Transcript CHAPTER 6

Unit 4: Chemical Periodicity
More About the Periodic Table
• Establish a classification scheme of the elements
based on their electron configurations.
Noble Gases
– All of them have completely filled electron shells.
• Since they have similar electronic structures, their
chemical reactions are similar.
–
–
–
–
–
–
He
Ne
Ar
Kr
Xe
Rn
1s2
[He] 2s2 2p6
[Ne] 3s2 3p6
[Ar] 4s2 3d10 4p6
[Kr] 5s2 4d10 5p6
[Xe] 6s2 4f14 5d10 6p6
Outer shell may be represented
as having the electron
configuration of ns2 np6
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More About the Periodic Table
Representative Elements
– Are the elements in A groups on
periodic chart.
• These elements will have
their “last” electron in an
outer s or p orbital.
• These elements have fairly
regular variations in their
properties.
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More About the Periodic Table
d-Transition Elements
– Elements on periodic chart in B
groups.
– Sometimes called transition
metals.
• Each metal has d electrons.
– ns (n-1)d configurations
– E.g. 21Sc through 30Zn have 4s
and 3d occupied but NOT 4p
• These elements make the transition from metals to
nonmetals.
• Exhibit smaller variations from row-to-row than the
representative elements.
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More About the Periodic
Table
f - transition metals
– Sometimes called inner transition
metals.
• Electrons are being added to
f orbitals.
– Lanthanides, 4f orbitals occupied
– Actinides, 5f orbitals occupied
• Electrons are being added
two shells below the valence
shell!
• Consequently, very slight
variations of properties from
one element to another.
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More About the Periodic Table
• Outermost electrons have the greatest
influence on the chemical properties of
elements.
– Adding an electron to an s or p orbital usually
causes dramatic changes in the physical &
chemical properties
– Adding an electron to a d or f orbital typically
has a much smaller effect on properties.
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Periodic Properties of the Elements
•
•
•
Knowledge of periodicity is valuable in understanding
bonding in simple compounds
Variations useful in predicting chemical behaviour
Changes in properties depend on:
•
•
•
•
•
•
•
electron configurations, especially configuration in outmost
occupied shell
How far away that shell is from the nucleus
Atomic Radii
Ionization Energy
Electron Affinity
Ionic Radii
Electronegativity
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Atomic Radii
• Effective nuclear charge, Zeff, experienced
by an electron in an outer shell is less than
the actual nuclear charge, Z.
– This is because the inner electrons block/
screen/shield the nuclear charge’s effect on the
outer electrons.
• The concept of shielding or screening
helps us to understand many periodic
trends in atomic properties.
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Atomic Radii
• Within a family (group) of representative
elements, atomic radii increase from the top to
bottom of the periodic table as electrons are
added to shells further from the nucleus.
• E.g. 3Li has a 1s2 2s1 configuration.
– The outermost 2s1 electron is not as effectively
shielded as an electron in a shell further from nucleus
• E.g. 11Na has 10 inner e-s 1s2 2s2 2p6 and one in
an outer shell, 3s1
– The 10 inner e-s shield the outer-shell electron from
most of the +11 nuclear charge
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Atomic Radii
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Atomic Radii
• Atomic radii decrease going from left to right
across the periodic table as a proton is
added to the nucleus and an electron is
added to a particular shell.
– Moving across a period, each element has an
increased nuclear charge and the electrons are
going into the same shell (2s and 2p or 3s and
3p, etc.).
• Consequently, the outer electrons feel a stronger
effective nuclear charge.
• For Li, Zeff ~ +1
For Be, Zeff ~ +2
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Atomic Radii
• Example 1: Arrange these elements in
order of increasing atomic radii.
– Se, S, O, Te
You do it!
O < S < Se < Te
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Ionization Energy
• First ionization energy (IE1)
– The minimum amount of energy required to remove the
