CHAPTER 6
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Transcript of CHAPTER 6
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CHAPTER 6
Chemical Periodicity
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Chapter Goals
1. More About the Periodic Table
Periodic Properties of the Elements
2. Atomic Radii
3. Ionization Energy
4. Electron Affinity
5. Ionic Radii
6. Electronegativity
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More About the Periodic TableEstablish 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 (full s and p orbitals), their chemical reactions are similar. He 1s2
Ne [He] 2s2 2p6
Ar [Ne] 3s2 3p6
Kr [Ar] 4s2 4p6
Xe [Kr] 5s2 5p6
Rn [Xe] 6s2 6p6
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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.
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More About the Periodic Table
d-Transition Elements Elements on periodic chart
in B groups.
Each metal has d electrons. ns (n-1)d configurations
These elements make the transition from metals to nonmetals.
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More About the Periodic Table
f - transition metals Sometimes called inner
transition metals.Electrons are being added to f orbitals.Consequently, very slight variations of properties from one element to another.Outermost electrons have the greatest influence on the chemical properties of elements.
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Periodic Properties of thePeriodic Properties of theElements – Atomic RadiiElements – Atomic Radii
Atomic radii describes the relative sizes of atoms.
Atomic radii increase within a column going from the top to the bottom of the periodic table.
Atomic radii decrease within a row going from left to right on the periodic table.
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Atomic Radii
The reason the atomic radii decrease across a period is due to shielding or screening effect. Effective nuclear charge, Zeff, experienced by an electron is less
than the actual nuclear charge, Z. The inner electrons block the nuclear charge’s effect on the outer
electrons.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: Arrange these elements based on their atomic radii. Se, S, O, Te
Example: Arrange these elements based on their atomic radii. P, Cl, S, Si
Example: Arrange these elements based on their atomic radii. Ga, F, S, As
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Redox Reactions
Why do metals losa electrons in their reactions?
Why does Mg form Mg2+ ions and not Mg3+?
Why do nonmetals take on electrons?
Ionization Energy
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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+ + e-
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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-
Mg+ has 12 protons and only 11 electrons. Therefore, IE for Mg+ > Mg
Atoms can have 3rd (IE3), 4th (IE4), etc. ionization energies.
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Ionization Energy
IE2 > IE1 It always takes more energy to remove a second electron from an ion than from a neutral atom.
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.
IE1 generally decreases moving down a family.
IE1 for Li > IE1 for Na, etc.
Periodic trends for Ionization Energy:
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Ionization Energy
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First Ionization Energies of Some Elements
0
500
1000
1500
2000
2500
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
Atomic Number
Ionization Energy (kJ/mol)
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
AlSi
P
S
Cl
Ar
K
Ca
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First Ionization Energies of Some Elements
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Ionization Energy
Example: Arrange these elements based on their first ionization energies (IE). Sr, Be, Ca, Mg
Example: Arrange these elements based on their first IE. Al, Cl, Na, P
Example: Arrange these elements based on their first IE. B, O, Be, N
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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
IA
Na
IIA
Mg
IIIA
Al
IVA
Si
IE1 (kJ/mol)
496 738 578 786
IE2
(kJ/mol)4562 1451 1817 1577
IE3
(kJ/mol)6912 7733 2745 3232
IE4
(kJ/mol)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+.
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Electron Affinity
Electron affinity is the amount of energy absorbed when an electron is added to an isolated gaseous atom to form an ion with a 1- charge.
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
Mg(g) + e- + 231 kJ/mol Mg-(g)
EA = +231 kJ/mol
Br(g) + e- Br-(g) + 323 kJ/mol
EA = -323 kJ/mol
Two examples of electron affinity values:
Sign conventions for electron affinity.
If electron affinity > 0 energy is absorbed.
If electron affinity < 0 energy is released.
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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.
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Electron Affinities of Some Elements
-400-350-300-250-200-150-100-50
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
Atomic Number
Ele
ctro
n A
ffin
ity
(kJ/
mo
l)
Electron Affinity
H
He
Li
Be B
C
N
O
F
NeNa
Mg Al
Si
P
S
Cl
ArK
Ca
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Electron Affinity
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Electron Affinity
Example: Arrange these elements based on their electron affinities. Al, Mg, Si, Na
Example: Arrange the following elements in order of increasing values of electron affinity, i.e., from most negative to least negative.
Cl, Se, S, Cs, Rb, Te
(a) Cl < S < Se < Rb < Te < Cs (b) Cl > Te > Se > S > Rb > Cs(c) Cl > Se > S > Te > Rb > Cs (d) Cl < S < Se < Te < Cs < Rb(e) Cl < S < Se < Te < Rb < Cs
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Ionic Radii
Cations are always smaller than their respective neutral atoms.
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Ionic Radii
Anions are always larger than their neutral atoms.
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Ionic RadiiCations 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 RadiiAnions 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
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Ionic Radii
Example: Arrange these elements based on their ionic radii. Ga3+, K+, Ca2+
Example: Arrange these elements based on their ionic radii. Cl-, Se2-, Br-, S2-
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Isoelectronic ions
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1- Which of the following statements is CORRECT with regard to
atomic or ionic size?
(1) S2- < Cl- (2) Br < Br- (3) Li- < Li (4) P < N
2- Select the largest species from the following group:
(A) Mg (B) Cl (C) S (D) Al
3- Select the smallest species from the following group:
(A) Fe3+ (B) Fe2+ (C) Fe+ (D) Fe
4- Select the element with the lowest ionization energy (the easiest
to ionize):
(A) Ga (B) In (C) B (D) Al
Problem
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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. Cesium and francium are the least electronegative
elements.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
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Electronegativity
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Electronegativity
Example: Arrange these elements based on their electronegativity. Se, Ge, Br, As
Example: Arrange these elements based on their electronegativity. Be, Mg, Ca, Ba
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Homework AssignmentHomework Assignment
One-line Web Learning (OWL):One-line Web Learning (OWL):
Chapter 6 Exercises and Tutors – Chapter 6 Exercises and Tutors – OptionalOptional
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End of Chapter 6