Download Topic 3: Periodicity

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Topic 3: Periodicity
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Chemistry is not study of a random collection of
elements.
Periodicity = repeating patterns of chemical and
physical properties.
3.1 The periodic table
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The development of the periodic
table took many years and involved
scientists from several countries.
Mendeleev grouped the known
elements into families on the basis
of their relative atomic masses and
their chemical properties.
Dimitri Mendeleev (1834-1907)
He left gaps where no known
elements fitted in and predicted the
physical and chemical properties of
these missing elements.
http://www.youtube.com/watch?v=157KVUQXRJA
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In the modern periodic table the elements are
arranged in order of increasing atomic number , Z.
http://education.jlab.org/itselemental/
http://www.rsc.org/periodic-table/
Groups
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Elements with similar chemical and physical properties
are placed underneath each other in vertical columns
called groups.
The groups are numbered 1-18.
The number of valence electrons of the elements in
group 1-2 and 13-18 can be found from the group
number .
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Certain groups have their own names:
- alkali metals
- alkaline earth metals
- pnictogens
- chalcogens
- halogens
- noble gases
Periods
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The period number is equal to the principal quantum
number, n, of the highest occupied energy level in the
elements of the period.
Metals, non-metals and metalloids
http://www.ptable.com/
Metals
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Shiny
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Malleable and ductile
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Good conductors of heat and
electricity
Non-metals
Sulphur
Carbon
Selenium
Phosphorus
Metalloids
B
Si
Ge
Both metallic and non-metallic properties.
● They resemble metals physically and non-metals
chemically.
● Silicon and germanium are semiconductors.
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Transition metals
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They have very similar chemical and physical
properties:
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Relatively high mp, high bp and high densities
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They form more than one stable cation
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They often have coloured compounds and coloured
solutions.
Lanthanoides
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Elements from Z=57 to Z=71
Ex. Europium (Eu) is a hard silvery-white metallic
element that is used in the security marking of
euronotes.
Actinoids
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Elements from Z=89 to Z=103
A collection of uranium glassware
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Ex. Naturally occuring uranium is a silvery-white metal.
It is a mixture of U-238 (99.27%) U-235 (0.72%) and U234 (0.006%).
Uranium decays slowly by emitting an alpha particle.
The half-life of U-238 is about 4.47 billion years and
that of U-235 is 704 million years.
3.2 Periodic trends
Effective nuclear charge
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In every atom there is a balance between the
attraction of the positively charged nucleus for the
negatively charged electrons and repulsion between
the electrons.
The outer electrons (valence electrons) are shielded
from the nucleus and repelled by the inner electrons.
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The outer electrons do not experience the full
attraction of the positive nucleus because of the
presence of inner electrons.
The effective charge experienced by the outer
electrons is less than the full nuclear charge.
Atomic radius
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Atomic radii
increase down a
group and
decrease across
a period.
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The effective charge increases as a period is crossed
from the left to the right:
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One proton is added to the nucleus and one
electron to the outermost electron shell.
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There is no change in the number of inner
electrons.
The effective charge remains almost the same down a
group, because:
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The increase in the nuclear charge is offset by the
increase in the number of inner electrons.
Ionic radius
Ionic radii
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Positive ions are smaller than their parent atoms
(because of loss of the outer shell).
Negative ions are larger than their parent atoms
(because of increased electron repulsion by addition of
electrons).
The ionic radii decrease as a period is crossed from
the left to the right (because of increased attraction
between the nucleus and the electrons).
The ionic radii increase down a group as the number
of electron shells increase.
Ionization energies
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The energy required to remove one mole of electrons
from one mole of gaseous atoms.
Electronegativity
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The ability of an atom to attract the shared pairs of
electrons in a covalent bond.
Electron affinity
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The energy change when an electron is added to an
isolated atom in the gaseous state.
X(g) + e- → X- (g)
Ex. F (g) + e- → F- (g)
Ea = -328 kJ mol-1
Melting points
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The nature of the bonding between the particles of an
element determines its melting point.
Strong bonds require higher energy
to break:
- The stronger the bonds, the
higher the melting point.
Chemical properties
http://www.webelements.com/
Metallic and non-metallic character
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Metals have a tendency to lose electrons during a
chemical reaction = they tend to be oxidized.
Non-metals have a tendency to gain electrons during a
chemical reaction = they tend to be reduced.
http://www.periodicvideos.com/
Group 18: the noble gases
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Colourless gases
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They exist as single atoms
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They are very unreactive
Group 1: the alkali metals
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Shiny, silvery, soft
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Good conductors of electricity and heat
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Low densities
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Too reactive to be found in nature
Reaction with water
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The alkali metals react with water to produce hydrogen
and a metal hydroxide = the resulting solution is
alkaline.
Reactivity increases down the group.
Li:
Na:
K:
https://www.youtube.com/watch?v=uixxJtJPVXk
Reaction with halogens
2 Na(s) + Cl2 (g) → 2 NaCl (s)
2 K(s) + Br2 (l) →2 KBr(s)
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The most vigorous reaction occurs between the most
reactive alkali metal and the most reactive halogen:
francium with fluorine.
