Transcript Chapter 1: Fundamental Concepts

Ch. 8 Periodic Properties of the Elements
Multielectron Atoms
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“Hydrogen-like” orbitals are used for all atoms
Energy levels are affected by other electrons
– Coulomb’s Law—electrostatic repulsion of like charges is proportional
to the amount of charge, and inversely proportional to the distance
between them (see text for eqn)
– Shielding—screening of one electron from the nuclear charge by other
electrons around the same atom
– Penetration—probability of the electron to be close to the nucleus
– Effective nuclear charge (Zeff)—the amount of nuclear charge an
electron experiences after taking shielding into account
– Degenerate—of equal energy
Order of Filling Subshells
Electron Spin and the Pauli Exclusion Principle
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Electrons have intrinisic angular momentum -- “spin” -- ms
– Possible values: ms = +1/2 and -1/2 (only two possible values)
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Pauli Exclusion Principle:
– No two electrons in an atom can have identical values of all 4 quantum
numbers -- maximum of 2 electrons per orbital!
– A single orbital can hold a “pair” of electrons with opposite “spins”
– e.g. the 3rd shell (n = 3) can hold a maximum of 18 electrons:
n=3
l =
0
1
2
subshell
3s
3p
3d
# orbitals
1
3
5
# electrons
2
6
10 = 18 total
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A single electron in an orbital is called “unpaired”
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Atoms with 1 or more unpaired electrons are paramagnetic,
otherwise they are diamagnetic
Electronic Configurations
• The Aufbau Principle -- Order of Filling Subshells
– Atomic # = # of protons = # electrons (in neutral atom)
– Add electrons to atomic orbitals, two per orbital, in the general
order of increasing principle quantum number n, for example:
#
Atom
Configuration
1
H
1s1
2
He
1s2
3
Li
1s22s1
4
Be
1s22s2
5
B
1s22s22p1
6
C
1s22s22p2
7
N
1s22s22p3
8
O
1s22s22p4
9
F
1s22s22p5
10
Ne
1s22s22p6
11
Na
1s22s22p63s1
Hund’s Rule
• Maximum number of unpaired electrons in orbitals of equal
energy
Orbital diagrams:
C __ __ __ __ __
1s
2s
2p
N __ __ __ __ __
1s
2s
2p
O __ __ __ __ __
1s
2s
2p
Relationship to Periodic Table
e.g. complete electronic configuration of Ge (#32, group IV)
Ge 1s22s22p63s23p64s23d104p2
or, Ge 1s22s22p63s23p63d104s24p2 (by values of n)
• Short-hand notation -- show preceding inert gas config.
– Ge [Ar]4s23d104p2 where [Ar] = 1s22s22p63s23p6
Valence Shell Configurations
• valence shell --
largest value of n (e.g. for Ge, n = 4)
plus any partially filled subshells
Ge 4s24p2 (valence shell electron configuration)
Ge
__ __ __ __ (valence shell orbital diagram)
4s
4p
Elements in same group have same valence shell e–
configurations
e.g. group V:
N
P
As
Sb
Bi
2s22p3
3s23p3
4s24p3
5s25p3
6s26p3
Sample Questions
• Write the complete electron configuration of gallium.
Answer:
• Write the short-hand electron configuration for zirconium.
Answer:
• Write the orbital diagram for the valence shell of tellurium.
Answer:
Sample Questions
• Write the complete electron configuration of gallium.
Answer:
Ga 1s22s22p63s23p64s23d104p1
• Write the short-hand electron configuration for zirconium.
Answer:
Zr [Kr]5s24d2
• Write the orbital diagram for the valence shell of tellurium.
Answer:
Te
___ ___ ___ ___
5s
5p
Sample Question
How many unpaired electrons does a ruthenium(II) ion, Ru2+,
have?
Show an appropriate orbital diagram to explain your answer.
Is the atom paramagnetic or diamagnetic?
Sample Question
How many unpaired electrons does a ruthenium(II) ion, Ru2+,
have?
Show an appropriate, valence-shell orbital diagram to explain
your answer.
