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Periodic Trends in Atomic Properties
UNIT 3: Periodicity
IB Topics 3 & 13
An Introduction to the
Periodic Table
DMITRI
MENDELEEV IS
OFTEN REGARDED
AS THE FATHER
OF THE PERIODIC
TABLE,
MENDELEEV
HIMSELF CALLED
HIS TABLE, OR
MATRIX, THE
PERIODIC SYSTEM.
2–3
Mendeleev’s Periodic Table
Dmitri Mendeleev
Modern Russian Table
Chinese Periodic Table
Stowe Periodic Table
This is Timothy Stowe's physicists’ periodic table. This table depicts periodicity
in terms of quantum numbers.
A Spiral Periodic Table
This example was devised by Theodor Benfey and depicts the elements as a seamless
series with the main group elements radiating from the center with the d- and felements filling around loops.
Triangular Periodic Table
Based on the work of Emil Zmaczynski and graphically reflects the process of the
construction of electronic shells of atoms.
“Mayan”
Periodic
Table
Giguere Periodic Table
This is a 3-D periodic table constructed by Paul Giguere based primarily on the
electronic structure of the atoms. The four main groups of elements are separated
according to the type of atomic orbitals being filled. This animated graphic may take
a few moments to load.
Periodic Table
Periodic Table
• Group: elements with same number of valence
electrons and therefore similar chemical and physical
properties; vertical column of elements (“family”)
• Period: elements with same outer shell; horizontal
row of elements
• Periodicity: regular variations (or patterns) of
properties with increasing atomic weight. Both
chemical and physical properties vary in a periodic
(repeating pattern) across a period.
Periodic Table
From the IB Data Booklet…
0 or 8
Transition metals
Lanthanides
Actinides
Another name for “metalloid” is “semi-metal”.
alkali metals
alkaline earth metals
lanthanides
actinides
halogens
noble gases
Transition metals
The Periodic Table
Copyright © Houghton Mifflin Company. All
rights reserved.
2–20
PERIODIC GROUPS
•
•
•
•
•
•
•
alkali metals
alkaline earth metals
transition metals
halogens
noble gases
lanthanides
actinides
Periodic Table
Metals
Conductors
Lose electrons
Malleable and ductile
Nonmetals
Brittle
Gain electrons
Covalent bonds
Semi-metals or Metalloids
Diatomic Elements
H2
He
Li Be
B C N2 O2 F2 Ne
Na Mg
Al Si P S Cl2 Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br2 Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I2 Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Alkali Metals
• Group 1 on periodic table
• Called alkali metals because they all react
with water to form an alkali solution of the
metal hydroxide and hydrogen gas.
– i.e. Na(s) + H2O(l)  NaOH(aq) + H2(g)
• React by losing outer electron to form the
metal ion
• Good reducing agents (because they can
readily lose an electron)
Chemical Properties
• Group 1 – The alkali Metals
Silvery metals
To reactive to be found in nature
Usually stored in oil to prevent contact with air or water
Good conductors of electricity
Have low densities
For ionic compounds with non-metals
Chemistry of the Groups
The1 alkali metals (Group 1)
H
1
Li
3
Na
11
K
19
- The alkali metals are lithium (Li), sodium (Na), potassium
(K), rubidium (Rb), cesium (Cs), and francium (Fr).
- Hydrogen is placed in Group 1 but is not a metal.
- The alkali metals react readily with nonmetals to give ions
with a +1 charge.
- Compounds of alkali metals are common in nature and
daily life.
Rb
37
Cs
55
Fr
87
Copyright 2007 Pearson Benjamin Cummings. All rights reserved.
Alkali Metals, Group 1
H
He
Li Be
B C N O F Ne
Na Mg
Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Alkaline Earth Metals
•
•
•
•
Group 2 on periodic table
Abundant metals in the earth
Not as reactive as alkali metals
Higher density and melting point than
alkali metals
Chemistry of the Groups
The alkaline earths (Group 2)
2
Be
4
Mg
12
- The alkaline earths are beryllium (Be), magnesium (Mg),
calcium (Ca), strontium (Sr), barium (Ba), and
radium (Ra).
- All are metals that react readily with nonmetals to give
ions with a 2 charge.
Ca
20
Sr
38
Ba
56
Ra
88
Copyright 2007 Pearson Benjamin Cummings. All rights reserved.
