Periodic Relationships
Among the Elements
Chapter 8
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How do we arrange and categorize
atoms?
19th Century – Did not know about protons,
neutrons and electrons
 Try organizing elements by their mass
1864 John Newlands noticed when atoms
are arranged by mass we see similar
properties every 8 elements
 Law of Octaves
1869 Dmitri Mendeleev & Lother Meyer independently
proposed more extensive tabulation based on regular,
periodic recurrence of properties
Grouped elements more accurately
Predicted that other elements should exist
Mendeleev’s periodic table had 66 known elements
By 1900 30 more elements were added
When the Elements Were Discovered
8.1
What happens if we arrange elements solely by
atomic mass?
Ar = 39.95 amu
>
K = 39.10 amu
Co = 58.93 amu
>
Ni = 58.69 amu
Te = 127.6 amu
>
I = 126.9 amu
Predict K is a noble gas and Ar is an alkali metal !!
1913 Henry Moseley proposed the idea of an atomic
number and predicted the frequency (ν) of x-rays emitted
from elements based on atomic number (Z)
 ν = a (Z-b)
where a & b are constants (same for all elements)
Atomic number = number of protons =
number of electrons (when neutral)
=> Electron configuration predicts chemical
properties
=> Look at outermost ground-state electron
configurations
=> valence electrons
Li
1s22s1
s1
Na 1s22s12p63s1
Electron involved in chemical bonding
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
4f
5f
8.2
Categorize elements by the type of subshell being filled
 Representative elements




 main group elements
 Groups 1A  7A
 incompletely filled s & p orbitals
Noble gases
 Group 8A
 filled s & p orbitals
Transition metals
 Group 1B, 3B  8B
 incompletely filled d orbitals
Group 2B - Zn, Cd, Hg
 filled d orbitals
 do not fit transition metals or representative elements
 no special name
Lanthanides & Actinides
 incompletely filled f orbitals
Classification of the Elements
8.2
Electron Configurations of Ions
Cations - Positive charge => Remove electrons
Li
1s22s1
Li+ 1s2
3 electrons
2 electrons
Anions - Negative charge => Add electrons
1s22s12p4
8 electrons
O2- 1s22s12p6
10 electrons
O
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. Inner orbitals shield outer orbitals.
1s orbital will shield 2s orbital, but 2s orbital will not shield 1s orbital.
Zeff = Z - s
0 < s < Z (s = shielding constant)
How many protons will the electron feel?
Zeff  Z – number of inner or core electrons
Core Zeff
Radius
Z
Na
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.
Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
8.3
Atomic radius = ½ distance between two nuclei of
adjacent atoms
2r
2r
Non-Metals that occur as
diatomic molecules
Metals
8.3
Increasing shell number
Increasing nuclear charge
Increasing effective charge
Atomic Radii
8.3
What about the size of ions ?
Cation is always smaller than atom from
which it is formed. Generally empty all valence electrons.
Anion is always larger than atom from
which it is formed. Add electrons to fill orbital.
Get more repulsive forces between electrons.
Shift
smaller
Shift
larger
8.3
NonMetals
Metals
8.3
Energy of orbitals in a multi-electron atom
Energy depends on n and l
E=0
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
It gets harder and
harder to remove
electrons !!
Where do we
see the
highest
ionization
energies?
8.4
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
Trend to smaller I1 due to increasing distance from nucleus
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
Reverse the sign
8.5
8.5
Orbitals
that are
almost
filled want
an electron
Noble
gases do
not want to
accept an
electron
Charge
density is
similar along
the diagonal
=> expect
similar
properties
Happens most with lighter elements.
8.6
Hydrogen (1s1)
Where should we put it in the periodic table?
- Hydrogen is in a class by itself
- has 1 electron is s orbital like alkali metals
=> can form H+
can also gain an electron to form H=> hydride ion
=> resembles halogens
Ionic hydrides react with water to produce H2 gas
+ metal hydroxides
NaH (s) + H2O (l)
CaH2 (s) + 2 H2O (l)
NaOH (aq) + H2 (g)
Ca(OH)2 (aq) + 2 H2 (g)
Hydrogen Displacement
Hydrogen can also burn in air
2 H2 (g) + O2 (g)
2 H2O (l)
Group 1A Elements (ns1, n  2)
Low first ionization energies
M+1 + 1e-
2M(s) + 2H2O(l)
4M(s) + O2(g)
2MOH(aq) + H2(g)
2M2O(s)
Metals lose their shiny
appearance to form oxides.
