Download Alkali metals

Chapter 7
Periodic Properties
of the Elements
熊同銘
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Contents
1.
2.
3.
4.
5.
6.
7.
Development of the PeriodicTable
Effective Nuclear Charge
Sizes of Atoms and Ions
Ionization Energy
Electron Affinity
Metals, Nonmetals, and Metalloids
Trends for Group 1A and Group 2A
Metals
8. Trends for Selected Nonmetals
1. Development of the Periodic Table
 Mendeleev
•
•
Ordered elements by atomic mass and saw a repeating
pattern of properties
Put elements with similar properties in the same column
and used pattern to predict properties of undiscovered
elements such as eka-aluminum (Gallium) and eka-silicon
(Germanium).
 Modern periodic table
• Ordered elements by atomic number. The number
of protons was considered the basis for the periodic
property of elements.
• For example, the atomic weight of Ar is greater than
that of K.
• Discovering the elements:
2. Effective Nuclear Charge
 Effective nuclear charge (Zeff): The net positive
charge experienced by an electron in a manyelectron atom; this charge is not the full nuclear
charge because there is some shielding of the
nucleus by the other electrons in the atom.
• Electrons are both
attracted to the
nucleus and repelled
by other electrons.
 The simplest method for Estimating Zeff of the valence
electron
Zeff = Z – S
[7.1]
Zeff: effective nuclear charge
Z: actual nuclear charge (atomic number)
S: screening constant (numbers of core electrons)
(A)
Examples for representative elements:
22s22p63s1
Na:
1s
Zeff = 11 – 10 = 1
11
2
2
6
2
Zeff = 12 – 10 = 2
12Mg: 1s 2s 2p 3s
**For same period: Z increase Zeff increase
(B)
Examples for transition elements:
64s2
Fe:
[Ar]3d
Zeff = 26 – 24 = 2
26
7
2
Zeff = 27 – 25 = 2
27Co: [Ar]3d 4s
8
2
Zeff = 28 – 26 = 2
28Ni: [Ar]3d 4s
**For 1st transition elements, Zeff are about the same
 The advanced method for Estimating Zeff of the
valence electron
11Na
for example, because there is a small probability that the
3s electron is close to the nucleus, the value of S in
Equation 7.1 changes from 10 to 8.5. The actual value of Zeff
for the 3s electron in Na is Zeff = 2.5+.
 Variations in effective
nuclear charge for
period 2 and period 3
elements.
*Zeff increases
across a period.
V
V
3. Sizes of Atoms and Ions
 Atomic Radius: The
average radius of an
atom based on
measuring large
numbers of elements
and compounds.
 Trends in bonding atomic radii
Continued



*****
Within each group, bonding atomic radius tends to
increase from top to bottom because the principal
quantum number (n) of the electrons in the outermost
principal energy level increases.
Within each period, bonding atomic radius tends to
decrease from left to right because the effective
nuclear charge (Zeff) experienced by the electrons in the
outermost principal energy level increases.
The radii of transition elements vary slightly across each
row because the Zeff are about the same.
Sizes of Ions
 Determined by
interatomic distances
in ionic compounds
 Ionic size depends on
– the nuclear charge.
– the number of
electrons.
– the orbitals in which
electrons reside.
Continued


Cations are smaller than their parent atoms:
– The outermost electron is removed and
repulsions between electrons are reduced.
Anions are larger than their parent atoms:
– Electrons are added and repulsions
between electrons are increased.
Isoelectronic Series
• In an isoelectronic series, ions have the same
number of electrons.
• Ionic size decreases with an increasing nuclear
charge.
 An Isoelectronic Series (10 electrons)
• Note increasing nuclear charge with decreasing
ionic radius as atomic number increases
O2–
1.26 Å
F–
1.19 Å
Na+
1.16 Å
Mg2+
0.86 Å
Al3+
0.68 Å
4. Ionization Energy

*****
Ionization Energy (IE): The minimum energy
required to remove an electron from the ground state
of a gaseous atom or ion.
 First ionization energy (I1)
X(g) + I1  X+(g) + e–
 Second ionization energy (I2)
X+(g) + I2  X2+(g) + e–
•
•
•
•
Endothermic process
The easier to remove valence electron, the
lower IE
For an particular element: I1 < I2 < I3.....
Removing a core electron takes much more
energy than removing a valence electron

Trends in First Ionization Energy (IE1)
Continued



*****
I1 generally decreases down a group
because electrons in the outermost valence
electron are farther away from the nucleus.
I1 generally increases across a period
because electrons in the outermost valence
electron experience a greater effective
nuclear charge (Zeff).
The s- and p-block elements show a larger
range of values for I1. (The d-block generally
increases slowly across the period; the f-block
elements show only small variations.)
 Irregularities in the General Trend
Why actually the I1 of Be (2A) > I1 of B (3A)?
Ans: The 2p electron (of B) is at the higher energy level than
2s electron (of Be), thus it is relatively easy to be
removed, therefore, I1 of Be > I1 of B.
Be 
1s

2s
2p
Be+ 
1s

2s
2p
To ionize Be, you must break up a full sublevel, which
costs extra energy.
  
