Transcript Periodicity

Periodicity
General info.
 Periodicity is concerned with trends or patterns
seen within elements on the periodic table.
 The patterns that will be covered are:
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Atomic radius
Ionization energy
Electronegativity
Melting points (IB/SL only)
Period 3 (IB/SL only
Atomic radius
 Remember, the radius is determined
directly by the electron cloud
 Atomic radii
 Going down a group (column)
 The radius increases going down a group.
 This increase is due to the addition of
energy levels. Each e.l. increases the
electron cloud significantly.
 Going across a period.
 The radius decreases going across a period.
 The nucleus increases, which means there is a
greater positive charge present. This increase in
positive charge pulls the electrons in closer to the
nucleus, thus, reducing the radius.
 No energy levels are being added.
Chart of atomic radii
Chart showing radius
trends across several
periods
Atomic radii with d block
elements
Ionic radii
 The radius of an atom increases when an
electron is added to make a (-) ion
(anion).
 Anions have more e- than protons so the
additonal e- is not pulled in close to the
nucleus. This causes the e- cloud to
expand.
Ionic radii continued
 The radius of an atom decreases when an
electron is lost to make a (+) ion (cation).
 Cations have more protons than electrons, so
the electrons now have a stronger pull on them
from the nucleus, thus shrinking the e- cloud.
 Also, usually when atoms lose any e-, it is all of
the valence e- they lose, so the entire energy
level is lost. This really decreases the e- cloud.
 Transitional metals do not lose an e.l.
Metal ionic radii
Nonmetal ionic radii
Ionic radii charts
Ionization energy (IE)
 Ionization energy- the energy required to remove an
electron from a gaseous atom.
 Since the electrons are attracted to the nucleus, it
takes energy to pull the e- from the atom.
 Electrons in the last e.l. are the electrons that are being
removed. These electrons are called valence
electrons. D-block elements have their valence ein the last two energy levels.
 First ionization energy. The energy required to remove
the first valence electron from a neutral atom.
Trends in the first
ionization energy
 Going down a group
 The first ionization energy decreases.
 The radius is increasing and the further away an e- is from
the nucleus, the less pull (attractive force) there is on the
electron.
 The further away an e- is from the nucleus, there are more
e- (core e-) between the nucleus and the valence e-, thus
weakening the pull from nucleus. (this is called shielding.)
 With the addition of each new energy level the valence
electrons become further away from the nucleus.
 Going across a period
 The first ionization energy increases
 The atomic radius decreases. The valence ebecome closer to the nucleus going from left to
right , therefore there is a stronger pull (attractive
force) from the nucleus on the valence e-. So,
more energy is required to remove an electron.
3-d chart on first
ionization energies
First ionization energy
trends across several
periods.
Successive ionization
energies (IE)
 The energy required to remove the first
e- from an atom is the 1st ionization
energy.
 The energy required to remove a 2nd eis the 2nd ionization energy and so on.
 The energy required to remove an eincreases with each successive e-. For
example: Al: IE1= 577 IE2= 1815,
IE3=2740 all in kJ/mole
Succesive ionization
energies cont.
 Successive ionization energies increase
because.
 After each e- is removed, the radius
becomes smaller making the remaining ecloser to the nucleus.
 There are more protons than e-, making the
attractive force from the nucleus stronger on
the core e-.
Successive ionization
energies cont.
 After all of the valence e- have been
removed from an atom, there is a huge
increase in the IE.
 IE for core e- are extremely high due to them
being in a lower e.l. that is closer to the
nucleus.
 There are many more protons than e-.
Chart of succesive
ionization energies ( IE)
Electronegativity
 A measurement of an element’s ability to
attract an electron from another atom
within a bond. When 2 atoms of different
elements bond, one atom is better
attracting the electrons in the bond. This
attraction is electronegativity.
Electronegativity
continued
 An atom is more electronegative if:
 The atom has a small e- cloud. The small ecloud allows its nucleus to get closer to
another atom’s electrons and pull off e Larger nucleus: more attractive force on
another atom’s electrons.
 In summary: volume/mass ratio should
be small.
Electronegativity
continued.
 The most electronegative element is Fluorine.
 Fluorine is given an electronegative rating of
4.0 which is the highest rating. All other
elements are compared to it and are given a
relative rating.
 The trends in electronegativity are the same as
the ionization energy: increases across a
period, decreases down a group.
Chart of
electronegativities
Electronegative trends for
the first 5 periods.