Group 17: the halogens
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Diatomic molecules
Reaction between halogens and alkali metals
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The outer electron moves from the alkali metal to the
halogen atom and an ionic halide is formed.
2 Na (s) + Cl2 (g) → 2NaCl (s)
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The most vigorous reaction occurs between elements
which are furthest apart in the Periodic Table (Fr and
F).
https://www.youtube.com/watch?v=Mx5JJWI2aaw
Reactions between halogens and halides
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Reactivity decreases down a group as the atomic radius
increases and the attraction for outer electrons decreases.
A solution of a more reactive halogen, X2(g), will react with
a solution of halide ions of a less reactive halogen, X-(g).
Ex. Which of the following chemical reactions are
possible?
Silver halide precipitate
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The halogens form insoluble salts (= precipitate) with
silver:
AgNO3 (aq) + NaI (aq) → AgI (s)
Metal oxides
Metal oxides (Na2O, MgO) are ionic compounds:
- solids in room temperature
- high mp & bp
- conduct electricity when molten (or in aqueous
solutions)
- basic in aqueous solutions:
Na2O (s) + H2O (l) → 2 NaOH (aq)
MgO (s) + H2O (l) → Mg(OH)2 (s)
Aluminium oxide
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Al2O3 is an ionic oxide with some covalent character.
It is amphoteric as it acts as a base when it reacts with
acids and acts as an acid when it reacts with bases:
Silicon oxide
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Silicon oxide, SiO2, has a giant covalent structure with
very high melting and boiling points.
It idoes not dissolve in water.
It is classified as an an acidic oxide, because it reacts
with NaOH at temperatures above 350° C.
Non-metallic oxides
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Nonmetal oxides are covalently bonded because of
the small difference in the elements' electronegativity
values.
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sulfur: SO2, SO3
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chlorine: Cl2O, Cl2O7
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phosphorus: P4O6, P4O10
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They have low mp and bp
and do not conduct electricity.
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Non-metallic oxides are acidic in aqueous solutions:
P4O10 + 6 H2O (l) → 4 H3PO4 (aq)
phosphoric(V)acid
SO3 (g) + H2O (l) → H2SO4 (aq)
sulfuric(VI)acid
Chemical properties of elements in the Period 3
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Na, Mg and Al are metals. They are shiny and good
conductors of heat and electricity.
Si is a semi-conductor and is called a metalloid since it has
some of the properties of a metal and some of a nonmetal.
P, S, Cl and Ar are all non-metals and do not conduct
electricity.
Formula of
oxide
Na2O(s)
MgO(s)
Al2O3(s)
SiO2(s)
P4O10(s)
SO3(l) and
SO2(g)
Nature of
oxide
basic
basic
amphoteric
acidic
acidic
acidic
Acidic rain
• Rain is naturally acidic, because the water molecules
react with the CO2 in the air and form the weak acid
H2CO3.
• Acid rain: precipitation (rain, snow) with pH lower than
5.6.
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The main acids present in
acid rain are sulfuric acid
nitric acid.
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The sulfuric acid in the rain reacts with calcium
carbonate (in limestone or marble) to create calcium
sulfate, which then flakes off.
CaCO3(s) + H2SO4(aq) → CaSO4(aq) + CO2(g) +
H2O(l)
13.1 First row d-block elements
d-block elements:
Transition elements
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An element whose atoms have an incomplete d subshell, or which can give rise to cations with an
incomplete d sub-shell.
Zinc is a d-block element but not a transition metal.
Physical properties
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Small atomic radii compared to the neighbouring sblock elements
Only a small increase in atomic radii across the period
Metallic bonding:high electrical and thermal
conductivity, high mp, malleable, ductile, high density
Magnetic properties
Trends in the first IE
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The rate of increase in the first IE across the period is
much lower for transition elements compared to that
for the main-group elements.
Chemical properties
Oxidation number (oxidation state)
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The oxidation number of an element keeps track of the
number of electrons it has lost or gained.
Some elements always have the same oxidation
number:
Variable oxidation states
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When transition metals lose electrons they lose the 4s
electrons first.
All transition elements (except for Cr and Cu) contain
two 4s electrons, which means that they all have an
oxidation state of +2
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In an positive ion all energy levels are closer to the
nucleus (more p than e) and as the energy levels
move down in energy the order of the 4s and 3d
sub-levels are reversed.
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3+
The M ion is the most stable for scandium to
chromium.
The M2+ ion is the most stable for Mn to Zn (the
increased nuclear charge makes it more difficult
to remove a third electron).
In the higher oxidation states the elements
usually not exist as a free metal ions, but
covalently bonded or as a oxyanions (MnO4-).
Complexes
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Because of their small size and relatively high charge,
the transition metal ions have a high charge density.
They attract species that are rich in electrons: ligands.
