four ionisation energies of
aluminium, for example, are given by
1st I.E. = 577 kJ mol-1
2nd I.E. = 1820 kJ mol-1
3rd I.E. = 2740 kJ mol-1
4th I.E. = 11600 kJ mol-1
• In order to form an Al3+(g) ion from Al(g) you
would have to supply:
577 + 1820 + 2740 = 5137 kJ mol-1
Why do successive ionisation energies get
larger?
• Once you have removed the first electron you
are left with a positive ion. Trying to remove a
negative electron from a positive ion is going
to be more difficult than removing it from a
neutral atom.
• Removing an electron from a 2+ or 3+ (etc)
ion is going to be progressively more difficult.
Using ionisation energies to work out which group
an element is in
• A big jump between two successive ionisation
energies is typical of suddenly breaking in to
an inner level. (new shell)
• You can use this to work out which group of
the Periodic Table an element is in from its
successive ionisation energies.
• Magnesium (2,8,2) is in group 2 of the
Periodic Table and has successive ionisation
energies:
• Here the big jump occurs after the second
ionisation energy.
• It means that there are 2 electrons which are
relatively easy to remove (the 2 electrons in
last shell), while the third one is much more
difficult (because it comes from an inner level
closer to the nucleus and with less screening).
• Silicon (2,8,4) is in group 4 of the Periodic
Table and has successive ionisation energies:
:
• Decide which group an atom is in if it has
successive ionisation energies?
Plotting graphs of successive
ionisation energies
• Chlorine has the electronic structure 2,8,7
• This graph plots the first eight ionisation
energies of chlorine.
• The green labels show which electron is being
removed for each of the ionisation energies.
.
• The seventeenth ionisation energy of chlorine is nearly
400,000 kJ mol-1, and the vertical scale has to be squashed to
accommodate this.
.
• Lets do some problems
Three Types of Bonding
• IONIC BONDING
• COVALENT BONDING
• METALLIC BONDING
IONIC STRUCTURES
• This type of bonding occurs only between a metal
and a non metal.
Example:
Ionic bonding in sodium chloride
• Sodium (2,8,1) has 1 electron more than a stable
noble gas structure (2,8). If it gave away that electron
it would become more stable. cation
• Chlorine (2,8,7) has 1 electron short of a stable noble
gas structure (2,8,8). If it could gain an electron from
somewhere it too would become more stable. anion
• The answer is obvious. If a sodium atom gives
an electron to a chlorine atom, both become
more stable.
.
Some other examples of ionic bonding
magnesium oxide
Covalent Bonding
• Occurs between two non metals
• Involves the sharing of electrons.
Some very simple covalent molecules
Chlorine
For example, two chlorine atoms could both achieve
stable structures by sharing their single unpaired
electron as in the diagram.
The two chlorine atoms are said to be joined by a covalent bond. The
reason that the two chlorine atoms stick together is that the shared
pair of electrons is attracted to the nucleus of both chlorine atoms.
• Hydrogen chloride
• The hydrogen has a helium structure, and
the chlorine an argon structure.
For example:
Metallic Bonding
• Metals tend to have high melting points and
boiling points suggesting strong bonds
between the atoms. Even a metal like sodium
(melting point 97.8°C) melts at a considerably
higher temperature than the element (neon)
which precedes it in the Periodic Table.
Metallic Bonding
Metallic Bonding
•atoms in metals are packed very closely in an orderly arrangement
•each atom loses its valence electrons to become a positive ion
• The electrons can move freely within these molecular
orbitals, and so each electron becomes detached from
its parent atom. The electrons are said to be
delocalized. The metal is held together by the strong
forces of attraction between the positive nuclei and
the delocalized electrons.
• Metallic bonding is the electrostatic attraction
between the positively charged ions and
negatively charged electrons
Trends in Electronegativity
• Electronegativity is a measure of the tendency of an
atom to attract a bonding pair of electrons.
• It is usually measured on the Pauling scale, on which
the most electronegative element (fluorine) is given
an electronegativity of 4.0 and values range down to
caesium and francium which are the least
electronegative at 0.7.
