7.1.
Introduction to the s block elements
The position of Group 1 Alkali Metals and Group 2 Alkaline Earth Metals
in the Periodic Table
Pd |
s block elements |
d blocks and f blocks of metallic
elements |
p block elements |
Gp1 |
Gp2 |
Gp3/13 |
Gp4/14 |
Gp5/15 |
Gp6/16 |
Gp7/17 |
Gp0/18 |
1 |
1H
|
2He |
2 |
3Li
lithium |
4Be
beryllium |
The modern Periodic Table of Elements
ZSymbol, z = atomic or proton
number
highlighting position of
Group
1
and
Group
2
elements
outer electrons: ns1
and ns2 |
5B |
6C |
7N |
8O |
9F |
10Ne |
3 |
11Na
sodium |
12Mg
magnesium |
13Al |
14Si |
15P |
16S |
17Cl |
18Ar |
4 |
19K
potassium |
20Ca
calcium |
21Sc |
22Ti |
23V |
24Cr |
25Mn |
26Fe |
27Co |
28Ni |
29Cu |
30Zn |
31Ga |
32Ge |
33As |
34Se |
35Br |
36Kr |
5 |
37Rb
rubidium |
38Sr
strontium |
39Y |
40Zr |
41Nb |
42Mo |
43Tc |
44Ru |
45Rh |
46Pd |
47Ag |
48Cd |
49In |
50Sn |
51Sb |
52Te |
53I |
54Xe |
6 |
55Cs
caesium |
56Ba
barium |
57-71 |
72Hf |
73Ta |
74W |
75Re |
76Os |
77Ir |
78Pt |
79Au |
80Hg |
81Tl |
82Pb |
83Bi |
84Po |
85At |
86Rn |
7 |
87Fr
francium |
88Ra
radium |
89-103 |
104Rf |
105Db |
106Sg |
107Bh |
108Hs |
109Mt |
110Ds |
111Rg |
112Cn |
113Nh |
114Fl |
115Mc |
116Lv |
117Ts |
118Og |
outer electrons: Group 1 ns1
and Group 2 ns2
-
The
first two vertical columns of the Periodic Table, i.e. Groups 1 and 2, are called the
s–block metals, because they only have 1 or 2 electrons in their outer shell.
-
These outer electrons are of an s–orbital type (s sub–shell or sub–quantum
level) and the chemistry of the metals, with their relatively low ionisation
energies, is dominated by the loss of these s electrons
to
form a cation and also accounts for their generally high chemical reactivity ...
-
the outer s1 electron loss
by the Group 1 Alkali Metals gives the M+
ion, and,
-
the outer s2 electrons lost by the Group 2 Alkaline Earth
Metals forms the M2+
ion,
-
and in each case the
cation has a residual very stable noble gas core of electrons.
-
The only chemically
stable
oxidation states are +1 for Group 1 metals and +2 for Group 2 elements,
governed by the relative ease of loss of the outer s electrons and
the subsequent very high ionisation energies required to remove a
2nd (for group 1) or 3rd electron (for group 2) from the inner noble
gas core of electrons left (i.e. from the next principal quantum
level or shell).
-
The relative ease of
delocalising the outer 1/2 electrons in the metal lattice makes them good conductors of
heat and electricity (bonding
model for metals).
-
The low ionisation
energies and low electronegativity means that when combined with
non–metals, most compounds of the Group 1–2 elements tend to be
ionic in nature.
-
Group 1and Group 2 ions
are their compounds are important in the natural world of living
systems and geology.
-
Calcium carbonate and
phosphate minerals are important components of skeletons, teeth as
well as bone!
-
For marine organisms the
skeletal remains form sedimentary rocks like limestone and chalk
which mainly consist of calcium carbonate and magnesium carbonate.
-
In biochemistry, at the
heart of the chlorophyll molecules involved with photosynthesis is
magnesium ion (Mg2+).
-
Sodium and potassium
ions (Na+ and K+) are important components of
a balanced electrolyte solutions in living systems and in nerve
impulse transmission systems.
7.2.
