|
7.5.
The reaction of s–block
metals and oxygen & their oxide (O2–) chemistry
-
The oxides and hydroxides
are white ionic solids.
-
The reaction of Group 1
metals with oxygen (a redox reaction)
-
Group 1
metals: 4M(s)
+ O2(g) ==> 2M2O(s)
(M = Li, Na, K, Rb, Cs)
-
shows the formation of the 'simple' oxide expected from their position in the
periodic table when the element is heated or burned in air.
-
Oxidation state changes: M
is 0 to +1, Oxygen is 0 to –2 in the oxide ion O2–.
-
The oxides are soluble in
water forming the strongly alkaline hydroxide:
-
Unfortunately, except for
lithium (an anomaly), 'higher' oxides can be formed e.g.
-
2M(s) + O2(g)
==> M2O2(s) [redox change, M
(0 to +1), O (0 to –1)]
-
shows the formation of
the yellow–orange peroxide by Na, K, Rb and Cs
-
each oxygen is in the –1
oxidation state in the peroxide ion O22–
-
they readily hydrolyse
with water forming hydrogen peroxide
-
M(s) + O2(g)
==> MO2(s) shows the formation of the 'superoxide'
by K, Rb and Cs
-
oxidation number
changes are M from 0 to +1 as expected, but on average each oxygen
changes from 0 to –1/2 in the superoxide ion
O2–
-
2MO2(s) + 2H2O(l)
==>
2MOH(aq) + H2O2(aq) + O2(g)
(redox change)
-
oxidation state changes:
M and H no change (+1), four O's change from –1/2
in superoxide ions to two of –1 in the peroxide molecule and two at
zero in the oxygen molecule.
-
The simple oxides
readily dissolve in acids and are neutralised to form salts.
-
M2O(s) +
2HCl(aq) ==> 2MCl(aq) +
H2O(l) (M = Li, Na, K, Rb, Cs)
-
M2O(s)
+ 2HNO3(aq) ==> 2MNO3(aq) + H2O(l)
to give the
soluble nitrate salt
-
M2O(s)
+ H2SO4(aq) ==> M2SO4(aq) + H2O(l)
to give the
soluble
sulphate salt
-
M2O(s)
+ 2CH3COOH(aq) ==> 2CH3COOM(aq) + H2O(l)
to give the
soluble ethanoate
salt
-
The reaction of Group 2
metals with oxygen
(a redox reaction)
-
Group 2
metals: 2M(s)
+ O2(g) ==> 2MO(s)
(M = Be, Mg, Ca, Sr, Ba)
-
shows the formation of the oxide expected from their position in the
periodic table when the element is heated or burned in air. Oxidation state changes: M
from 0 to +2, and oxygen from 0 to –2.
-
The oxide, apart from
beryllium, is slightly soluble in water forming the alkaline hydroxide,
which increases in strength of basic character down the group.
-
MO(s) + H2O(l)
==> M(OH)2(s/aq) (not a
redox change, M = Be, Mg, Ca, Sr, Ba)
-
All the oxides are basic and readily neutralised by
acids (not a
redox change).
-
MO(s) +
2HCl(aq) ==> MCl2(aq) + H2O(l)
(M = Be, Mg, Ca, Sr, Ba)
-
to give the soluble chloride salt
-
ionically: M2+O2–(s) +
2H+(aq) ==> M2+(aq) + H2O(l)
-
This applies to all
four acid reactions examples in this section, acid proton donation to the oxide ion base.
-
In each case the chloride Cl–,
nitrate NO3– and sulphate SO42–
are spectator ions.
-
MO(s) +
2HNO3(aq)
==> M(NO3)2(aq) + H2O(l)
(M = Be, Mg, Ca, Sr, Ba)
-
MO(s) + H2SO4(aq)
==> MSO4(aq/s) + H2O(l)
(M = Be, Mg, Ca, Sr, Ba)
-
MO(s) +
2CH3COOH(aq)
==> (CH3COO)2M(aq) + H2O(l)
(M = Be, Mg, Ca, Sr, Ba)
-
Beryllium oxide BeO
is amphoteric
(another Be Gp 2 anomaly) and dissolves in strong bases
like sodium hydroxide.