What happens if two atoms of equal electronegativity
bond together?
Consider a bond between two atoms, A and B.
•If the atoms are equally electronegative, both have the
same tendency to attract the bonding pair of electrons, and
so it will be found on average half way between the two
atoms.
What happens if B is slightly more electronegative than A?
• B will attract the electron pair rather more than A
does.
What happens if B is a lot more electronegative than A?
In this case, the electron pair is dragged right over to B's end of the
bond.
A has lost control of its electron, and B has complete control over
both electrons.
Patterns of electronegativity in the Periodic
Table
• The most electronegative element is fluorine. If you remember
that fact, everything becomes easy, because electronegativity
must always increase towards fluorine in the Periodic Table.
Trends in electronegativity across a period
• As you go across a period the electronegativity
increases. The chart shows electronegativities from
sodium to chlorine - you have to ignore argon. It
doesn't have an electronegativity, because it doesn't
form bonds.
Trends in electronegativity down a group
• As you go down a group, electronegativity
decreases. (If it increases up to fluorine, it
must decrease as you go down.) The chart
shows the patterns of electronegativity in
Groups 1 and 7.
Trends in Melting and Boiling Points
Trends in Melting Point (METALS)
.
Trends in Melting Point and Boiling Point (NON METALS)
Melting and boiling points across a period
The chart shows how the melting and boiling points of
the elements change as you go across the period. The
figures are plotted in Kelvin rather than °C to avoid
having negative values.
Table of physical data
Symbol
Melting
point
(K)
Boiling
point
(K)
11
Na
371
1156
magnesium
12
Mg
922
1380
aluminium
13
Al
933
2740
silicon
14
Si
1683
2628
phosphorus
15
P
317
553
sulphur
16
S
392
718
chlorine
17
Cl
172
238
argon
18
Ar
84
87
Element
Proton
number
sodium
back to top
Explanation of trend across a Period
1. The Period always begins with metals
Example:
Sodium, magnesium and aluminium
• Sodium, magnesium and aluminium are all metals. They have
metallic bonding, in which positive metal ions are attracted
to delocalized electrons. Going from sodium to aluminium:
• the charge on the metal ions increases from +1 to +3 (with
magnesium at +2) ...
•
the number of delocalized electrons increases ...
• so the strength of the metallic bonding increases and
the melting points and boiling points increase.
2. The Period then comes across the
Metalloids.
Example Silicon
• Silicon is a metalloid (an element with
some of the properties of metals and
some of the properties of non-metals).
• Silicon has giant covalent bonding. It
has a giant lattice structure, in which
each silicon atom is covalently-bonded to
four other silicon atoms in a tetrahedral
arrangement. A giant covalent structure
•
• The structure is held together by strong
covalent bonds in all three dimensions.
Now the Period reaches the non-metals.
Example
• Phosphorus, sulphur, chlorine and argon
• These are all non-metals, and they exist as small,
separate molecules. Phosphorus, sulphur and
chlorine exist as simple molecules, with strong
covalent bonds between their atoms.
• Their melting and boiling points are very low because
little energy is needed to overcome their bonds.
• Sulphur has a higher melting point and boiling point than the other three
because:
phosphorus exists as P4 molecules ...
sulphur exists as S8 molecules ...
chlorine exists as Cl2 molecules ...
argon exists individual Ar atoms ...
• the strength of the forces decreases as the size
of the molecule decreases ...
• so the melting points and boiling points
decrease in the order
S8 > P4 > Cl2 > Ar
IONIC RADII
IONIC RADIUS
• Ions aren't the same size as the atoms they
come from. Compare the sizes of sodium and
chloride ions with the sizes of sodium and
chlorine atoms.
Positive ions
• Positive ions are smaller than the atoms they come from.
Sodium is 2,8,1; Na+ is 2,8.
• You've lost a whole layer of electrons, and the remaining 10
electrons are being pulled in by the full force of 11 protons.
Negative ions
Negative ions are bigger than the atoms they come from.
