Transcript Unit2_Periodicity_vs3

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Unit 2 – 2014-2015
Periodicity
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Lesson 1: Review of Periodic
Table Structure
Thursday, October 2nd
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Periodicity
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
IB Understandings

The periodic table is arranged into four blocks associated with the four
sublevels – s, p, d and f

The periodic table consists of groups (vertical columns) and periods
(horizontal columns)

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
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3.1 Applications and Skills

Applications and skills:

Deduction of the electron configuration of an atom from the element’s
position on the periodic table, and vice versa.

Guidance:

The terms alkali metals, halogens, noble gases, transition metals,
lanthanoids and actinoids should be known.

The group numbering scheme from group 1 to group 18, as
recommended by IUPAC, should be used.
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Periodic Table
•
The development of the periodic table brought a
system of order to what was otherwise an collection of
thousands of pieces of information.
•
The periodic table is a milestone in the development
of modern chemistry. It not only brought order to the
elements but it also enabled scientists to predict the
existence of elements that had not yet been
discovered .
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Early Attempts to Classify Elements
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 Dobreiner’s Triads
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•
(1827)
Classified elements in sets of three having similar
properties.
Found that the properties of the middle element were
approximately an average of the other two elements
in the triad.
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Dobreiner’s Triads
Element
Cl
Br
I
Atomic
Mass
(amu)
35.5
79.9
126.9
Ca
Sr
Ba
40.1
87.6
137.3
Average
Density
(g cm-3)
Average
81.2
1.56
3.12
4.95
3.25
88.7
1.55
2.6
3.5
2.53
Note: In each case, the numerical values for the
atomic mass and density of the middle element are
close to the averages of the other two elements 7
+ Newland’s Octaves -1863
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 John
Newland attempted to
classify the then 62 known
elements of his day.
 He
observed that when classified
according to atomic mass, similar
properties appeared to repeat for
about every eighth element
 His
attempt to correlate the
properties of elements with
musical scales subjected him to
ridicule.
 In
the end his work was
acknowledged and he was
vindicated with the award of the
Davy Medal in 1887 for his work.
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Dmitri Mendeleev
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Dmitri Mendeleev is
credited with creating
the modern periodic
table of the elements.
He gets the credit
because he not only
arranged the atoms, but
he also made
predictions based on his
arrangements His
predictions were later
shown to be quite
accurate.
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Mendeleev’s Periodic Table
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Mendeleev organized all of the elements into
one comprehensive table.
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Elements were arranged in order of increasing
mass.
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Elements with similar properties were placed in
the same row.
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.1
+ Mendeleev’s Periodic Table
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+ Mendeleev’s Periodic Table
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Mendeleev left some blank spaces in his periodic table. At the
time the elements gallium and germanium were not known. He
predicted their discovery and estimated their properties.
+ Periodic Properties
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 Elements
show gradual changes in certain
physical properties as one moves across a
period or down a group in the periodic table.
These properties repeat after certain intervals.
In other words they are PERIODIC
Periodic properties
include:
-- Ionization Energy
-- Electronegativity
-- Electron Affinity
-- Atomic Radius
-- Ionic Radius
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REVIEW: Periodic Table
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 Elements
are arranged by increasing atomic
number, Z
 Groups, or
vertical columns, group elements with
the same number of valence electrons and
therefore similar chemical properties
 Periods, or
rows, are horizontal groups; the period
number is equal to the principal quantum number,
n, of the highest occupied energy level of the
elements in that period
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Important Group To Remember
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Group Number Recommended Characteristics
name
1
Alkali metals
Very reactive metals
2
Alkaline metals
Less reactive metals
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Pnictogens
16
Chalcogens
17
Halogens
Very reactive non-metals
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Noble Gases
Non-reactive non-metals; stable octet
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Arrangement
 Metals
16
make up most of the periodic table and are
found on the left side
 Metalloids
include B, Si, Ge, As, Sb and Te and
separate metals and non-metals (Po and At are
sometimes considered metalloids); have
characteristics of both metals and nonmetals
 Non-Metals
are on the right side of the periodic
table after the metalloids
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17
Metals and Nonmetals
Metalloids
Transition
metals
+Additional Groupings in the
Periodic Table

Nonmetals, Metals, Metalloids, Noble gases
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IB Goodness!
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 You
have a large number of periodic tables in your
data booklet
 All
groups are numbered 1-18
 The
position of an element is related to its electron
configuration
 Know
the s-block, p-block, d-block and f-block;
those blocks represent which valence electrons
are getting filled
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Blocks
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Let’s Practice!
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Let’s Practice
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Lesson #2 – Review of
Periodic Table Families
Friday, October 3rd
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3.2 Periodic Trends
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IB Understandings

Vertical and horizontal trends in the Periodic Table exist for atomic
radius, ionic radius, ionization energy, electron affinity, and
electronegativity.
Guidance

Only examples of general trends across periods and down groups are
required. For ionization energy the discontinuities in the increase
across a period should be covered.
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Trends in metallic and non-metallic behaviour are due to the trends above.

Oxides change from basic through amphoteric to acidic across a period.
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Topic 3.2 Applications and Skills
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Applications and skills:

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).
Guidance

Group trends should include the treatment of the reactions of alkali
metals with water, alkali metals with halogens and halogens with halide
ions.

