Name: transition metals in the periodic table
Uploaded: 2020-09-24T15:06:53.000Z
Duration: 2 min 58 s
Channel: JoVE

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Properties of Transition Metals

Transition metals are a group of elements located in the d-block. They are placed between the main group elements and exhibit unique characteristics, such as electrical conductivity, colors, hardness, high melting points, and magnetism.
Many of these characteristics are attributed to the filling of d orbitals. Moving across the period, electrons are added following the Aufbau principle. However, the outermost s orbital is filled first, before the electron enters the (n &#8722; 1) d subshell, with some exceptions. Chromium and copper have a half-filled s orbital because it is energetically more favored to have a half-filled or full d subshell.
The atomic size of transition metals decreases slightly across a period. As the electrons enter the d subshell, the number of valence electrons in the outermost s orbital remains steady, leading to an overall constant nuclear charge. Down the column, the trend increases from the fourth to the fifth period, but not below. 
The 6th period contains 14 additional elements of the lanthanide series, where electrons enter the (n &#8722; 2)f subshell before entering the (n &#8722; 1)d subshell. Electrons located in the f orbitals do not sufficiently shield the valence electrons in the outermost s orbital, resulting in an increased effective nuclear charge. Thus, the valence electrons are attracted more strongly, leading to smaller than expected atomic radii. This is known as the Lanthanide contraction.
Similarly, the transition elements&#8217; electronegativity is influenced by its atomic size and nuclear charge. Moving across a period, it increases only slightly. Down the column, the electronegativity increases from the fourth to the fifth period but remains constant in the sixth period. This is attributed to a decrease in atomic size and an increase in the nuclear charge.
The ionization energy of transition metals increases slightly across a period. Interestingly, down the column, it is higher for the 6th period than the 4th and 5th periods. Here, the valence electrons are held more tightly as the nuclear charge increases with a negligible increase in atomic size.
Lastly, transition metals can have a range of stable oxidation states by losing outermost s orbital electrons before the unpaired d orbital electrons. The partially filled d orbitals enable oxidation states ranging from +7 to +1, characteristic colors, and magnetic properties.

