Topic 13 Transition Elements

Answer these questions.

– List out all the elements in 3d block?

– List out all the transition elements?

-Why Sc and Zn were not include in transition elements?

A transition metal is one that forms one or more stable ions which have incompletely filled d orbitals. On the basis of this definition, scandium and zinc do not count as transition metals, even though they are members of the d block.

Scandium has the electronic structure [Ar] 3d1 4s2. When it forms ions, it always loses the 3 outer electrons and ends up with an argon structure. The Sc3+ ion has no d electrons and so does not meet the definition.

Zinc has the electronic structure [Ar] 3d104s2. When it forms ions, it always loses the two 4s electrons to give a 2+ ion with the electronic structure [Ar] 3d10. The zinc ion has full d levels and does not meet the definition either.

– Why is the energy of 4s orbital less than that of 3d orbital? Why write 4s orbital lower energy level than 3d?

The reason the 4s orbital is usually filled first is because the 3d orbital electrons feel a lot of repulsion between the other 3rd level electrons (those in 3s and 3p). The 4s orbital, while farther away, requires less energy for the electron to occupy because it is not repelled as much by the coexisting electrons.

– Why is the electron configuration for copper is 1s2 2s2 2p6 3s2 3p6 3d10 4s1 instead of 1s2 2s2 2p6 3s2 3p6 3d9 4s2 ?

A full-filled sublevel is more stable than a half-filled sublevel.  The arrangement of electrons with the same spin is as small as possible. The 4s orbital has higher energy than the 3d orbital.

– Why does copper included into transition elements even though it has fully filled d sub-shell?
Copper, [Ar] 3d10 4s1, forms two ions. In the Cu+ ion the electronic structure is [Ar] 3d10. However, the more common Cu2+ ion has the structure [Ar] 3d9. Copper is definitely a transition metal because the Cu2+ ion has an incomplete d sub-shell.
Cu2+ ion has the structure [Ar]3d9. Copper is a transition metal because the Cu2+ ion has an incomplete d orbital.
– What are the transition elements’ properties?
Properties of transition elements include:
1. they are good conductors of heat and electricity.
2. they can be hammered or bent into shape easily.
3. they have high melting and boiling points (but mercury is a liquid at room temperature).
4. they are usually hard and tough.
5. they have high densities.
6. have large charge/ radius ratio.
7. form compounds which are often paramagnetic.
8. show variable oxidation states.
9. form coloured ions and compounds: the colour of aqueous potassium manganate(VII) is purple.
10. form compounds with profound catalytic activity: iron filing is used as a catalyst in Haber process, copper reacts with ethane to form tris(ethanedioato)cuprate(II) ion.

– Why transition elements formed coloured ions?

The reason why transition metal in particular are colorful is because they have unfilled or either half filled d orbitals. There is Crystal field theory which explains the splitting of the d orbital, which splits the d orbital to a higher and lower orbital. Now, the electrons of the transition metal can “jump”. d-d transition can occur.

– Why some of the elements in d-block form white compounds but not in coloured?

Compounds where the 3d sub-shell of the transition elements is either empty or fully filled are white or colourless because there is no d-d transition can take place.

– Why Scandium(III) ion, titanium(IV) ion, copper(I) ion and zinc ion form white compounds?

– Why is Cu2+ blue, but Zn2+ is colourless?

– Why is Cu+ not Coloured?

 

– Explain the trends of atomic radius of transition elements?

Moving left to right across the periodic table, there is a trend of decreasing atomic radius. Going from vanadium to copper, the nuclear charge increases faster than the increase in the screening effect (due to the ineffective screening effect of the inner 3d electrons), causing the atomic size to decrease. In chromium and copper, the 4s orbital is half-filled. For V and Cr, the 3+ state (V3+, Cr3+) is more stable. In general, the first ionisation energy increases across a Period.

However, in the transition metals, moving left to right, there is a trend of increasing atomic radius which levels off and becomes constant.

In the transition elements, the number of electrons are increasing but in a particular way (cara yang istimewa/ khas).

The number of electrons increase going across a period, thus, there is more pull of these electrons towards the nucleus. However, with the d−electrons, there is some added electron-electron repulsion.

For example, in chromium, there is a promotion of one of the 4s electrons to half fill the 3d sublevel; the electron-electron repulsions are less and the atomic size is smaller. The opposite holds true for the latter part of the row.