Chlorine is 2,8,7; Cl- is 2,8,8.
Although the electrons are still all in the 3-level, the extra
repulsion produced by the incoming electron causes the atom to
expand.
There are still only 17 protons, but they are now having to hold
18 electrons.
Trend
• Both cations and anions increase in
size down a group.
PHYSICAL AND CHEMICAL
PROPERTIES OF THE
GROUP I ELEMENTS
•
•
•
•
•
•
•
The Alkali Metals are:
Lithium
Sodium
Potassium
Rubidium
Cesium
Francium
ATOMIC AND PHYSICAL PROPERTIES OF THE
GROUP 1 ELEMENTS
Trends in Atomic Radius
You can see that the atomic radius increases as you go down the
Group.
Trends in First Ionisation Energy
Notice that first ionisation energy falls as you go down the group.
Trends in Electronegativity
Trends in Melting and Boiling Points
Chemical Properties of Group I
• REACTIONS OF THE GROUP 1 ELEMENTS WITH
WATER
All of these metals react vigorously or even explosively with
cold water. In each case, a solution of the metal hydroxide is
produced together with hydrogen gas.
This equation applies to any of these metals and water -
Details for the individual metals
• Lithium
• Lithium's density is only about half that of water so it
floats on the surface, gently fizzing and giving off
hydrogen. It gradually reacts and disappears, forming
a colourless solution of lithium hydroxide.
• Sodium
• Sodium also floats on the surface, but enough heat is
given off to melt the sodium (sodium has a lower
melting point than lithium and the reaction produces
heat faster) and it melts almost at once to form a
small silvery ball that dashes around the surface.
• A white trail of sodium hydroxide is seen in the water
under the sodium, but this soon dissolves to give a
colourless solution of sodium hydroxide.
• Potassium
• Potassium behaves rather
like sodium except that the
reaction is faster and enough
heat is given off to set light
to the hydrogen.
• This time the normal
hydrogen flame is
contaminated by potassium
compounds and so is
coloured lilac (a faintly
bluish pink).
• Rubidium
• Rubidium is denser than water and so sinks. It
reacts violently and immediately, with
everything spitting out of the container again.
Rubidium hydroxide solution and hydrogen
are formed.
• Caesium
• Caesium explodes on contact with water, quite
possibly shattering the container. Caesium
hydroxide and hydrogen are formed
Reactions with Oxygen
• Group I elements combine with oxygen to
form metal oxides which are basic.
GENERAL EQUATION
Metal + Oxygen
Metal oxide
However, each of the metals has its own
behaviour as you move down the group.
Lithium
• Lithium forms the normal oxide
Li+ O2Sodium
Forms an oxide and the peroxide
Sodium Oxide
When there is a lot of oxygen available then the following
reaction occurs.
Reactions with the Halogens
• Group I react with the halogens to for
metal halides.
• General Equation
Metal + Halogen
Metal Halide
The halides formed are usually a white solid
• Metal + Fluorine
Metal Fluoride
• Metal + Bromine
Metal Bromide
• Metal + Iodine
Metal Iodide
Reaction with dilute acids
• Group I metals react violently with dilute acids to
produce hydrogen gas and a salt.
For example, sodium will react with dilute hydrochloric
acid to give sodium chloride solution and hydrogen gas.
Reaction with Hydrogen
When heated with hydrogen they react to form hydrides
(metal hydride)
The oxidation number of hydrogen is -1
Reactions with Hydrogen
• Halogens react with Hydrogen to produce Hydrogen
Halides
Examples:
Physical properties
• The hydrogen halides are colourless gases at room
temperature, producing steamy fumes in moist air.
Hydrogen fluoride has an abnormally high boiling
point for the size of the molecule (293 K or 20°C),
and could condense to a liquid on a cool day.
• The hydrogen Halides react with water and form
solutions of strong acids
• Hydrogen chloride gas is very soluble in water, reacting with it
to produce hydrochloric acid.
• The familiar steamy fumes of hydrogen chloride in moist air
are caused by the hydrogen chloride reacting with water
vapour in the air to produce a fog of concentrated hydrochloric
acid.
