Transcript File - Mr Weng`s IB Chemistry

IB CHEMISTRY
Topic 3 Periodicity
Higher level
3.1 The periodic table
OBJECTIVES
• The periodic table is arranged into four blocks associated with the
four sub-levels—s, p, d, and f.
• The periodic table consists of groups (vertical columns) and periods
(horizontal rows).
• The period number (n) is the outer energy level that is occupied by
electrons.
• The number of the principal energy level and the number of the
valence electrons in an atom can be deduced from its position on the
periodic table.
• The periodic table shows the positions of metals, non-metals and
metalloids.
• Deduction of the electron configuration of an atom from the
element’s position on the periodic table, and vice versa.
Periodic Table and orbitals
• Atomic number (Z) – the number of protons in the
nucleus of an atom of that element
• The atomic number of each element increases left to
right across each period
• The s,p,d,f atomic
orbitals are
arranged in blocks
of the periodic table.
Group vs. Period
• Group – vertical columns of the periodic table
which contain elements having similar chemical
and physical properties
• The groups to be known are 1 alkali metals, 17
halogens, 18 noble gases, transition metals,
lanthanoids and actinoids
• Period – horizontal rows of the periodic table
PT and groups
PT and electron shells
Metals vs nonmetals
Metals
• Conductors of heat and electricity
• Malleable (bent into shapes)
• Ductile (drawn into wires)
• Lustre (shiny)
• Oxidized (lose electrons)
Nonmetals
• Insulators of heat and electricity
• Brittle
• Dull
• Reduced (gain electrons)
PT metals vs nonmetals
3.2 Periodic trends
OBJECTIVES
• Vertical and horizontal trends in the periodic table exist for atomic
radius, ionic radius, ionization energy, electron affinity and
electronegativity.
• Trends in metallic and non-metallic behaviour are due to the trends
above.
• Oxides change from basic through amphoteric to acidic across a period.
• Prediction and explanation of the metallic and non-metallic behaviour
of an element based on its position in the periodic table.
• Discussion of the similarities and differences in the properties of
elements in the same group, with reference to alkali metals (group 1) and
halogens (group 17).
• Construction of equations to explain the pH changes for reactions of
Na2O, MgO, P4O10, and the oxides of nitrogen and sulfur with water.
Atomic Radii
• Decrease across a period, increase down a group.
Ionic
Radii
• cations
get
smaller
• anions
get
larger
First Ionization Energy
First Ionization Energy – The energy required to
remove one mole of electrons from a mole of
atoms or ions in the gaseous phase
X(g) X+(g) + e-
Explaining IE
• IE increases across a period because the nuclear charge
increases, attracting the electrons and the number of
electrons in the shell are increasing.
• IE decreases down a group because the electrons further
away from the nucleus and electrons in lower shells are
blocking the attraction causing electron shielding.
Anomaly type 1:
• B has a lower IE than Be because the 2p electrons are
slightly higher in energy than the 2s electrons, and so the
ionization for B is lower than for Be
Anomaly type 2:
• O has a lower IE than N because the px, py, and pz only
contain one electron. The extra electron in O causes a pair
and repulsion making it easier to remove, hence giving it a
lower IE than N.
Electron affinity
Electron affinity – is the energy released when 1
mol of electrons is attached to 1 mole of neutral
atoms or molecules in the gas phase
X(g) + e-  X-(g)
Increasing Eea
Electronegativity
Electronegativity – is a measure of the ability of an
atom to attract bonded electron pairs to itself
when in a covalent bond
Electronegativity and metallic nonmetallic character
Metals have small electronegativity values, nonmetals have high values. Differences greater than
1.8 will form an ionic bond rather than a covalent
