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Ionization energy
TOPIC 3: PERIODICITY
Physical properties chapters 8.8, 9.5
Periodicity = regularly repeating pattern of physical and chemical properties in the Periodic Table.
electro­
negativity
ionization
energy
The energy (in J) required for the removal of 1 e­ from an atom in gaseous state.
The relative ability to attract a bonding pair of e­ to itself in a covalent bond (Pauling's scale: 0 ­ 4,0).
melting
point
The temperature where the forces holding particles together in a crystal structure are overcome.
Electronegativity
physical
properties
ionic
radius
Cations' ionic radii are ca. ½ of the atomic radii, whereas anions' radii are nearly twice as large as atomic radii.
atomic
radius
Half the distance between 2 identical atoms bonded together (measured by X­
ray diffraction). The bond determines whether the radius is covalent or metallic.
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Melting point
Atomic radius
Ionic radius (cations)
Ionic radius (anions)
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Chemical properties Group 7: Halogens
chapters 9.6, 17.4, 24.1, 24.5 ­ 24.7
Group 1: Alkali metals
­ diatomic molecules
­ shiny, silvery, soft
Br2
Rb
Li
Na
Cs
K
­ reaction w/ O2 produces a metal oxide
4 M (s) + O2 (g) 2 M (s) + 2 H2O (l) ­ reaction w/ halogens:
2 M (s) + X2 (g/l/s) ­ oxidizing agents = accept readily electrons
therefore "force" other elements to lose theirs
This oxidizing ability decreases down the group.
Youtube: Alkali metals in water
In that case, which of the following chemical reactions are possible?
2 MOH (aq) + H2 (g)
Youtube: Sodium and chlorine
2 MX (s)
Cl2
F2 > Cl2 > Br2 > I2
2 M2O (s)
­ reaction w/ H2O greatly exothermic
F2
I2
a.
b.
c.
d.
Cl­ (aq) + Br2 (l)
F­ (aq) + Br2 (l)
I­ (aq) + Br2 (l)
Cl­ (aq) + F2 (g)
Br­ (aq) + Cl2 (g)
Br­ (aq) + F2 (g)
Br­ (aq) + I2 (s)
F­ (aq) + Cl2 (g)
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Silicon's oxide, SiO2, has a so called giant covalent
structure with very high melting and boiling points.
It is insoluble in water. It is however classified as an acidic oxide, because it reacts with NaOH at temperatures above 350 oC:
Period 3:
Oxides are products of reactions where an element
reacts with oxygen.
Metal oxides (Na2O, MgO) are ionic compounds, i.e. they
­ are solids in room temperature
­ have high mp & bp
­ conduct electricity when molten (or in aqueous
solution)
­ are basic in aqueous solutions:
Na2O (s) + H2O (l) 2 NaOH (aq)
MgO (s) + H2O (l)
Mg(OH)2 (s)
Aluminium oxide, Al2O3 is all of the above except
for the latter. It is amphoteric, which means it acts as a base when it reacts with acids and acts as an acid
when it reacts with bases.
SiO2 (s) + NaOH (aq)
Na2SiO3 (aq) + H2O (l)
Non­metal oxides are covalently bonded. This is due to the small difference in the elements' electro­
negativity values.
sulfur: SO2, SO3 chlorine: Cl2O, Cl2O7
phosphorus: P4O6, P4O10
Non­metal oxides:
­ have low mp & bp
­ do not conduct electricity in liquid state
­ are acidic in aqueous solutions
P4O10 + 6 H2O (l)
4 H3PO4 (aq)
phosphoric(V)acid
SO3 (g) + H2O (l)
H2SO4 (aq)
sulfuric(VI)acid
Remember: noble gases do not form oxides!
SL ends HERE.
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Cl2O (g) + H2O (l)
2 HClO (aq) chloric(I)acid
(hypochlorous acid)
Cl2O7 (l) + H2O (l)
2 HClO4 (aq)
chloric(VII)acid
perchloric acid
P4O6 (s) + 6 H2O (l) 4 H3PO3 (aq)
phosphoric(III)acid
SO2 (g) + H2O (l)
H2SO3 (aq)
sulfuric(IV)acid
(sulfurous acid)
Chlorides are compounds of chlorine and one other element. Metal chlorides are ionic and thereby conduct electricity when molten or aqueous and have high melting points.
