Download Chapter 7. Periodic Properties of the Elements.

Chapter 7. Periodic Properties
of the Elements.
7.1 Development of The Periodic Table.
The number of known
elements in 1800 was 31, but
had increased to 63 by 1865.
In 1869 Mendeleev and Lothar
Meyer published almost
identical classification
schemes. Both noticed that
similar chemical properties
recurred when elements were
placed in order of increasing
atomic weight. Mendeleev
advanced his ideas more
vigorously, and had the
insight to leave blank spaces
and predict the existence of
yet undiscovered elements.
Dimitri Mendeleev
(1834-1907)
Group 4b
Group 5b
Group 6b
Missing
Element
=eka-silicon
Group 1b
Group 2b
Group 3
Group 4
Group 5
Group 6
Group 7
Group 1
Group 2
Thus, Mendeleev predicted the existence of what he
called eka-silicon, which we today refer to as
Germanium.
Mendeleev predictions.
At. Wt.
Density:
Sp. Heat:
M.Pt.
Color
Oxide:
Density oxide:
Chloride:
B.Pt. chloride:
72
5.5
0.305
high
dark gray
MO2
4.7
XCl4
<100
Observed props. of Ge
72.59 g/mol
5.35 g/cc
0.309 J/g-K
947 oC
grayish white
GeO2
4.70
GeCl4
84 oC
The shells and sub-shells being filled across
the periodic table:
Because they have the same valence electron
configuration, elements in the same group have
similar chemistry:
Appearance:
Group I A
Li
Na
Group II A
Be
Mg
All the elements in group 1 react violently with water:
2 M (s)
+
2 H2O (l)  2 MOH (aq) + H2(g)
Reactivity:
Sodium in water
Potassium in water
Atomic Numbers (Z):
Henry Moseley (1887-1915)
developed the concept of
atomic numbers. Each element
produces X-rays of a distinct
frequency, which increased
with atomic number. Atomic
number = number of protons
in the nucleus. The concept of
At. No. allowed correct
placement of e.g. Ar (Z = 18)
before K (Z = 19) even though
At. Wt. Ar > K. A similar
problem occurs with Co (At.
Wt = 58.9) and Ni (58.7), even
though Z = 27 for Co and 28
for Ni.
Henry Moseley
(1887-1915)
The K x-ray spectrum of germanium:
Energy (keV) (kilo-electron-volts)
Relationship between the frequency of the K x-ray
peak and atomic number of the element (Z)
Kr
Z
K
Al
Ar
determining Z from energy
of x-ray peak now places
Ar and K in the right order
The core electrons and the valence
electrons:
Nucleus = 12+
valence
electrons
in 3s shell
experience
small net
positive
charge
the core
electrons shield
the valence
electrons from
the nucleus
The Mg
atom
e
e
core
Electrons
= Ne 10 e-
7.2 Effective Nuclear Charge.
In a many-electron atom, each electron is
simultaneously attracted to the nucleus, and repelled
by the other electrons. Generally, the valence
electrons are screened from the nucleus by the core
electrons (~S, where S is the screening constant =
number of core electrons). The effective nuclear charge
is the approximate charge experienced by the valence
electrons. Thus, the effective nuclear charge Zeff,
experienced by the valence electrons is given by:
Zeff
=
Z
-
S
number of
core electrons
Thus, for Na, Zeff = 11 – 10 ~ 1+. However, because
electrons in higher shells are less good at shielding
the nucleus, the charge experience by Na+ is actually
about +2.5.
Which would you expect to experience a greater Zeff, a 2p electron
of an F atom or a 3s electron of an Na atom?
Zeff (F, 2p) =
+9 – 2 = 7+
much more tightly bound !
Zeff (Na, 3s) = +11 – 10 = 1+
Which electron do you think would be easier to remove?
The Na 3s electron is easily removed => metals easily lose electrons
The F 2p electron is NOT easily removed => it is much harder to remove an
electron from a halogen
7.3 Sizes of Atoms and Ions.
We can picture atoms as having size in
various ways. The bonding atomic radius, or
covalent radius is half the distance in
homonuclear diatomic molecules, e.g. H2,
Cl2. The C-C distance in diamond gives us
the covalent radius of C as 1.54/2 = 0.77 Å.
Bond lengths are given quite closely as the
sum of covalent radii, e.g. C-Cl = 0.77 + 0.99
= 1.76 Å.
The non-bonding atomic radius = van der
Waals radius.
Atomic non-bonding (van der Waals) and
bonding (covalent) radii of atoms:
Non-bonding atomic radii of I and H atoms
The bonding atomic radius
(‘covalent radius’) of I is ½
the I-I distance in I2. The H-I
internuclear distance in HI
is the sum of the bonding
atomic radii of H and I.
iodine atom
2 x I-I bonded
Atomic radius
I2 molecule
hydrogen atom
2 x H-H bonded
Atomic radius
H2 molecule
HI molecule
Ionic radii:
The ionic radii are the radii of the ions. For cations the ionic radius is
smaller than the atomic radius, because there are fewer electrons
around the nucleus, while for anions the ionic radii are larger than the
atomic radii because there are now more electrons to repel each
other:
Lithium atom
radius = 1.34 Å
fluorine atom
radius = 0.71 Å
Li
Li+
F
F-
lithium cation
ionic radius =
0.68 Å
fluorine anion
ionic radius =
1.33 Å
The sizes of atoms (gray) cations (pink) and anions (blue)
(textbook p. 263):
Periodic trends in Atomic and
ionic radii.
