Development of the Periodic Table
The periodic table is a methodical arrangement of the chemical elements, organized on the basis of their electron configurations.
Discuss the origins and history of the periodic table.
- Although the work of alchemists was originally a misguided effort to convert lead into silver and gold, their studies laid a foundation that aided a later fundamental understanding of matter.
- The modern periodic table was devised by Dmitri Mendeleev and is a useful framework for organizing and analyzing chemical and physical behavior of the elements.
- The notation in the periodic table includes references to atomic mass and atomic number.
- proton: A positively charged subatomic particle forming part of the nucleus of an atom and determining the atomic number of an element; the nucleus of the most common isotope of hydrogen, composed of two up quarks and a down quark.
- element: Any one of the simplest chemical substances that cannot be decomposed in a chemical reaction or by any chemical means, made up of atoms all having the same number of protons.
- alchemy: The ancient search for a universal panacea and for the philosopher’s stone. The process eventually developed into chemistry.
The modern periodic table organizes the known elements in several ways: it lists them in order of patterns of atomic weight, electron configuration, reactivity, and electronegativity. It is such a good method of organizing and presenting the known elements that it has been used to successfully predict the existence of certain elements. Today, it is applied not only by chemists but also in all related sciences to understand the properties and reactivity of atoms and molecules. The table has recognizable origins in the 17th century and draws on knowledge and experience of medieval and earlier eras.
History of the Periodic Table
Atomic theory dates back to the ancient Greek philosophers and those of Hellenistic Egypt. They theorized that all substances were made of fundamental building blocks; however, the nature of those blocks was the object of fierce debate.
The fundamental blocks were called atoms, derived from the Greek “atmos,” meaning “indivisible.” Early atomic theory attempted to explain properties of matter by assigning attributes to atoms that might match the attributes of the various matter they combined to form, such as slipperiness, liquidity, color, and cohesiveness. Philosophers categorized the world around them by property and function, a type of approach that later led to the development of the periodic table of elements.
In the Middle Ages, practitioners of alchemy sought to make gold and silver from lead. Although their efforts were in vain, their investigation has ultimately led to a systematic understanding of the chemical world. It also established the mindset that gave us the periodic table of elements.
Alchemists were influenced by international trade, especially along the Silk Road between China and Europe. Chemical knowledge spread across cultures, and by about the middle of the 18th century, there were already 33 known elements. At the beginning of the 19th century, Joseph Proust and others were demonstrating the Law of Definite Proportions experimentally. This provided fundamental evidence that matter existed in pure compounds as opposed to just mixtures of any proportion. These observations strengthened the atomic theory and demanded a systematic method of organizing the elements.
The Modern View of the Periodic Table
Scientists began to notice similarities and patterns among known elements, and a great research interest of the 19th century was to develop a systematic method to report and classify them. Russian chemistry professor Dmitri Mendeleev and German chemist Julius Meyer independently presented their own versions of the periodic table in 1869 and 1870. Mendeleev’s approach was ultimately adopted for several reasons: For one, he left gaps for elements that had yet to be discovered. In doing so, he predicted the elements gallium and germanium. He also placed atoms based principally on their chemical properties, not atomic mass. As it turns out, organizing by chemical family correctly sorts most of the elements by their atomic number; atomic mass is not perfectly correlated with atomic number.
The modern version of Mendeleev’s periodic table now contains some 118 different elements. In the periodic table, the number above the element’s symbol is the atomic number, which represents the number of protons in the nucleus. The atomic mass is given by the sum of the neutrons and protons.
Periods 1 through 3
Elements of the same period have the same number of electron shells.
Discuss the relationship between an atom’s electron structure and its period (row) on the periodic table.
- As you move through a period (across the table to the right), the electron shells of the elements in that period are filling up, approaching the stable configuration of the noble gas at the end of that row.
- For any element in periods 1, 2, and 3, the elements directly above and below it are members of the same group and have similar chemical properties based on similar arrangements of valence electrons.
- The Aufbau principle describes how electrons are put into orbitals in a particular order for filling.
- atom: The smallest possible amount of matter that still retains its identity as a chemical element, now known to consist of a nucleus surrounded by electrons.
- electron: The subatomic particle having a negative charge and orbiting the nucleus; the flow of electrons in a conductor constitutes electricity.
- electron shell: The collective states of all electrons in an atom having the same principal quantum number (visualized as an orbit in which the electrons move).