properties of transition metals

Transition metals are defined as those elements that have partially filled d orbitals. As shown in Figure 1, the d-block elements in groups 3&#8211;12 are transition elements. The f-block elements, also called inner transition metals (the lanthanides and actinides), also meet this criterion because the d orbital is partially occupied before the f orbitals.
<img alt="Periodic table diagram with element blocks: s, p, d, f; essential chemistry reference." src="/files/ftp_upload/11454/11454_Image_1.png" style="text-align: center; height: 500px;" data-altupdatedat="1754904545000" data-alt="Periodic table diagram highlighting s, p, d, f blocks; educational chemistry resource." />
Figure 1: Periodic Table. The transition metals are located in groups 3&#8211;11 of the periodic table. The inner transition metals are in the two rows below the body of the table.
The d-block elements are divided into the first transition series (the elements Sc through Cu), the second transition series (the elements Y through Ag), and the third transition series (the element La and the elements Hf through Au). Actinium, Ac, is the first member of the fourth transition series, which also includes Rf through Rg.
The f-block elements are the elements Ce through Lu, which constitute the lanthanide series (or lanthanoid series), and the elements Th through Lr, which constitute the actinide series (or actinoid series). Because lanthanum behaves very much like the lanthanide elements, it is considered a lanthanide element, even though its electron configuration makes it the first member of the third transition series. Similarly, the behavior of actinium means it is part of the actinide series, although its electron configuration makes it the first member of the fourth transition series.
The transition elements have many properties in common with other metals. They are almost all hard, high-melting solids that conduct heat and electricity well. They readily form alloys and lose electrons to form stable cations. In addition, transition metals form a wide variety of stable coordination compounds, in which the central metal atom or ion acts as a Lewis acid and accepts one or more pairs of electrons. Many different molecules and ions can donate lone pairs to the metal center, serving as Lewis bases.
Properties of the Transition Elements
Transition metals demonstrate a wide range of chemical behaviors. Some transition metals are strong reducing agents, whereas others have very low reactivity. For example, the lanthanides all form stable 3+ aqueous cations. The driving force for such oxidations is similar to that of alkaline earth metals such as Be or Mg, forming Be2+ and Mg2+. On the other hand, materials like platinum and gold have much higher reduction potentials. Their ability to resist oxidation makes them useful materials for constructing circuits and jewelry.
Ions of the lighter d-block elements, such as Cr3+, Fe3+, and Co2+, form colorful hydrated ions that are stable in water. However, ions in the period just below these (Mo3+, Ru3+, and Ir2+) are unstable and react readily with oxygen from the air. The majority of simple, water-stable ions formed by the heavier d-block elements are oxyanions such as MoO42&#8722; and ReO4&#8722;.
Ruthenium, osmium, rhodium, iridium, palladium, and platinum are the platinum metals. With difficulty, they form simple cations that are stable in water, and, unlike the earlier elements in the second and third transition series, they do not form stable oxyanions.
Both the d- and f-block elements react with nonmetals to form binary compounds; heating is often required. These elements react with halogens to form a variety of halides ranging in oxidation state from +1 to +6. On heating, oxygen reacts with all of the transition elements except palladium, platinum, silver, and gold. The oxides of these latter metals can be formed using other reactants, but they decompose upon heating. The f-block elements, the elements of group 3, and the elements of the first transition series except copper react with aqueous solutions of acids, forming hydrogen gas and solutions of the corresponding salts.
Transition metals can form compounds with a wide range of oxidation states. Some of the observed oxidation states of the elements of the first transition series are shown in Table 1. Moving from left to right across the first transition series, the number of common oxidation states increases at first to a maximum towards the middle of the table, then decreases. The values in the table are typical values; there are other known values, and it is possible to synthesize new additions. For example, in 2014, researchers were successful in synthesizing a new oxidation state of iridium (+9).
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			<td style="color: black; border: 1px solid black; text-align: center;">21 Sc</td>
			<td style="color: black; border: 1px solid black; text-align: center;">22 Ti</td>
			<td style="color: black; border: 1px solid black; text-align: center;">23 V</td>
			<td style="color: black; border: 1px solid black; text-align: center;">24 Cr</td>
			<td style="color: black; border: 1px solid black; text-align: center;">25 Mn</td>
			<td style="color: black; border: 1px solid black; text-align: center;">26 Fe</td>
			<td style="color: black; border: 1px solid black; text-align: center;">27 Co</td>
			<td style="color: black; border: 1px solid black; text-align: center;">28 Ni</td>
			<td style="color: black; border: 1px solid black; text-align: center;">29 Cu</td>
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			<td style="color: black; border: 1px solid black; text-align: center;">+1</td>
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			<td style="color: black; border: 1px solid black; text-align: center;">+2</td>
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			<td style="color: black; border: 1px solid black; text-align: center;">+3</td>
			<td style="color: black; border: 1px solid black; text-align: center;">+3</td>
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			<td style="color: black; border: 1px solid black; text-align: center;">+4</td>
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			<td style="color: black; border: 1px solid black; text-align: center;">+5</td>
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			<td style="color: black; border: 1px solid black; text-align: center;">+6</td>
			<td style="color: black; border: 1px solid black; text-align: center;">+6</td>
			<td style="color: black; border: 1px solid black; text-align: center;">+6</td>
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			<td style="color: black; border: 1px solid black; text-align: center;">+7</td>
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Table 1. Transition metals of the first transition series can form compounds with varying oxidation states.
For the elements scandium through manganese (the first half of the first transition series), the highest oxidation state corresponds to the loss of all of the electrons in both the s and d orbitals of their valence shells. The titanium(IV) ion, for example, is formed when the titanium atom loses its two 3d and two 4s electrons. These highest oxidation states are the most stable forms of scandium, titanium, and vanadium. However, it is not possible to continue to remove all of the valence electrons from metals as we continue through the series. Iron is known to form oxidation states from +2 to +6, with iron(II) and iron(III) being the most common. Most of the elements of the first transition series form ions with a charge of 2+ or 3+ that are stable in water, although those of the early members of the series can be readily oxidized by air.
The elements of the second and third transition series generally are more stable in higher oxidation states than are the elements of the first series. In general, the atomic radius increases down a group, which leads to the ions of the second and third series being larger than are those in the first series. Removing electrons from orbitals that are located farther from the nucleus is easier than removing electrons close to the nucleus. For example, molybdenum and tungsten, members of group 6, are limited mostly to an oxidation state of +6 in aqueous solution. Chromium, the lightest member of the group, forms stable Cr3+ ions in water and, in the absence of air, less stable Cr2+ ions. The sulfide with the highest oxidation state for chromium is Cr2S3, which contains the Cr3+ ion. Molybdenum and tungsten form sulfides in which the metals exhibit oxidation states of +4 and +6.
The variety of properties exhibited by transition metals is due to their complex valence shells. Unlike most main group metals where one oxidation state is normally observed, the valence shell structure of transition metals means that they usually occur in several different stable oxidation states. In addition, electron transitions in these elements can correspond with absorption of photons in the visible electromagnetic spectrum, leading to colored compounds. Because of these behaviors, transition metals exhibit a rich and fascinating chemistry.
This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/19-1-occurrence-preparation-and-properties-of-transition-metals-and-their-compounds">Openstax, Chemistry 2e, Chapter 19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds.</a>