Applying periodic trends
to chemical reactivity
 Metals are reactive if:
 (lose e- in
reactions)
 Larger radius
 Lower ionization
energy
 Fewer valence e-
 Nonmetals are
reactive if:
 (gain e- in
reactions.)
 Smaller radius
 High
electronegativity
 More valence e-
Melting point trends
 Down a group: The melting points tend to
decrease. The large radius decreases the
attraction between atoms, thus, making a
substance easier to melt.
 Across a period: the melting points
increase until group 14. Then they
decrease.
Melting points for the
group 1 elements
States of matter for all the
elements at 3507 C
Melting and boiling point
trends for elements 1-95
Chemical and physicial
properties
 Alkali metals:
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Very soft metals
Not very dense and so float on water
Low melting points: Li=454K to Cs=302 K
One valence e- with a low ionization energy.
Readily react with nonmetals
React with water to form a strong base.
Example:2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)
Continued:
 Halogens
Halogen
Fluorine Chlorine Bromine Iodine
Color
Pale
yellow
Yellow
green
Red
brown
Black or
purple
State at
room
temp.
Gas
Gas
Liquid
solid
Continued: Halogens
 All exist as diatomic molecules in their pure state. F2,
Cl2 Br2 I2
 Are nonpolar and have only van der waal forces when
in their pure form.
 Slightly soluble in water.
 When in water they dissociate slightly:
 X2 + H2O  H+ + X- + HOX,
 The HOX is a weak acid but a strong oxidant due to the
oxygen which reacts with other materials and bleaches them.
Example is HOCl is in bleach. These acids are also toxic to
microbes, and so make disinfectants and are used in water
treatment.
Continued: halogens
 They combine with metals to produce
ionically bonded salts called hallides. Ex.
NaCl, Most are soluble in water,
exception, AgCl, PbCl2
 The reactivity decreases going down the
group. (This is the opposite of the Alkali
metals.)The radius increases which
reduces the ability of the elements to
attract electrons.
Trends in period 3
 Left hand side: groups 1 and 2 alkali and alkaline
metals.
 Large radii, low IE, and so react as metals, losing e React with nonmetals to make solid ionic
compounds.
 These solids are crystaline solids with high melting
and boiling points. Ex. NaCl
 The solids easily dissolve in water to form ions.
 Oxides, such as MgO, K2O dissolve in water to form
bases. Ex. MgO + H2O  Mg(OH)2
 and can neutralize acids. ex MgO + 2HCl  MgCl2
+ H2O
 Middle of period 3
 Group 3:Aluminum is a metal
 Its oxide is amphoteric, that is it can dissolve in either an acid
or a base. Al2O3 + 6HCl  2AlCl3 + 3H2O,
 Al2O3 + 2OH- + 3H2O 2Al(OH)4 Makes crystalline solids like AlCl3 with high melting/boiling
points.
 Group4: Si, the ionization energy too great to behave as a metal
and so the first nonmetal appears.
 Its oxide will react with water to produce a weak acid.
 Ex. CO2 + H2O  H2CO3 (H+ + HCO3-)
 Make large crystalline solid networks with very high melting
a boiling points. Ex. Graphite, diamond, silicon chips.
 Group 5: Phosphorus: Makes primarily
covalently bonded molecules.
 Weak forces are between its molecules,
with chlorine or anything else.
 Very low melting and boiling points.
 Oxides form acid in water (versions of
phosphoric acid.) H3PO4
 Group 6 Sulfur.
 Makes covalent compounds with weak
intermolecular forces, low melting/boiling
points.
 Oxides dissolve to form acid in water:
 SO2 + H2O  H2SO3
 or SO3 + H2O  H2SO4
 Group 7:Halogens
 Are nonmetals that exist as diatomic
molecules.
 F2, Cl2, are gases, Br2 liquid, I2 solid. Only van
der waal forces exist between molecules.
 These elements combine with metals to make
water soluble ionic compounds that have high
melting and boiling points. (Silver compounds
are not water soluble.)
 Reactions are in previous slide.
Summary of period 3:
trends seen.
 Metal or nonmetal: Metallic  nonmetallic
 Physical state in pure form: solidgas.
 Pure element combines with water: bases  to
acids
 Oxides react with water to form: bases
amphoteric  acids, oxidants or bleaches.
 Compounds: crystalline solids (groups 1,2) 
strong networks (group 4)  weak molecular
solids (groups 5,6)  gases (group7)
exception is iodine