From titanium to copper, the number of protons in the nucleus increases. However, each additional electron is added to an inner 3d sub-shell. These additional inner electrons shields the outer electrons from the nucleus and to a large extent and cancel out most of the effect of the increase in the nucleus charge. As a result, the effective nuclear charge increases only slightly resulting in only a small change in the atomic radius of the elements.

The sudden increase in the atomic radius of manganase is due to the greater repulsion involving a half-filled 3d sub-shell (3d5, when the electron density is evenly distributed) repulse with the electrons in the inner shells of 3s2 and 3p6.

Similarly, the slight increase in the atomic radius of copper is due to the greater repulsion involving a completely filled third shell (3s2 3p6 3d10) and the inner shells.

– Why there is sudden increase atomic radius of manganese?

The sudden increase in the atomic radius of manganase is due to the greater repulsion involving a stable half-filled 3d sub-shell (3d5, when the electron density is evenly distributed) repulse with the electrons in the inner shells of 3s2 and 3p6.

– List out all of the transition elements’ electronic configuration?

*Argon, Ar = 18 proton number

– What were the electronic configuration of potassium and calcium? 

Proton number of potassium is 19, the electronic configuration is 2.8.8.1,

1s2 2s2 2p6 3s2 3p6 4s1

Proton number of calcium is 20, the electronic configuration is 2.8.8.2,

1s2 2s2 2p6 3s2 3p6 4s2

*You can say that for potassium and calcium, the 3d orbitals have a higher energy than the 4s, and so for these elements, the 4s levels fill before than the 3d. 

– What were the electronic configuration of chromium and copper?

Proton number of chromium is 24, the electronic configuration is 2.8.8.6,

1s2 2s2 2p6 3s2 3p6 3d5 4s1

Proton number of copper is 29, the electronic configuration is 2.8.18.1,

1s2 2s2 2p6 3s2 3p6 3d10 4s1

*4s orbitals filled before 3d orbitals except chromium and copper.

– Why Zn is bigger atomic radius than Cu?

Proton number copper is 29, the electronic configuration is 2.8.18.1

1s2 2s2 2p6 3s2 3p6 3d10 4s1 and

Proton number zinc is 30, the electronic configuration is 2.8.18.2,

1s2 2s2 2p6 3s2 3p6 3d10 4s2

Both copper and zinc have the same number of inner electrons. Zinc has more protons than copper. The effective nuclear charge of zinc is higher than that of copper, which would cause the size of zinc to be smaller than copper.

However, this is not the case. The larger size of zinc is due to the greater repulsion between the inner shell electrons and a fully filled 3s orbital as compared to the repulsion between the inner shell electrons and a half-filled 3s orbital of copper.

Size of zinc is greater than copper. In zinc all the 10 d electrons are available which provide very large screening effect and decreases the effect of nuclear charge on the electrons which results in repulsion in the electrons of zinc which lead to large atomic radius.

– Explain the trends of melting point and boiling point of transition elements?

Melting points and bp are a measure of the foces that hold the particles together in the solid state and liquid state respectively.

The stronger the forces of attraction, the higher the melting point and boiling point.

The transitional elements are all metals. The ‘intermolecular force’ is the metallic bond.

All transition elements have very high melting points (in excess of 1000oC) and very high boiling points (in excess of 2000oC) compared to other main group metals such as calcium (melting point = 838oC, boiling point = 1440oC). This indicates the presence of strong metallic bonds.

The high melting point and boiling point of the elements are due to the very small difference in the energy of the 3d subshell and 4s sub-shell.

As a result, atoms of the transition elements can make use of the outer 4s as well as the inner 3d electrons to form metallic bonds. The large number of electrons available accounts for the strong metallic bonds formed.

On the other hand, potassium (a Group 1 metal) is soft with a low melting point (64.2 °C), because potassium has only one valence electron (4s’) available for the formation of the metallic bonds. Similarly for calcium which has only two valence electrons (4s2).

Due to the large energy difference between the 3p and 4s sub-shells, potassium and calcium cannot make use of the inner 3p electrons to form metallic bonds.

The ‘sudden’ drop in the melting point and boiling point of manganese is attributed to the electronic configuration of the manganese atom:

Mn: [Ar]3d54s2

The half-filled 3d5 configuration has extra stability. As a result, the d electrons are less available for metallic bond formation.

 

– Why very high melting and boiling point (physical property) of transition elements?

Presence of strong metallic bonds with very small difference in the energy of 3d and 4s sub-shell which available to make use to form metallic bonds except Mn (half-filled that has extra stability then less available for make use as metallic bonds).

– Why is it potassium (group 1 metal) is soft with low melting point?

