Transcript No Slide Title

Periodic Relationships Among
the Elements
Chapter 8
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
1. Explain the basis of the periodic table as described by
Mendeleev and Meyer and indicate the shortcomings of
their
method.
2. Explain the basis of the periodic table as described by
Moseley
and how it predicted properties of “missing” elements.
3. Identify elements that correspond to each of the
following
groups:
·
* representative elements
·
* noble gases
·
* transition metals
·
* lanthanides
·
* actinides
4. Describe the electron configuration of cations and
5. Apply the concept of effective nuclear charge and
shielding
constants (screening constants) to justify why the first
ionization energy is always smaller than the second
ionization
energy of a given atom.
6.
Predict the trends from left to right and top to bottom
of the
periodic table for each of the following:
·
atomic radius
·
ionic radius
·
ionization energy
·
electron affinity
·
metallic character
7.
Relate why hydrogen could be placed in a class by
itself
when reviewing its chemical properties.
8.
Provide examples of Group 1A elements reacting with
11.
Relate how strontium-90 could lead to human illness.
12.
Compare the reactivity of boron, a metalloid, to
aluminum.
13.
4A.
Identify the metals, nonmetal, and metalloids of Group
14.
Recall the reactions that form nitric acid, phosphoric
acid and
sulfuric acid.
15.
List the halides (halogens)
16.
Indicate the three hydrohalic acids that are strong acid
and the
one hydrohalic acid that is a weak acid.
17.
Explain why the name for Group 8A has changed from
19.
Rationalize the characteristics of the properties of
oxides of
the third period elements.
20.
Classify oxides as acidic, basic, or amphoteric.
21. Explain why concentrated bases such as NaOH
should not
be stored in glass containers
8.1
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
4f
5f
8.2
8.2
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
Atoms gain electrons
so that anion has a
noble-gas outer
electron configuration.
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
8.2
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
8.2
Na+: [Ne]
Al3+: [Ne]
O2-: 1s22s22p6 or [Ne]
F-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He
8.2
Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal,
electrons are always removed first from the ns orbital and
then from the (n – 1)d orbitals.
Fe:
[Ar]4s23d6
Fe2+: [Ar]4s03d6 or [Ar]3d6
Mn:
[Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5
8.2
Effective nuclear charge (Zeff) is the “positive charge” felt
by an electron.
Zeff = Z - s
0 < s < Z (shielding constant)
Zeff  Z – number of inner or core electrons
Na
Z
Core
Zeff
Radius
11
10
1
186
Mg
12
10
2
160
Al
13
10
3
143
Si
14
10
4
132
For a Period
as Zeff increases
radius decreases
8.3
8.3
Atomic Radii
8.3
8.3
Cation is always smaller than atom from
which it is formed.
Anion is always larger than atom from
which it is formed.
8.3
8.3
Ionization energy is the minimum energy (kJ/mol) required
to remove an electron from a gaseous atom in its ground
state.
I1 + X (g)
X+(g) + e-
I1 first ionization energy
I2 + X (g)
X2+(g) + e-
I2 second ionization energy
I3 + X (g)
X3+(g) + e-
I3 third ionization energy
I1 < I2 < I3
8.4
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
8.4
General Trend in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
8.4
Electron affinity is the negative of the energy change that
occurs when an electron is accepted by an atom in the
gaseous state to form an anion.
X (g) + e-
X-(g)
F (g) + e-
X-(g)
DH = -328 kJ/mol
EA = +328 kJ/mol
O (g) + e-
O-(g)
DH = -141 kJ/mol
EA = +141 kJ/mol
8.5
8.5
8.6
Group 1A Elements (ns1, n  2)
M+1 + 1e-
2M(s) + 2H2O(l)
2M(s) + O2(g)
2MOH(aq) + H2(g)
M2O(s)
Increasing reactivity
M
8.6
Group 2A Elements (ns2, n  2)
M+2 + 2e-
Be(s) + 2H2O(l)
Mg(s) + 2H2O(g)
M(s) + 2H2O(l)
No Reaction
Mg(OH)2(aq) + H2(g)
M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba
Increasing reactivity
M
8.6
Group 3A Elements (ns2np1, n  2)
4Al(s) + 3O2(g)
2Al(s) + 6H+(aq)
2Al2O3(s)
2Al3+(aq) + 3H2(g)
8.6
Group 7A Elements (ns1np5, n  2)
X2(g) + H2(g)
X-1
2HX(g)
Increasing reactivity
X + 1e-
8.6