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Transition Metals!
Part 1:
1. Why do elements in groups have similar chemical properties?
Chemical properties depend on ____________________________
2. Why do elements physical properties within a period (consider melting points of Na, Mg, Al)
change?
Physical properties depend on _____________________________
3. Define oxidation and reduction
4. Define oxidizing agent and reducing agent.
5. How are atomic radius, ionization energy and ability to be oxidized related?
6. Which is the most powerful reducing agent amongst Li, Na, K?
7. Which is the most powerful oxidizing agent amongst Cl 2, Br2, I2?
Transition Metals
o Transition metals are strictly defined as metals or metal ions with partially filled d and/or f orbitals
o Transition metals are different than metals in other groups because their properties are very similar
across a period as well as in columns.
 The number of valence electrons and the atomic radius affect the chemical properties of a
particular atom.
 As you go across the period for transition metals you are adding inner electrons (3d) rather than
valence electrons. These inner electrons don’t as
readily participate in bonding because they are
“inner electrons”.


The above isn’t quite correct. Technically you have to
consider wave mechanics. The 4s orbital is lower in
energy than the 3d before Sc because the electrons in a
4s orbital can be closer to the nucleus. Remember that S
orbitals look like a series of rings. There is a ring relatively
close to the nucleus where an electron might be found.
We say that the s orbital electrons are penetrating. With a
higher probability of being close to the nucleus, an
increased nuclear charge can change the energy of the
orbital significantly compare to other orbitals (3d) which
do not penetrate the nucleus as much.
But, once the 4s orbital is filled, the 3d orbitals have a
significant drop in energy (below that of 4s electrons)
because they penetrate the 4s orbital and therefore, are
not shielded by the 4s electrons. Without shielding from
the 4s electrons, electrons in the 3d orbitals feel a greater
effective nuclear charge.

As you go across a period in the d-block the radius also stays about the same because only
inner electrons are added. These inner electrons shield the 4s electrons from the nuclear
charge. With effectively the same valence electrons and atomic radius, transition metals have
very similar chemical properties.
o

Common characteristics
They are practically all hard, strong, high-melting point, high boiling-point metals which conduct
heat and electricity well. In short, they are “typical” metals of the sort we meet in ordinary
circumstances.
Multiple oxidation states
 Results in different colors
 Results in usefulness as catalysts
Form complex ions
Colored
Paramagnetic




Ion Formation of Transition Metals
1. Why would transition metals lose their 4s electrons before their 3d electrons?
2. Write the electron configurations for Mn2+, Cu2+, Ti3+,
Answer 3-5 in the table below
3. What ion would result from the loss of all s electrons for the metals (Sc to Zn)?
4. What ion would result form the loss of all the s and d electrons for all the metals (Sc to Zn)?
5. Now look at the most common oxidation states for each metal. (Use the big periodic table I gave
you or the one posted in the front of the classroom. The most common oxidation state is bolded)
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ion that
results from
the loss of
all s
electrons
Ion that
results from
the loss of
all s and d
electrons
Most
common
oxidation
state
6. Do some transition metals lose all of their s and d electrons?
7. For elements that lose both their s and d electrons, would you say that the energy difference
between 4s and 3d is large or small? Explain your answer.
8. Is there a general pattern? Is there a point at which, most often, only the s electrons are lost?
9. What increases as you go across the periodic table? Why would this effect the ability to lose d
electrons?
10. Look at chart on page 991 (fig 21.2) You can see that that first ionization energy changes gradually
but the third ionization energy changes much more significantly for each step in the row. How does
this support the idea that the energy level of the 3d orbitals decreases across a row?
Open the IB Book to page 87.
11. What is the charge of each element on MnO4-1 and in Cr2O72-?
12. Why would MnO4-1 and Cr2O72- make good oxidizing agents?
13. An element has 8 electrons in its d shell, will a +3 ion of that element be a better oxidizing agent or
reducing agent? Explain.
14. An element has 3 electrons in its d shell, will a +4 ion of that element be a better oxidizing agent or
reducing agent? Explain.
15. Why would Fe’s +3 be more stable than +2?
15. Given that usually, the physical properties of groups varies slightly, but that the 4d and 5d periods
have very similar chemical and physical properties, what can you infer about the electronic structures
of the 4d and 5d elements?