most loosely bound electron from an isolated gaseous
atom to form a 1+ ion.
• Symbolically:
Atom(g) + energy  ion+(g) + e-
Mg(g) + 738kJ/mol  Mg+ + e13
Ionization Energy
• Second ionization energy (IE2)
– The amount of energy required to remove the
second electron from a gaseous 1+ ion.
• Symbolically:
– ion+ + energy  ion2+ + e-
Mg+ + 1451 kJ/mol Mg2+ + e-
•Atoms can have 3rd (IE3), 4th (IE4), etc.
ionization energies.
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First Ionization Energies of Some Elements
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Ionization Energy
•
Periodic trends for Ionization Energy:
1. IE2 > IE1
It always takes more energy to remove a second electron
from an ion than from a neutral atom.
2. IE1 generally increases moving from IA elements to
VIIIA elements.
Important exceptions at Be & Mg, N & P, etc. due to filled and
half-filled subshells.
3. IE1 generally decreases moving down a family.
IE1 for Li > IE1 for Na, etc.
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Ionization Energy
• Example 2: Arrange these elements based
on their first ionization energies.
– Sr, Be, Ca, Mg
You do it!
Sr < Ca < Mg < Be
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Ionization Energy
• First, second, third, etc. ionization energies
exhibit periodicity as well.
• Look at the following table of ionization
energies versus third row elements.
– Notice that the energy increases enormously
when an electron is removed from a completed
electron shell.
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Ionization Energy
Group
and
element
IE1
(kJ/mol)
IE2
(kJ/mol)
IE3
(kJ/mol)
IE4
(kJ/mol)
IA
Na
IIA
Mg
IIIA
Al
IVA
Si
496
738
578
786
4562
1451
1817
1577
6912
7733
2745
3232
9540
10,550
11,580
4356
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Ionization Energy
• The reason Na forms Na+ and not Na2+ is that the
energy difference between IE1 and IE2 is so large.
– Requires more than 9 times more energy to remove the
second electron than the first one.
• The same trend is persistent throughout the
series.
– Thus Mg forms Mg2+ and not Mg3+.
– Al forms Al3+.
Attaining a noble gas configuration favours an atom of
a representative element in forming a monoatomic ion
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Ionization Energy
• The relative values of IE helps in predicting
whether an element would form ionic or
covalent compounds
– Elements with low IE  ionic compounds by
losing e-s (cations)
– Elements with intermediate IE  covalent
compounds
– Elements with very high IE  ionic compounds
by gain e-s (anions)
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Electron Affinity
• Electron affinity (EA) is the amount of energy
absorbed when an electron is added to an
isolated gaseous atom to form an ion with a 1charge.
• Sign conventions for electron affinity.
– If electron affinity > 0 energy is absorbed.
– If electron affinity < 0 energy is released.
• Electron affinity is a measure of an atom’s ability
to form negative ions.
• Symbolically:
atom(g) + e- + EA ion-(g)
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Electron Affinity
Two examples of electron affinity values:
Mg(g) + e- + 231 kJ/mol  Mg-(g)
EA = +231 kJ/mol
Br(g) + e-  Br-(g) + 323 kJ/mol
EA = -323 kJ/mol
Elements with very –ve
electron affinities gain
electrons easily to form
negative ions (anions)
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Electron Affinity
• General periodic trend for electron affinity is
– the values become more negative from left to right
across a period on the periodic chart.
– the values become more negative from bottom to
top up a row on the periodic chart.
Noteworthy exceptions:
Group 2A – very difficult to
add an e- because these
elements have their outer s
subshell filled
Group 5A – an additional ewould have to be added to a
half-filled set of np orbitals
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Electron Affinity
• Example 3: Arrange these elements based
on their electron affinities.
– Al, Mg, Si, Na
You do it!
Si < Al < Na < Mg
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Ionic Radii
• Cations (+ve ions)
are always smaller
than their
respective neutral
atoms.
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Ionic Radii
• Anions (negative ions) are always larger
than their neutral atoms.
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Ionic Radii