Information and Data
Table GROUP 1 ALKALI METALS
(from left to right is down the group!)
property\Z
symbol, name |
3Li
Lithium |
11Na
Sodium |
19K
Potassium |
37Rb
Rubidium |
55Cs
Caesium |
87Fr
Francium |
melting
point/oC |
181 |
98 |
64 |
39 |
29 |
27 |
boiling
point/oC |
1347 |
883 |
774 |
688 |
679 |
677 |
density/gcm–3 |
0.53 |
0.97 |
0.86 |
1.48 |
1.87 |
>1.87 |
1st
IE/kJmol–1 |
513 |
496 |
419 |
403 |
376 |
400 |
2nd
IE/kJmol–1 |
7298 |
4562 |
3051 |
2632 |
2420 |
2100 |
atomic metallic
radius/pm |
152 |
186 |
231 |
244 |
262 |
270 |
M+
ionic radius/pm |
78 |
98 |
133 |
149 |
165 |
180 |
electronegativity |
0.98 |
0.93 |
0.82 |
0.82 |
0.79 |
0.70 |
simple electron
configuration |
2,1 |
2,8,1 |
2,8,8,1 |
2,8,18,8,1 |
2,8,18,18,8,1 |
2,8,18,32,18,8,1 |
electron
configuration |
[He]2s1 |
[Ne]3s1 |
[Ar]4s1 |
[Kr]5s1 |
[Xe]6s1 |
[Rn]7s1 |
Electrode
potential M/M+ |
–3.04V |
–2.71V |
–2.92V |
–2.92V |
–2.92V |
–2.92V |
Symbol
– flame colour |
Li
– red/crimson |
Na
– yellow |
K
– lilac/purple |
Rb
– red |
Cs
– blue |
Fr
– na |
Note the outer electron
configurations (n = principal quantum number = period number)
Group 1 is
ns1
and
Group 2 is
ns2
-
Typical metals in some ways
e.g. silvery grey lustrous solids*, very good conductors of heat and electricity,
relatively high boiling points.
-
*
When freshly cut they are quite shiny,
but they rapidly tarnish by reaction with oxygen to form an oxide layer,
which is why they are stored under oil. Ask your teacher to show you the
bottle from which we extract the lumps for doing our demonstrations –
not a pretty sight if its a bit old!
-
The 1st ionisation energies
are the lowest of any group of elements, but note the jump up to a very
high 2nd ionisation energy.
-
Untypical in other ways e.g. relatively soft, low density (Li, Na and K float on water before reacting ...), and
very low
melting points.
-
The reasons for the low melting
points, densities and physical hardness lies in looking at the
bonding model for metals.
-
The more electrons that can be
delocalised and the closer the atoms (actually ions in a metallic lattice)
can approach each other the stronger the bond, thereby raising melting
points, densities and tensile strength.
-
However, Group 1 metals can only
release one electron per atom to contribute towards the pool of bonding
delocalised electrons. Group 2, 3 and transition metals etc. can contribute
more electrons for metallic bonding. This why the Group 2 elements have
higher melting points, densities and are harder metals.
-
Also, being the first element on
a given period, they have the highest relative atomic/ionic radii because
you have the minimum nuclear attractive force on the outer electrons.
-
Any metal
flame colour is due
to electronic transitions in the atom or cation.
-
Electrons are promoted to
higher quantum levels via collisions of the high thermal kinetic energy
particles in the hot flame. When the promoted electron 'relaxes' or
'falls' back to its more stable electronic level, energy is emitted
(this is the basis of an emission spectrum). If the
frequency/wavelength/energy of the photons emitted is in the visible
region of the electromagnetic spectrum, a 'flame colour' results e.g. as
observed in fireworks.
-
The
set of quantum levels and associated energies are unique for each atom
which means the
quantum level differences varies from atom to atom, therefore the
frequency/energy of emitted photons is different, hence you see different flame
colours in the visible region of light from each Group 1/2 metal – each
element has a 'finger print' emission spectrum.
-
The s–block groups 1/2 have
the lowest ionisation energies of all the elements in the periodic
table, so the high temperature of the flame (>1000oC) means
that the kinetic energy of the flame particles is quite sufficient to
promote electrons to a higher quantum level when the metal atoms collide
with other high KE atoms or molecules.