Construction of equations to explain the pH changes for reactions of
Na2O, MgO, P4O10, and the oxides of nitrogen and sulfur with water.
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Periodicity
Properties
of elements repeat themselves
periodically
The
Periodic Table is arranged to show
these trends
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Effective Nuclear Charge
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 Nuclear
charge is the number of protons in the
nucleus of an atom and increases steadily from left
to right on the Periodic Table
 HOWEVER, the
outer electrons are shielded from
the nuclear charge by the inner electrons and feel
less of this positive pull
 The
effective nuclear charge experience by the
outer electrons is less than the full nuclear charge
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Effective Nuclear Charge - Trends
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 Effective
nuclear charge INCREASES from left to
right across the Periodic Table (no change in
number of inner electrons)
 Effective
nuclear charge REMAINS THE SAME as
we go down a group
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Atomic Radius
 Atomic
radius is measured as ½ the distance between
neighboring nuclei
 Atomic
radii DECREASE from left to right across the
period and INCREASE down a group
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Atomic Radius (cont.)
Atomic
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radii increase down a group because
we are adding Principle Energy Levels
Atomic
radii decrease from left to right as
we increase effective nuclear charge
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Ionic Radius
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 Positive
ions (cations) have a smaller radius than the
parent atom; lose outer valence shell
 Negative
ions (anions) have a larger radius than the
parent atom; adding electrons to outer shell which
increases electron repulsion
 Atomic
radii decrease from Groups 1 to Group 14 for the
positive ions as we increase effective nuclear charge
 Atomic
radii decrease from Groups 14 to Group 17 for the
negative ions as we increase effective nuclear charge
 Ionic
radii increase down the group as number of energy
shells increases
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Let’s Practice
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 Describe
and explain the trend in radii of the
following atoms and ions: O2–, F–, Ne, Na+, and
Mg2+.

The ions and the Ne atom have 10 electrons and the electron
configuration 1s22s22p6. The nuclear charges increase with atomic
number:

O: Z = +8, F: Z = +9, Ne: Z = +10, Na: Z = +11, Mg: Z = +12

The increase in nuclear charge results in increased attraction between
the nucleus and the outer electrons. The ionic radii decrease as the
atomic number increases.
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Ionization Energies
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Ionization Energy Trends
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Ionization
energy increases across a period
as we increase effective nuclear charge
Ionization
energy decreases down a group
as we increase the shielding effect
Exceptions
to these trends can be explained
in terms of electron configuration stability
(i.e. it takes more energy to remove an
electron from a full sub-level or shell)
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Electron Affinity
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 The
first electron affinity of an element is the energy
change when one mole of electrons is added to one mole of
gaseous atoms to form one mole of gaseous ions:
X(g) + e– → X–(g)
 Values
 1st
are tabulated in Table 8 of the IB data booklet.
electron affinity is general exothermic as the electron is
attracted to the positive nucleus; 2nd and 3rd electron
affinities become endothermic as the atom already has a
negative charge!
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Electron Affinity Trends
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
Group 17 have the highest electron affinity

Group 1 has the lowest electron affinity

Group 2 and Group 15 have the highest electron affinities; here you are
adding the first electron to a half-filled sub-level; electrostatic
repulsion
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Electronegativity
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
The electronegativity of an element is a measure of the ability of its
atoms to attract electrons in a covalent bond

Electronegativity increases from left to right across a period owing to
the increase in nuclear charge, resulting in an increased attraction
between the nucleus and the bond electrons.

Electronegativity decreases down a group. The bonding electrons are
furthest from the nucleus and so there is reduced attraction.
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Metals vs. Non-Metals
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The ability of metals to easily conduct electricity is due to their low
ionization energy and ability to move their electrons away from the
nucleus

As you move from left to right on the Periodic Table, there is a slow
transition from metal to semi-metal to non-metal
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Lesson #3 – Review of
Periodic Table Families
Tuesday, October 7, 2014
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Nota Bene

Remember when studying that the group of an element on the
Periodic Table also tells you the number of valence electrons for that
element!
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+The Electron Shielding Effect
 Electrons
between the
nucleus and the
valence
electrons repel
each other
making the atom
larger.
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How is atomic radius measured?
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Half the distance between neighboring nuclei.
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
OR
Distance from the nucleus to the outermost electron.
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Atomic Radius Across a Period

Why does atomic radius decrease across a period?
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Think about the effective nuclear charge!
How many occupied energy levels does each atom have?
Fluorine’s radius is almost half of Lithium’s.
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Atomic Radius
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Atomic Radius
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
Extra Info:
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We typically measure atomic radius at bonding atomic radius where
we look at atoms in chemical bonds with other atoms of the same
element and take ½ the diameter between the two nuclei

There is also nonbonding atomic radius or van der Waals’ radius for
things like Noble Gases that do not bond; take a look at these atoms in
the solid phase and measure the distance between two nuclei
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Ionic Radius – Some More Info
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
The factors that effect ionic radius include nuclear charge, number of
filled energy shells, electrostatic repulsion and overall charge

At the HL level, the IB likes to ask questions about ionic radius that are
not clear-cut and require you to think

Which ion would have a larger ionic radius, Na+ or Mg2+?
Let’s think about this…we cannot directly compare these thinking
about trends from left to right because they have different charges!
But, we can recognize that both now have the electron configuration
of Neon. However, Mg has one more proton in the nucleus meaning
electrons are pulled in tighter!
+ Atomic Radius
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Trends in Ion Sizes
Radius in pm
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Effective Nuclear Charge

Extra Info (way beyond IB!)

The formula for effective nuclear charge is:
Zef = Z – S
Where s= screening or shielding constant
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Melting Points
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
Comparing melting points is complex because it depends on bonding
as well as nuclear charge

Trends down Groups 1 through 17 can be explained as elements in
each group bond in the same way!

