Download Topic 3 - periodicity

TOPIC 3 - PERIODICITY
IB CHEMISTRY SL
3.1.1 DESCRIBE THE ARRANGEMENT OF ELEMENTS IN THE
PERIODIC TABLE IN ORDER OF INCREASING ATOMIC NUMBER
The elements in a
periodic table are
placed in order of
increasing atomic
number (Z), which we
know is a
fundamental property
of the element number of protons in
the nucleus of its
atom. To read the
periodic table, simply
start from the top left
and read across. You
will find that the
atomic number
increases.
3.1.2 DISTINGUISH BETWEEN THE TERMS GROUP AND
PERIOD
Groups are the
vertical columns
of elements.
Periods are the
horizontal rows of
elements.
3.1.3 APPLY THE RELATIONSHIP BETWEEN THE ELECTRON ARRANGEMENT OF
ELEMENTS AND THEIR POSITION IN THE PERIODIC TABLE UP TO Z=20.
• Groups state how many valence
electrons there are
Periods state how many electron
shells there are.
• Example,
-Carbon has an electron arrangement of 2, 6
It is in the 2nd period and 6th group.
-Calcium has an electron arrangement of 2, 8, 8, 2
Valence electron is an electron in
the outer shell of an atom
It is in the 4th period and 2nd group.
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
• An atom with one valence electron
will be in which group?
• An atom with the electron
arrangement 2,8,7 is located where
on the periodic table?
• An atom with a full outer shell will
be in which group?
• An atom with the electron
arrangement 2,8,8,1 is located
where on the periodic table?
• http://www.youtube.com/watch?v=
0RRVV4Diomg&list=PL8dPuuaLjXtP
HzzYuWy6fYEaX9mQQ8oGr&index=
4
3.2 PHYSICAL PROPERTIES
•
3.2.1 DEFINE THE TERMS FIRST IONIZATION ENERGY
AND ELECTRONEGATIVITY
• First Ionization Energy is the
minimum energy required to
remove an electron from a neutral
gaseous atom in its ground state.
• Electronegativity defined as the
relative attraction that an atom
has for the shared pair of
electrons in a covalent bond.
3.2.2 DESCRIBE AND EXPLAIN THE TRENDS IN ATOMIC RADII, IONIC RADII, FIRST IONIZATION ENERGIES,
ELECTRONEGATIVITIES AND MELTING POINTS FOR ALKALI METALS AND THE HALOGENS
• Atomic Radii
Since the outer electrons are
difficult to locate, in practice, the
atomic radius is measured as half
the distance between two bonded
atoms. For this reason noble gases
are given no value as they do not
bond with other atoms.
On descending a group, the atomic
radius increase. This is because the
outer electrons are getting further
from the nucleus. This applies for
both alkali metals and halogens.
• Ionic Radii
It is defined by the distance between nucleus
and outer most electrons of a positive metal
cation or negative non-metal anion.
Both metal cation and non-metal anion
increase in size on descending the group
• Metal cations tend to be smaller than their
atom as they have lost their outer most
electrons
• Non-metal anions tend to be larger than
their atom as they have gained electrons into
their outer energy levels
• 1st Ionization Energy
It is defined by the minimum energy
required to remove an electron
from a neutral gaseous atom in its
ground state.
Ionization energy decreases on
descending a group
• The outer most electrons are easier
to remove as they are a greater
distance from the nucleus and are
shielded from the positive nucleus
charge by the core electrons.
• Electronegativity
Electronegativity is defined
as the relative attraction
that an atom has for the
shared pair of electrons in
a covalent bond.
On descending a group,
electronegativity
decreases. While the
atomic radius and the
nuclear charge increases,
the level of shielding
increases and the effective
nuclear charge decrease.
• The most electronegative
elements are in the top
right of the periodic table.
(F, O and N). Which are
involved in hydrogen
bonding.
MELTING POINT TRENDS
• Page 57
• What is the melting point trend for
group 1 and group 7?
• Why is this the trend?
MELTING POINT TREND FOR ALKALI METALS AND HALOGENS
3.2.3 DESCRIBE AND EXPLAIN TRENDS IN ATOMIC RADII, IONIC RADII, FIRST
IONIZATION ENERGIES AND ELECTRONEGATIVITIES FOR ELEMENTS ACROSS
PERIOD 3
• First Ionization Energies
• Atomic radii
Across a period, the atomic radius
decreases. This is because of the
increase in effective nuclear charge
and no increase shielding, therefore
the electrons get pulled closer
towards the nucleus.
Ionization energy generally
increases across a period. This is
due to the increase effective
nuclear charge as the number of
protons increases with no
increase in shielding.
Ionic radii
Electronegativity
Across a period, the metal cation
and the non-metal anion get smaller.
This is because of an increase in
protons (Effective Nuclear Charge)
and all the ions have the same
electronic configuration.