Is the atom paramagnetic or diamagnetic?
Answer:
4 unpaired electrons, so paramagnetic
Orbital diagram:
Ru2+ ___ ___ ___ ___ ___
4d
Variation of Atomic Properties
Atomic Size (atomic radius, expressed in pm -- picometers)
(10–12 m!)
e.g. group 1 metals:
Atom
Radius in pm
Valence Shell
Li
152
2s1
Na
186
3s1
K
227
4s1
Cs
248
5s1
e.g. some elements in 2nd period:
Atom
B
C
N
O
F
radius
88
77
70
66
64
e– config
2p1
2p2
2p3
2p4
2p5
General Trend in Atomic Size
Relative sizes of ions
cations are smaller than parent atoms
e.g. Na
Na+
186 pm
95 pm
2s22p63s1
2s22p6
anions are larger than parent atoms
e.g. Cl
Cl–
99 pm
181 pm
3s23p5
3s23p6
Ionization Energy
I.E. = energy required to remove an electron from an atom or
ion (always endothermic, positive values)
e.g.
Li(g) --> Li+(g) + e–
I.E. = 520 kJ/mole
Exceptions: special stability of filled subshells, and of half-filled
subshells
Electron Affinity
• E. A. = energy released when an electron is added to an
atom or ion (usually exothermic, negative EA values)
e.g.
Cl(g) + e– --> Cl–(g)
E. A. = -348 kJ/mol
• The general trends in all these properties indicate that
there is a special stability associated with filled-shell
configurations.
• Atoms tend to gain or lose an electron or two in order to
achieve a stable “inert gas configuration” -- many important
consequences of this in chemical bonding.
Types of Elements
Metals:
Shiny, malleable, ductile solids with
high mp and bp
Good electrical conductors
Metal character increases to lower left of periodic table
Nonmetals:
Gases, liquids, or low-melting solids
Non-conductors of electricity
Diatomic elements: H2, O2, N2, F2, Cl2, Br2, I2
Metalloids:
Intermediate properties, often semiconductors
Sample Questions
Of the following atoms, circle the one with the highest electron
affinity.
K
Cl
P
Br
Na
Write a balanced chemical equation that corresponds to the
electron affinity of the element that you selected above.
Sample Question
Of the following atoms, circle the one with the highest electron
affinity.
K
Cl
P
Br
Na
Write a balanced chemical equation that corresponds to the
electron affinity of the element that you selected above.
Answer:
Cl(g) + e– --> Cl–(g)
Alkali Metals
• They want to be +1!
• Easily oxidized, low EA, low IE.
• Density increases moving down the group. (mass rises
faster than atomic radius)
• Reactions
– With halogens to form salts, e.g.
2 Na(s) + Cl2(g)  2 NaCl(s)
– With water to make base + hydrogen, e.g.
2 K(s) + 2 H2O(l)  2 K+(aq) + 2 OH–(aq) + H2(g)
• Reactions are more vigorous as you get lower in the group
(why?)
http://www.youtube.com/watch?v=9bAhCHedVB4&feature=rel
mfu
http://www.youtube.com/watch?v=rtNaEFXOdAc&feature=rel
mfu
Halogens
• They want to be –1!
• Easily reduced, high EA, high IE.
• Density increases moving down the group. (mass rises
faster than atomic radius)
• Reactions
– With metals to form metal halides, e.g.
2 Al(s) + 3 Cl2(g)  2 AlCl3(s)
– With hydrogen to form hydrogen halides (binary acids!), e.g.
H2(g) + I2(s)  2 HI(g)
– With other halogens to form interhalogen compounds, e.g.
Br2(l) + F2(g)  2 BrF(g)
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http://www.youtube.com/watch?v=F4IC_B9i4Sg
Noble Gases
• Closed-shell electron configuration; very unreactive!
• Used for lights, airtanks for divers, cryogens
• Few reactions! Fluorides, oxides can be made under severe
conditions.
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Helium--helios (sun)
Krypton--kryptos (hidden)
Neon--neos (new)
Xeno--xenos (stranger)