Early philosophers and scientists appreciated that matter was composed of atoms and that many elements reacted in predictable proportions to each other. The periodic table was constructed in order to organize those observations and measurements. The principle of valence emerged, attributable to the presence or absence of electrons and the energy of those electrons in the volume around an atom’s nucleus. Electrons, negatively charged subatomic particles, define an atom’s chemical reactivity. Electron are organized in energy levels or electron shells, which correspond to the periods on the periodic table.
The Bohr Atom
Neils Bohr proposed a simplified picture of an atom, with a central nucleus surrounded by electrons in specific energy levels (n). The periodic table codifies the energy levels in periods, the rows on the table. The simplest atoms, hydrogen and helium, are found in row 1, or the first period. These atoms have electrons occupying the energy level n=1. Moving down, row 2, or period 2, contains the elements Li (lithium) through Ne (neon). The elements in period 2 have their level n=1 energy completely filled; they proceed to fill their n=2 level moving across the table to the right. In a similar fashion, moving down one period to row 3, there are the elements Na (sodium) through Ar (argon). The period-3 atoms have levels n=1 and n=2 filled; they are populating the n=3 level moving across the table.
It is important to remember that the periodic table is a representation of atoms with zero net charge; they have as many electrons around the nucleus as they have protons in the nucleus.
The Aufbau Principle
In the n=1, n=2, and n=3 energy levels, electrons are organized in orbitals, designated as s, p, d, and f. For example, the atomic number of Ne (neon) is 10 and contains 5 orbitals (1s, 2s, 2px, 2py, and 2pz). In each full orbital, there are 2 electrons, giving a total of 10 to balance the positive charge provided by the 10 protons in the nucleus.
In the periodic table, there are 2 electrons in period 1, while both periods 2 and 3 have 8 electrons in the filled level. For atoms with atomic numbers less than about 20, the octet rule of electron addition and orbital filling applies. This simply states that the n=2 and n=3 levels, in particular, are full when there are 8 electrons. The Aufbau principal describes how electrons are put into orbitals in a particular order for filling.
The d-block elements are commonly known as transition metals or transition elements.
Identify the distinctive and characteristic properties of the transition metals.
- Transition metals are elements in the ten middle groups of the fourth, fifth, sixth, and seventh periods of the periodic table.
- Transition metals and their compounds can exhibit color due to internal d-d electron transfers.
- Transition metals and their compounds can exhibit ferromagnetism, paramagnetism, and diamagnetism.
- Transition metals and their compounds are well known for their catalytic activities.
- paramagnetic: Exhibiting paramagnetism; the tendency of magnetic dipoles to align with an external magnetic field.
- Oxidation State: The state of an atom having a particular oxidation number.
- diamagnetic: Exhibiting diamagnetism; repelled by a magnet.
The d-Block of the Periodic Table
The transition metals are also known as thetransition elements or the d-block elements. As the name implies, the chemistry of this group is determined by the extent to which the d-electron suborbital levels are filled. Chemical similarities and periodicities can be easily seen horizontally across the d-block of the periodic table.
The chemistry is far from simple, however, and there are many exceptions to the orderly filling of the electron shell. The Aufbau principle provides an methodical framework for predicting the order in which most atoms will populate their electron shells.
Chemical properties in the periodic table are organized vertically, by group, for similar chemical and physical properties. For example, the metals in group 11 have similar characteristics of electrical conductivity, luster, crystal structure, ductility, and tensile strength. Moving horizontally across the periodic table trends in properties such as atomic radius, electronegativity, and electron affinity are observed.
Characteristic Properties of Transition Metals
Transition metals can be said to possess the following characteristics generally not found in the main grouping of the periodic table. They can be mostly attributed to incomplete filling of the electron d-levels:
- The formation of compounds whose color is due to d–d electronic transitions.
- The formation of compounds in many oxidation states due to the relatively low reactivity of unpaired d electrons.
- The formation of many paramagnetic compounds due to the presence of unpaired d electrons. A few compounds of main group elements are also paramagnetic (e.g., nitric oxide, oxygen).
Ligand -to-Metal Charge-Transfer (LMCT) Transition
Color in transition-series metal compounds is generally due to the electronic transitions of two principal types of charge transfer transitions. An electron may jump from a predominantly ligand orbital to a predominantly metal orbital, giving rise to a ligand-to-metal charge transfer (LMCT) transition. These can most easily occur when the metal is in a high oxidation state. For example, the color of chromate, dichromate, and permanganate ions is due to LMCT transitions. Another example is that mercuric iodide (HgI2) is red because of a LMCT transition.