Aluminium chloride, however, possesses a covalent
character:
Al2Cl6
AlCl3
AlCl3
Transition elements chapters 6.1, 23.3 ­23.9, 15.9
Transition elements are elements that contain an
incomplete d level of electrons in one or more of
their oxidation states. This is why neither Zn nor Sc are considered transition elements:
Zn: [Ar]4s23d10
Sc: [Ar]4s23d1
Zn2+: Sc3+: Transition elements show characteristic properties:
The acidity of aqueous solutions of chlorides
increases across the period:
NaCl :
MgCl2 :
AlCl3 :
SiCl4 :
PCl3 :
PCl5 :
Cl2 :
neutral aqueous solution
weakly acidic aq. sol.
acidic
acidic
reaction with water produces HCl
acidic
and the element's oxoacid
acidic
acidic
variable
oxidation states
complex
ion formation
transition
elements
colored
compounds
catalytic
properties
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Complex ion formation:
+1
+2
­1
­2
H
H (most often) and alkali metals: earth alkali metals:
halogens, especially F: oxygen: These electron pairs form coordinate covalent
bonds with the metal ion to form complex ions.
H
H
H
When cations are formed, 4s electrons are lost first.
All 1st row transition elements (apart from Cr & Cu) contain two 4s electrons meaning: O
O
H
3+
Fe
H
O
O
H
H
O
all show +2 oxidation state.
Even higher oxidation states occur, but those are found in covalent bonds or the polyatomic ions:
oxyanions. Examples: Cr2O72­, MnO4­
H
H
H
Some form the +3 or +4 ion, the latter being rare due to small size & high charge (which easily leads to
covalent bonding). Examples: CrCl3, Fe2O3, MnO2
H
The oxidation state of an element keeps track of
the number of electrons being lost or gained. Some
elements always have the same oxidation state in
its compounds:
Because of their small size, transition element ions
attract species that are rich in electrons: ligands.
They are neutral molecules or negative ions that
contain a non­bonding pair of electrons.
O
Variable oxidation states:
structural
formula:
chemical
formula:
name:
shape:
coordination number:
Other possible ligands are e.g. NH3, Cl­ and CN­.
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Colored complexes:
When white light falls on a solution containing a
complex ion of a transition element, some of the
light corresponding to a particular wavelength (or
energy) is absorbed.
Catalytic properties:
A catalyst increases the rate of a reaction without
participating in the reaction itself. Therefore, a catalyst does not become chemically changed at the end of the reaction and can be reused.
light with λ ≈ 700 nm
is absorbed, complementary
green color is shown
white
light
Heterogenous catalysts are in a different phase from
the reactants and products.
light with λ ≈ 400 nm
is absorbed, complementary
yellow color is shown
H ­
H
ΔE
excited
state
Four main factors affect the size of the splitting, hence the color:
­ nature of transition element (Cu or Ni?)
­ oxidation state
(+2 or +3?)
­ ligand
(H2O or NH3?)
­ shape of complex ion
(linear or
­ H
H H ­ H
­ H
H ΔE
ground
state
When the light is absorbed, its energy raises an electron from one energy level to another. Since bonding in complex ions involves always d orbitals, the electron transition occurs within the split d orbital.
nickel surface
Transition metals are good heterogenous catalysts because they are adsorb small molecules particularly well.
Other examples:
Fe in the Haber process (NH3)
Rh, Pt, Pd in vehicle catalysts
square planar?)
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Homogenous catalysts are in the same phase as the reactants and products. The two reacting species
bond chemically to the transition metal forming
an intermediate, and then leave.
Transition metals can be relatively easily oxidized
and reduced due to their variable oxidation states.
This is why the intermediate can be formed easily.
Examples:
2 H2O2 (aq)
MnO2 (s)
2 SO2 (g) + O2 (g)
2 H2O (l) + O2 (g)
V2O5 (s)
2 SO3 (g)
cobalt in
vitamin B12
more efficient
reaction
more product
more quickly
catalyst
N
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EM
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less money
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environment
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