Across a period, electronegativity
increases. This is the result of an
increased number of protons and
thus an increased effective
nuclear charge with no greater
level of shielding.
Atomic radii
Across a period, the atomic radius decreases. This is because of the increase in effective
nuclear charge and no increase shielding, therefore the electrons get pulled closer towards
the nucleus.
Ionic radii
Across a period, the cations and the anions get smaller. This is because of an
increase in protons (Effective Nuclear Charge) and all the ions have the same
electronic configuration.
IONIZATION ENERGY GENERALLY INCREASES ACROSS A PERIOD. THIS IS DUE TO THE
INCREASE EFFECTIVE NUCLEAR CHARGE AS THE NUMBER OF PROTONS INCREASES WITH NO
INCREASE IN SHIELDING.
ELECTRONEGATIVITY
ACROSS A PERIOD, ELECTRONEGATIVITY INCREASES. THIS IS THE RESULT OF AN
INCREASED NUMBER OF PROTONS AND THUS AN INCREASED EFFECTIVE
NUCLEAR CHARGE WITH NO GREATER LEVEL OF SHIELDING.
3.2.4 COMPARE THE RELATIVE ELECTRONEGATIVITY VALUES
OF TWO OR MORE ELEMENTS BASED ON THEIR POSITIONS
IN THE PERIODIC TABLE.
• By looking at the difference in electro-negativity
values
We can tell whether a molecule has
covalent or ionic bonds
3.3 CHEMICAL PROPERTIES
3.3.1 DISCUSS THE SIMILARITIES AND DIFFERENCES IN THE
CHEMICAL PROPERTIES OF ELEMENTS IN THE SAME GROUP
Group 0: Noble Gases
-Colorless
-Monatomic
-Unreactive
-Their lack of reactivity can be
explained by the inability of their
atoms to lose or gain electrons.
Group 1: Alkali Metals
• Good conductors of electricity
• The alkali metals react with water to
produce hydrogen and the metal
hydroxide.
• Low Density
• Shiny surface
• Reactive metals
• Forms ionic compounds with non-metals
• They form single charged ions M+. Their
low ionization energies give an indication
of the ease with which the outer electron
is lost.
• Reaction with Water
• The reaction becomes more
vigorous as we descend the group.
• http://periodicvideos.com/
Group 7: Halogens
-Diatomic molecules
-Colored
-Gradual change from gases (Fluorine
and Chlorine) to Liquid (Bromine) to
Solids (Iodine and Astatine)
-Reactive non-metals
-Reactivity decrease down the group
-Form ionic compounds with metals or
covalent compounds with other nonmetals
• http://periodicvideos.com/
HALOGEN REACTION WITH GROUP 1 METALS
• Halogens react with Group 1 metals
to form ionic halides. The halogen
atoms gains one electron from the
Group 1 elements to form a halide
ion. The electrostatic force of
attraction between the alkali metal
ions and halides bonds the ions
together.
Once the transfer is complete, the ions are
pulled together by the mutual attraction of
their opposite charges.
This would be the equation
DISPLACEMENT REACTIONS WITH HALOGENS
• The relative reactivity of the
elements can also be seen by
placing them in direct competition
for an extra electron.
A chlorine nucleus has a stronger attraction
for an electron than a bromine nucleus
because of its smaller atomic radius and so
takes the electron from the bromide ion.
The chlorine has gained an electron and so
forms the chloride ion. The bromide ion
loses an electron to form bromine.
HALOGEN REACTION WITH SILVER
The halogens form insoluble salts with
silver. Adding a solution containing the
halide to a solution containing silver
ions produces a precipitate which is
useful in identifying the halide.
Silver Chloride = White
Silver Bromide = Cream
Silver Iodide = Yellow
Silver Nitrate = Colorless
3.3.2 DISCUSS THE CHANGES IN NATURE, FROM IONIC TO COVALENT AND FROM
BASIC TO ACIDIC, OF OXIDES ACROSS PERIOD 3
The acid-base properties of the oxides are closely linked to their bonding.
Metallic elements, which form ionic oxides, are basic; non-metal oxides,
which are covalent are acidic. Aluminum oxide, which can be considered as
an ionic oxide with some covalent character, shows amphoteric properties
reacting with both acids and bases
• Sodium and Magnesium Oxide
are bases
Na2O + HCl →2NaCl + H2O
MgO + H2SO4 → MgSO4 + H2O
Phosphorous Oxide, Sulfur Oxide,
Dichlorine Oxide reacts with water to
form Acid Solutions
P4O10 + 6H2O → 4H3PO4
• Aluminium oxide
is amphoteric (base and acid) so
it reacts with both bases and
acids
Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O
Al2O3 + 2NaOH +3H2O → 2NaAl(OH)4
SO3 + H2O → H2SO4
Cl2O7 + H2O → 2HClO4