Hydrobromic acid and hydriodic acid as strong acids
• Hydrogen bromide and hydrogen iodide dissolve in
(and react with) water in exactly the same way as
hydrogen chloride does. Hydrogen bromide reacts to
give hydrobromic acid; hydrogen iodide gives
hydriodic acid. Both of these are also strong acids.
Hydrofluoric acid as an exception
• By contrast, although hydrogen fluoride dissolves
freely in water, hydrofluoric acid is only a weak acid
Reactions with water
• Fluorine will react with water to produce
O2 gas
• The other halogens will behave differently
in water forming an acid
Reactions with Metals
• The halogens combine with metals to give a
salt containing the halide ion:
Example:
The salt formed is usually white and when
dissolved in water gives a colourless solution.
•The only metal halides that are insoluble in water are
usually halides of lead and silver
TESTING FOR HALIDE IONS
• Since the salt formed are colourless when aqueous
there must be a way to test for the presence of Cl- Brand I• Using silver nitrate solution
By adding a few drops of this solution and looking for
the formation of a precipitate.
• Silver nitrate solution is added to give:
observation
ion present
:
F-
no precipitate
Cl-
white precipitate
Br-
very pale cream precipitate
I-
very pale yellow precipitate
The chloride, bromide and iodide precipitates are shown in the photograph:
.
The chemistry of the test
• The precipitates are the insoluble silver
halides - silver chloride, silver bromide or
silver iodide.
Silver fluoride is soluble, and so you don't get a precipitate
The chemistry of the test
• The precipitates are the insoluble silver
halides - silver chloride, silver bromide or
silver iodide.
Silver fluoride is soluble, and so you don't get a precipitate
Reactions of halogens with one another
• Since the reactivity of these elements
decreases going down the group then a
halogen positioned higher in the group will
react with a lower halogen halide.
Bromide added to Chlorine
• Bromide ion is added to chlorine water.
We see that the reaction has produced
bromine.
2Br-(aq ) + Cl2(aq ) --> 2Cl-(aq ) + Br2(aq )
Bromide ion is added to chlorine water
and is shaken
The reaction has produced bromine
Iodide added to Chlorine
• Iodide ion is added to chlorine water. we
see that the reaction has produced iodine.
2I-(aq ) + Cl2(aq ) --> 2Cl-(aq ) + I2(aq )
Iodide ion is added to chlorine water
Chloride added to Bromine
• Chloride ion is added to bromine water.
• we see that no reaction occurrs.
2Cl-(aq ) + Br2(aq ) --> No reaction
Chloride ion is added
to bromine water
... and is shaken.
No reaction took
place.
Iodide added to Bromine
• Iodide ion is added to bromine. we see that the
reaction has produced iodine.
• 2I-(aq ) + Br2(aq ) --> 2Br-(aq ) + I2(aq )
Iodide ion is added to
bromine water
... and is shaken.
The reaction has
produced iodine.
Chloride added to Iodine
2Cl-(aq ) + I2(aq ) --> No chemical reaction took place
Chloride ion is
added to iodine
water
... and is shaken.
No chemical
reaction took
place.
Bromide added to Iodine
2Br-(aq ) + I2 --> No chemical reaction took place
Trends Across a Period.
• Reducing strength decreases (ability to donate
electrons)
• Oxidizing strength increases (ability to accept
electrons)
• as atoms become less able to release electrons they
have a greater tendency to acquire additional
electrons.
• Elements on the left of the table form positive ions
and those to the right form negative ions.
