Download chemical periodicity

CHEMICAL
PERIODICITY
Development of Periodic Table
Dmitri Mendeleev, Russian Chemist in the mid-1800’s
arranged the 70 known elements into a systematic way.
– Put the names of elements on a card
(Just like you arranged the pieces)
– He also put atomic masses, physical and chemical
properties.
– When arranged in the order of atomic mass, he
discovered a repeated, or periodic, pattern.
– He left blank spaces where he didn’t know the element
for the spot.
– By doing this, he constructed the first periodic table for
elements.
Mendeleev’s Problems
1. Why could most of the elements be
arranged in the order of increasing
atomic mass but a few could not?
2. What was the reason for chemical
periodicity?
Henry Moseley’s Solution
1. Instead of arranging in atomic mass,
in1913 Henry Moseley arranged in
increasing order of nuclear charge.
His work led to modern definition and
recognition of atomic number.
2. He discovers that it is the number of
electrons that affect an atoms’ reactivity,
or chemical property.
Parts of the Periodic Table
• Periods:
– The horizontal rows
• Periodic Law:
– When the elements are arranged in order of
increasing atomic number, there is a periodic
repetition of their physical and chemical
properties.
• Groups: (a.k.a. Families)
– The vertical columns
Parts of the Periodic Table, cont.
• Representative Elements:
– Group A elements. (The tall parts of the
periodic table.)
– These can be divided into 3 groups:
• Metals
• Nonmetals
• Metalloids
Metals
• Malleable-Very easy to be hammered or beaten into thin
sheets
• Ductile-Very easy to produce a wire
• High luster-De-excitation of electrons causes shiny
appearance of metal surface
• High electrical conductivity - Due to the free movement of
electrons
• High thermal conductivity
• Most are solid at room temperature
• Metallic Bond Strength
– Directly proportional to the heat of vaporization
• 80% of all elements are metals
Nonmetals
• Upper right portion of the periodic table
• Generally nonlustrous
• Generally poor conductors of electricity
• Some are gases at room temperature and
some brittle solids
• One is even a liquid at room temperature
Metalloids
• These border the stair step, except for
aluminum
• These have properties of both metals and
nonmetals
• Si, Ge, As, Sb, Te
Group 1 – Alkali Metal
• Very reactive!Not found in nature as free elements
• Combine vigorously with most nonmetals
• React strongly with water to produce hydrogen gas and
aqueous solutions of substances known as “alkalis.”
• Usually stored in kerosene (paraffin oil)
Group 2 – Alkali Earth Metal
• Harder, denser, and stronger than alkali metal
• Higher melting points
• Not so reactive compared to alkali metal;
Still very reactive! Not found as free element
Group 3 – 12 Transition Metal
• Good conductors of electricity and have a
high luster
• Less reactive than alkali and alkali-earth
metal
• Some (e.g. platinum and gold) are so
unreactive that they do not form compounds
easily.
• Some are found as free element.
Group 17 - Halogens
• Most reactive nonmetal
• React vigorously with most metals to
form “salts”
• Need one electron to obtain stable
noble-gas configuration
• F and Cl are gaseous.
• Br is liquid.
• Iodine is solid.
Group 18 – Noble Gases
• Lack chemical reactivity
• Very stable
• Octet electronic configuration in the
outermost energy level (or “shell”)
Categories of electrons
• Core (inner) electrons – all those shared by the
previous noble gas
• Outer electrons – those in the highest occupied
energy level
– Similar chemical properties of elements in groups is a result
of similar outer electron configurations
– In the main group, the group number equals the number of
outer electrons
• Valence electrons – those involved in bonding
– In the main group, the outer electrons are valence
– In the transition metals can include some d electrons
Trends in metallic behavior
Metallic Behavior
• Metallic behavior increases left and
down
• Metals tend to lose electrons in
reactions
• Highly metallic elements, are likely to
make positive ions
• Least metallic elements are likely to
make negative ions
• In middle are more likely to make
covalent bonds
Ion (Cation and Anion)
Neutral atom
F
Neutral atom
Ca
Add electron
+ 1eRemove electron
-2e-
Negative ion (Anion)
F-1 (Anion)
Positive ion (Cation)
Ca+2 (Cation)
# of electrons added/removed= # of charge
Main-group ions and the noble
gas configurations
Octet (Duplet) Rule
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
8.2
Periodic Law
All the elements in a group have the same electron configuration in
their outermost shells
Example:
Group 2
Be 1S2 2S2
2, 2
Mg
1S2 2S2 2p6 3S2
2,
8,
2
Ca
1S2 2S2 2p6 3S2 3p6 4S2
2,
8,
8,
2
Periodic Table and Electron Configuration
Groups 1-2
Groups 3-8
Transition
Lantanides
s1 s2
1
2
3
4
5
6
=
=
=
=
s level
p level
d level
f level
p1 p2 p3 p4 p5 p6
d1 - d10
f1 - f14
Figure 8.13
The relation between orbital filling and the periodic table
Figure 8.12
Periodic Patterns
s
p
1
2
3
4
5
6
7
f (n-2)
d (n-1)
6
7
© 1998 by Harcourt Brace & Company
Trends in Some Periodic
Properties
• The physical and chemical behavior of the
elements is based on the electron
configurations of their atoms.