Transition metals are defined as those elements that have partially filled d orbitals. As shown in Figure 1, the d-block elements in groups 3&#8211;12 are ...

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Transition Metals and Coordination Complexes

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KEY TERMS AND CONCEPTS

transition metals in the periodic table

Transition metals are the d-block elements and the inner transition metals in the f-block. They are defined by partially filled d orbitals. In the periodic table, the transition elements are in groups 3 through 11, and the inner transition metals appear in the two rows below the main body of the table.

The d-block is often grouped into transition series. The first transition series runs from Sc through Cu. The second series runs from Y through Ag. The third series includes La and Hf through Au, and the fourth series begins with Ac and continues through Rg. The f-block includes Ce through Lu, which make up the lanthanide series, and Th through Lr, which make up the actinide series. Lanthanum and actinium behave like the lanthanides and actinides, so they are included with those series even though their electron configurations place them at the start of the third and fourth transition series.

Transition metals share several common metal properties. Most are hard, high-melting solids that conduct heat and electricity well. They also form alloys easily and lose electrons to make stable cations. Another major feature is their ability to form stable coordination compounds, where the metal atom or ion acts as a Lewis acid and accepts electron pairs from Lewis bases such as molecules or ions with lone pairs.

These elements show a wide range of chemical behavior. Some transition metals are strong reducing agents, while others are much less reactive. The lanthanides, for example, form stable 3+ aqueous cations, and this oxidation is similar in driving force to the formation of Be2+ and Mg2+ by alkaline earth metals. At the other extreme, platinum and gold have high reduction potentials, so they resist oxidation and are useful in circuits and jewelry.

Ion stability also changes across the d-block. Lighter ions such as Cr3+, Fe3+, and Co2+ often form colorful hydrated ions that remain stable in water. The next heavier ions, such as Mo3+, Ru3+, and Ir2+, are unstable and react quickly with oxygen in air. For the heavier d-block elements, many of the water-stable simple ions are oxyanions, such as MoO4 2− and ReO4−. Ruthenium, osmium, rhodium, iridium, palladium, and platinum are called the platinum metals. They form simple cations only with difficulty, and they do not form stable oxyanions like the earlier elements in the second and third transition series.

Transition metals and inner transition metals also react with nonmetals to form binary compounds, and heating is often needed. They react with halogens to make halides in oxidation states from +1 to +6. When heated, oxygen reacts with all transition elements except palladium, platinum, silver, and gold. The oxides of those metals can be made by other reactions, but they break down when heated. The f-block elements, group 3 elements, and the first transition series except copper react with aqueous acids to form hydrogen gas and the matching salts.

A key reason transition metals are so versatile is that they can form compounds in many oxidation states. In the first transition series, scandium through manganese show a rise in the number of common oxidation states, with a maximum near the middle of the series, and then the number falls again. For scandium through manganese, the highest oxidation state usually comes from removing all valence s and d electrons. Titanium(IV), for example, forms when titanium loses its two 3d and two 4s electrons. Iron can form oxidation states from +2 to +6, but +2 and +3 are the most common. Most first-series transition metals form 2+ or 3+ ions that are stable in water, although the early members are easily oxidized by air.