Chemical properties
• Group 2 – Alkaline Earth Metals
• The alkaline earth metals are silvered colored, soft
metals. Elements classified as Alkaline Earth Metals
are all found in the Earth’s crust, but not in the
elemental form as they are so reactive. Alkaline metals
are usually shiny solids that conduct heat or electricity
and can be formed into sheets.
Alkaline Earth Metals, Group 2
H
He
Li Be
B C N O F Ne
Na Mg
Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Transition Metals
•
•
•
•
Variable valency & oxidation state
Forms colored compounds
Forms complex ions
Catalytic behavior
– Need to know these well, but more on all
this later…
HL Transition Elements
•
•
•
•
•
•
Show lull in periodic patterns
Small Radii
Small Ionization energy
More then one oxidation number
Forms complex ions
Zn is not transition element
Transition metals
Chemistry of the Groups
Group 13
– Of the Group-13 elements, only the lightest, boron, lies on the
diagonal line that separates nonmetals and metals, it is a
semimetal and possesses an unusual structure.
– The rest of Group 13 are metals (aluminum, gallium, indium, and
thallium) and are typical metallic solids.
– Elements of Group 13 are highly reactive and form stable
compounds with oxygen.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Chemistry of the Groups
Group 14
– Group 14 elements straddle the
diagonal line that divides nonmetals from
metals.
– Carbon is a nonmetal, silicon and
germanium are semimetals, and tin and
lead are metals.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Chemistry of the Groups
Group 15, the Pnicogens
15
N
– The pnicogens are nitrogen, phosphorus, arsenic, antimony, and
bismuth.
7
P
15
As
33
Sb
1. –3, in which three electrons are added to give the closedshell electron configuration of the next noble gas
2. +5, in which all five valence electrons are lost to give the
closed-shell electron configuration of the preceding noble
gas
51
Bi
83
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Pnicogens, Group 15
H
He
Li Be
B C N O F Ne
Na Mg
Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Chemistry of the Groups
Group 16, the Chalcogens
16
O
8
– The chalcogens are oxygen, sulfur, selenium, tellurium, and
polonium.
1. –2, in which two electrons are added to achieve the
closed-shell electron of the next noble gas.
S
16
Se
34
2. +6, in which all six valence electrons are lost to give the
closed-shell electron configuration of the preceding noble
gas.
Te
52
Po
84
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Chalcogens, Group 16
H
He
Li Be
B C N O F Ne
Na Mg
Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Halogens
• Group 7 on periodic table
– Referred to as group 17 or group VIIA on
other periodic tables
• React by gaining one more electrons to
from halide ions
• Good oxidizing agents (because they
readily gain an electron)
Halogens, Group 17
H
He
Li Be
B C N O F Ne
Na Mg
Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Chemistry of the Groups
The halogens (Group 17)
17
F
9
Cl
- The halogens are fluorine (F), chlorine (Cl), bromine (Br)
iodine (), and astatine (At).
- They react readily with metals to form ions with a 1
charge.
17
Br
35
Element
At. Mass
Normal Form at STP
Fluorine
F
19.0
F2
Chlorine
Cl
35.5
Cl2 greenish-yellow gas
53
Bromine
Br
79.9
Br2 red-brown liquid
At
Iodine
I
126.9
85
Astatine
At
(210)
I
Copyright 2007 Pearson Benjamin Cummings. All rights reserved.
I2
pale-yellow gas
black solid (m.p.113oC)
b.p., oC
-187.0
-34.5
58.0
184.0
Noble Gases
• Group 0 (or 8) on periodic table
– Referred to as group 18 or group VIIIA on other
periodic tables
• Sometimes called rare gases or inert gases
• Relatively nonreactive
• Gases at room temperature
Chemistry of the Groups
The noble gases (Group 18)
18
He
2
Ne
10
Ar
- are helium (He), neon (Ne), argon, (Ar), krypton (Kr), xenon
(Xe), and radon (Rn);
- are monatomic;
- are unreactive gases at room temperature and pressure;
- are called inert gases.