Melting points DECREASE down Group 1 as these metallic bonds have
delocalized electrons and as you move down the group the attraction
decreases

Melting points INCREASE down Group 17 because these diatomic
elements are held together by London Dispersion forces which get
stronger as the electron cloud gets bigger

Melting points generally rise from left to right until Group 14 and then
fall from Group 14 to Group 18
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Melting Points
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Let’s Practice
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1 (a) Explain what is meant by the atomic radius of an element.
(b) The atomic radii of the elements are found in Table 9 of the IB data
booklet.
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(i) Explain why no values for ionic radii are given for the noble gases.

(ii) Describe and explain the trend in atomic radii across the Period 3
elements.
2 Si4+ has an ionic radius of 4.2 × 10–11 m and Si4– has an ionic radius of
2.71 × 10–10 m. Explain the large difference in size between the Si4+ and
Si4– ions.
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Answers
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
1 (a)Half the distance between the nuclei of neighbouring atoms of the
same element.

(b) (i) The noble gases do not form stable ions and engage in ionic
bonding so the distance between neighbouring ions cannot be
defined.

(ii) The atomic radii decrease from Na to Cl. This is because the
number of inner, shielding, electron is is constant (10) but the nuclear
charge increases from +11 to +17. As we go from Na to Cl, the
increasing effective nuclear charge pulls the outer electrons closer.

2. Si4+ has an electronic configuration of 1s22s22p6 where Si4– has an
electronic configuration of 1s22s22p63s23p6. Si4+ has two occupied
energy levels and Si4− has three and so Si4− is larger.
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Let’s Practice

More practice!
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Even more practice

Yay!
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Answers
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
(a) The electron in the outer electron energy level (level 4) is removed
to form K+. The net attractive force increases as the electrons in the
third energy level experience a greater effective nuclear charge.

(b) P3− has electronic configuration of 1s22s22p63s23p6 whereas Si4+
has an electronic configuration of 1s22s22p2. P3– has one more
principal energy level than Si4+ so its valence electrons will be further
from the nucleus and it will have a larger ionic radius.

(c) The ions have the same electron configuration, 1s22s22p63s23p6:
both have two complete shells; the extra proton in Na+ attracts the
electrons more strongly
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More Practice
Phosphorus exists as molecules with four
atoms: P4. Sulfur exists as molecules with
eight atoms: S8. There are stronger London
dispersion forces between the larger S8
molecules as there are
more electrons.
Cl- < Cl < Cl+
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More
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More Practice
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Lesson #4 – Vertical and
Horizontal Periodic Trends
Wednesday, October 8, 2014
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Warm-Up

Why does I2 have a higher melting point than F2?

Why does aluminum have a higher melting point than sodium?
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Metals
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
Metallic properties are related to ionization energy; the lower the
ionization energy, the more metallic an element is

A metallic structure consists of a regular lattice of positive ions in a sea
of delocalized electrons

Metallic character decreases from left to right and increases from top
to bottom
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Metals
1.
Are good conductors of heat and electricity
2.
Are malleable (capable of being hammered into thin sheets)
3.
Are ductile (capable of being draw into wires)
4.
Have lustre (they are shiny)
5.
Typically lose electrons; i.e. they like to be oxidized
** Note: Mercury is the only metal that is liquid at room temperature! It
can actually dissolve other metals and solutions formed this way are
called amalgams
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Non-Metals + Metalloids
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Non-Metals

Are poor conductors of heat and electricity

Typically gain electrons; i.e. they like to be reduced
Metalloids

Are semiconductors meaning they are able to conduct electricity only
at high temperature which has widespread applications in electronics
(e.g. Silicon Valley)
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Group 1 Alkali Metals

These metals are highly reactive, soft and have low melting points

Tend to form +1 ions

Held together by metallic bonds; as you go down Group 1 melting
point decreases because electrons are held less tightly meaning you
can disrupt the lattice more easily
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Group 1 Reactions
Reaction With Oxygen

All alkali metals react vigorously with oxygen to form an oxide
4Na(s) + O2(g)  2Na2) (s)

As you go down the group, the reactions become more vigorous
because of their lower ionization energy.
Reaction With Water

All alkali metals react rapidly with water; more reactive down the
group
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)

The products are highly basic!
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Group 1 Reactions With Water
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
Lithium floats and reacts slowly. It releases hydrogen but keeps its
shape.

Sodium reacts with a vigorous release of hydrogen. The heat produced
is sufficient to melt the unreacted metal, which forms a small ball that
moves around on the water surface.

Potassium reacts even more vigorously to produce sufficient heat to
ignite the hydrogen produced. It produces a lilac coloured flame and
moves excitedly on the water surface.

Cesium just EXPLODES!
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Group 17 Elements

Group 17 are known as the halogens

They all exist as diatomic elements; Cl2, I2, etc.

Reactivity decreases down the group (opposite of Group 1)
Reactions with Group 1:

All halogens react with Group I to form a salt (halides)
2Na(s) + Cl2(g)  2NaCl(s)
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Salts
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
Just as a reminder, salts are ionic compounds, positive and negative
ions held together by the electrostatic attraction

Salts have high melting points and only conduct electricity in the liquid
phase, gas phase or dissolved in water

More on this in the future!
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Displacement Reactions With
Halogens

More reactive halogen with displace a less reactive halogen
(REMEMBER from last year – single displacement reactions!)