The second and third transition series are usually more stable in higher oxidation states than the first series. Their larger atomic radius makes electron removal easier because the electrons are farther from the nucleus. For example, molybdenum and tungsten, both group 6 elements, are mostly limited to +6 in aqueous solution. Chromium, the lightest member of the group, forms stable Cr3+ ions in water and less stable Cr2+ ions when air is absent. The highest chromium sulfide is Cr2S3, which contains Cr3+, while molybdenum and tungsten form sulfides with +4 and +6 oxidation states.

The many properties of transition metals come from their complex valence shells. Unlike most main group metals, which usually show one oxidation state, transition metals often appear in several stable oxidation states. Their electron transitions can also absorb visible light, which gives many of their compounds color. These features make transition metal chemistry especially varied and important.

Transition Metals in the Periodic Table

Transition metals are the d-block elements and the inner transition metals in the f-block. They are defined by partially filled d orbitals. In the periodic table, the transition elements are in groups 3 through 11, and the inner transition metals appear in the two rows below the main body of the table. The d-block is often grouped into transition series. The first transition series runs from Sc through Cu. The second series runs from Y through Ag. The third series includes La and Hf through...

coordination compounds and nomenclature

How Transition Metal Complexes Are Named

Transition metal complexes are named by following rules for ligands, metal ions, and oxidation state. These compounds form when a central metal ion or atom, often a transition metal or inner transition metal, accepts electron pairs from ligands. A ligand is a Lewis base that can be an atom, molecule, or ion with one or more lone pairs to donate. In these complexes, the bond is a coordinate covalent bond. That means both electrons in the bond come from the donor, such as ammonia, and are...

metal-ligand bonds

Coordination Complexes and Chelation

Coordination complexes form when a metal ion bonds to ligands. These compounds appear in familiar examples such as hemoglobin in blood, chlorophyll in green plants, vitamin B-12, and the catalyst used to make polyethylene. Transition metal ions are especially likely to form these complexes. In a coordination complex, the metal acts as a Lewis acid. The ligand acts as a Lewis base and donates both electrons in the coordinate covalent bond. A ligand can be an atom, molecule, or ion, as long as...

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Transition Metal Shapes from Coordination Number

Transition metal complexes are often identified by their coordination number and the shape around the metal ion. The coordination number is the number of donor atoms bonded to the central metal. In many complexes, the most common shapes are octahedral, tetrahedral, and square planar. Coordination number and molecular geometry are linked in a clear pattern. A coordination number of 2 gives a linear shape, as in [Ag(NH3)2]+. A coordination number of 3 gives a trigonal planar shape, as in...

structural isomerism

Linkage and Ionization Isomers

Isomers are different chemical species that share the same chemical formula. In coordination compounds, structural isomerism includes linkage isomers and coordination-sphere isomers. These two forms show how the same atoms can be arranged in different ways around a metal center. Linkage isomers form when a ligand can attach to a transition metal through two different atoms. The cyanide ligand, CN−, can bind through carbon or nitrogen. The thiocyanate ligand, SCN−, can bind through sulfur or...

stereoisomerism

Isomer Shapes in Coordination Complexes

Isomer shapes in coordination complexes can change a compound’s properties even when the chemical formula stays the same. Isomers are different chemical species with the same formula. In transition metal complexes, the atoms stay connected through the same types of bonds, but their arrangement in space can differ. One common form is geometric isomerism. In these complexes, ligands (the ions or molecules attached to the metal) can sit in cis or trans positions. The octahedral [Co(NH3)4Cl2]+ ion...

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Metal Ligand Bonds in Coordination Complexes

Coordination complexes form metal-ligand bonds through valence bond theory. In this model, Linus Pauling explained bonding with hybridization and overlap of atomic orbitals. The central metal atom or ion acts as a Lewis acid. It uses empty orbitals of suitable energy to accept electron pairs from ligands, which act as Lewis bases. These coordinate covalent bonds help explain the color, geometry, and magnetic behavior seen in coordination compounds. The type of hybridization helps determine the...

crystal field theory - octahedral complexes

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Octahedral d-Orbital Splitting in Crystal Fields