All Group 1 elements have one electron in their outermost shell which is held very weakly by the nucleus. This electron can drift further from the nucleus than in most atoms of other elements. This results in Group 1 elements having larger atomic radii than those elements that follow them in their respective periods. The large atomic size results in weaker forces between neighbouring atoms. It is these weaker attractive forces due to the large atomic radii between neighbouring atoms of Group 1 elements that result in lower melting and boiling points when compared to other metals.

Potassium (a Group 1 metal) is soft with a low melting point (64.2 °C), because potassium has only one valence electron (4s’) available for the formation of the metallic bonds. Similarly for calcium which has only two valence electrons (4s2).

Alkali metals are soft metals that can be cut with a knife and silvery coloured. Freshly cut alkali metals are shiny but tarnish rapidly due to reaction with oxygen in the air. They are generally stored under oil.

It is the reduced interatomic forces in these elements that make them relatively soft.

– Why is it calcium (group 2 metal) is soft with low melting point?

Calcium (a Group 2 metal) is soft with a low melting point (?? °C), because potassium has only two valence electron (4s2) available for the formation of the metallic bonds. Similarly for potassium (a Group 1 metal) which has only one valence electrons (4s’).

– Explain why sudden drop in melting point and boiling point of manganase?

The ‘sudden’ drop in the melting point and boiling point of manganese is attributed to the electronic configuration of the manganese atom:

Mn: [Ar] 3d5 4s2

The d5 configuration has extra stability. As a result, the d electrons are less available for metallic bond formation.

– Why the density of transition elements increase across period?

Since the number of protons is also increasing, the effective nuclear charge increases across a period. This causes the atomic radius to decrease. As radius decreases across a period and the atomic number increases so density increases across a period.

Due to the presence of strong metallic bonds, atoms of the transition elements are very tightly packed in the solid lattice.

As a result, all the d-block elements have high densities (compared to aluminium = 2.70 g cm-3).

The density increases gradually across the period because, although the atomic size does not change much, the relative atomic mass increases gradually. This results in a gradual increase in the Mass over Volume ratio.

– Why chromium has higher density than potassium?

chromium (density = 7.2 g cm-3)

potassium (density = 0.86 g cm’)

Both elements from the 4th Period. Each element has one electron in their outer shells. There were a big difference in their densities. The size of the chromium atom (0.129 nm) is smaller than potassium (0.227 nm). However, it is heavier than potassium, the chromium atoms are more closely packed in the solid lattice (due to the strong metallic bond in 3d and 4s sub-shell). As a result, Mass the ratio of Mass/ Volume(= density) for chromium is larger than potassium.

……

 

The following are high-level thinking skills (KBAT) questions designed based on the learning outcomes provided:

13 Transition Elements.

13.1 Physical properties of first row transition elements.

1. Describe the camparisons between transition elements and d-block elements?

2. Describe the trends of the first row of transition elements’ physical properties such as:

a. atomic radius,
b. first ionisation energy,
c. melting point,
d. density,
e. electrical conductivity.

3. Explain the variation in successive ionisation energies in transition elements?

13.2 Chemical properties of first row transition elements.

1. Explain why variable oxidation states in terms of the energies of 3d and 4s orbitals?

2. Explain the colours of transition metal ions in terms of a partially filled 3d orbitals?

3. State the principal oxidation numbers of these elements in their common cations, oxides and oxo ions?

4. Explain qualitatively the relative stabilities of these oxidation states?

5. Explain the uses of standard reduction potentials in predicting the relative stabilities of aqueous ions?

6. Explain the terms complex ion and ligand?

7. Explain the formation of complex ions and the colour changes by exchange of ligands?

8. Explain the use of first row transition elements in homogeneous catalysis?

9. Explain the use of first row transition elements in heterogeneous catalysis?

13.3 Nomenclature and bonding of complexes.

1. Discuss coordinate bond formation between ligands and the central metal atom/ion, and state the types of ligands, i.e. monodentate, bidentate and hexadentate?

13.4 Uses of first row transition elements and their compounds.

1. Describe the use of chromium (in stainless steel), cobalt, manganese, titanium (in alloys) and TiO2 (in paints)?

13 Transition Elements

13.1 Physical properties of first row transition elements

Candidates should be able to:

(a) define a transition element in terms of incomplete d orbitals in at least one of its ions;

(b) describe the similarities in physical properties such as atomic radius, ionic radius and first ionisation energy;

(c) explain the variation in successive ionisation energies;

(d) contrast qualitatively the melting point, density, atomic radius, ionic radius, first ionisation energy and conductivity of the first row transition elements with those of calcium as a typical s-block element.