Properties of the Period 3
Elements
Atomic Properties
• Electronic structures
• In Period 3 of the Periodic Table, the 3s and 3p
orbitals are filling with electrons. The electronic
structures for the eight elements are:
Na
[Ne] 3s1
Mg
[Ne] 3s2
Al
[Ne] 3s2 3px1
Si
[Ne] 3s2 3px1 3py1
P
[Ne] 3s2 3px1 3py1 3pz1
S
[Ne] 3s2 3px2 3py1 3pz1
Cl
[Ne] 3s2 3px2 3py2 3pz1
Ar
[Ne] 3s2 3px2 3py2 3pz2
First ionisation energy
The pattern of first ionisation energies across Period 3
Explaining the pattern
• First ionisation energy is governed by:
• the charge on the nucleus;
• the distance of the outer electron from the nucleus;
• the amount of screening by inner electrons;
• whether the electron is alone in an orbital or one of a
pair.
The upward trend
• In the whole of period 3, the outer electrons are in 3level orbitals. These are all the same sort of distances
from the nucleus, and are screened by the same
electrons in the first and second levels.
• The major difference is the increasing number of
protons in the nucleus as you go from sodium across
to argon. This causes greater attraction between the
nucleus and the electrons and so increases the
ionisation energies.
The fall at aluminium
• You might expect the aluminium value to be
more than the magnesium value because of the
extra proton.
• But the fact that aluminium's outer electron is in
a 3p orbital rather than a 3s.
• The 3p electron is slightly more distant from the
nucleus than the 3s, and partially screened by
the 3s electrons as well as the inner electrons.
Both of these factors offset the effect of the extra
proton.
The fall at sulphur
• The screening is identical in phosphorus and sulphur
(from the inner electrons and, to some extent, from
the 3s electrons), and the electron is being removed
from an identical orbital.
• The difference is that in the sulphur case the electron
being removed is one of the 3p2 pair. The repulsion
between the two electrons in the same orbital means
that the electron is easier to remove than it would
otherwise be.
Atomic radius
• The trend
The diagram shows how the atomic radius changes as you
go across Period 3.
• The figures used to construct this diagram are based
on:
• metallic radii for Na, Mg and Al;
• covalent radii for Si, P, S and Cl;
• the van der Waals radius for Ar because it doesn't form
any strong bonds.
• The general trend is atomic size decreases across the
period (ignoring the noble gases)
Electronegativity
The trend
The trend across Period 3 looks like this:
Physical Properties
Structures of the elements
• The structures of the elements change as you go across
the period. The first three are metallic, silicon is giant
covalent, and the rest are simple molecules.
Electrical conductivity
• Sodium, magnesium and aluminium are all
good conductors of electricity.
• Silicon is a semiconductor.
• None of the rest conduct electricity.
Melting and boiling points
• The chart shows how the melting and
boiling points of the elements change as
you go across the period.
.
The sizes of the melting and boiling points are governed entirely
by the sizes of the molecules. Remember the structures of the
molecules:
CHEMICAL REACTIONS OF
THE PERIOD 3
ELEMENTS
Reactions with water
Sodium
• Sodium has a very exothermic reaction with cold water
producing hydrogen and a colourless solution of sodium
hydroxide.
.
Magnesium
• Magnesium has a very slight reaction with cold water,
but burns in steam.
• A very clean coil of magnesium dropped into cold
water eventually gets covered in small bubbles of
hydrogen which. float it to the surface. Magnesium
hydroxide is formed as a very thin layer on the
magnesium and this tends to stop the reaction.
•Magnesium burns in steam with its typical white flame to produce
white magnesium oxide and hydrogen.
Aluminium
• Aluminium powder heated in steam produces hydrogen and
aluminium oxide. The reaction is relatively slow because of
the existing strong aluminium oxide layer on the metal, and the
build-up of even more oxide during the reaction.
Silicon
•The common shiny grey lumps of silicon with a rather metal-like
appearance are fairly un-reactive. Most sources suggest that this form of
silicon will react with steam at red heat to produce silicon dioxide and
hydrogen.
Phosphorus and sulfur
These have no reaction with water.
Chlorine
• Chlorine dissolves in water to some extent to give a
green solution. A reversible reaction takes place to
produce a mixture of hydrochloric acid and chloric(I)
acid (hypochlorous acid).
Argon
There is no reaction between argon and water.
Reactions with oxygen
Sodium
Sodium burns in oxygen with an orange flame to produce a white
solid mixture of sodium oxide and sodium peroxide.