• Cation (positive ions) radii decrease from left to
right across a period.
– Increasing nuclear charge attracts the electrons and
decreases the radius.
Ion
Rb+
Sr2+
In3+
Ionic
Radii(Å)
1.66
1.32
0.94
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Ionic Radii
• Anion (negative ions) radii decrease from left
to right across a period.
– Increasing electron numbers in highly charged
ions cause the electrons to repel and increase the
ionic radius.
Ion
N3-
O2-
F1-
Ionic
Radii(Å)
1.71
1.26
1.19
– Example: O2- is larger than the isoelectric F- because
the oxide ion contains 10 e-s held by a nuclear
charge of 8+, whereas the F- ion has 10 e-s held by 29
a nuclear charge of 9+
Ionic Radii
• Both cation and anion sizes increase going
down a group
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Ionic Radii
• Example 4: Arrange these elements based
on their ionic radii.
– Ga, K, Ca
You do it!
K1+ > Ca2+ > Ga3+
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Electronegativity
• Electronegativity is a measure of the relative
tendency of an atom to attract electrons to itself
when chemically combined with another
element.
– Electronegativity is measured on the Pauling scale.
– Fluorine is the most electronegative element.
• E.g. EN value for F is 4.0  when F is chemically bonded to
other elements, it has a greater tendency to attract electron
density to itself than any other element
– Cesium and francium are the least electronegative
elements.
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Electronegativity
•
For the representative elements, electronegativities usually increase from
left to right across periods and decrease from top to bottom within groups.
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Electronegativity
• Example 5: Arrange these elements based
on their electronegativity.
– Se, Ge, Br, As
You do it!
Ge < As < Se < Br
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Periodic Trends
• It is important that you understand and
know the periodic trends described in the
previous sections.
– They will be used extensively in Chapter 7 to
understand and predict bonding patterns.
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Chemical Reactions & Periodicity
• In the next sections periodicity will be
applied to the chemical reactions of
hydrogen, and oxygen.
– They form the most kinds of compounds with
other elements
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Hydrogen
• Colourless, odourless ,
tasteless gas,
• Lowest molecular weight &
density
• Flammable
– Combustion reaction is
exothermic enough to provide
the heat needed to sustain the
reaction
– 2H2 (g) + O2 (g) 2H2O(l) + heat
At 803.8 feet in length and 135.1
feet in diameter, the German
passenger airship Hindenburg (LZ129) was the largest aircraft ever
to fly. The very flammable
hydrogen was responsible for the
Hindenburg disaster in 1937 37
Preparation of Hydrogen
• Hydrogen gas, H2, can be made in the
laboratory by the reaction of a metal with a
nonoxidizing acid.
Mg + 2 HCl MgCl2 + H2
• Hydrogen is commercially prepared by the
thermal cracking of hydrocarbons.
C4H10  2 C2H2 + 3 H2
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Preparation of Hydrogen
• Hydrogen may also be prepared by steam
cracking:
– CH4 (g) + H2O (g)
Ni catalyst
8300C
CO(g) + 3 H2 (g)
– The mixture of H2 and CO gases are referred to
“synthesis gas” and can be used to produce a variety of
organic compounds e.g. methanol, and hydrocarbon
mixtures for gasoline, kerosene
2002 prototype car from Chrysler that uses
methanol for fuel. A small reactor converts
methanol, H2O and O2 into H2 and CO2.
the H2 then reacts further with O2 to
produce electricity to power the car.
Methanol easier and safer to store than H2
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Reactions of Hydrogen & the Hydrides
• Hydrogen reacts with active metals to yield
solid ionic hydrides.
– Example:
2 K(l) + H2 (g)  2 KH (s)
Ba + H2  BaH2
• In general for IA metals, this reaction can be represented as:
2 M + H2  2 MH
•In general this reaction for IIA metals can be represented as:
M + H2  MH2
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Reactions of Hydrogen & the Hydrides
• The ionic hydrides are basic.
– The H- reacts with water to produce H2 and
OH-.
H- + H2O  H2 + OH• For example, the reaction of LiH with water proceeds
in this fashion.