-
The equation below shows the formation of a hydroxo
beryllate complex ion
(not a redox change).
-
BeO(s) +
2NaOH(aq) + H2O(l) ==> Na2[Be(OH)4](aq) (beryllate
salt)
-
ionically: Be2+O2–(s)
+ 2OH–(aq) + H2O(l)
==> [Be(OH)4]2–(aq)
TOP OF PAGE and
sub-index
7.6.
Reaction of s–block
metals and water & their hydroxide (OH–) chemistry
-
The oxides and hydroxides
are usually white ionic solids.
-
 The
reaction of group 1 metals with water
(a redox reaction)
-
Group 1 metal hydroxide formation
-
2M(s)
+ 2H2O(l) ==> 2M+OH–(aq) + H2(g)
(M = Li, Na, K, Rb, Cs)
-
shows the formation of the
alkaline metal hydroxide and hydrogen.
-
M = Li (slow at first), Na
(fast), K (faster – may ignite hydrogen to give a lilac coloured flame*
from hot potassium atoms), Rb, Cs, Fr (very explosive) i.e. the
reactivity increases down the group.
-
The hydroxides, MOH, are
white ionic solids, all very soluble (except LiOH), strong
bases, getting stronger down the group.
-
All Group 1 hydroxides
are soluble in water giving strongly alkaline solutions,
-
and their aqueous
solutions readily neutralised by acids (not a
redox change) e.g.
-
MOH(aq) + HCl(aq)
==> MCl(aq) +
H2O(l) (M = Li, Na, K, Rb, Cs)
-
to give the soluble chloride salt*
-
ionically: OH–(aq) +
H+(aq) ==> H2O(l)
-
an acid–base reaction, same for all four
examples in this section
-
MOH(aq) + HNO3(aq)
==> MNO3(aq) + H2O(l)
(M = Li, Na, K, Rb, Cs)
-
2MOH(aq) + H2SO4(aq)
==> M2SO4(aq) + 2H2O(l)
(M = Li, Na, K, Rb, Cs)
-
MOH(aq) + CH3COOH(aq)
==> CH3COOM(aq) + H2O(l)
(M = Li, Na, K, Rb, Cs)
-
*
The hydroxide solutions
are readily titrated with standardised hydrochloric acid (burette)
using phenolphthalein indicator, the colour change is from pink to
colourless.
-
The
reaction of group 1 metals with water
(a redox reaction)
-
Group 2 metal hydroxide
formation
-
M(s)
+ 2H2O(l) ==> M(OH)2(aq/s) + H2(g)
(M = Mg, Ca, Sr, Ba)
-
shows the formation of the hydroxide and hydrogen with
cold water.
-
ionically: M(s)
+ 2H2O(l) ==> M2+(aq) + 2OH–(aq) + H2(g)
-
oxidation number changes, M
is 0 to +2, for one H per water it changes from +1 to 0 in H2.
-
M = Be (no reaction, anomalous), Mg
(very slow reaction), Ca, Sr, Ba (fast to very fast).
-
i.e. the reactivity
increases down the group.
-
The reactivity trend
for Group 2, and its explanation, are similar to that above for the
Group 1 Alkali Metals.
-
The reactivity
trend for s–block metals is explained below
-
Magnesium hydroxide and calcium hydroxide
(limewater) are
sparingly soluble, but the solubility increases down the group, so
barium hydroxide is moderately soluble.
-
As previously
mentioned, a mixture of magnesium oxide/hydroxide and water
is sometimes called milk of magnesia and the
saturated aqueous solution of calcium hydroxide is called
limewater.
-
If the metal is heated in
steam the oxide is formed:
-
REACTIVITY TREND THEORY: The Group 1/2 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.
-
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 as you go down from one period to the
next one.
-
This means the
outer electron is further and further from the nucleus.
-
This also
means the outer electron is also shielded by the extra full electron shell of negative charge.
-
Due to this shielding
the effective nuclear charge on the external electron is ~ +1 (~
proton number – number of noble gas inner core electrons).
-
Further more, the
effective nuclear charge of ~+1 is acting over a larger 'surface
area' as the atomic radius increases.
-
Therefore both of these
factors combine to make the outer electron less
and less strongly held by the positive nucleus as the atomic number
increases (down the group).
