Transcript Document

Periodic Table – Ionization Energies
• Energy is required to remove an electron from
an isolated gaseous atom. This energy, the
ionization energy, can be determined very
accurately. The energy is often supplied in the
form of light. High frequency/short
wavelength light is needed to ionize atoms.
The energy needed to ionize one gaseous atom
can be obtained from Ephoton = hνPhoton.
Ionization Energies – cont’d:
• The H atom is unique in that H has only one
electron. After this electron is removed no
further ionization processes are possible. For
all other atoms several ionization steps are
possible. Increasing amounts amounts of
energy are required to remove successive
electrons. Why? There are clear periodic trends
for ionization energies. Ionization energies will
be used eventually in discussions of chemical
bonding.
Ionization Process
Ionization Energy
Mg(g) → Mg+(g) + e-
I1 = 738 kJ
Mg+(g) → Mg2+(g) + e-
I2 = 1451 kJ
Zeff2
I = RH 2
n
Ionization energies decrease as atomic radii increase.
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Ionization Energies – Group Trends
• For chemical families (Groups) ionization
energies drop as atomic radii increase. We
need to consider both atomic radii and
effective nuclear charge. In alkali metals the
single valence electron (ns1) is removed first.
As n increases the valence e- is located on
average further from the nucleus. The effect of
growing nuclear charge is again offset by inner
or core electron screening.
Ionization Energies – Group Trends
(I1Values in kJ∙mol-1)
Alkali Metals & H
Alkaline Earth
Metals
The Halogens
H
1312
Noble Gases
He
2372
Li
520
Be
899
F
1681
Ne
2080
Na
496
Mg
738
Cl
1251
Ar
1520
K
419
Ca
590
Br
1140
Kr
1351
Rb
403
Sr
549
I
1008
Xe
1170
Cs
376
Ba
503
At
Rn
1037
Ionization Energies – Period Trends
• As we move from left to right across a period
values of first ionization energies increase and,
at the same time, atomic radii decrease.
Moving across the 2nd and 3rd periods the
atomic number, the nuclear charge and the
total number of valence electrons increases
steadily. The number of core electrons remains
constant and causes screening to produce an
effective nuclear charge for the outermost
electrons. Zeffective ≈ Z – S (Very roughly!)
Third Period – Effective Charges
Trends in Atomic Radii (pm)and
Ionization Energies
Na
Mg
Al
Si
P
S
Cl
Ar
6.8
ZEffective
2.5
3.3
4.1
4.3
4.9
5.5
6.2
Atomic
Radius(pm)
186
160
143
117
110
104
99
I1 (kJ∙mol-1)
496
738
578
786
1012
1000
1251
1520
FIGURE 9-10
First ionization energies as a function of atomic number
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Successive Ionization Energies Trends
• Ionization energies increase as successively
more electrons are removed from an atom. The
relative values of the ionization energies show
surprising jumps or “discontinuities”. This is
illustrated by data for Mg where the first
second and third ionization energies are:
• I1 = 738 kJ∙mol-1, I2 = 1451 kJ∙mol-1 and I3 =
7733 kJ∙mol-1 respectively. Let’s see if electron
configurations help us understand these data.
Magnesium – Successive Ionizations
• Mg(g) → Mg+(g) + eI1 = 738 kJ
1s22s22p63s2 1s22s22p63s1
• Mg+(g) → Mg2+(g) + eI2 = 1451 kJ
1s22s22p63s1 1s22s22p6
• Mg2+(g) → Mg3+(g) + eI3 = 7733 kJ
1s22s22p6
1s22s22p5
• The first large jump in ionization energy
corresponds to the removal of the first nonvalence electron. Why is this reasonable?
577.6
1012
999.6
1451
7733
I2 (Mg) vs. I3 (Mg)
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I1 (Mg) vs. I1 (Al)
General Chemistry: Chapter 9
I1 (P) vs. I1 (S)
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Electron Affinities
• In the gas phase most nonmetal atoms will
pick up an electron (or two!) to form a
negatively charged monatomic ion. The
process is usually exothermic and the term
electron affinity tells us the size of the energy
change associated with this process.
Neglecting the Noble Gases the EA values are
generally most exothermic for the non-metals
with the smallest atomic radii.
Electron Affinities – cont’d:
• Surprisingly, most gaseous metal atoms can
also pick up an electron (usually an exothermic
process)! (Discuss this again when consider
chemical bonding.) There are surprises with
nonmetals! E.g., the second EA value for
oxygen is +ve (an endothermic process). The
O2- ion is stable in binary ionic compounds,
though, due to the lattice energy of compounds
such as MgO(s).