For the simple oxide:
For the peroxide:
Magnesium
Magnesium burns in oxygen with an intense white flame to give
white solid magnesium oxide.
• Aluminium
• Aluminium will burn in oxygen if it is powdered,
otherwise the strong oxide layer on the aluminium
tends to inhibit the reaction. If you sprinkle
aluminium powder into a Bunsen flame, you get
white sparkles. White aluminium oxide is formed.
Silicon
Silicon will burn in oxygen if heated strongly enough. Silicon dioxide
is produced.
• Phosphorus
• White phosphorus catches fire spontaneously in air, burning
with a white flame and producing clouds of white smoke - a
mixture of phosphorus(III) oxide and phosphorus(V) oxide.
• The proportions of these depend on the amount of oxygen
available. In an excess of oxygen, the product will be almost
entirely phosphorus(V) oxide.
For the phosphorus(III) oxide:
For the phosphorus(V) oxide:
• Sulphur
• Sulphur burns in air or oxygen on gentle heating with
a pale blue flame. It produces colourless sulphur
dioxide gas.
With an excess of oxygen produces sulfur trioxide
2S
+
3O2
2SO3
Chlorine and argon
Despite having several oxides, chlorine will not react directly
with oxygen. Argon doesn't react either.
PHYSICAL PROPERTIES
OF THE
PERIOD 3 OXIDES
A quick summary of the trends
The oxides
• The oxides we'll be looking at are:
Na2O MgO Al2O3 SiO2 P4O10 SO3
P4O6
SO2
Cl2O7
Cl2O
Oxides of the Period 3
Na2O
MgO
Al2O3
SiO2
P4O10
and
P4O6
ionic
ionic
ionic
SO2
and
SO3
Cl2O
and
Cl2O7
Covalent Covalent Covalent Covalent
Electrical conductivity
• None of these oxides has any free or mobile
electrons. That means that none of them will conduct
electricity when they are solid.
• The ionic oxides can, however, undergo electrolysis
when they are molten. They can conduct electricity
because of the movement of the ions towards the
electrodes and the discharge of the ions when they get
there.
The trend in acid-base behaviour
• The trend is from strongly basic oxides on the lefthand side to strongly acidic ones on the right, via an
amphoteric oxide (aluminium oxide) in the middle.
An amphoteric oxide is one which shows both acidic
and basic properties.
Chemistry of the individual oxides
Sodium oxide
• Sodium oxide is a simple strongly basic oxide. It is
basic because it contains the oxide ion, O2-, which is
a very strong base with a high tendency to combine
with hydrogen ions.
Reaction with water
• Sodium oxide reacts exothermically with cold water
to produce sodium hydroxide solution. Depending on
its concentration, this will have a pH around 14.
Reaction with acids
• As a strong base, sodium oxide also reacts with acids.
For example, it would react with dilute hydrochloric
acid to produce sodium chloride solution.
Magnesium oxide
Magnesium oxide is a simple basic oxide, because it
also contains oxide ions. However, it isn't as strongly
basic as sodium oxide because the oxide ions aren't
so free.
• Reaction with water
• If you shake some white magnesium oxide powder
with water, nothing seems to happen - it doesn't look
as if it reacts. However, if you test the pH of the
liquid, you find that it is somewhere around pH 9 showing that it is slightly alkaline
• Some magnesium hydroxide is formed in the
reaction, but this is almost insoluble - and so not
many hydroxide ions actually get into solution.
• Reaction with acids
• Magnesium oxide reacts with acids as you would
expect any simple metal oxide to react. For example,
it reacts with warm dilute hydrochloric acid to give
magnesium chloride solution.
Aluminium oxide
• Aluminium oxide is amphoteric. It has reactions as
both a base and an acid.
• Reaction with water
• Aluminium oxide doesn't react in a simple way with
water in the sense that sodium oxide and magnesium
oxide do, and doesn't dissolve in it. Although it still
contains oxide ions, they are held too strongly in the
solid lattice to react with the water.