18
Kr
36
Xe
54
Rn
86
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Noble Gases, Group 18
H
He
Li Be
B C N O F Ne
Na Mg
Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Lanthanides
•
•
•
•
•
Part of the “inner transition metals”
Soft silvery metals
Tarnish readily in air
React slowly with water
Difficult to separate because all have two
effective valence electrons
Lanthanide Series
H
He
Li Be
B C N O F Ne
Na Mg
Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
La
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Actinides
• Radioactive elements
• Part of the “inner transition metals”
Actinide Series
H
He
Li Be
B C N O F Ne
Na Mg
Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
La
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
La
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
PERIODIC TRENDS
• Properties that have a definite trend as
you move through the Periodic Table
– Valence Electrons
– Atomic radii
– Ionic radii
– Electronegativity
– Ionization energy
– Melting points
Many periodic trends can be explained by…
1. How many protons are in the nucleus
(nuclear charge)
2. How far the outer electrons are from
the nucleus (number of shells)
3. How many core electrons are between
the outer electrons and the nucleus
(amount of shielding)
The periodic table is full of
repeating patterns.
Hence the name periodic table.
The electrons in the outermost electron
shell (highest energy level) are called
valence electrons.
“vale” =
Latin to be strong
Most chemical reactions occur as valence
electrons seek a stable configuration.
(full energy level, and to a lesser extent ½ full)
Non Valence electrons are called “core
electrons”. Core electrons are relatively
stable.
Consider Hydrogen v. Helium…
+1
Hydrogen
electron
Consider Hydrogen v. Helium…
+1
electron
Hydrogen
+2
Helium
Greater (+) Charge in
the nucleus produces
stronger attraction.
All orbitals get smaller as protons
are added to the nucleus.
Also…
Electrons are attracted to
the nucleus, but are repelled
by the “screening” core
electrons.
+
If the electron is further
away, the attraction force is
much less.
+
Electron attraction to the nucleus depends on…
1. How many protons are in the nucleus
2. How far the electron is from the nucleus
3. How many core electrons are between
the electron and the nucleus
Atomic Radii
• Trend: increases down a group
• WHY???
– The atomic radius gets bigger because
electrons are added to energy levels farther
away from the nucleus.
– Plus, the inner electrons shield the outer
electrons from the positive charge (“pull”) of
the nucleus; known as the SHIELDING
EFFECT
Atomic Radii
• Trend: decreases across a period
• WHY???
– As the # of protons in the nucleus
increases, the positive charge increases
and as a result, the “pull” on the electrons
increases.
Atomic Size
• First problem where do you start measuring
• The electron cloud doesn’t have a definite edge.
• They get around this by measuring more than 1 atom
at a time
Atomic Size
}
Radius
• Atomic Radius = half the distance between two
nuclei of a diatomic molecule
Atomic Radii
The atomic radius is the distance from the nucleus
to the outermost electron.
Since the position of the outermost
electron can never be known
precisely, the atomic radius is
usually defined as half the distance
between the nuclei of two bonded
atoms of the same element.
Thus, values not listed in IB data
booklet for noble gases.
Trends in Atomic Size
•
•
•
•
•
Influenced by two factors
Shielding
More shielding is further away
Charge on nucleus
More charge pulls electrons in closer
Group trends
• As we go down a
group
• Each atom has
another energy
level
• So the atoms get
bigger
H
Li
Na
K
Rb
Periodic Trends
• As you go across a period the radius
gets smaller.
• Same energy level
• More nuclear charge
• Outermost electrons are closer
Na
Mg
Al
Si
P
S Cl Ar
Ionic size
•
•
•
•
Anions form by gaining electrons
Anions are bigger than the atom they come from
Nonmetals form anions
Anions of representative elements have noble gas
configuration.
Ionic Radii
• Cations (+) are always
smaller than the metal
atoms from which they
are formed. (fewer
electrons than protons
& one less shell of e’s)
• Anions (-) are always
larger than the
nonmetal atoms from
which they are formed.
(more electrons than
protons)
© 2002 Prentice-Hall, Inc.
Ionic Radii
• Trend: For both cations and anions, radii increases
down a group
• WHY???
– Outer electrons are farther from the nucleus (more shells/
energy levels)
Atomic
Radii for
Selected
Atoms
Copyright © Houghton Mifflin Company. All rights reserved.
7–76
For the tall groups the column number
is the number of valence electrons.
The Orbitals Being Filled for Elements in Various
Parts of the Periodic Table
The Values of First Ionization Energy for the
Elements in the First Six Periods
Copyright © Houghton Mifflin Company. All rights reserved.