A metal-to-ligand charge transfer (MLCT) transition will be most likely when the metal is in a low oxidation state and the ligand is an easily reduced d–d transition. An electron jumps from one d-orbital to another. In complexes of the transition metals, the d orbitals do not all have the same energy.
Paramagnetic and Diamagnetic Compounds
Transition metal compounds are paramagnetic when they have one or more unpaired d electrons. Some compounds are diamagnetic. These include octahedral, low-spin, d6 and square-planar d8complexes. In these cases, crystal field splitting is such that all the electrons are paired up. Ferromagnetism occurs when individual atoms are paramagnetic and the spin vectors are aligned parallel to each other in a crystalline material. Metallic iron and the alloy alnico are examples of ferromagnetic materials involving transition metals. Anti-ferromagnetism is another example of a magnetic property arising from a particular alignment of individual spins in the solid state.
The transition metals and their compounds are known for their homogeneous and heterogeneous catalytic activity. This activity is attributed to their ability to adopt multiple oxidation states and to form complexes.
The Bottom of the Periodic Table
The periodic table currently contains 7 periods, but theorists predict that two additional periods may exist.
Identify the key characteristics of the f-block elements.
- The lanthanide and actinide series derive properties from the f-block electrons.
- Four elements in the actinide series are naturally occurring, five isotopes of others are occasionally produced by decay of uranium, while the rest of the transuranics have been synthetically produced.
- The highest atomic number synthesized to date is 118, the element ununoctium (Uuo).
- lanthanide: Any of the 14 rare earth elements from lanthanum to lutetium in the periodic table. Because their outermost orbitals are empty, they have very similar chemistry. Below them are the actinides.
- actinide: Any of the 14 radioactive elements of the periodic table that are positioned under the lanthanides, with which they share similar chemistry.
- transuranium: Transuranic. A transuranium element is any synthetic element having an atomic number greater than that of uranium (92).
There are a few ways to approach this particular topic, and they all refer to how the elements on the table itself are presented.
The most classic representation of the periodic table shows the relative positions of the known elements in the table. The table itself is comprised of 7 periods and 18 groups, with the latest known element being number 118, ununoctium. However, there is a glaring discontinuity apparent in the table. In Row 6, Column 3, an empty space appears between Ba and Hf. The atomic number that should be here, 57, is located at the bottom of the table in the row called the Lanthanides. Directly below the space in Row 6, in Row 7, is another empty space, which is filled by a row called the Actinides, also seen at the bottom of the chart.
Expanding the Dimensions of the Periodic Table
By expanding the horizontal dimensions of the table, the actinide and lanthanide rows can be put into their correct relative positions. Since the chemistries of this group rest largely on the f-shell electrons and the interactions at this energy level, this is called the f-block. This representation, clumsy as it is, correctly shows the elements known to date, up to z=118, unonoctium. In fact, this representation is predictive in that it shows chemical families (groups) and the periodicities (periods) in their correct relative positions.
Taking the extension of the periodic table even further, consider an element with atomic number 92 in the actinide series, called uranium. Once elements of this atomic number range were discovered at the end of the 19th century, isotopes of uranium were the largest and heaviest elements known in nature. In 1934, Enrico Fermi predicted the existence of transuranium elements—those elements with atomic number (z) greater than or equal to 93. In 1934 only 4 actinides were known, all smaller than uranium, so it was not known that they formed a period or family like the lanthanides. The first transuranium element, Np (neptunium), was synthetically produced in 1940 by bombarding uranium with slow neutrons. Over the next two decades, a great many actinide isotopes were produced, generally by bombardment with either other atoms or subatomic particles. The actinides were added along with the lanthanides.
Two New Periods
By using the predictive properties of the periodic table, along with a growing expertise in atomic and subatomic theory, two entirely new periods were predicted. On the advice of Glenn Seaborg and others, Periods 8 and 9 were added to the periodic table, comprising the g-block. The positioning of the g-block in the table (to the left of the f-block, to the right, or in between) is speculative. The positions in the table correspond to the assumption that the Madelung rule (that orbitals with lower value of the sum of n and l quantum numbers will be filled before those of higher n+l values) will continue to hold for higher atomic numbers. At element 118, the orbitals 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f, 5s, 5p, 5d, 5f, 6s, 6p, 6d, 7s, and 7p are assumed to be filled, with the remaining orbitals unfilled. The orbitals of the eighth period are predicted to be filled in the order 8s, 5g, 6f, 7d, 8p. However, after approximately element 120, the proximity of the electron shells makes placement in a simple table problematic.