Electron Affinity
F(g) + e- → F-(g)
EA = -328 kJ
F(1s22s22p5) + e- → F-(1s22s22p6)
Li(g) + e- → Li-(g)
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EA = -59.6 kJ
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Values are in kilojoules per mole for the process
X(g) + e-
X-(g).
FIGURE 9-11
Electron affinities of main-group elements
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Second Electron Affinities
O(g) + e- → O-(g)
EA = -141 kJ
O-(g) + e- → O2-(g)
EA = +744 kJ
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Magnetic Properties
• Diamagnetic atoms or ions:
– All e- are paired.
– Weakly repelled by a magnetic field.
• Paramagnetic atoms or ions:
– Unpaired e-.
– Attracted to an external magnetic field.
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Magnetic Properties
• For both atoms and monatomic ions we can
readily identify paramagnetic species using
orbital diagrams for the atoms and monatomic
ions. Often one writes electron configurations
first. With some practice one can jump to
considering the valence electrons and write
partial orbital diagrams for the valence
electron subshells.
Paramagnetism
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Periodic Properties of the Elements
FIGURE 9-12
•Atomic properties and the periodic table – a summary
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“Periodic Physical Properties”
• The next slide shows pictures of three
molecular halogens – chlorine, bromine and
iodine. “Obviously”, from the pictures alone!,
one can see that chlorine has the lowest
melting and boiling point. Why? Let’s estimate
values for the missing melting point and
boiling point of bromine and compare the
estimates to the experimental values.
266
?
332
?
FIGURE 9-13
Three halogen elements
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FIGURE 9-14
Melting points of the third-period elements
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Reducing Abilities of Group 1 and 2
Metals
2 K(s) + 2 H2O(l) → 2 K+ + 2 OH- + H2(g)
I1 = 419 kJ
I1 = 590 kJ
I2 = 1145 kJ
Ca(s) + 2 H2O(l) → Ca2+ + 2 OH- + H2(g)
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Oxidizing Abilities of the Halogen
Elements
(Group 17)
2 Na + Cl2 → 2 NaCl
Cl2 + 2 I- → 2 Cl- + I2
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Acid-Base Nature of Element Oxides
Basic oxides or base anhydrides:
Li2O(s) + H2O(l) → 2 Li+(aq) + 2 OH-(aq)
Acidic oxides or acid anhydrides:
SO2 (g) + H2O(l) → H2SO3(aq)
Na2O and MgO yield basic solutions
Cl2O, SO2 and P4O10 yield acidic solutions
SiO2 dissolves in strong base, acidic oxide.
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Class Examples:
• 1. Which of the following atoms and ions are
paramagnetic (i.e. have unpaired electrons).
Note: An even number of electrons does not
indicate that all electrons are paired. (a) He
atom, (b) F atom, (c) As atom, (d) F- ion (e)
Al3+ ion and (f) Fe atom.
• 2. Arrange the following in order of increasing
atomic radius: (a) Mg, Ba, Be, Sr (b) Rb+, Se2-,
Br- and Sr2+ (c] Ca, Rb, F, S (d) Fe, Fe3+, Fe2+.
Class Examples – cont’d:
• 3. Write balanced chemical equations to
represent the reactions of the following oxides
with water: (a) SO3(g), (b) P4O10(s), (c) BaO(s)
and (d) Li2O(s).
• 4. Arrange the following atoms in order of
increasing first ionization energy: (a) Fr, He,
K, Br (b) P, As, N, Sb and (c) Sr, F, Si, Cl.
• 5. Why are transition metal atoms and ions so
often paramagnetic?
Class Examples
• 6. What information does the term “degenerate
orbitals” convey?
• 7. How do a ground state and an excited state
electron configuration differ?
• 8. How many electrons are described using the
notation 4p6? How many orbitals does this
notation include? What is the shape of the
orbitals described using the 4p6 notation?
Class Examples
• 9. Are all (neutral) atoms having an odd atomic
number paramagnetic? Are all atoms having an
even atomic number necessarily diamagnetic?
Explain.
• 10. The following electron configurations do
not correspond to the ground electronic state of
any atom. Why? (a) 1s22s22p64s1, (b)
1s22s22p63s23p63d2 (c)
1s22s22p63s23p63d84s24p2.