7–79
Successive Ionization Energies (KJ per Mole)
for the Elements in Period 3
Copyright © Houghton Mifflin Company. All rights reserved.
7–80
Ionic Size
•
•
•
•
Cations form by losing electrons
Cations are smaller than the atom they come from
Metals form cations
Cations of representative elements have noble gas
configuration.
Electron Affinity
• The energy change associated with adding
an electron to a gaseous atom
• High electron affinity gives you energy• exothermic
• More negative
• Increase (more - ) from left to right
– greater nuclear charge.
• Decrease as we go down a group
– More shielding
Electronegativity
• Definition: a relative measure of the attraction
that an atom has for a shared pair of electrons
when it is covalently bonded to another atom.
Electronegativity
• Trend: decreases down a group
• WHY???
– Although the nuclear charge is increasing,
the larger size produced by the added
energy levels means the electrons are
farther away from the nucleus; decreased
attraction, so decreased electronegativity;
plus, shielding effect
Effective Nuclear Charge
• http://www.youtube.com/watch?v=IvSmfgxCSNQ
Electronegativity
• Trend: increases across a period
(noble gases excluded!)
• WHY???
– Nuclear charge is increasing, atomic radius is
decreasing; attractive force that the nucleus
can exert on another electron increases.
Copyright © Houghton Mifflin Company. All rights reserved.
7–87
Copyright © Houghton Mifflin Company. All rights reserved.
7–88
Electronegativity
• Is the measure of the attraction that an atom has for a
shared pair of electrons when it is covalently bonded
to another atom
• When the size of the atom decreases the
electronegativity increases
• Value increases from left to right, and decreases down
a group
89
Melting points
•
•
•
•
Melting points decrease down Group 1
Melting points increase down Group 17
Melting points generally rise across a period.
All the Period 3 elements are solids at room temp
except Chlorine and Argon
Melting Point
C
B
Si
Be
Ca
Mg Al
Li
NO F
H
He
Na
Ne
P S
Cl
K
Ar
Melting Point
Melting points depend on both…
1.The structure of the element
2.Type of attractive forces holding the atoms
together
Melting Point
Trend (using per 3 as an example):
• Elements on the left exhibit metallic bonding
(Na, Mg, Al), which increases in strength as
the # of valence electrons increases.
Melting Point
Trend (using per 3 as an example):
• Silicon in the middle of the period has a
macromolecular covalent structure (network) with
very strong bonds resulting in a very high melting
point.
Melting Point
Trend (using per 3 as an example):
• Elements in groups 5, 6 and 7 (P4, S8 and Cl2) show
simple molecular structures with weak van der Waals’
forces (more on that next unit) of attraction between
molecules (which decrease with molecular size).
Melting Point
Trend (using per 3 as an example):
• The noble gases (Ar) exist as single
individual atoms with extremely weak forces
of attraction between the atoms.
Melting Point
Within groups there are also clear trends:
element Li
m.p. (K) 454
Na
371
K
336
Rb
312
Cs
302
In group 1 the m.p. decreases down the
group as the atoms become larger and the
strength of the metallic bond decreases.
Melting Point
Within groups there are also clear trends:
element F2
m.p. (K) 53.53
Cl2
171.60
Br2
265.80
I2
386.85
In group 7 the van der Waals’ attractive
forces between the diatomic molecules
increase down the group so the melting
points increase.
Reactivity
• Group 1 – increase reactivity from top to bottom
• Group 17 – reactivity decreases down a group
PART 4: d-block elements (first row)
The first row transition elements:
electron configurations are [Ar]…
(Sc)
4s23d1
Ti
V
Cr
4s23d2 4s23d3 4s13d5
Fe
4s23d6
Co
Ni
Mn
4s23d5
Cu
(Zn)
4s23d7 4s23d8 4s13d10 4s23d10
Note: for Cr and Cu it is more energetically favorable to half-fill or
completely fill the d sub-level so they have only one 4s electron
Oxidation states of the first row transition series: (you
need to be familiar with the ones in bold)
(Sc) Ti
V
Cr
Mn
Fe
Co
Ni
Cu (Zn)
+1
+3
+2
+2
+2
+2
+2
+2
+2
+2
+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
Transition elements
• Transition element: an element that possesses an
incomplete d sublevel in one or more of its oxidation states.