Reactivity: F2 > Cl2 > Br2 > I2

Example:
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Cl2(g) +2KBr(aq)  2KCl(aq) + Br2(g)

All of these reactions are redox reactions, meaning one atom gains
elements and another atom loses electrons

You need to know that solutions of chlorine are green, bromine are
orange and iodine are violet (red-brown); solutions of the halides
are colorless
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Silver + Halides

Halogens form an insoluble (i.e. not able to be dissolved in water)
solution when combined with silver

Adding these together produces a precipitate
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Group 18 Noble Gases


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Group 18 consists of the Noble Gases which are all extremely
stable because they have a stable octet

They are colorless

They are monatomic; they exist as single atoms

They are very unreactive
Other elements all want to be the Noble Gases!

Elements in Groups 1, 2, and 13 lose electrons to adopt the arrangement of
the nearest noble gas with a lower atomic number.

Elements in Groups 15 to 17 gain electrons to adopt the electron
configuration of the nearest noble gas on their right in the Periodic Table.

The metalloids in the middle of the table show intermediate properties.
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Chemical Properties
1.
Why are alkali metals called alkali?
2.
How are halides formed?
3.
Why does Cl- displace Br- in a
displacement reaction?
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Chemical Properties in Groups
1,2,and 7
4. How do the reactivities of the alkali metals
and the halogens vary down a group?
5. Which property of the halogens increases
from fluorine to iodine?
A. ionic charge
B. electronegativity
C. melting point of the element
D. chemical reactivity with metals
+
Let’s Practice
75
+
Let’s Practice
76
+ Lesson 5 – Construction of
Oxide Equations and
Explaining pH Changes
Thursday, October 9th
77
Warm up: Explain why each trend occurs below in terms of effective nuclear charge:
+
+
Oxides
 What
is an oxide?
 Oxides
are formed from the combination of an
element with oxygen
+
pH of Oxides

Metal oxides are basic; they react with water to form metal
hydroxides
CaO(s) + H2O(l)  Ca(OH)2(aq)

Non-metallic oxides are acidic: they react with water to form acidic
solutions
CO2(g) + H2O(l)  H2CO3(aq) Carbonic Acid

If an oxide can act both as an acid and a base, it is classified as
amphoteric!
Al2O3(s) + 2NaOH(aq) + 3H2O(l)  2NaAl(OH)4 (acts as an acid)
Al2O3(s) + 6HCl(aq)  2AlCl3(aq) + 3H2O(l) (acts as a base)
80
+
Group 3 Oxides – Know the Trends!
81
Formula
of Oxide
Na2O
MgO
Al2O3
SiO2
P4O10
SO3 and
SO2
Metal/No
n-Metal
Metal
Metal
Metal
Nonmetal
Nonmetal
Nonmetal
Nature of
Oxide
Basic
Basic
Amphoter Acidic
ic
Acidic
Acidic
+
Metal Oxides
 The
transition from metallic to non-metallic character is
illustrated by the bonding of the Period 3 oxides. Ionic
compounds are usually formed between metals and
non-metal elements, so the oxides of elements Na to Al
have what we call super ionic structures.
+
Metalloid Oxides
 Covalent
 Giant
Bonds
Molecular Structure
 Ex: SiO2
Silicon dioxide does
not dissolve in water
but it is still classified
as acidic because it
can react with NaOH to
form Na2SiO3(aq)
+
Non-metal Oxides
Molecular
Ex: P4O10,
covalent bonds
SO3,
Cl2O7
+
Oxides of Period 3
Oxide
Formula
Na2O
(s)
MgO (s) Al2O3
(s)
SiO2 (s)
P4O10(s)
/P4O6
(s)
SO3(l)/S
O2(g)
Cl2O7(l)
/Cl2O
(g)
Oxidation
number
+1
+2
+3
+4
+5/+3
+6/+4
+7/+1
Electrical
conductivity
in molten
state
high
high
high
very low
none
none
none
Structure
giant ionic
giant
covalent
molecular covalent
+
Oxide Trends
 As
you go across period 3, the ionic character
decreases.
 As
you go down a group of oxides, the ionic character
increases.
 How
can you explain these trends in terms of
electronegativity differences within the oxide?
Li2O
Na2
O
K2O
Rb2
O
Cs2O
MgO
Al2O3
SiO2
P4O10
SO3
Cl2O7
87
+
Acid-Base Character
Formula
of oxide
Na2(s)
Acid-Base
character
basic
MgO(s)
Al2O3(s)
SiO2(s)
amphoteric
P4O10 (s) SO3(l)/S
/P4O6
O2(g)
(s)
acidic
Cl2O7(l)
/Cl2O
(g)
+
Basic Character
 Sodium
and magnesium oxides will dissolve in
water to form alkaline solutions:
Na2O (s) + H2O (l)  2NaOH (aq)
MgO (s) + H2O (l)  Mg(OH)2 (aq)
+
Basic Oxides
 Basic
oxides react with an acid to form a salt and
water:
MgO (s) + HCl (aq)  MgCl2 (aq) + H2O (l)
Li2O (s) + HCl (aq)  LiCl (aq) + H2O(l)
+
Acidic Oxides
 Non-metallic
oxides react with water to produce
acidic solutions:
P4O10(s) + 6H2O  4H3PO4(aq)
P4O6(s) + 6H2O  4H3PO3 (aq)
Acidic
Oxides
+
 Sulfur
trioxide reacts with water to produce
sulfuric acid:
SO3 (l) + H2O (l)  H2SO4 (aq)
 Sulfur
dioxide reacts with water to produce sulfic
acid
 Dichlorine
heptoxide reacts with water to produce
chloric acid
 Dichlorine
monoxide reacts with water to produce
chlorous acid

Silicon dioxide does not react with water, but will react with alkalis to
form silicates.
+
+
95
96
+
Lesson 7 – Intro to Transition
Metals
Tuesday, October 14
97
+
98
Mole Day Is Coming!
Let’s plan a party!
+
IB Understandings
99

Transition elements have variable oxidation numbers, form complex
ions with ligands, have coloured compounds, and display catalytic and
magnetic properties.