bond.
Comparing electronegativies
Summary of trends in the Periodic Table
Alkali metals
• low melting and boiling points
• melting and boiling points decrease
due to increase shielding and less
nuclear attraction
• very reactive due to need to lose just
one electron to have full electron shell
• large atoms so metals are soft and not
dense
Alkali metals and water
• Alkali metals react vigourously with water to
create hydrogen and a base
Na(s) + H2O(l) → NaOH(aq) + H2(g)
Alkali metals with halogens
• Alkali metals react vigourously with halogens to
form salts
2Na(s) + Cl2(g) → 2NaCl(s)
Halogens
• very reactive due to need to gain just
one electron to have full electron shell
• very electronegative as just have to
gain one electron
• melting and boiling points increase due
to increased London dispersion forces
(IMF) between the simple covalent
molecules, and increased molecular
weight
• Halogens become darker are move from gas to
solid down the group
• A halogen higher on up the group will displace
(is more reactive) than one lower down
Halogens with halide ions
The more reactive halogen (further up the group)
will will take an electron from a halide ion to itself
become a halide ion.
Cl2(aq) + 2Br-(aq)  2Cl-(aq) + Br2(aq)
Oxide reactions and pH
• The oxides of elements have increasing acidity
across a period (Al is amphoteric being both acidic
and basic).
• Metal oxides are basic, non-metal oxides are acidic.
Na2O(s) + H2O(l)  2NaOH(aq)
MgO(s) + H2O(l)  Mg(OH)2(aq)
P4O10(s) + 6H2O(l)  4H3PO4(aq)
SO2(g) + H2O(l)  H2SO3(aq)
SO3(l) + H2O(l)  H2SO4(aq)
Increasing
acidity across
a period
(All these and previous equations must be learnt.)
OBJECTIVES
• Transition elements have variable oxidation states, form complex
ions with ligands, have coloured compounds, and display catalytic and
magnetic properties.
• Zn is not considered to be a transition element as it does not form
ions with incomplete d-orbitals.
• Transition elements show an oxidation state of +2 when the selectrons are removed.
• Explanation of the ability of transition metals to form variable
oxidation states from successive ionization energies.
• Explanation of the nature of the coordinate bond within a complex
ion.
• Deduction of the total charge given the formula of the ion and
ligands present.
• Explanation of the magnetic properties in transition metals in terms
of unpaired electrons.
Higher level
13.1 First-row d-block elements
• A transition element is defined as an element that
possesses an incomplete d-sublevel in one or more
oxidation states (ie. as an ion)
• All the elements in group 12, Zn, Cd, Hg, and Cn are
not transition metals as they contain full d-sublevels
with 10 d-electrons as ions (lose s electrons).
Higher level
Transition metals
• Transition metals have an empty d orbital. The d orbital
splits into two energy sublevels and electrons moving
between these gives them their properties
• Note: for Cr and Cu it is more energetically favourable
to half-fill and completely fill the d sub-level respectively
so they contain only one 4s electron
Higher level
Transition metals
• Produces colours and allows for complex ion formation
• It gives them variable oxidation numbers and makes
them good catalysts
• Most materials are diamagnetic (repelled by a
magnet), some transition metals are paramagnetic
(ferromagnetic – attracted by a magnet) due to unpaired
electrons allowing spin in one direction to form poles.
• When transition metals lose electrons they lose the
4s electrons first
• All transition metals can show an oxidation state of +2
and occurs when they lose both s orbital electrons
Higher level
Properties
Variable oxidation states come from the fact that they
have relatively small differences in their successive
ionization energies. Cf. 1st and 2nd IE of Na vs one of
the transition metals.
Common oxidation states from the data booklet:
Higher level
Oxidation numbers
Element
Z
3d
4s
Sc
21
[Ar]