-
Planck's Equation:
ΔE = hν, where
-
ΔE = E2 –
E1, the energy difference between e.g. the outer s level
E1 and a higher level E2,
-
h = Planck's
constant and ν = frequency of light of the emitted photons.
-
For more examples–details
see electronic structure and
emission/absorption spectra notes.
-
Oxidation state or oxidation
number is always +1
in Group 1 Alkali Metal compounds.
-
Only
the single outer s–electron is easily
lost, the 2nd, and subsequent ionisation energies are far too high to
form chemically stable cations of 2+ etc. i.e. the energy required will
not be compensated by ionic bond formation.
-
The stable Group 1 cation has the electron
configuration of a noble gas,
-
e.g. the sodium atom, Na, is
2,8,1 or
1s22s22p63s1 or [Ne]3s1
-
so the sodium ion, Na+, is
2,8 or
1s22s22p6 or [Ne]
-
See also section 4. on Group
trends and comparison with Group 2 metals.
PLEASE
NOTE that Francium is highly radioactive and therefore difficult and
dangerous to study BUT all its known physical and chemical properties fit in
with it being at the foot of Group 1 and other properties could be inferred
from the properties and group trends of Li to Cs.
TOP OF PAGE and
sub-index
7.3.
Information and Data
Table GROUP 2 ALKALINE EARTH METALS
(from left to right is down the group!)
property\Zsymbol,
name |
4Be
Beryllium |
12Mg
Magnesium |
20Ca
Calcium |
38Sr
Strontium |
56Ba
Barium |
88Ra
Radium |
melting
pt./oC |
1278 |
649 |
839 |
769 |
729 |
700 |
boiling
pt./oC |
2487 |
1090 |
1484 |
1384 |
1637 |
1140 |
density/
gcm–3 |
1.85 |
1.74 |
1.55 |
2.54 |
3.51 |
5.0 |
1st
IE/ kJmol–1 |
900 |
738 |
590 |
550 |
503 |
509? |
2nd
IE/kJmol–1 |
1757 |
1451 |
1145 |
1064 |
965 |
979 |
3rd
IE/kJmol–1 |
14848 |
7733 |
4910 |
4210 |
3600 |
3300 |
atomic
radius/ pm |
111 |
160 |
197 |
215 |
217 |
223 |
M2+
ionic radius/ pm |
34 |
78 |
106 |
127 |
143 |
152 |
electronegativity |
1.57 |
1.31 |
1.00 |
0.95 |
0.89 |
0.89 |
electron
config. |
2,2 |
2,8,2 |
2,8,8,2 |
2,8,18,8,2 |
2,8,18,18,8,2 |
2,8,18,32,18,8,2 |
electron
config. |
[He]2s2 |
[Ne]3s2 |
[Ar]4s2 |
[Kr]5s2 |
[Xe]6s2 |
[Rn]7s2 |
Electrode
potential M/M2+ |
–1.97V |
–2.36V |
–2.84V |
–2.89V |
–2.92V |
–2.92V |
Symbol
– flame colour (see
chemical tests
new window) |
Be
– na |
Mg
– na |
Ca
– brick red |
Sr
– crimson |
Ba
– apple green |
Ra
– na |
Note the outer electron
configurations (n = principal quantum number = period number)
Group 1 is
ns1
and
Group 2 is
ns2
-
Very typical metals, silvery grey
lustrous solids, relatively high melting and boiling points, good conductors
of heat and electricity.
-
Compared to adjacent Group 1
metal on same period:
-
The melting and boiling points
are higher, and they are harder, stronger and more dense than the adjacent Group 1
metal on the same period. This is
because their are two delocalised electrons per ion in the crystal lattice
giving an overall stronger electrical attraction with the more highly
charged M2+ ions.
-
Chemically very similar
e.g. form mainly ionic compounds but different formulae and less reactive
because the 1st ionisation energies are higher (due to extra nuclear
charge) and a 2nd ionisation energy input to form the stable M2+
ion.