• Based on the above definition, which of the elements above
are not transition elements? Explain.
Transition elements
• Transition element: an element that possesses
an incomplete d sublevel in one or more of its
oxidation states.
• Scandium and zinc are not considered
transition metals because…
– Sc3+ has no d-electrons
– Zn2+ has a full d sublevel
• None of the possible oxidation states have
incomplete d sublevels
Characteristic properties of transition elements:
•
•
•
•
Variable oxidation states
Formation of complex ions
Colored complexes
Catalytic behavior
Oxidation-Reduction Reactions
• All oxidation reduction reactions have one element
oxidized and one element reduced
• Occasionally the same element may undergo both
oxidation and reduction. This is known as an autooxidation reduction
LEO says
GER!
LEO
says
Loss of Electrons = Oxidation
Gain of Electrons = Reduction
OIL RIG
• Oxidation Is Loss
• Reduction Is Gain
Redox

Oxidation-Reduction Reactions
a. Defined
1. Oxidation-reduction (“redox”) reactions involve
the transfer of electrons from one substance
to another.
2. Oxidized substances lose electrons and
reduced substances gain electrons.
111
Oxidation Numbers
• Oxidation is the loss of electrons;
Reduction is the gain of electrons
• Oxidation and reduction go together.
Whenever a substance loses electrons
and another substance gains electrons
• Oxidation Numbers are a system that
we can use to keep track of electron
transfers
Assigning Oxidation Numbers
• Assigning Oxidation Numbers
What is the general rule for assigning
oxidation numbers?
Assigning Oxidation Numbers
An oxidation number is a positive or
negative number assigned to an atom
to indicate its degree of oxidation or
reduction.
Assigning Oxidation Numbers
As a general rule, a bonded atom’s
oxidation number is the charge that it
would have if the electrons in the
bond were assigned to the atom of
the more electronegative element.
Assigning Oxidation Numbers
In binary ionic compounds, the oxidation
numbers of the atoms equal their ionic
charges.
• The compound sodium chloride is composed
of sodium ions (Na1+) and chloride ions (Cl1–).
• The oxidation number of sodium is +1.
• That of chlorine is –1.
– Notice that the sign is put before the oxidation
number.
Assigning Oxidation Numbers
Because water is a molecular compound,
no ionic charges are associated with its
atoms.
• Oxygen is reduced in the formation of water.
• Oxygen is more electronegative than hydrogen.
• The two shared electrons in the H–O bond are
shifted toward oxygen and away from hydrogen.
– The oxidation number of oxygen is –2.
– The oxidation number of each hydrogen is +1.
Assigning Oxidation Numbers
Oxidation numbers are often written above
the chemical symbols in a formula.
+1 –2
H 2O
Oxidation Numbers
Oxidation numbers always refer to
single atoms
The oxidation number of an
uncombined element is always 0
O2, H2, Ne
Zn
The oxidation number of Hydrogen
is usually +1 Hydrides are an
exception They are -1
HCl, H2SO4
The oxidation number of Oxygen is H2O, NO2, et
usually -2 Peroxides are an exception
They are –1 In OF2 oxygen is a +2
Oxidation numbers of monatomic
ions follow the charge of the ion
O2-, Zn2+
The sum of oxidation numbers is
zero for a neutral compound. It is
the charge on a polyatomic ion
LiMnO4
SO42-
Practice Assigning Oxidation Numbers
NO2
N2O5
HClO3
HNO3
Ca(NO3)2
KMnO4
Practice Assigning Oxidation Numbers
Fe(OH)3
K2Cr2O7
CO32CNK3Fe(CN)6
Oxidation numbers
• Oxidation number is the number of electrons
gained or lost by the element in making a
compound
• Metals are typically considered
more 'cation-like' and would
possess positive oxidation
numbers, while nonmetals are
considered more 'anion-like' and
would possess negative
oxidation numbers.