Crystal field theory explains how ligands affect the d orbitals in transition metal complexes. This electrostatic model helps predict the colors, magnetic behavior, and some structures of coordination compounds. It focuses on the nonbonding electrons on the central metal ion, not on the metal-ligand bonds. In crystal field theory, the metal ion and ligands are treated as very small point charges. Because electrons are negative, electrons donated by the ligands repel the electrons in the...

crystal field theory - tetrahedral and square planar complexes

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Crystal Field Splitting in Non-Octahedral Shapes

Crystal field theory also explains d orbital splitting in tetrahedral and square planar complexes. These shapes do not follow the same pattern as an octahedral complex, so the d orbitals respond to ligand placement in different ways. In tetrahedral complexes, the four ligands sit between the axes. None of the d orbitals points directly at the ligands, but the d x2−y2 and d z2 orbitals lie along the Cartesian axes and overlap less with the ligands than the d xy, d xz, and d yz orbitals. The d...

colors and magnetism

Color and Spin in Coordination Complexes

Color in coordination complexes comes from how metal electrons absorb light. When atoms or molecules absorb light at the right frequency, electrons move to higher-energy orbitals. For many main-group atoms and molecules, the absorbed light is ultraviolet and cannot be seen. In coordination compounds, the energy gap between d orbitals often matches visible light, so the compound can absorb and emit colors that the human eye detects. Small changes in orbital energy can cause large shifts in the...

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Learn where transition metals and inner transition metals are found, why they form many oxidation states, and how their ions, compounds, and colors differ.

Transition Metals in the Periodic Table | JoVE K12

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Ions and Ionic Charges

In ordinary chemical reactions, the nucleus &mdash; which contains the protons and neutrons of each atom and thus identifies the element &mdash; remains ...

Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outermost energy level and the composition of the nucleus of atoms.  Clarification Statement: Chemistry: Examples of properties that could be predicted from patterns could include reactivity of metals, types of bonds formed, numbers of bonds formed, atomic radius, atomic mass, or reactions with oxygen. Emphasis is on main group elements and qualitative understanding of the relative trends of ionization energy and electronegativity.

HS-PS1-1

The Atomic Theory of Matter

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers. Leucippus and Democritus argued that all matter ...

Chemical Reactions

A chemical reaction is a process by which the bonds in the atoms of substances are rearranged to generate new substances. Matter cannot be created or ...

Chemical Equations

Chemical equations represent the identities and relative quantities of substances involved in a chemical reaction. The substances undergoing reaction are ...

Introduction to Chemical Reactions

All chemical reactions begin with a reactant, the general term for one or more substances entering the reaction. Sodium and chloride ions, for example, ...

Temperature Dependence on Reaction Rate

The Collision Theory
Atoms, molecules, or ions must collide before they can react with each other. Atoms must be close together to form chemical bonds. ...

stoichiometry-product-yield-and-limiting-reactants

Stoichiometry, Product Yield, and Limiting Reactants

Source: Smaa Koraym at Johns Hopkins University, MD, USA

	Preparation of 25% v/v Ethylenediamine in Deionized WaterExpand
	Here, we show the laboratory ...

physical-and-chemical-properties-of-matter

Physical and Chemical Properties of Matter

The characteristics that enable us to distinguish one substance from another are called properties.

Physical Properties of Matter

A physical property is ...

Dynamic Equilibrium

A reversible chemical reaction represents a chemical process that proceeds in both forward (left to right) and reverse (right to left) directions. When ...

Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.  Clarification Statement: Chemistry: Examples of chemical reactions could include the reaction of sodium and chlorine, carbon and oxygen, or carbon and hydrogen. Reaction classification aids in the prediction of products (e.g. synthesis, decomposition, single displacement, double displacement, and acid-base).

Intermolecular vs Intramolecular Forces

Intermolecular forces (IMF) are electrostatic attractions arising from charge-charge interactions between molecules. The strength of the intermolecular ...

comparison-between-electrical-and-gravitational-forces

Comparison Between Electrical And Gravitational Forces

There are four fundamental forces in nature: the gravitational force, the electromagnetic force, the strong nuclear force, and the weak nuclear force. To ...

intermolecular-forces-and-physical-properties

Intermolecular Forces and Physical Properties



Intermolecular forces are attractive forces that exist between molecules. They dictate several bulk properties, such as melting points, boiling points, ...

electric-field-of-two-equal-and-opposite-charges

Electric Field of Two Equal and Opposite Charges

Atoms generally contain the same number of positively and negatively charged particles, protons, and electrons. Hence, they are electrically neutral. ...