Guidance

Common oxidation numbers of the transition metal ions are listed in
the IB Data booklet in sections 9 and 14.

Zn is not considered to be a transition element as it does not form ions
with incomplete d orbitals.

Transition elements show an oxidation number of +2 when the s
electrons are removed.
+
Applications and Skills
100

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.
+
Periodic Trends
101

The d-block elements show a lull in the periodic patterns we have seen
in the s and p block

The 10 d-block elements all show similar characteristics to each other
(as an example look at the small range of atomic radii)
+
Electron Configuration

102
Any of the d-block elements in the same group as chromium (Cr)
and Copper (Cu) have electron configurations that are exceptions to
Aufbau’s principal because s electrons get promoted to either half-fill
or fully fill the d-sublevel
+
Effective Nuclear Charge
103

The small decrease in atomic radii across the d-block is due to the fact
that there is only a small increase in the effective nuclear charge as you
move left to right

Why is this?

Remember that the d-sublevel is not the valence shell! It is actually an
inner shell so each time you add a proton, you are also adding an
electron to an inner shell which provides some shielding effect!

Why is this amazing?

Since most of the d-block metals have similar radii, they can easily form
alloys or metal mixtures because mixing two different elements does
not disrupt the metallic bond structure!
+
Let’s Practice
104
+
Answer!
105
+
Physical Properties

High electrical and thermal conductivity

High melting point

Malleable – they are easily beaten into shape

High tensile strength – they can hold large loads without breaking •
ductile – they can be easily drawn into wires

Iron, cobalt, and nickel are ferromagnetic.
106
Properties can be explained by the fact that the s and d sublevel
electrons all delocalize to form a sea of electrons! Sea of electrons
hold metal lattice together.
+
Chemical Properties
107

Form compounds with more than one oxidation number • form a variety
of complex ions

Form coloured compounds

Act as catalysts when either elements or compounds.
+
Zinc

Why is zinc not considered to be a transition element? Discuss.
+
Zinc
109

Zinc only forms a +2 ion AND in both its natural state and as an ion it
has a completely filled d-sublevel

Scandium3+ (Sc3+) also forms a colorless solution because it has NO d
sublevel electrons but it is still a transition metal because it has an
incomplete d sublevel. It can also sometimes for the +2 ion!
+
Oxidation States
110

One of the key features of transition metals is their ability to form
multiple ions with various oxidation states!

This is different than alkali metals which only form +1 ions and alkaline
metals which only form +2 ions

Look at the different in ionization energies between Calcium, which
only forms a +2 ion, and Titanium, a transition metal. Ti is lacking the
big jump Ca has from the 2nd to 3rd Ionization Energy
+
Multiple Oxidation States
111

This is because in Calcium, once you remove 2 electrons, you get to a
noble gas configuration; removing one more electron makes that ion
very unstable!

In Titanium, the 3d and 4s electrons are very close in energy so it can
lose from both sublevels until they are all gone; the jump doesn’t
happen until the next electron lost would have to be from the 3p
sublevel!
+
Oxidation States

What is the most common oxidation state of the transition
elements? Look at the red!
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
+1
+2
+2
+2
+2
+2
+2
+2
+2
+2
+3
+3
+3
+3
+3
+3
+3
+3
+3
+4
+4
+4
+4
+4
+4
+4
+5
+5
+5
+5
+5
+6
+6
+6
+7
+2
+
Important Points For IB!!
113
 All
the transition metals show both the +2 and +3 oxidation
states. The M3+ ion is the stable state for the elements from
scandium to chromium, but the M2+ state is more common
for the later elements. The increased nuclear charge of the
later elements makes it more difficult to remove a third
electron.
 The
maximum oxidation state of the elements increases in
steps of +1 and reaches a maximum at manganese. These
states correspond to the use of both the 4s and 3d electrons
in bonding. Thereafter, the maximum oxidation state
decreases in steps of –1.
+
Important Points For IB!!
 Oxidation
states above +3
generally show covalent
character. Ions of higher charge
have such a large charge
density that they polarize
negative ions and increase the
covalent character of the
compound (see Figure 3.13).
 Compounds
with higher
oxidation states tend to be
oxidizing agents. The use of
potassium dichromate(VI)
(K2Cr2O7), for example, in the
oxidation of alcohols.
114
+
Let’s Practice!
115
+
Answers!
116
+
Next Week
Quiz
on Tuesday!!! Short + review
Complex
Ions – Wednesday
D-Sublevel
Colored
Exam
117
Split – Thursday
Ions – Friday
the following week!!! Thursday,
October 23rd tentatively unless we need one
more day of review!
+
Lesson 8 – Complex Ions and
Coordinate Complexes
Tuesday, October 14th, 2004
118
+
Complex Ions
Transition
119
metal ions in solution have a high
charge density
These
ions attract polar water molecules
which form coordinate bonds
REMBEMBER: A
coordinate bond is a bond
where one atom provides both of the
electrons to the covalent bond
+
Ligands
120

When a complex is formed where a central ion is surrounded by
molecules or ions which possess a lone pair of electrons that can enter
into a coordinate bond, the surrounding species are called ligands

Ligands have to have a lone pair of electrons to donate!