• e- configurations can be used to explain
many of the repeating or “periodic” properties
of the elements
Figure 8.14
Defining metallic and covalent radii
Atomic Radius - size of atom
8.3
Definition of Some Periodic Properties
Atomic Radius/Ionic Radius- size of atom/Ion
Electronegativity is a measure of an atom’s
attraction for another atom’s electrons.
Ionization energy: the energy required to
remove an electron from an atom is ionization
energy. (measured in kilojoules, kJ)
Electron affinity is the energy change that
occurs when an atom gains an electron (also
measured in kJ).
Atomic radii of the main-group and
transition elements.
Vertically
•The trend for atomic radius in a vertical
column is to go from smaller at the top to
larger at the bottom of the family.
•With each step down the family, we add
an entirely new energy level, making the
atoms larger with each step.
Horizontally
•Each step adds a proton and an electron
(and 1 or 2 neutrons).
•Electrons are added to existing energy
level.
The effect is that the more positive
nucleus has a greater pull on the electron
cloud.
The nucleus is more positive and the
electron cloud is more negative.
The increased attraction pulls the cloud
in, making atoms smaller as we move
from left to right across a period.
Electron Shells and Sizes of
Atoms
• Size in main group elements follow 2 general
rules
– 1. AR increases going down in a group
– 2. AR decreases going left to right in a period
• Why?
• Moving down in a group:
– Zeff is constant because Z and S increase equally
– Electrons in a higher n and therefore larger
• Moving across:
– Z increases and S does not (electrons in same
shell don't shield each other well) so e- attracted
more strongly and are closer/smaller
• In transition metals there is an initial decrease in size
but then remains relatively constant
Figure 8.16
Periodicity of atomic radius
SAMPLE PROBLEM 8.3
PROBLEM:
Using only the periodic table (not Figure 8.15), rank each set of
main group elements in order of decreasing atomic size:
(a) Ca, Mg, Sr
PLAN:
Ranking Elements by Atomic Size
(b) K, Ga, Ca
(c) Br, Rb, Kr
(d) Sr, Ca, Rb
Elements in the same group increase in size and you go down;
elements decrease in size as you go across a period.
SOLUTION:
(a) Sr > Ca > Mg
These elements are in Group 2A(2).
(b) K > Ca > Ga
These elements are in Period 4.
(c) Rb > Br > Kr
Rb has a higher energy level and is far to the left.
Br is to the left of Kr.
(d) Rb > Sr > Ca
Ca is one energy level smaller than Rb and Sr.
Rb is to the left of Sr.
Effective nuclear charge (Zeff) is the “positive charge” felt
by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Z
Core
Zeff
Radius (pm)
Na
11
10
1
186
Mg
12
10
2
160
Al
13
10
3
143
Si
14
10
4
132
Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
Ions
• Here is a simple way to remember which
is the cation and which the anion:
+
This is Ann Ion.
She’s unhappy and
negative.
+
This is a cat-ion.
He’s a “plussy” cat!
Cation Formation
Effective nuclear
charge on remaining
electrons increases.
Na atom
1 valence electron
11p+
Valence elost in ion
formation
Result: a smaller
sodium cation, Na+
Remaining e- are
pulled in closer to
the nucleus. Ionic
size decreases.
Anion Formation A chloride ion is
produced. It is
larger than the
original atom.
Chlorine
atom with 7
valence e17p+
One e- is added
to the outer
shell.
Effective nuclear charge is
reduced and the e- cloud
expands.
Cation is always smaller than atom from
which it is formed.
Anion is always larger than atom from
which it is formed.
The Radii (in pm) of Ions of Familiar Elements
Ionic Size
• Cations are smaller than the parent atom
• Anions are larger than the parent atom.
• Ionic size increase down in a group
• In an isoelectronic series, the most negative ion
is largest, the most positive is smallest.
• Atoms making more than one ion, most positive
is smallest.
Ionization Energy
• This is the second important periodic trend.
• The atom has been “ionized” or charged.
• The number of protons and electrons is no longer
equal.