-
So, the outer electron is more easily lost,
and 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 (i.e. ease of ion formation) and a similar argument
applies to the Group 2 metals, but two electrons are removed to
form the cation.
-
The enthalpy change
in forming the hydrated cation from the solid metal does not
appear to be as important here.
-
At a more advanced and detailed
level, this change can be theoretically split into the
-
enthalpies of
(i) atomisation, (ii) ionisation, (iii) hydration of gaseous ion
... (BUT not here!).
-
The reactivity trend is
also paralleled by the
increasingly negative half–cell potential (EθM/M+ and EθM/M2+) down
groups, 1 and 2 i.e. increasing potential to acts as a reducing agent
– an electron donor.
-
As with water, the
reaction of a group 1/2 metal with oxygen or halogens gets more
vigorous as you descend the group.
-
All the hydroxides are basic
with increasing strength down the group and readily neutralised by acids
(not redox reactions). Magnesium hydroxide is sparingly soluble in
water but the solubility increases down the group.
-
M(OH)2(aq/s) +
2HCl(aq) ==> MCl2(aq) + 2H2O(l)
(M = Be, Mg, Ca, Sr, Ba)
-
to give the
soluble chloride salt*
-
all base (OH–)
... acid (H+)
reactions
-
ionically if
soluble the reaction is: OH–(aq) + H+(aq)
==>
H2O(l)
-
ionically if
insoluble: M2+(OH–)2(s) + 2H+(aq)
==>
M2+(aq) + 2H2O(l)
-
M(OH)2(aq/s)
+
2HNO3(aq)
==> M(NO3)2(aq) + 2H2O(l)
(M = Be, Mg, Ca, Sr, Ba)
-
M(OH)2(aq/s) + H2SO4(aq)
==> M2SO4(aq/s) + 2H2O(l)
(M = Be, Mg, Ca, Sr, Ba)
-
M(OH)2(aq/s) +
2CH3COOH(aq)
==> (CH3COO)2M(aq) + 2H2O(l)
(M = Be, Mg, Ca, Sr, Ba)
-
*
Saturated calcium
hydroxide solution (limewater) can be titrated with standardised
hydrochloric acid (burette, low molarity) to determine its
solubility. You normally use phenolphthalein indicator and the
end–point colour change is from pink to colourless.
-
The Group 2 hydroxides,
M(OH)2, get more
soluble down the group:
-
Beryllium hydroxide is
amphoteric
(an anomaly in the group), because apart from the reactions
above, it dissolves in strong alkalis like sodium hydroxide to form a
hydroxo–complex ion salts called 'beryllates' e.g.
-
For
the reaction of Group 1 and 2 hydroxides with carbon dioxide
to form the carbonates and hydrogen carbonates,
see
section 7.9
-
For the thermal
decomposition of nitrates
see section 7.11
-
See also
VOLUMETRIC
TITRATION QUESTIONS involving acids and group1 /2 alkaline
hydroxides
TOP OF PAGE and
sub-index
7.7.
The reaction of s–block
metals with acids
-
Group 1 metals are far too
reactive to contemplate adding them to acids in a school laboratory!
-
Group 2
metals, apart from
beryllium (another anomaly), readily react with acids, with increasing vigour down the group (explanation
in section 7.4). A redox reaction to form the
soluble chloride salt.
-
M(s) + 2HCl(aq)
==> MCl2(aq) + H2(g) (M
= Mg, Ca, Sr, Ba)
-
M(s) + 2HNO3(aq)
==> M(NO3)2(aq) + H2(g)
-
to
form the soluble nitrate salt
-
Looks ok in principle,
and does this with Mg and very dilute nitric acid, but rarely this simple, the nitrate(V) ion can get reduced to nasty
brown nitrogen(IV) oxide gas (nitrogen dioxide, NO2) and
other products, NO gas?, NO2– ion?
-
M(s) + H2SO4(aq)
==> MSO4(aq/ s) + H2(g)
-
to form
soluble ==> insoluble sulphate salt
-
The reaction from
magnesium to barium becomes increasingly slower as the sulphate
becomes less soluble, it coats the metal, inhibiting the reaction.