The number of coordinate bonds from the ligands to the central atom is
called the coordination number
6 ligands attached
to the central ion!
Coordination
number = 6
+
Shapes of Complex Ions
121

These complex ions form a variety of 3-dimensional shapes

Think back to VSEPR (electron cloud repulsion) and the shapes of
covalent molecules; electron clouds want to be as far apart as possible
+
Shapes of Complex Ions
Coordination
Number
122
Oxidation
Shape
Number of
Central Ion
6
+3
Octahedral
4
+7
Tetrahedral
4
+2
Square Planar
4
+2
Tetrahedral
2
+1
Linear
Yikes! Why?! There is no easy way to predict
whether it will be square planar or tetrahedral. Sad!
Most are tetrahedral…that’s all I can tell ya!
+
Oxidation Number of Metal and
Charge of Ion

123
Don’t cheat! Let’s try to figure out the oxidation state of the metal
in each of these ions…
Let’s look at Fe(H2O)63+ first… The overall charge is +3. H2O does
not have a charge itself so the charge of the iron must be +3.
+3 + 0 = +3
Now, let’s look at [CuCl4]2-. The chlorine here is chlorine ion
which always has a charge of -1. We have four of them so they
contribute an overall charge of -4. The total charge on the ion is 2 so…
X + (-4) = -2, X= +2, the oxidation state of the copper
Let’s try the rest!
HINT: CN-1 is a polyatomic ion with a charge of -1.
+
Oxidation Number of Central Ion
1.
Ligands can either be neutral (H2O, NH3, etc.) or negatively
charged (Cl-, CN-, etc.). Decide the ligands’ charges
2.
Add up all the charges on the ligands
3.
Set up the following equation:
124
X (oxidation # of trans. Metal) + Y (sum of charges of ligands) = Z (overall
charge)

HINT: The complex ion will always be kept in brackets [ ]!



[Fe(H2O)6]2+
[Ni(CN)4]2-
 Fe has a +2 oxidation #
 Ni has a +2 oxidation #
Hint 2: These can enter into ionic compounds! Reverse criss-cross

Li2[Ni(CN)4]  I know Lithium forms a +1 ion so the complex ion has to have
a -2 charge (+2 + -2 = 0)
+
Examples
125
+
Determining Charge of Complex Ion
126

When a complex ion bonds with another ion, we can use the
reverse criss-cross method to determine the charge on that ion!

EXAMPLE: What is the charge for the complex ion in compound?
[Cr(H2O)6]Cl3?

So, we know that chlorine always has a charge of -1. Since the overall
charge for this compound is zero (I don’t see a charge on the outside), I
know that the charge of the ion has to be +3 to balance out the 3 -1
charges from the chlorine. Let’s try two more!

What is the charge for the complex ion in the following compounds?
[CrCl(H2O)5]Cl2.H2O and [CrCl2(H2O)4]Cl.2H2O

+2 and +1
+
Let’s Practice

Platinum (II) can form a complex ion with 1 ammonia and 3
chloride ligands. What is the overall charge and formula for this
complex ion?

Platinum forms a +2 ion. Ammonia is neutral and chloride ions each
have a -1 charge. So, +2 + (3*-1) = -1 overall charge

[Pt(NH3)Cl3]-1
127
+
Monodentate vs. Polydentate

Monodentate ligands contain a single donor atom and have one lone
pair contributing to the coordinate bond in a complex


128
E.g. Cl- and H2O
Polydentate (chelate) ligands contain two or more donor atoms that
form coordinate bonds with a transition metal center.

E.g. 1,2-ethanediamine (en)

Ethanedioate (ox)
+
Chelating Agents
129

If something has more than one lone pair of electrons and can
coordinate bond with a central metal ion in more than one place,
it is called a polydentate ligand

EDTA4– (old name ethylenediaminetetraacetic acid) is an example of a
polydentate ligand as it has six atoms (two nitrogen atoms and four
oxygen atoms) with lone pairs available to form coordinate bonds.
1
3
2
6
4
5
+
Chelating Agents
130

Since compounds like EDTA4+ can grip the central atom in more
than one place, it is called a chelate

The resulting complex ions formed with chelates are very energetically
stable, more stable than if individual ligands are bound to the central
ion
+
EDTA Usage
Removal of heavy metals
1.

Example: lead poisoning
Chelating Therapy
2.

Reduces calcium ions in the blood and can remove gunk from hardened
arteries
Water Softening
3.

Getting rid of metal ions in water
Food preservation
4.