• Reaction with acids
• Aluminium oxide contains oxide ions and so reacts
with acids in the same way as sodium or magnesium
oxides. That means, for example, that aluminium
oxide will react with hot dilute hydrochloric acid to
give aluminium chloride solution.
• Reaction with bases
• Aluminium oxide has also got an acidic side to its nature, and
it shows this by reacting with bases such as sodium hydroxide
solution.
• Various aluminates are formed - compounds where the
aluminium is found in the negative ion. This is possible
because aluminium has the ability to form covalent bonds with
oxygen.
.
• With hot, concentrated sodium hydroxide solution, aluminium
oxide reacts to give a colourless solution of sodium
tetrahydroxoaluminate.
Silicon dioxide (silicon(IV) oxide)
• Silicon dioxide has no basic properties - it doesn't
contain oxide ions and it doesn't react with acids.
Instead, it is very weakly acidic, reacting with strong
bases.
• Reaction with water
• Silicon dioxide doesn't react with water, because of
the difficulty of breaking up the giant covalent
structure.
• Reaction with bases
• Silicon dioxide reacts with sodium hydroxide
solution, only if it is hot and concentrated. A
colourless solution of sodium silicate is
formed.
The phosphorus oxides
• Phosphorus(III) oxide (P4O6)
• Phosphorus(III) oxide reacts with cold water to give
a solution of the weak acid, H3PO3 - known
variously as phosphorous acid, orthophosphorous
acid or phosphonic acid.
• P4O6
+
6 H2 O
4H3PO3
Phosphorus(V) oxide
• Phosphorus(V) oxide reacts violently with
water to give a solution containing a mixture
of acids, the nature of which depends on the
conditions. We. usually just consider one of
these, phosphoric(V) acid, H3PO4 - also known
just as phosphoric acid.
The sulphur oxides
• We are going to be looking at sulphur dioxide, SO2,
and sulphur trioxide, SO3.
• Sulphur dioxide
• Sulphur dioxide is fairly soluble in water, reacting
with it to give a solution of sulphurous acid
.
(sulphuric(IV) acid),
H2SO3. This only exists in
solution, and any attempt to isolate it just causes
sulphur dioxide to be given off again.
• Sulphur trioxide
• Sulphur trioxide reacts violently with water to
produce a fog of concentrated sulphuric acid
droplets.
The chlorine oxide
• Chlorine(VII) oxide
• Chlorine(VII) oxide is the highest oxide of chlorine the chlorine is in its maximum oxidation state of +7.
It continues the trend of the highest oxides of the
Period 3 elements. towards being stronger acids.
• Chlorine (VII) oxide reacts with water to give the
very strong acid, chloric (VII) acid - also known as
perchloric acid. The pH of typical solutions will, like
sulphuric acid, be around 0.
PROPERTIES OF THE
PERIOD 3 CHLORIDES
• The chlorides
• The chlorides we'll be looking at are:
NaCl MgCl2 AlCl3 SiCl4 PCl5 S2Cl2
PCl3
Reactions with water
• As an approximation, the simple ionic
chlorides (sodium and magnesium chloride)
just dissolve in water.
• The other chlorides all react with water in a
variety of ways.
• The reaction with water is known as
hydrolysis.
Sodium chloride, NaCl
• Sodium chloride is a simple ionic
compound consisting of a giant array of
sodium and chloride ions.
• A small representative bit of a sodium
chloride lattice looks like this:
This is normally drawn in an exploded form as:
• The strong attractions between the positive and
negative ions need a lot of heat energy to break, and
so sodium chloride has high melting and boiling
points.
• It doesn't conduct electricity in the solid state because
it hasn't any mobile electrons and the ions aren't free
to move. However, when it melts it undergoes
electrolysis.
• Sodium chloride simply dissolves in water to give a
neutral solution.
Magnesium chloride, MgCl2
• Magnesium chloride is also ionic, but with a
more complicated arrangement of the ions to
allow for having twice as many chloride ions
as magnesium ions
• Again, lots of heat energy is needed to
overcome the attractions between the ions, and
so the melting and boiling points are again
high.