B+  
1s 2s
2p
1s 2s
2p
When you ionize B, you get a full sublevel, which costs
less energy.
B
Continued
Why actually the I1 of N (5A) > I1 of O (6A)?
Ans: Repulsion between paired electrons in a 2p orbital (of
oxygen), lead the electron relatively easy to be removed
therefore, IE1 of N > IE1 of O.
N

1s

2s
  
2p
N+ 
1s

2s
 
2p
To ionize N, you must break up a half-full sublevel,
which costs extra energy.
O 
1s
   
2s
2p
O+ 
1s
   
2s
2p
When you ionize O, you get a half-full sublevel, which
costs less energy.
*****
 Successive Values of Ionization Energies
•
•
Removal of each successive electron costs more
energy. due to having more protons than electrons.
When all valence electrons have been removed, it
takes a great deal more energy to remove the next
electron.
 Electron Configurations of Ions
(A) Anions examples:
Br(Z=35) [Ar]4s23d104p5
N(Z=7) [He]2s22p3
Br–: [Ar]4s23d104p6 or [Kr]
N3–: [He]2s22p6 or [Ne]
(B) Cations examples (2nd and 3rd period
metals):
Na(Z=11) [Ne]3s1
Al(Z=13) [Ne]3s23p1
*****
Na+: [Ne] or 1s22s22p6
Al3+: [Ne]
*****
(C) Cations examples (4th period and beyond
metals)
ELevel
[Ar]3d104s24p1
(a) Ga(Z=31):
Ga3+: [Ar]3d10
(b) Sn(Z=50): [Kr]4d105s25p2
Sn2+: [Kr]4d105s2
Sn4+: [Kr]4d10
(c) Cr(Z=24):[Ar]4s23d4
Cr2+:[Ar]3d4
Cr3+:[Ar]3d3
(As the (n-1)d
orbitals begin to fill
in the first transition
series)
5. Electron Affinities
*****
 Electron affinity (EA): The energy change
associated with the gaining of an electron by an
atom in its gaseous state.
M(g) + e– → M–(g) ΔE = EA
 The more negative
(exothermic) the
value, the larger
the EA.
 Trend of EA is
much more
irregular than IE
*****
 Trends in Electron Affinity
 Alkali metals decrease electron affinity down
the column.
 “Generally” increases across period
(becomes more negative from left to right).
 Group 5A generally has lower EA than
expected because extra electron must pair.
 Groups 2A and 8A generally have very low EA
(many of these elements is positive) because
added electron goes into higher energy level or
sublevel.
 Halogen have highest EA in any period.
6. Metal, Nonmetals, and Metalloids
Metals Differ from Nonmetals
• Metals tend to form cations.
• Nonmetals tend to form anions.
Metals
Properties of metals:
Shiny luster
Conduct heat and electricity
Malleable and ductile
Solids at room temperature (except mercury)
Low ionization energies/form cations easily
Metal Chemistry
• Compounds formed between metals and
nonmetals tend to be ionic.
• Metal oxides tend to be basic.
Nonmetals
• Nonmetals are found on the right hand side of the
periodic table.
• Properties of nonmetals include the following:
Solid, liquid, or gas (depends on element)
• Solids are dull, brittle, poor conductors
• Large negative electronegativity/form anions
readily
Nonmetal Chemistry
• Substances containing only nonmetals are
molecular compounds.
• Most nonmetal oxides are acidic.
Recap of a Comparison of the
Properties of Metals and Nonmetals
Metalloids
• Metalloids have some characteristics of metals
and some of nonmetals.
• Several metalloids are electrical semiconductors
(computer chips).
Group Trends
• Elements in a group have similar properties.
• Trends also exist within groups.
• Groups Compared:
 Group 1A: The Alkali Metals
 Group 2A: The Alkaline Earth Metals
 Group 6A: The Oxygen Group
 Group 7A: The Halogens
 Group 8A: The Noble Gases
Alkali Metals
• Alkali metals are soft,
metallic solids.
• They are found only in
compounds in nature,
not in their elemental
forms.
• Typical metallic
properties (luster,
conductivity) are seen
in them.
Alkali Metal Properties
• They have low densities and melting points.
• They also have low ionization energies.
Alkali Metal Chemistry
Their reactions with water are famously exothermic.
Differences in Alkali
Metal Chemistry
• Lithium reacts with oxygen to make an oxide:
4 Li + O2  2 Li2O
• Sodium reacts with oxygen to form a peroxide:
2 Na + O2  Na2O2
• K, Rb, and Cs also form superoxides:
M + O2  MO2
*****
Flame Tests
• Qualitative tests for alkali metals include
their characteristic colors in flames.
Alkaline Earth Metals—Compare to
Alkali Metals
• Alkaline earth metals have higher densities
and melting points than alkali metals.
• Their ionization energies are low, but not as
low as those of alkali metals.
Alkaline Earth Metals
• Beryllium does not
react with water, and
magnesium reacts only
with steam, but the
other alkaline earth
metals react readily
with water.
• Reactivity tends to
increase as you go
down the group.
Group 6A—Increasing in Metallic
Character down the Group
• Oxygen, sulfur, and selenium are nonmetals.
• Tellurium is a metalloid.
• The radioactive polonium is a metal.
Group 7A—Halogens
• The halogens are typical nonmetals.
• They have highly negative electron affinities, so
they exist as anions in nature.
• They react directly with metals to form metal
halides.
Group 8A—Noble Gases
• The noble gases have very large ionization energies.
• Their electron affinities are positive (can’t form stable
anions).
• Therefore, they are relatively unreactive.
• They are found as monatomic gases.
End of Chapter 07