Gets rid of metal ions that would otherwise spoil food
5.
Restoring metal scultures
6.
Cosmetics - presevative
131
+
Use In Medicine and Food
132

Chelating agents are use widely in food because they remove
transition metals by forming very energetically stable complex ions;
these transition metals would otherwise catalyze oxidation reactions
and spoil the food

Chelating agents are also used in medicine to remove toxic metals in
the body
+
Let’s Practice
133
+
Answers
134
+
Lesson 9 – d-Block Elements
as Catalysts and Magnets
Wednesday, October 15, 2014
135
+
Catalysts
136

A catalyst alters the rate of reaction by providing an alternate
pathway with a lower activation energy

Catalysts are extremely important as they allow reactions to proceed
faster

In a homogeneous catalyst, the catalyst is in the same state as the
reactants

In a heterogeneous catalyst, the catalyst is in a different state as the
reactants

Transition metals are super effective heterogeneous catalysts since
they can use their s and d electrons to weakly interact with reactant
molecules and arrange them in the correct orientation (remember that
in order for a reaction to occur, the reactants must combine with
enough energy in the correct orientation!)
+
Examples of Transition Metals as
Heterogeneous Catalysts

Iron – Used in the Haber process to make ammonia
Fe
N2(g) + 3H2(g)  2NH3(g)

Nickel – Converts alkenes to alkanes

Palladium and Platinum (and Nickel) – Used in cars’ catalytic
Pd or Pt or Ni
converters
2CO(g) + 2NO(g)  2CO2(g) + N2(g)

MnO2 – Decomposition of H2O2 (we did this last year in lab!)
MnO2
2H2O2(l)  2H2O(l) + O2(g)

V2O5
V2O5 – Used in Contact process - 2SO2(g) + O2(g)  2SO3(g)
137
+
Hydrogenation of Oil

138
Nickel can be used to add hydrogen to unsaturated oils
RCH=CHR + H2  RCH2CH2R

This allows these oils to be solid at room temperature and is useful for
cooking

However, this process forms unhealthy trans fats which are not
metabolized correctly in the body and leads to cardiovascular
problems!
+
Examples of Transition Metals as
Homogeneous Catalysts
139

The ability of transition metals to show various oxidation states
makes them effective homogeneous catalysts in redox reactions

Many of these reactions are important in Biochemistry!

Fe2+ in heme  Oxygen gets transported throughout the body by
forming a weak bond the the iron ion held in the center of the
hemoglobin protein! Why would you want a weak bond here?

Co3+ in Vitamin B12  Part of the vitamin B12 molecule consists of an
octahedral Co3+ complex. Vitamin B12 is needed for the production of
healthy red blood cells and for a healthy nervous system

In a biological setting, catalysts are called enzymes!
+
Catalytic Converters in Cars
140

In a running car engine, nitrogen gas in the air and oxygen can
react in the high heat to form nitrogen monoxide

When NO is released into the air, it combines with more oxygen to form
NO2

This NO2 pollutes the air forming smog and can also react with water to
form acid rain! (Think about our oxide trends!!!)

Also, carbon monoxide can be formed – CO

Therefore these gases pass through a catalytic converter – often
Pd, Pt or Rd, which helps convert these dangerous high energy
molecules into more stable ones!
+
Magnetism
141

Every spinning electron can behave like a tiny magnet

However, when electrons are paired and have opposite spins, their
spins cancel out this magnetic effect

Some transition metals are unusual because of the number of unpaired
electrons they have and when aligned create highly magnetic
substances!
+
Types of Magnetism

Diamagnetism is a property of all materials and produces a very
weak opposition to an applied magnetic field.

Paramagnetism, which only occurs with substances which have
unpaired electrons, is stronger than diamagnetism. It produces
magnetization proportional to the applied field and in the same
direction.

Ferromagnetism is the largest effect, producing magnetizations
sometimes orders of magnitude greater than the applied field.
142
* Note: all substances with paired electrons exhibit diamagnetism but the
effect is much smaller than paramagnetism and certainly much smaller
than ferromagnetism!**
+
Transition Metals
143

Iron, nickel, and cobalt are ferromagnetic; the unpaired d electrons in
large numbers of atoms line up with parallel spins in regions called
domains.

Paramagnetism increases with the number of unpaired electrons so
generally increases from left to right across the Periodic Table,
reaches a maximum at chromium, and decreases. Zinc has no unpaired
electrons and so is diamagnetic.
+
Let’s Practice
144
Fe2+
+
Let’s Practice
145
+
Quiz
Thursday, October 16, 2014
146
+
Lesson 10 – d-Sublevel Splits
Friday, October 17, 2014
147
+
IB Understandings
148
Understandings:

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.
Guidance

The relation between the colour observed and absorbed is illustrated
by the colour wheel in the IB Data booklet in section 17.
+
Applications and Skills

149
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.
Guidance

Students are not expected to recall the colour of specific complex ions.

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.

The spectrochemical series is given in the IB data booklet in section 15. A
list of polydentate ligands is given in the data booklet in section 16.
Students are not expected to know the different splitting patterns and their
relation to the coordination number. Only the splitting of the 3-d orbitals in
an octahedral crystal field is required.
+
Crystal Field Theory (CFT)

The d-sublevel consists five orbitals – dxy, dyz, dxz, dx2-y2, dz2
150
+
Colored Ions

The color of the ions is related to the presence of partially filled dsublevel orbitals

Any ions that do not have any electrons in the d sublevel will not be
colored!
151
+
Visible Spectrum

152
The visible spectrum
ranges from about 400nm to 700nm. The
color we see depends on the wavelength
The color
depends on
which colors
are absorbed
and which
colors are
transmitted or
reflected!!
+
Light
 White
light contains all
wavelengths in the visible
spectrum
 Transition
metals absorb some
light and transmits others
 The
light emitted is the composite
of the wavelengths
153
+
D-Level Splits
154

But, why do these transition metals absorb light?!