Oxidation - reduction
• Oxidation is loss of electrons
• Reduction is gain of electrons
– Oxidation is always accompanied by reduction
• The total number of electrons is kept constant
• Oxidizing agents oxidize and are themselves reduced
• Reducing agents reduce and are themselves oxidized
Predicting oxidation numbers
•
•
•
•
•
•
•
Oxidation number of atoms in element is zero in all cases
Oxidation number of element in monatomic ion is equal to the charge
sum of the oxidation numbers in a compound is zero
sum of oxidation numbers in polyatomic ion is equal to the charge
F has oxidation number –1
H has oxidn no. +1; except in metal hydrides where it is –1
Oxygen is usually –2. Except:
–
–
–
O is –1 in hydrogen peroxide, and other peroxides
O is –1/2 in superoxides KO2
In OF2 O is +2
Variable oxidation states
• All transition metals can show an oxidation
number of +2
• Some transition metals can form the +3 or +4
ion
Fe3+
Mn4+
• The M4+ ion is rare and in higher oxidation
states the element is generally found not as the
free metal ion, but either covalently bonded or
as the oxyanion, such as MnO4-.
Variable oxidation states
• Transition metals lose their s-electrons
first
Cu: [Ar]4s13d10
Cu1+: [Ar]3d10
Variable oxidation states
• 3d and 4s sublevels are similar in energy
Variable oxidation states
The gradual increase in ionization energy from losing the last 4s
election to the first 3d election explains the existence of additional
oxidation states
Succissive Ionzation Energies for Ca and Ti
14000
12000
Ionization Energy (kJ/mol)
10000
8000
Ti’s I.E. increases
gradually
Ca
(because energy
of 4s and
Ti
3d are similar) and does not
jump until the removal of
the 5th electron, which
explains why Ti can be +2,
+3, or +4, but not +5…
Ca has a huge “jump” in
6000
I.E. after removing just
2 electrons
4000
2000
0
0
1
2
3
4
Ionization Number
5
6
7
Some common examples of variable oxidation states in
addition to +2
Transition Oxidation Formula
element
#
Cr
Name
+3
CrCl3
Cr
+6
Cr2O72-
Mn
+4
MnO2
manganese (IV) oxide
Mn
+7
MnO4-
manganate (VII) ion
Fe
+3
Fe2O3
iron (III) oxide
Cu
+1
Cu2O
copper (I) oxide
chromium (III) chloride
dichromate (VI) ion
Complex Ions
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7–130
Formation of complex ions
• Ions of the d-block elements attract species that
are rich in electrons (ligands) because of their
small size.
Formation of complex ions
• Ligand: a neutral molecule or anion
which contains a nonbonding pair of
electrons.
– H2O is a common ligand
– The word “ligand” is derived from ligandus,
the Latin word for “bound”
– Most (but not all) transition metal ions exist as
hexahydrated complex ions in aqueous
solutions (i.e. [Fe(H2O)6]3+)
– Ligands can be replaced by other ligands
(such as NH3 or CN-).
Coordination complex: A structure containing a metal (usually a metal ion)
bonded (coordinated) to a group of surrounding molecules or ions.
Ligand (ligare is Latin, to bind): A ligand is a molecule or ion that is directly
bonded to a metal ion in a coordination complex
A ligand uses a lone pair of electrons to bond to the metal ion
Coordination sphere: A metal and its surrounding ligands
133
Formation of complex ions
• The electron pair from a ligand can form
coordinate covalent bonds with the metal ion
to form complex ions.
– A “coordinate covalent bond” is also known as a
“dative” - a bond in which both shared electrons are
supplied by one species…
ligand
coordinate covalent bond
Formation of complex ions
• Ions of the d-block elements attract species that
are rich in electrons (ligands) because of their
small size.
• The electron pair from a ligand can form
coordinate covalent bonds with the metal ion
to form complex ions.
– A Coordinate covalent bond is also known as a
“dative” - a bond in which both shared electrons are
supplied by one species
Formation of complex ions
Ligands:
Formation of complex ions
• Coordination number: the number of lone
pairs bonded to the metal ion.
L:
Mn+
Shape: linear
Coordination # = 2
:L
Formation of complex ions
• Coordination number: the number of lone
pairs bonded to the metal ion.
:L
L:
Mn+
L:
:L
Shape: square planar
Coordination # = 4
Formation of complex ions
• Coordination number: the number of lone
pairs bonded to the metal ion.
L
..
Mn+
L:
..
L
Shape: tetrahedral
Coordination # = 4
:L
Formation of complex ions
• Coordination number: the number of lone
pairs bonded to the metal ion.
L:
L:
L
..
:L
Mn+
Shape: octahedral
Coordination # = 6
..
L
:L
Coordination number: the number of lone pairs bonded to
the metal ion.