Increasing reactivity
M
8.6
4 Li (s) + O2 (g)
2 Li2O (s)
Some alkali metals can form peroxides
(O22- ion)
2 Na (s) + O2 (g)
Na2O2 (s)
K, Rb & Cs can form superoxides (O2- ion)
K (s) + O2 (g)
KO2 (s)
As we move down the periodic table we
can for other oxide forms.
Group 2A Elements (ns2, n  2)
Less reactive than alkali metals.
M+2 + 2e-
Be(s) + 2H2O(l)
Mg(s) + 2H2O(g)
M(s) + 2H2O(l)
cold
Form divalent ions
No Reaction
Mg(OH)2(aq) + H2(g) Reacts slowly with steam
M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba Much
more
reactive
Increasing reactivity
M
Increasing metallic character as you go
down the group.
BeH2, BeCl2, and MgH2 are more molecular
than ionic in nature
Strontium - 90 is radioactive
- Major product in atomic bomb explosions
- Since Sr2+ and Ca2+ are chemically similar, Sr2+
can replace Ca2+ in our bones
=> high energy radiation
=> can cause anemia, leukemia and other
chronic illnesses
Group 3A Elements (ns2np1, n  2)
Boron is a metalloid and the rest are metals
Boron does not react with O2 (g) & water
4Al(s) + 3O2(g)
2Al2O3(s) Forms unreactive oxide coating
Increasing metallic
character
2Al(s) + 6H+(aq)
2Al3+(aq) + 3H2(g)
Al only forms 3+ ions
Other 3A elements form + & 3+ ions
8.6
Group 4A Elements (ns2np2, n  2)
C, Si & Ge do not form ionic compounds
Sn2+(aq) + H2 (g)
Pb(s) + 2H+(aq)
Pb2+(aq) + H2 (g)
Increasing metallic
character
Sn(s) + 2H+(aq)
Sn & Pb will not
react with water, but
will react with HCl
non-metal
metalloids
metals
8.6
Group 4A elements form compounds with
2+ & 4+ oxidation states
4+ is generally the most stable
CO2 is more stable than CO
SiO2 is very stable and SiO does
not exist under normal conditions
For Pb the stability trend is reversed
2+ is more stable than 4+
Pb [Xe] 6s24f145d106p2
Lose only 6p2 electrons
=> Pb2+ [Xe] 6s24f145d10
Group 5A Elements (ns2np3, n  2)
Larger variation in properties
Increasing metallic
character
Nitrogen
forms diatomic gas - N2
forms oxides - NO, N2O, NO2, N2O4 & N2O5 => NOx
all oxides are gases except N2O5 (s)
accepts 3 electrons to become isoelectric with Ne
compounds w/ metals are ionic – Li3N & Mg3N2
non-metals
metalloids
metal
Phosphorus
 exists as P4 molecules
forms solid oxides as P4O6 & P4O10
N & P form important oxoacids
N2O5(s) + H2O(l)
2 HNO3(aq)
P4O10(s) + 6 H2O(l)
Arsenic
Antimony
Bismuth
4 H3PO4(aq)
Form
extensive
3-d
networks
Group 6A Elements (ns2np4, n  2)
No metals in this group
Oxygen form diatomic gas – O2
Sulfur forms molecular S8
Selenium forms molecular Se8
Tellurium & Polonium – form extensive 3-d networks
Polonium is radioactive
non-metals
metalloids
8.6
Oxygen tends to form O2- ion in many compounds
S2-, Se2-, and Te2- ions are common
S forms important oxides – SO2, SO3 => SOx
SO3(g) + H2O(l)
H2SO4(aq)
Group 7A Elements (ns2np5, n  2)
All halogens are non-metals => X2
=> Want to gain 1 electron => halide ions
X-1
=> acids
2HX(g)
Halogens are very reactive 2 F2 (g) + H2O (l)
4 HF (aq) + O2 (g)
Increasing reactivity
X + 1eX2(g) + H2(g)
8.6
Group 8A Elements (ns2np6, n  2)
Completely filled ns and np subshells.
Highest ionization energy of all elements.
No tendency to accept extra electrons.
Occur as monatomic gases.
Also called inert gases.
8.6
1963 – Neil Bartlett made the first Xe compounds
Xe (g) + PtF6 (g)
XePtF6 (s)
Since then XeF4, XeO3, XeO4, XeOF4, KrF2
have been prepared.
Properties of Oxides Across a Period
basic
acidic
8.6