1) Atomic and ionic radius increases as
we move down a group.
2) Atomic and ionic radius tends to
decrease along a row from left to right
Variation in atomic radius with Z (atomic number)
Rb
Cs
K
Na
Li
Xe
Ar
Kr
Ne
He
Atomic number (Z)
Rn
Periodic trends in ionic radii.
Ionic radii are based on the distance
between ions in ionic compounds. In
general, cations are smaller than the
atoms from which they are derived, but
anions are larger.
O21.40
F1.33
Na+
0.97
Mg2+
0.66
Al3+
0.51Å
The sizes of atoms (gray) cations (pink) and anions (blue):
Ionization energy.
The ionization energy is the minimum
energy required to remove an electron
from the ground state of the isolated
gaseous atom or ion. E.g. First (I1)and
second (I2) ionization energies of Na+
are given by ΔH for the processes:
Na(g)
→ Na+(g)
+
eNa+(g)
→ Na2+(g) +
e-
Trends in 1st Ionization Energies
Variations in successive ionization energies:
The value of In increases as n increases for any one
element. It increases sharply as we start to remove
inner-shell electrons. Core (inner-shell) electrons are
much harder to remove that valence electrons:
Why does the ionization energy increase sharply after removal
of the 5th electron for Nitrogen?
What would a similar graph look like for beryllium (Be)?
core electrons are much harder to remove than valence electrons !
Nitrogen
Beryllium
Periodic Trends in First
Ionization Energies:
If we look at how I1 varies across the periodic
table, we find that:
1. I1 increases across a row with increasing Z.
2. I1 decreases down a group, e.g.
Li>Na>K>Rb>Cs, or He>Ne>Ar>Kr>Xe.
3. The representative elements (‘main group
elements’) show a larger variation in I1 than do
transition elements. The f-block elements
show only a small variation in I1.
Variation in first ionization energy with Z along the
periodic table:
He
Ne
Noble gases occur at the peaks
Ar
Kr
Xe
Rn
Li Na K
Rb
Cs
Atomic number (Z)
Alkali metals occur
at the troughs
Electron Configuration of Ions:
When electrons are lost from an atom
to form a cation, they are always lost
from the shell having the highest
principal quantum number, i.e. the
highest energy.
e.g. Li(1s22s1)
loses
electron
→
Li+ (1s2)
Electronic configurations of
transition metals (the d-block)
For transition metal ions one finds that
in e.g. Fe2+, the two electrons lost are
lost from the 4s2 level, even though in
the neutral atom this level is lower in
energy than the 3d level.
Fe ([Ar]3d64s2) → Fe2+([Ar]3d6).
Addition of electrons:
When an electron is added, it goes into
the empty or partially filled orbital
having the lowest value of n.
F (1s22s22p5)
→
F-(1s22s22p6)
7.5 Electron affinities:
This is the energy for adding electrons to the atom.
Cl(g) +
e-
→
Cl-(g) ΔE = -349 kJ/mol
The halogens have the most negative electron
affinities, followed by the chalcogens. Group 2 has
very small EA’s, because they have a filled s orbital.
Similarly, group 5 has small EA’s because they have
a half-filled p subshell, which is particularly stable.
For the noble gases, EA’s are all positive,
corresponding to the lack of chemistry, except for
Xe.
Electron affinities (kJ/mol) for some elements:
7.6 Metals, non-metals, and
metalloids.
Metals are lustrous (shiny) and malleable
(can be pounded into thin sheets) and ductile
(can be drawn into wires), and conduct heat
and electricity. All are solids at room
temperature except Hg, and Cs and Ga melt
at slightly above room temperature.
Compounds of metals with non-metals tend
to be ionic substances. Most metal oxides
are basic, i.e. dissolve in water to give basic
solutions.
Na2O + H2O = 2 NaOH
Non-metals.
These tend not to be lustrous, and
usually are poor conductors of heat
and electricity. Generally low-melting,
although diamond melts at 3570 oC.
Because of their electron affinities,
non-metals tend to gain electrons when
they react with metals, e.g.
2 Mg(s)
+
O2(g) → 2 MgO(s)
Compounds composed entirely of nonmetals tend to be molecular substances.
Most non-metal oxides dissolve in water to
form acids:
SO3(s) + H2O(l)
→
H2SO4(aq)
CO2(g) + H2O(l)
→
H2CO3(aq)
acid + base =
H2SO4 + CaO =
salt +
CaSO4 +
water
H2O
7.7. Trends for the active metals.
Students read 7.7 through 7.9 on own
time.