In an isolated transition metal, the d-orbitals are said to be degenerate
or have the same energy

However, when ligands bond with the central atom, an electric field is
produced which causes the d sublevel to split into two – a higher and a
lower energy level

When light passes through these d sublevels, the d electrons jump from
the lower energy d orbitals to the higher energy d orbitals
155
+
Energy Level Splits
156
The energy separation between the orbitals is ∆E and hence the
color of the complex depends on the following factors:
1.
the nuclear charge and the identity of the central metal ion
2.
the charge density of the ligand
3.
the geometry of the complex ion (the electric field created by the
ligand’s lone pair of electrons depends on the geometry of the
complex ion)
4.
the number of d electrons present and hence the oxidation number of
the central ion.
+
Nuclear Charge
157

The strength of the bond between the ligand and the ion depends on
the electrostatic attraction

Ligands interact more efficiently the higher the nuclear charge

Example: [Mn(H2O)6]2+and [Fe(H2O)6]3+ both have the same electron
configuration but the iron nucleus has a higher nuclear charge and so
has a stronger interaction with the water ligands

Mn absorbs in the green region and therefore look pink in solution

Fe absorbs in the higher energy blue region and therefore look
yellow/brown
+
Charge Density of Ligand
158

The greater the charge density of the ligand, the greater the energy
level split between the d orbitals and the higher energy that is
absorbed

The spectrochemical series arranges the ligands according to the
energy separation, ∆E, between the two sets of d orbitals

The wavelength at which maximum absorbance occurs, λmax,
decreases with the charge density of the ligand, as shown in the table
below and in Section 15 in your IB Booklet!!
+
Geometry of the Complex

159
The coordination number and geometry of the complex ion also affects
the color
+
Number of d electrons and oxidation
state of central ion
160

The strength of the interaction between the ligand and the central
metal ion and the amount of electron repulsion between the ligand and
the d electrons depends on the number of d electrons and hence the
oxidation state of the metal.

EXAMPLE: [Fe(H2O)6]2+ absorbs violet light and so appears
green/yellow, whereas [Fe(H2O)6]3+ absorbs blue light and appears
orange/brown.
+
Let’s Practice!

State the formula and the shape of the complex ion formed in the
following reactions.
(a) Some iron metal is dissolved in sulfuric acid and then left exposed to air
until a yellow solution is formed.
(b) Concentrated hydrochloric acid is added to aqueous copper sulfate
solution to form a yellow solution.
(c) A small volume of sodium chloride is added to aqueous silver nitrate
solution. The white precipitate dissolves to form a colourless solution when
ammonia solution is added.
161
+
Answer
162

(a) [Fe(H2O)6]3+
The oxidation number is +3 as the complex is left exposed to air. The
shape is octahedral as the coordination number = 6 (see left).

(b) The complex [CuCl4]2– is yellow.
The shape is tetrahedral as the coordination number = 4 (see left).

(c) NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)
The complex [Ag(NH3)2]+ is linear as the coordination number is 2.
[H3N—Ag—NH3]+
163
164
165
+
Lesson 11 – Colored Ions
Wednesday, October 22nd
166
+The D Block Colored Compounds
 In
an isolated atom all of the d sublevel electrons have
the same energy.
 When
an atom is surrounded by charged ions or polar
molecules, the electric field from these ions or molecules
has a unequal effect on the energies of the various d
orbitals and d electrons.
 The
colors of the ions and complex ions of d block
elements depends on a variety of factors including:
 The particular element
 The oxidation state
 The kind of ligands bound to the element
Various oxidation
states of Nickel (II)
+
Colors in the D Block

The presence of a partially filled d sublevels in a transition
element results in colored compounds.

Elements with completely full or completely empty subshells
are colorless,

For example Zinc which has a full d subshell. Its compounds are
white

A transition metal ion exhibits color, if it absorbs light in the
visible range (400-700 nanometers)

If the compound absorbs a
particular wavelengths of light its
color is the composite of those
wavelengths that it does not absorb.

It shows the complimentary color.
.16
+Colors and d Electron Transitions

The d orbitals may split into two groups so that two orbitals are
at a lower energy than the other three

The difference in energy of these orbitals varies slightly with
the nature of the ligand or ion surrounding the metal ion

When white light passes through a compound of a transition
metal, light of a particular frequency is absorbed as an electron
is promoted from a lower energy d orbital to a higher one.

When the energy of the transition: ∆E =hn may occur in the
visible region, the compound is colored
+ Magnetic Properties

Paramagnetism --- Molecules with one or
more unpaired electrons are attracted to a
magnetic field. The more unpaired
electrons in the molecule the stronger the
attraction. This type of behavior is called

Diamagnetism --- Substances with no
unpaired electrons are weakly repelled by
a magnetic field.

Transition metal complexes with unpaired
electrons exhibit simple paramagnetism.

The degree of paramagnetism depends
on the number of unpaired electrons
.17
+ Catalytic Behavior

Many D block elements are catalysts for
various reactions

Catalysts speed up the rate of a
chemical reaction with out being
consumed.

The transition metals form complex ions
with species that can donate lone pairs of
electrons.

This results in close contact between the
metal ion and the ligand.

Transition metals also have a wide variety
of oxidation states so they gain and lose
electrons in redox reactions
.
+Some Common D Block Catalysts
 Examples
of D block elements that are
used as catalysts:
1. Platinum or
rhodium is used in a
catalytic converter
2. MnO2 catalyzes the
decomposition
of hydrogen peroxide
3. V2O5 is a catalyst for
the contact process
4. Fe in Haber process
5. Ni in conversion of
alkenes to alkanes
.17
+
Lesson 12 – Review
Tuesday, October
173
+
Hints

Know the EXACT definition of electronegativity

Remember that silicon forms network solids; these have VERY high
melting points!

Any elements in Group 12 (where Zinc lives) are not considered
transition metals
174