• Examples: state the coordination
numbers of the species below.
[Fe(CN)6]3-
6
[CuCl4]2-
4
[Ag(NH3)2]+
2
Follow the electrons
Colored complexes
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7–143
Colored Complexes
• In the free ion the five d-orbitals are degenerate (of
equal energy). However, in complexes the d orbitals
are split into two distinct energy levels.
E
Colored Complexes
• The energy difference between the levels
corresponds to a specific frequency and wavelength
in the visible region of the electromagnetic spectrum.
E=hf
Colored Complexes
• When the complex is exposed to light, energy of a
specific wavelength is absorbed and electrons are
excited from the lower level to the higher level.
In the example above, [Ti(H2O)6]3+ contains a single d-electron in lower
energy orbitals. 500 nm light absorption promotes the d-electron.
Colored Complexes
• Cu2+(aq) appears blue because it is the
complementary color to the wavelengths that have
been absorbed.
When yellow light is
subtracted out of
white light, blue light
is transmitted
Colored Complexes
• The observed color is across the color wheel
from the absorbed color.
Color Wheel
white light in
violet transmitted
Colored Complexes
• The energy separation between the orbitals and hence the
color of the complex depends on the following factors:
1) Nuclear charge (based on identity of the central metal ion)
Colored Complexes
2) Charge density of the ligand
[Ni(NH3)6]2+
[Ni(H2O)6]2+
[Ni(en)3]2+
Colored Complexes
Ex: NH3 has a higher charge density than H2O and so produces
a larger split in the d sublevel.
• [Cu(H2O)6]2+ absorbs red-orange light and appears pale blue
• [Cu(NH3)4(H2O)2]2+ absorbs the higher energy yellow light and
appears deep blue
I- < Br- < Cl- < F- < OH- < H2O < NH3 < en < NO2- < CNWeak-field
ligands
(small Δ)
Strong-field
ligands
(large Δ)
Colored Complexes
3) Number of d electrons present (and hence the
oxidation # of the central ion)
Mn2+
Mn3+
Mn4+
Mn6+
Mn7+
Colored Complexes
4) Shape of the complex ion
• Electric field created by the ligand’s lone pair of electrons depends
on the geometry of the complex ion
Colored Complexes
• If the d sublevel is completely empty, as in Sc3+, or
completely full, as in Cu+ or Zn2+, no transitions within
the d sublevel can take place and the complexes are
colorless.
Colored Complexes
NOTE: it is important to distinguish between the words “clear” and
“colorless.” Neither AP, nor IB, will give credit for use of the word clear
(which means translucent) when colorless should have been used.
Think about it, something can be pink and clear… colorless means
something else.
Both are “clear.” Only
the beaker on the left
is “colorless.”
Catalytic Behavior
• Many transition elements and their compounds are
very efficient catalysts.
• Catalysts increase the rate of chemical reactions.
Catalytic Behavior
• Examples: (need to be familiar with examples --- which metal or ion helps with
each process --- plus economic significance of those in bold, but mechanisms will not
be assessed)
–
–
–
–
–
–
–
Fe in the Haber Process
V2O5 in the Contact Process
Ni in the conversion of alkenes to alkanes
Pd and Pt in catalytic converters
MnO2 in the decomposition of hydrogen peroxide
Fe2+ in heme
Co3+ in vitamin B12
Catalytic Behavior
*Fe in the Haber Process
Ammonia (NH3) is the raw material for a
large number of chemical products including
fertilizers, plastics, drugs and explosives.
Catalytic Behavior
*V2O5 in the Contact Process
Sulfur trioxide (SO3) is
used in the manufacture
of sulfuric acid, the
manufacturing world’s
most important chemical.
Catalytic Behavior
Ni in the conversion of alkenes to alkanes
This rxn allows unsaturated vegetable oils with a double
C-C bond to be converted to margerine.
Catalytic Behavior
Pd and Pt in catalytic converters
This rxn removes harmful
primary pollutants from a
car’s exhaust gases.
Catalytic Behavior
MnO2 in the decomposition of hydrogen peroxide
Catalytic Behavior
Fe2+ in heme
O2 is transported
through bloodstream by
forming a weak bond
with the heme group of
hemoglobin.
The O2 - Fe2+ bond is
easily broken when
oxygen needs to be
released.