-
M(s) + 2CH3COOH(aq)
==> (CH3COO)2M(aq) + H2(g)
to form soluble ethanoate salt
-
 In
aqueous solutions the metal cations formed are hydrated to aqa–complex
ions.
-
As described above, The soluble
groups 1/2 salt solutions contain the hydrated cations derived from
the metal:
-
tetra–aqua cations [Li(H2O)4]+(aq)
and [Be(H2O)4]2+(aq)
-
or the hexa–aqua ions [M(H2O)6]+(aq)
M = Na, K etc. for Group 1
-
and [M(H2O)6]2+(aq)
where M = Mg, Ca etc. for Group 2
-
The tetraaqua beryllium ion and
the hexaaqua magnesium ions generate a slight acidity in their salt
solutions due to the significant polarising power of the ions (Be2+
very small and double charged, Mg2+ double charged) e.g.
-
for beryllium:
[Be(H2O)4]2+(aq)
+ H2O(l)
[Be(H2O)3(OH)]+(aq)
+ H3O+(aq)
-
or magnesium:
[Mg(H2O)6]2+(aq)
+ H2O(l)
[Mg(H2O)5(OH)]+(aq)
+ H3O+(aq)
TOP OF PAGE and
sub-index
7.8.
The reaction of s–block
metals with chlorine & halide (X–) salts
-
The salts are
white or colourless crystalline solids
-
Group 1 metals readily react
with halogens
(a redox reaction)
-
e.g. heating the metal in
chlorine will cause it to burn forming the chloride
-
2M(s) + Cl2(g)
==> 2MCl(s) (redox reaction, M = Li, Na,
K, Rb, Cs)
-
Oxidation state changes:
M from 0 to +1, X = F, Cl, Br & I from 0 to –1
-
The salt products, M+X–,
are
white–colourless crystalline ionic solids that dissolve in water to give
neutral solutions of about pH 7. The crystalline solids have high
melting and boiling points.
-
The solids do not
conduct electricity (no mobile ions or electrons) but will conduct
and undergo electrolysis when molten or dissolved in water when ions
are free to move to electrodes.
-
The halogen is in the –1
oxidation state in the halide ion X–
-
The halides of groups
1–2 are important raw materials e.g.
-
sodium chloride
==> sodium hydroxide from rock salt by electrolysis of aqueous
solution
-
potassium
bromide/iodide ==> elemental bromine/iodine from seawater by
oxidation
-
calcium
chloride ==> calcium metal by electrolysis of molten chloride
-
Group 2 metals (except Be)
readily react on heating with halogens
(a redox reaction)
-
e.g. heating in chlorine the
chloride is formed
-
M(s) + Cl2(g)
==> MCl2(s) (M = Mg, Ca, Sr,
Ba)
-
Oxidation state changes:
M from 0 to +2, X = F, Cl, Br & I from 0 to –1
-
The salt products, M2+(X–)2,
are similar in properties to the Group 1 M+X– compounds.
-
However, beryllium chloride has a
polymeric covalent structure, due to the high polarising influence of
beryllium in its +2 oxidation state and the smaller difference in
electronegativity between Be–Cl compared to chlorine and the other
group 1 and 2 metals.
keywords equations:
CaCl2(aq) + 2NaOH(aq) ==> 2NaCl(aq) + Ca(OH)2(s) * MgSO4(aq) + 2KOH(aq)
==> K2SO4(aq) + Mg(OH)2(s) * Ba(NO3)2(aq) + 2NaOH(aq) ==>
2NaNO3(aq) + Ba(OH)2(s) * CaCl2 + 2NaOH ==> 2NaCl + Ca(OH)2 * MgSO4 +
2KOH ==> K2SO4 + Mg(OH)2 * Ba(NO3)2 + 2NaOH ==>
2NaNO3 + Ba(OH)2
|
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.
|
| Website content © Dr Phil Brown
2000+. All copyrights reserved on revision notes, images, quizzes,
worksheets etc. Copying of website material is NOT permitted. Doc
Brown's Chemistry theoretical-physical chemistry revision notes for
pre-university level students on s-block metals group 1 alkali
metals group 2 alkaline earth metals physical and chemical
properties reactions equations and trends explained |
|
(I will replace
with proper diagrams later) 