• Solid magnesium chloride is a non-conductor of
electricity because the ions aren't free to move.
However, it undergoes electrolysis when the ions
become free on melting.
• Magnesium chloride dissolves in water to give a
faintly acidic solution (pH = approximately 6).
Aluminium chloride, AlCl3
• At room temperature, solid aluminium chloride has an
ionic lattice with a lot of covalent character.
• Solid aluminium chloride doesn't conduct electricity at
room temperature because the ions aren't free to move.
• The reaction of aluminium chloride with water is
dramatic. If you drop water onto solid aluminium
chloride, you get a violent reaction producing clouds of
steamy fumes of hydrogen chloride gas.
• The aluminium chloride reacts with the water rather
than just dissolving in it.
Silicon tetrachloride, SiCl4
• Silicon tetrachloride is a colourless liquid at room
temperature which fumes in moist air.
• It doesn't conduct electricity because of the lack of
ions or mobile electrons.
• It fumes in moist air because it reacts with water in
the air to produce hydrogen chloride
• It is a violent reaction to produce silicon dioxide and
fumes of hydrogen chloride.
The phosphorus chlorides
• There are two phosphorus chlorides phosphorus(III) chloride, PCl3, and
phosphorus(V) chloride, PCl5.
Phosphorus(III) chloride (phosphorus
trichloride), PCl3
• This is another simple covalent chloride again a fuming liquid at room temperature.
• It doesn't conduct electricity because of the
lack of ions or mobile electrons.
• Phosphorus(III) chloride reacts violently with
water. You get phosphorous acid, H3PO3, and
fumes of hydrogen chloride (or a solution
containing hydrochloric acid if lots of water is
used).
Disulphur dichloride, S2Cl2
• Disulphur dichloride is a simple covalent liquid orange and smelly!
• Disulphur dichloride reacts slowly with water to
produce a complex mixture of things including
hydrochloric acid, sulphur, hydrogen sulphide
and various sulphur-containing acids and anions
(negative ions).
• There is no way that you can write a single
equation for this - and one would never be
expected in an exam.
PHYSICAL AND CHEMICAL
PROPERTIES OF THE
GROUP VII ELEMENTS
The Halogens are:
Fluorine
Chlorine
Astatine
Iodine
Bromine
• The name Halogen is derived from the Greek
word ‘halos’ meaning salt former
• All halogens have the outer shell configuration
s2p5
• The halogens all occur as diatomic molecules
by forming a single covalent bond to get a
stable electron configuration.
Trends in Atomic Radius
Explaining the increase in atomic radius
• The radius of an atom is governed by
- the number of layers of electrons around
the nucleus
- the pull the outer electrons feel from the
nucleus.
• Compare fluorine and chlorine:
F 2,7
Cl 2,8,7
• In each case, the outer electrons feel a net
pull from the nucleus.
Trends in Electronegativity
Explaining the decrease in electronegativity
• This is easily shown using simple dotsand-crosses diagrams for hydrogen
fluoride and hydrogen chloride.
.
Trends in Melting Point and Boiling Point
• The bonding pair of electrons between the hydrogen
and the halogen feels the same net pull of 7+ from
both the fluorine and the chlorine.
• (This is exactly the same sort of argument as you
have seen in the atomic radius.)
• However, in the chlorine case, the nucleus is further
away from that bonding pair. That means that it won't
be as strongly attracted as in the fluorine case.
• The larger pull from the closer fluorine nucleus is
why fluorine is more electronegative than chlorine is.
• Explaining the trends in melting point and boiling point
• All of the halogens exist as diatomic molecules - F2, Cl2, and
so on.
• The intermolecular attractions between one molecule and its
neighbours are van der Waals dispersion
• As the molecules get bigger there are obviously more electrons
which can move around and set up the temporary dipoles
which create these attractions.
• The stronger intermolecular attractions as the molecules get
bigger means that you have to supply more heat energy to turn
them into either a liquid or a gas - and so their melting and
boiling points rise.