(aq)
LiH (s)  H 2O( )  H 2(g)  OH
 Li

(aq)
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Reactions of Hydrogen & the Hydrides
• Hydrogen reacts with nonmetals to
produce covalent binary
compounds  molecular hydrides
– One example are the haloacids
produced by the reaction of hydrogen
with the halogens.
H2 + X2 
2 HX
• For example, the reactions of F2 and
Br2 with H2 are:
H2 + F2  2 HF
H2 + Br2  2 HBr
Hydrogen burns in an
atmosphere of pure Cl2 to
produce hydrogen chloride, HCl
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Reactions of Hydrogen & the Hydrides
• Hydrogen reacts with oxygen and other VIA
elements to produce several common binary
covalent compounds.
– Examples of this reaction include the production
of H2O, H2S, H2Se, H2Te.
2 H2 + O2  2 H2O
8 H2 + S8  8 H2S
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Reactions of Hydrogen & the Hydrides
• The hydrides of Group VIIA and VIA
hydrides are acidic.

(aq)

(aq)
HCl  H  Cl



H 2S  H (aq)  HS(aq)
(a strong acid)
(a weak acid)
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Reactions of Hydrogen & the Hydrides
• The primary industrial use of Hydrogen is in the synthesis of
ammonia, a molecular hydride, by the Haber process
• Most of the NH3 produced is used as a fertilizer or to make
other fertilizers e.g. ammonium nitrate NH4NO3 and
ammonium sulfate NH4SO4
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Reactions of Hydrogen &the Hydrides
•
There is an important periodic trend
evident in the ionic or covalent character
of hydrides.
1. Metal hydrides are ionic compounds
and form basic aqueous solutions.
2. Nonmetal hydrides are covalent
compounds and form acidic aqueous
solutions.
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Oxygen and the Oxides
• Joseph Priestley discovered oxygen in 1774
using this reaction:
2 HgO(s) 2 Hg() + O2(g)
Red powder
colourless gas
•A common laboratory preparation method for
oxygen is:
2 KClO3 (s)  2 KCl(s) + 3 O2(g)
•Commercially, oxygen is obtained from the
fractional distillation of liquid air.
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Oxygen and the Oxides
• Ozone (O3) is an allotropic form of oxygen
which has two resonance structures.
• Ozone is an excellent UV light absorber in the
earth’s atmosphere.
2 O3(g)  3 O2(g)
in presence of UV
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Reactions of Oxygen & the Oxides
• Oxygen is an extremely reactive element.
– O2 reacts with most metals to produce normal
oxides having an oxidation number of –2.
4 Li(s) + O2(g)  2 Li2O(s)
However, oxygen reacts with sodium to
produce a peroxide having an oxidation
number of –1.
2 Na(s) + O2(g)  Na2O2(s)
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Reactions of Oxygen & the Oxides
• Oxygen reacts with heavier members of
group 1  K, Rb, and Cs to produce
superoxides having an oxidation number of
-1/2.
K(s) + O2(g)  KO2(s)
Oxygen reacts with IIA metals to give
normal oxides.
2 M(s) + O2(g)  2 MO(s)
2 Sr(s) + O2(g)  2 SrO(s)
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Reactions of Oxygen & the Oxides
• At high oxygen pressures the 2A metals can form
peroxides.
Ca(s) + O2(g)  CaO2(s)
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Reactions of Oxygen & the Oxides
• Metals that have variable oxidation states,
such as the d-transition metals, can form
variable oxides.
– For example, in limited oxygen:
2 Mn(s) + O2(g)  2 MnO(s)
– In excess oxygen:
4 Mn(s) + 3 O2(g)  2 Mn2O3(s)
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Reactions of Oxygen & the Oxides
• Oxygen reacts with nonmetals to form
covalent nonmetal oxides.
• For example, carbon reactions with oxygen:
– In limited oxygen
2 C(s) + O2(g)  2 CO(g)

In excess oxygen
C(s) + O2(g)  CO2(g)
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Reactions of Oxygen & the Oxides
• Phosphorous reacts similarly to carbon
forming two different oxides depending on
the oxygen amounts:
– In limited oxygen
P4(s) + 3 O2(g)  P4O6(s)

In excess oxygen
P4(s) + 5 O2(g)  P4O10(s)
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Reactions of Oxygen & the Oxides
• Similar to the nonmetal hydrides, nonmetal
oxides are acidic.
– Sometimes nonmetal oxides are called acidic
anhydrides.
– They react with water to produce ternary acids.
• For example:
CO2(g) + H2O ()  H2CO3(aq)
Cl2O7(s) + H2O ()  2 HClO4(aq)
As2O5(s) + 6 H2O()  4 H3AsO4(aq)
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Reactions of Oxygen & the Oxides
• Similar to metal hydrides, metal oxides are
basic.
– These are called basic anhydrides.
– They react with water to produce ionic metal
hydroxides (bases)
Li2O(s) + H2O()  2 LiOH(aq)
CaO(s) + H2O ()  Ca(OH)2(aq)
Metal oxides are usually ionic and basic.
Nonmetal oxides are usually covalent and acidic.
An important periodic trend.
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Reactions of Oxygen & the Oxides
• Nonmetal oxides react with metal oxides to
produce salts.
Li2O(s) + SO2(g)  Li2SO3(s)
Cl2O7(s) + MgO(s)  Mg(ClO4)2(s)
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Combustion Reactions
• Combustion reactions are exothermic redox
reactions
– Some of them are extremely exothermic.
• One example of extremely exothermic
reactions is the combustion of
hydrocarbons.
– Examples are butane and pentane combustion.
2 C4H10(g) + 13 O2(g)  8 CO2(g) + 10 H2O(g)
C5H12(g) + 8 O2(g)  5 CO2(g) + 6 H2O(g)
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Fossil Fuel Contaminants
• When fossil fuels are burned, they frequently have
contaminants in them.
• Sulfur contaminants in coal are a major source of
air pollution.
– Sulfur combusts in air.
S8(g) + 8 O2(g)  8 SO2(g)
Next, a slow air oxidation of sulfur dioxide occurs.
2 SO2(g) + O2(g)  2 SO3(g)
Sulfur trioxide is a nonmetal
oxide, i.e. an acid anhydride.
SO3(g) + H2O()  H2SO4(aq)
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Fossil Fuel Contaminants
• Nitrogen from air can also be a source of
significant air pollution.
• This combustion reaction occurs in a car’s
cylinders during combustion of gasoline.
N2(g) + O2(g)  2 NO(g)
After the engine exhaust is released, a slow
oxidation of NO in air occurs.
2 NO(g) + O2(g)  2 NO2(g)
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Fossil Fuel Contaminants
• NO2 is the haze that we
call smog.
– Causes a brown haze in air.
• NO2 is also an acid
anhydride.
– It reacts with water to form
acid rain and, unfortunately,
the NO is recycled to form
more acid rain.
3 NO2(g) + H2O()  2 HNO3(aq) + NO(g)
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Group Question
• What do the catalytic converters that are
attached to all of our cars’ exhaust systems
actually do? How do they decrease air
pollution?
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