13.2 Chemical properties of first row transition elements

Candidates should be able to:

(a) explain variable oxidation states in terms of the energies of 3d and 4s orbitals;

(b) explain the colours of transition metal ions in terms of a partially filled 3d orbitals;

(c) state the principal oxidation numbers of these elements in their common cations, oxides and oxo ions;

(d) explain qualitatively the relative stabilities of these oxidation states;

(e) explain the uses of standard reduction potentials in predicting the relative stabilities of aqueous ions;

(f) explain the terms complex ion and ligand;

(g) explain the formation of complex ions and the colour changes by exchange of ligands. (Examples of ligands: water, ammonia, cyanide ions, thiocyanate ions, ethanedioate ions, ethylenediaminetetraethanoate, halide ions; examples of complex ions: [Fe(CN)6]4, [Fe(CN)6]3, [Fe(H2O)5(SCN)]2+);

(h) explain the use of first row transition elements in homogeneous catalysis, as exemplified by Fe2+ or Fe3+ in the reaction between I and S2O82;

(i) explain the use of first row transition elements in heterogeneous catalysis, as exemplified by Ni and Pt in the hydrogenation of alkenes.

13.3 Nomenclature and bonding of complexes

Candidates should be able to:

(a) name complexes using International Union of Pure and Applied Chemistry (IUPAC) nomenclature;

(b) discuss coordinate bond formation between ligands and the central metal atom/ion, and state the types of ligands, i.e. monodentate, bidentate and hexadentate.

13.4 Uses of first row transition elements and their compounds

Candidates should be able to:

(a) describe the use of chromium (in stainless steel), cobalt, manganese, titanium (in alloys) and TiO2 (in paints).

 

13 Transition Elements

13.1 Physical properties of first row transition elements

Candidates should be able to:

(a) define a transition element in terms of incomplete d orbitals in at least one of its ions;

Another definition of a Transition metal is that it has an incomplete d subshell. This excludes Cu and Zn. Atomic No. Symbol 4 s 3 d Configuration. NB 4s is of lower energy so fills before 3d. Cr and Cu are exceptions to Aufbau rule.

(b) describe the similarities in physical properties such as atomic radius, ionic radius and first ionisation energy;

atomic radius,

ionic radius,

first ionisation energy,

(c) explain the variation in successive ionisation energies;

(d) contrast qualitatively the melting point, density, atomic radius, ionic radius, first ionisation energy and conductivity of the first row transition elements with those of calcium as a typical s-block element.

melting point,

density,

atomic radius,

ionic radius,

first ionisation energy,

electrical conductivity,

13.2 Chemical properties of first row transition elements

Candidates should be able to:

(a) explain variable oxidation states in terms of the energies of 3d and 4s orbitals;

(b) explain the colours of transition metal ions in terms of a partially filled 3d orbitals;

(c) state the principal oxidation numbers of these elements in their common cations, oxides and oxo ions;

 

common cations,

oxides,

oxo ions,

(d) explain qualitatively the relative stabilities of these oxidation states;

(e) explain the uses of standard reduction potentials in predicting the relative stabilities of aqueous ions;

(f) explain the terms complex ion and ligand;

 

(g) explain the formation of complex ions and the colour changes by exchange of ligands.

Examples of ligands:

water,

ammonia,

cyanide ions,

thiocyanate ions,

ethanedioate ions,

ethylenediaminetetraethanoate,

halide ions;

examples of complex ions:

[Fe(CN)6]4,

[Fe(CN)6]3,

[Fe(H2O)5(SCN)]2+);

(h) explain the use of first row transition elements in homogeneous catalysis, as exemplified by Fe2+ or Fe3+ in the reaction between I and S2O82;

(i) explain the use of first row transition elements in heterogeneous catalysis, as exemplified byNi and Pt in the hydrogenation of alkenes.

 

13.3 Nomenclature and bonding of complexes

Candidates should be able to:

(a) name complexes using International Union of Pure and Applied Chemistry (IUPAC) nomenclature;

(b) discuss coordinate bond formation between ligands and the central metal atom/ ion, and state the types of ligands, i.e. monodentate, bidentate and hexadentate.

monodentate,

bidentate,

hexadentate,

 

13.4 Uses of first row transition elements and their compounds

Candidates should be able to:

(a) describe the use of chromium (in stainless steel), cobalt, manganese, titanium (in alloys) and TiO2 (in paints).

 

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