Ti
22
[Ar]


V
23
[Ar]



Cr
24
[Ar]






Mn
25
[Ar]






Fe
26
[Ar]






Co
27
[Ar]






Ni
28
[Ar]






Cu
29
[Ar]






Zn
30
[Ar]








Higher level
Electron configurations
A complex consists of a central atom, which
is usually a metal atom or ion, and attached
groups called ligands
The
coordination
number is the
total number of
points at which a
central atom or
ion attaches
ligands
Higher level
Complex ions
A substance
consisting of one
or more
complexes is
called a
coordination
compound
33
Higher level
The region surrounding the central atom or
ion and containing the ligands is called the
coordination sphere
Coordinati
on
number
Shape
6
octahedral
4
tetrahedral
or square planar
2
linear
Square planar compounds are rare, but usually d8 configurations with strong field ligands.
Higher level
• The number of lone pairs bonded to the metals
ion is known as the coordination number
Higher level
• Co-ordination number examples
Coordination
number
6
4
2
Examples
[Fe(CN)6]3-
[CuCl4]2-
[Ag(NH3)2]+
[Fe(OH)3(H2O)3] [Cu(NH3)4]2+
All of these complex ions are bonded to monodentate ligands
which means they all consist of one type of ligand.
4 lobes is normally tetrahedral, but with full d8 (Cu) and strong
ligands it is square planar.
• A ligand is a neutral molecule or anion which
contains a non-bonding pair of electrons, these
electron pairs, from the ligand, form coordinate
bonds with the metal ion to form complex ions
• Complex ions form with transition metals
because of their small size d-block ions attract
species that are rich in electrons
Higher level
Ligands
A common ligand is
water and most (but not
all) transition metal ions
exist as hexahydrated
complex ions in
aqueous solution, e.g.
[Fe(H2O)6]3+
(1) The lone pair of electrons from the water
molecules (ligands) form the coordinate bonds.
(2) Iron is forming 6 bonds, so the coordination
number of the iron is 6
Higher level
Coordinate bonds and numbers
• Ethylenediaminetetraacetate (EDTA) is a
chelate or polydentate ligand as it grabs onto
the metal with more than one donor atom.
• It has several important uses including the
removal of heavy metal ions such as
treatment for lead poisoning, as well food
preservation in preventing transition metals
catalyzing food rancidity (going off).
Higher level
EDTA
OBJECTIVES
• The d sub-level splits into two sets of orbitals of different
energy in a complex ion.
• Complexes of d-block elements are coloured, as light is
absorbed when an electron is excited between the d-orbitals.
• The colour absorbed is complementary to the colour observed.
• Explanation of the effect of the identity of the metal ion, the
oxidation number of the metal and the identity of the ligand on
the colour of transition metal ion complexes.
• Explanation of the effect of different ligands on the splitting of
the d-orbitals in transition metal complexes and colour observed
using the spectrochemical series.
Higher level
13.2 Coloured complexes
CFT suggest that the dxy, dyz, dxz orbitals are of a
lower energy state (more stable) than the 𝑑𝑥 2−𝑦2
and 𝑑𝑧 2 orbitals creating a split d-sublevel.
The x, y, z axis is
where the atomic axis
where ligands join.
Higher level
Crystal field theory (CFT)
Strong field ligands cause spin paired splitting
which has higher energy levels as the electrons in
the ligands are repelling the electrons in the
metal which are on the same axis.
These orbitals
are not
directly in line
with the
ligands.
Higher level
Ligands and energy levels
The free electron pair orbitals of
the ligands are attracted to the
nucleus of the metal cation. The
overlapping orbitals with the
𝑑𝑥 2 −𝑦2 and 𝑑𝑧 2 can cause
electrons to move to the split
lower energy orbitals of dxy, dyz,
dxz which do not have direct
overlap with the ligands. This
difference in energy is the ∆0 and
varies in size, hence movement
of electrons between these split
d orbitals will produce different
wavelengths (seen as different
colours).
Higher level
Crystal field splitting energy (∆0)
More negative
charge density in
ammonia makes ∆0
greater than than
water.
Similarly, as fluorine
has more electron
density than
chlorine, so ∆0
increases.
Higher level
1. ∆0 and ligand strength
Descending down a
group the ∆0
increases due to
more orbital
overlap.
Higher level
2. ∆0 and metal ions
The higher the
oxidation state of
the cation the more
∆0 increases due to
greater attraction of
the nucleus to the
ligand and hence
more orbital
overlap.
Higher level
3. ∆0 and oxidation state
The crystal field splitting energy is the difference
in energy between these two split sublevels.
Higher level
Crystal field splitting energy (∆0)
The strength of the ligand determines the
amount of splitting. Splitting can be spin free or
spin paired.
HIGH energy split, low spin
LOW energy split, high spin
Higher level
Ligands and ∆0
Eg. CN is a strong-field ligand creating
diamagnetism, H2O is a weak field ligand creating
paramagnetism.
diamagnetic
paramagnetic
Higher level
Diamagnetism and paramagnetism
The lower energy state is written as 𝑡2𝑔 and the
higher energy state is written as 𝑒𝑔 .
The configurations are written as follows:
5 0
𝑡2𝑔
𝑒𝑔
4 2
𝑡2𝑔
𝑒𝑔
Higher level
Electron configurations and ∆0
The spectrochemical series (from data booklet)
shows the splitting strength of the ligands:
Strong field ligands
Weak field ligands
M3+
metal
ion
M2+
metal
ion
Weak field ligands cause spin free splitting, <∆0.
Strong field ligands cause spin paired splitting, >∆0.
Higher level
Spectrochemical series
d8 configurations with strong field
ligands (low spin) will look like this 
Hence with the greater repulsion on the lower energy orbitals, the
ligands will best line up on the x and y axis of the dx2-y2(square
planar) to be as far apart from each other as possible (rather than
dz2) or in a tetrahedral shape.
Higher level
Incomplete square planar explanation
4 lobes
2 lobes
6 lobes
Linear as z2 is slightly closer to nucleus
Square planar – d8 with strong ligand (Cu)
Tetrahedral as visually looks like the easiest acess to the nucleus with a 3D model
Octahedral – only symmetrical place for 6 ligands
Higher level
Other configurations
• Many complex ions are colored because the energy
differences between d orbitals match the energies of
components of visible light
• The colour absorbed can be determined by taking
the colour transmitted (observed) and finding the
opposite wavelength on the colour wheel (from data
booklet).
Yellow light transmitted
Violet light absorbed
Higher level
Colours
Corundum Al2O3
Higher level
Why?
Ruby Al2O3: 1% Cr3+
Higher level
Beryl Be3Al2Si6O18
Higher level
Why?
Emerald Be3Al2Si6O18: 1% Cr3+
Higher level
Higher level
Different colours of chromium – induce the
strength of the ligand…