-
Oxidation state or oxidation
number is always +2
in Group 2 Alkaline Earth Metal compounds.
-
The two outer s–electrons are readily lost.
The 3rd, and subsequent ionisation energies are far too high to
form chemically stable cations of 3+ etc. i.e. the energy required will
not be compensated by ionic bond formation.
-
The stable Group 2 cation has electron
configuration of noble gas,
-
e.g. the calcium atom,
Ca, is 2,8,8,2 or
1s22s22p63s23p64s2
or [Ar]4s2
-
so the calcium ion, Ca2+, is
2,8,8 or
1s22s22p63s23p6
or [Ar]
-
See also section 7.4 below on
Group trends and comparison with Group 1 metals.
-
PLEASE NOTE
that Radium is highly
radioactive and therefore difficult and dangerous to study BUT all its known
physical and chemical properties fit in with it being at the foot of Group 2
and other properties could be inferred from the properties of Mg to Ba.
TOP OF PAGE and
sub-index
7.4.
General Trends down groups 1 & 2 with increasing atomic number and formula
patterns
IONISATION
ENERGY PATTERNS
-
Because from one element to
the next, an extra shell of electrons is added, increasing the electron
'bulk' and the outer electrons are increasingly less strongly held (see
above).
-
The radii of the adjacent
Gp 2 atom is smaller than Gp 1 atom on the same
period, because the
nuclear charge has increased by one unit (L to R in PT), but is
attracting electrons in the same shell.
-
Similarly the radii of a Gp 2
M2+ ion is smaller than the adjacent Gp 1 M+ ion on the same
period, because the
nuclear charge has increased by one unit (L to R in PT), but is
attracting the same number of electrons in the same shells. (see data
tables in section 7.2 and section
7.3)
-
REACTIVITY TREND THEORY – relate to atomic radius and ionisation
energy
-
The metal gets more
reactive down the group because
...
-
When an alkali metal
atom reacts, it loses an electron to form a singly positively
charged ion e.g. Na ==> Na+ + e– (in terms of
electrons 2.8.1 ==> 2.8 and so forming a stable ion with a noble
gas electron arrangement).
-
As
you go down the group from one element down to the next the
atomic radius gets bigger due to
an extra filled electron shell.
-
The
outer electron is further and further from the nucleus and is also
shielded by the extra full electron shell of negative charge.
-
Therefore the outer electron is less
and less strongly held by the positive nucleus.
-
This
combination of factors means the outer electron is more easily lost,
the M+ ion more easily formed, and so the element is more
reactive as you go down the group – best seen in the laboratory with
their reaction with water.
-
The
reactivity argument mainly
comes down to increasingly lower ionisation energy down the
group and a similar argument
applies to the Gp 2 metals, but two electrons are removed to
form the cation.
-
The reaction of a group 1/2
metal with oxygen, water or halogens gets more vigorous as you descend the
group.
-
Generally (but not always),
the melting and boiling points fall steadily:
-
This is because the ionic
radii increase down the group increasing charge separation between the
metal cations of the lattice and the free delocalised electrons.
-
This
weakens the electrical attractive bonding force and so less thermal KE
is needed to weaken the lattice to the 'collapse point' i.e. melting.
-
BUT
the situation is not as simple as might be expected, e.g. the metal ions
do not always have the same crystal lattice packing arrangement.
s–block cation \ anion |
oxide |
hydroxide |
carbonate |
hydrogencarbonate
|
O2– |
OH– |
CO32– |
HCO3– |
formula
derived from Group1 cation M+ |
M2O |
MOH |
M2CO3 |
MHCO3 |
(M+)2O2– |
M+OH– |
(M+)2CO32– |
M+HCO3– |
formula
derived from Group
2 cation M2+ |
MO |
M(OH)2 |
MCO3 |
M(HCO3)2 |
M2+O2– |
M2+(OH–)2 |
M2+CO32– |
M2+(HCO3–)2 |
-
There are also Group 1
hydrogen sulfates (hydrogensulfates/hydrogensulfates) of formula MHSO4, i.e. half neutralised
sulfuric acid which are rarely encountered but can be crystallised.