A ligand has at least one atom with a lone pair of
electrons, e.g. H2O, NH3, Cl-, CN-
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These electron pairs form co-ordinate covalent bonds
with the metal ion and a complex ion is formed.
The number of co-ordinate covalent bonds from
ligands to the central metal ion is called the
coordination number.
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In a co-ordinate bond the shared pair of electrons
orginates from the same atom.
The charge on a complex ion is the sum of the charge
of the d-block metal ion and the charge of the ligand (if
they are ions).
Ligand exchange (replacement)
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In aqueous solutions the water molecule usually act as
a ligand, but it can be replaced with another ligand.
Catalytic behaviour
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A catalyst increases the rate of a reaction without
participating in the reaction itself.
A catalyst does not become chemically changed at the
end of the reaction and can be reused.
Many transition elements and their compounds are
very efficient catalysts.
Heterogeneous catalysts
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The catalyst is in a different state from the reactants.
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Ex. Iron (Fe) in the Haber-Bosch process:
N2(g) + 3H2 (g)
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2NH3 (g)
Nickel (Ni) in the conversion of alkenes to alkanes
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Vanadium (V) oxide in the contact process
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Manganese (IV) oxide with hydrogen peroxide
Homogenous catalysts
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Homogenous catalysts are in the same state of matter
as the reactants and products.
Transition metals are effective homogenous catalysts in
redox reactions since they can relatively easily be
oxidized and reduced due to their variable oxidation
states.
Magnetic properties
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Every spinning electron can behave as a tiny magnet (=
a magnetic dipole).
Electrons with opposite spins cancel each other out.
Most substances have all electrons paired up and so
are non-magnetic.
Materials ca be classified as diamagnetic, paramagnetic
or ferromagnetic based on their behaviour in an external
magnetic field.
a) Diamagnetism
A property of all materials that do not contain
unpaired electrons.
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In an external magnetic field the paired electrons
orientate themselves such that the field created by
their spin opposes the applied field → They are
weakly repelled by an external magnetic field.
Ex. Ne
b) Paramagnetism
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Stronger than diamagnetism.
Increases with the number of unpaired electrons (from
left to right in the periodic table).
The spins of unpaired electrons in an atom or ion can
temporarily be aligned in an external magnetic field.
These unpaired electrons behave as tiny magnets and
are attracted by an external magnetic field.
https://www.youtube.com/watch?v=hK_Yi5nxuKQ
c) Ferromagnetism
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The largest effect.
Only occurs in materials where unpaired d electrons in
a large number of atoms can line up with parallell
spins to form regions or domains of parallell spins.
In an external magnetic field the unpaired electrons in
the domains align themselves such that the magnetic
field created by their spin is aligned with the applied
field.
After the external magnetic field is removed, the
domains remain aligned making a permanent magnet.
Colour of transition metal complexes
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White light is the combination of equal
brightness of red, green and blue light (they are
primary colours since all other colours can be formed
from them).
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When white light is shone on a substance, some light
is absorbed and some is reflected.
www.youtube.com/watch?
v=EHMH0uQDEOU
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If all light is absorbed, the substance appears black.
If only certain wavelengths are absorbed, the
compound appears coloured.
If all light is reflected, the compound appears white.
Transition metals absorb visible light
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When white light falls on an aqueous solution of a
transition metal complex, the ions absorb some
colours.
The Fe3+ ion appears yellow because it absorbs light in
the blue region of the spectrum. Yellow is the
complementary colour to blue (= they are opposite
each other in the colour wheel.)
The colour of transition metal ions
Splitting of the d-orbitals
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In an isolated gaseous transition metal atom, the five
3 d orbitals are degenerate since they all have the
same energy.
However, since the 3 d sub-shells all have different
orientations in space they will be orientated differently
relative to the ligands in a complex ion.
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The 3d electrons close to a ligand will experience
repulsion and be raised in energy.
The 3d electrons further from the ligand will be
reduced in energy.
→ The 3d sub/shell splits into two energy levels.
http://www.chemguide.co.uk/inorganic/complexions/whatis.html
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The amount the orbitals are split (∆E) and hence the
colour of the complex depends on:
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The nature of the transition metal
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The oxidation state
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The shape of the complex
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The nature of the ligand
If the orbital is completely empty (Sc3+) or completely
full (Cu+, Zn2+), the complexes are colourless.
Nuclear charge
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Ions with a higher nuclear charge form stronger
coordinate bonds with the ligands lone pair of
electrons (= because of stronger electrostatic
attraction).
This stronger attraction causes a larger split in the 3d
orbitals (larger ΔE = higher energy light is absorbed).
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Mn2+ is pale pink (green light absorbed)
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Fe3+ is yellow (blue light is absorbed)
Charge density of the ligand
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Ligands can be arranged in order of their ability to split
the d orbitals in octahedral complexes.
Spectrochemical series:
Ammonia has greater charge density than water and
produces a larger split in the d-orbitals.
UV-vis absorbtion spectrum of
some copper complexes