Electric Field

Consider two point charges, each exerting Coulomb force on the other. It is possible to describe the Coulomb interaction via an intermediate step by ...

Ionic Bonding and Electron Transfer

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence ...

Electric Dipoles and Dipole Moment

Consider two charges of equal magnitude but opposite signs. If they cannot be separated by an external electric field, the system is called a permanent ...

Plan and conduct an investigation to gather evidence to compare the structure of substances at the macroscale to infer the strength of electrical forces between particles.  Clarification Statement: Emphasis is on understanding the strengths of forces between particles, not on naming specific intermolecular forces (such as dipole-dipole). Examples of particles could include ions, atoms, molecules, and network solids (such as graphite). Examples of macro-properties of substances could include the melting point and boiling point, vapor pressure, and surface tension.

HS-PS1-3

bond-dissociation-energy-and-activation-energy

Bond Dissociation Energy and Activation Energy

Bond energy is the energy required to break a bond homolytically. These values are usually expressed in units of kcal/mol or kJ/mol and are referred to as ...

Bond Energies and Bond Lengths

Stable molecules exist because covalent bonds hold the atoms together. The strength of a covalent bond is measured by the energy required to break it, ...

Coupled Reactions

Cellular processes such as building and breaking down complex molecules occur through stepwise chemical reactions. Some of these chemical reactions are ...

energy-diagrams-transition-states-and-intermediates

Energy Diagrams, Transition States, and Intermediates

Free-energy diagrams, or reaction coordinate diagrams, are graphs showing the energy changes that occur during a chemical reaction. The reaction ...

What is Metabolism?

Overview

Metabolism represents all of the chemical activity in a cell, including reactions that build molecules (anabolism) and those that break ...

types-of-chemical-reactions-anabolic-and-catabolic

Types of Chemical Reactions: Anabolic and Catabolic

The first law of thermodynamics holds that energy can neither be created nor destroyed&#8212;it can only change form. An organism&#39;s essential function ...

Non-equilibrium in the Cell

An important concept in studying metabolism and energy is that of chemical equilibrium. Most chemical reactions are reversible. They can proceed in both ...

Activation Energy

Activation energy is the minimum amount of energy necessary for a chemical reaction to move forward. The higher the activation energy, the slower the rate ...

Energy Transfer in Chemical Reactions

Chemical reactions require sufficient energy to cause the matter to collide with enough precision and force that old chemical bonds can be broken and new ...

factors-influencing-the-rate-of-chemical-reactions

Factors Influencing the Rate of Chemical Reactions

A variety of factors influence the rate of chemical reactions. For a chemical reaction to happen, atoms must collide with enough energy to overcome the ...

Predicting Reaction Outcomes

Kinetics describes the rate and path by which a reaction occurs. In contrast, thermodynamics deals with state functions and describes the properties, ...

States of Matter and Phase Changes

The internal energy of a substance&#8212;the total kinetic energy of all its molecules and the potential energy of their associated forces&#8212;depends ...

Reaction Mechanisms

Chemical reactions often occur in a stepwise fashion, involving two or more distinct reactions taking place in a sequence. A balanced equation indicates ...

Temperature Dependence

Chemical Kinetics 
The reaction rate is the speed at which a chemical reaction occurs. The reaction rate is defined as the change in concentration of a ...

Molecular Kinetic Energy

The word &#34;gas&#34;&#160;comes from the Flemish word meaning &#34;chaos,&#34; first used to describe vapors by the chemist J. B. van Helmont. Consider ...

Develop a model to illustrate that the release or absorption of energy from a chemical reaction system depends upon the changes in total bond energy.  Clarification Statement: Emphasis is on the idea that a chemical reaction is a system that affects the energy change.  Examples of models could include molecular-level drawings and diagrams of reactions, graphs showing the relative energies of reactants and products, and representations showing energy is conserved.

HS-PS1-4

Transition Metals in the Periodic Table

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