-
The oxidation
numbers–states in the compounds listed in the two tables above and below:
-
+1 for metal cation of group 1, +2 for
metal cation of group 2
-
oxygen –2, hydrogen +1, carbon +4 in
table above and below
-
halogens e.g. Cl –1, nitrogen +5,
sulfur
+6
-
you need to be able to analyse an anion
to understand the relationship between the constituent oxidation states
and the charge on the anion, sum of oxidation states = overall charge
on ion e.g.
-
carbonate, CO32–,
C is +4, 3 O's at –2, sum of ox. states = +4 and –6 = –2 = charge on
anion
-
hydrogencarbonate,
HCO3–,
H is +1, C is +4, 3 O's at –2, sum of ox. states = +1 and +4 and –6
= –1 = charge on anion
-
nitrate(V),
NO3–,
N is +5, 3 O's at –2, sum of ox. states = +5 and –6 = –1 = charge on
anion
-
sulfate(VI),
SO42–,
S is +6, 4 O's at –2, sum of ox. states = +6 and –8 = –2 = charge on
anion
cation\anion |
halide |
nitrate(V) |
sulfate(VI) |
ethanoate |
X– |
NO3– |
SO42– |
CH3COO– |
formula
derived from Group1 cation M+ |
MX |
MNO3 |
M2SO42– |
CH3COOM |
M+X– |
M+NO3– |
(M+)2SO42– |
CH3COO–M+ |
formula
derived from Group
2 cation M2+ |
MX2 |
M(NO3)2 |
MSO4 |
(CH3COO)2M |
M2+(X–)2 |
M2+(NO3–)2 |
M2+SO42– |
(CH3COO–)2M2+ |
WHAT NEXT?
GCSE Level
periodic table notes (for the basics)
and
GCSE Level alkali
metal notes (for the basics)
INORGANIC Part
7 s–block Gp 1 Alkali Metals/Gp 2 Alkaline Earth Metals sub–index:
7.1 Introduction to s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals * 7.2
Group 1 data and graphs * 7.3
Group 2 data and graphs *
7.4 General trends down groups I & II and formulae
*7.5
Oxygen reaction & oxides of s–block
metals *
7.6 Water reaction & hydroxides of
group 1/2 metals
* 7.7 Acid reaction & salts of group1/2
metals * 7.8
chlorine
reaction & halides of group I/II metals * 7.9
carbonates & hydrogen carbonates
of s–block metals
* 7.10 Solubility trends of groups 1/2 OH, NO3,SO4,CO3
compounds
* 7.11 Thermal
decomposition and stability of group 1 and group 2 carbonates & nitrates * 7.12
Uses of
s–block Group 1 Alkali Metals and Group 2 Alkaline Earth Metals and their compounds
Each page has a matching sub-index
Advanced
Level Inorganic Chemistry Periodic Table Index:
Part 1
Periodic Table history
Part 2
Electron configurations, spectroscopy,
hydrogen spectrum,
ionisation energies *
Part 3
Period 1 survey H to He *
Part 4
Period 2 survey Li to Ne * Part
5 Period 3 survey Na to Ar *
Part 6
Period 4 survey K to Kr AND important
trends down a group *
Part 7
s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals *
Part 8
p–block Groups 3/13 to 0/18 *
Part 9
Group 7/17 The Halogens *
Part 10
3d block elements & Transition Metal Series
*
Part 11
Group & Series data & periodicity plots
All
11 Parts have
their own sub-indexes near the top of the pages
Group numbering and the modern periodic
table
The original group numbers of
the periodic table ran from group 1 alkali metals to group 0
noble gases. To account for the d block elements and their
'vertical' similarities, in the modern periodic table, groups 3
to group 0 are numbered 13 to 18. So, the p block elements are
referred to as groups 13 to group 18 at a higher academic level,
though the group 3 to 0 notation are still
used, but usually at a lower academic level. The 3d block
elements (Sc to Zn) are now considered the head (top) elements
of groups 3 to 12. The s-block elements of the Groups 1 and 2
metals retain their original numbers.
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