• The energy required to remove an electron from an
atom is ionization energy. (measured in kilojoules, kJ)
• The larger the atom is, the easier its electrons are to
remove.
• Ionization energy and atomic radius are inversely
proportional.
• Ionization energy is always endothermic, that is
energy is added to the atom to remove the electron.
Ionization Energy (Potential)
• Draw arrows on your help sheet like this:
Ionization energy
• Trends:
– Generally as size decreases, IE1 increases.
– 1. IE1 generally increases from left to right, some
exceptions
– 2. IE1 generally decreases going down in a group
– 3. transition and f-block elements have much
smaller variances in IE1
• But Why?
– Across - increasing Zeff and smaller size(e- closer
to nucleus)
– exceptions - Be to B because s shields p and
lowers Zeff
– N to O because repulsions in first paired electron
– Down - Zeff constant and larger so easier to ionize
Figure 8.17
Periodicity of first ionization
energy (IE1)
Figure 8.18
First ionization
energies of the
main-group
elements
SAMPLE PROBLEM 8.4
PROBLEM:
Using the periodic table only, rank the elements in each of the
following sets in order of decreasing IE1:
(a) Kr, He, Ar
PLAN:
Ranking Elements by First Ionization Energy
(b) Sb, Te, Sn
(c) K, Ca, Rb
(d) I, Xe, Cs
IE decreases as you proceed down in a group; IE increases as
you go across a period.
SOLUTION:
(a) He > Ar > Kr
Group 8A(18) - IE decreases down a group.
(b) Te > Sb > Sn
Period 5 elements - IE increases across a period.
(c) Ca > K > Rb
Ca is to the right of K; Rb is below K.
(d) Xe > I > Cs
I is to the left of Xe; Cs is furtther to the left and
down one period.
Ionization energy
• Second ionization energy IE2 – to remove 2nd
e– Always larger than IE1
– as e- are removed, Z remains constant and
remaining e- are harder to remove
• A huge increases occurs in ionization
energies when core electrons are reached
because of the lower shielding/higher Zeff.
• Result - Rx only involve outer e-
Figure 8.19
The first three ionization energies
of beryllium (in MJ/mol)
SAMPLE PROBLEM 8.5
PROBLEM:
PLAN:
Identifying an Element from Successive
Ionization Energies
Name the Period 3 element with the following ionization energies
(in kJ/mol) and write its electron configuration:
IE1
IE2
IE3
IE4
IE5
1012
1903
2910
4956
6278
IE6
22,230
Look for a large increase in energy which indicates that all of the
valence electrons have been removed.
SOLUTION:
The largest increase occurs after IE5, that is, after the 5th valence
electron has been removed. Five electrons would mean that the
valence configuration is 3s23p3 and the element must be
phosphorous, P (Z = 15).
The complete electron configuration is 1s22s22p63s23p3.
Electronegativity
• Electronegativity is a measure of an atom’s
attraction for another atom’s electrons.
• It is an arbitrary scale that ranges from 0 to
4.
• The units of electronegativity are Paulings.
• Generally, metals are electron givers and
have low electronegativities.
• Nonmetals are are electron takers and have
high electronegativities.
• What about the noble gases?
Electronegativity
• Your help sheet should look like this:
0
Overall Reactivity
• This ties all the previous trends together in
one package.
• However, we must treat metals and
nonmetals separately.
• The most reactive metals are the largest
since they are the best electron givers.
• The most reactive nonmetals are the
smallest ones, the best electron takers.
Overall Reactivity
• Your help sheet will look like this:
0
Electron Affinity
• What does the word ‘affinity’ mean?
• Electron affinity is the energy change that
occurs when an atom gains an electron
(also measured in kJ).
• Where ionization energy is always
endothermic, electron affinity is usually
exothermic, but not always.
Electron Affinity
• Electron affinity is exothermic if there is an
empty or partially empty orbital for an
electron to occupy.
• If there are no empty spaces, a new orbital
or PEL must be created, making the
process endothermic.
• This is true for the alkaline earth metals
and the noble gases.
Electron Affinity
• Your help sheet should look like this:
+
+
electron affinities
• energy required to put an e- on an atom
(making it negative or more negative)
• EA1 usually negative
• EA2 always positive
• Irregular trend: increases going right and up
– Remember: Full shells best, full subshells good,
and half full is kinda cool
– Cl and other halogens need only 1 e- so most
exothermic reaction
– noble gases have endo because new electron in
higher n
– Mg, Be endo because new e in higher l
– N endo because new e paired
Figure 8.20
Electron affinities of the main-group elements
Figure 8.21
Trends in three atomic properties
Atomic radii increases
Atomic radii decreases
Electronegativity Increases
Electronegativity Increases