The existence of elements with these high atomic numbers is speculative, and isotopes are expected to have fleetingly short half-lives. Various experts predict that z = approximately 130 is a maximum, while others feel that there is no effective upper limit. Experiments in the synthesis of transuranium elements continue.
Periodic Table Position and Electron Configuration
The position of elements on the periodic table is directly related to their electron configurations.
Use the periodic table to identify atom properties such as groups and electron configurations.
- Elements are organized by period and group, with the period corresponding to the principle energy level, and the group relating to the extent the subshells are filled.
- The properties of an atom relate directly to the number of electrons in various orbitals, and the periodic table is much like a road map among those orbitals such that chemical properties can be deduced by the position of an element on the table.
- The electrons in the outermost or valence shell are especially important because they can engage in the sharing and exchange that is responsible for chemical reactions.
- quantum number: One of certain integers or half-integers that specify the state of a quantum mechanical system (such as an electron in an atom).
- orbital: A specification of the energy and probability density of an electron at any point in an atom or molecule.
- electron shell: The collective states of all electrons in an atom having the same principal quantum number (visualized as an orbit in which the electrons move).
Major Divisions of the Periodic Table
The periodic table is a tabular display of the chemical elements organized on the basis of their atomic numbers, electron configurations, and chemical properties. Elements are presented in increasing atomic number. The main body of the table is a 18 × 7 grid. Elements with the same number of valence electrons are kept together in groups, such as the halogens and the noble gases. There are four distinct rectangular areas or blocks. The f-block is usually not included in the main table, but rather is floated below, as an inline f-block would often make the table impractically wide. Using periodic trends, the periodic table can help predict the properties of various elements and the relations between properties. It therefore provides a useful framework for analyzing chemical behavior and is widely used in chemistry and other sciences.
The electrons in the partially filled outermost shell (or shells) determine the chemical properties of the atom; it is called the valence shell. Each shell consists of one or more subshells, and each subshell consists of one or more atomic orbitals.
The properties of an atom depend ultimately on the number of electrons in the various orbitals, and on the nuclear charge which determines the compactness of the orbitals. In order to relate the properties of the elements to their locations in the periodic table, it is often convenient to make use of a simplified view of the atom in which the nucleus is surrounded by one or more concentric spherical “shells,” each of which consists of the highest-principal quantum number orbitals that contain at least one electron; these are s- and p-orbitals and can include d- or f-orbitals, which is atom dependent. The shell model, as with any scientific model, is less a description of the world than a simplified way of looking at it that helps us to understand and correlate diverse phenomena.
We will look at several visualizations of the periodic table. First, however, it would be instructive to see how it is constructed from a logical viewpoint. The table today is the result of an ongoing effort of more than 100 years of observation, measurement, prediction and proof of the relationships of chemical and physical phenomena to electron configurations and charges.
Periods 1, 2, & 3
Starting with simple elements, the first three rows of the periodic table, called Periods 1, 2 and 3, correspond to the n=1, n=2 and n=3 levels.
Hydrogen has 1 electron in the 1s level, and to the right, helium, in Group 18, has 2 electrons in the 1s level, a completely filled shell, the duet rule. Helium is the first in the series of noble gases. Moving down to Period 2, lithium is the first element in the row, with a filled 1s configuration. Across the period, first the 2s and then the 2p orbitals fill, arriving at the configuration for neon, following the octet rule. Period 3 follows a similar pattern. Please note that the number of outer-shell electrons is the major determinant of the element’s valence.
Electron Configuration of Cations and Anions
The elements on the periodic table exhibit different levels of reactivity based on the number of electrons in their highest energy shells.
Predict whether an atom will undergo ionization to provide an anion or cation based on its valence shell electron configuration.
- The electronic configuration of many ions is that of the closest noble gas to them in the periodic table.
- An anion is an ion that has gained one or more electrons, acquiring a negative charge.
- A cation is an ion that has lost one or more electrons, gaining a positive charge.
- anion: A negatively charged ion, as opposed to a cation
- ionization: Any process that leads to the dissociation of a neutral atom or molecule into charged particles (ions).
- cation: A positively charged ion, as opposed to an anion.
Cations and Anions form from Neutral Atoms
Every atom in its ground state is uncharged. It has, according to its atomic number, the same number of protons and electrons. Electrons are rather labile, however, and an atom will often gain or lose them depending on its electronegativity. The driving force for such gain or loss of electrons is the energetically optimal state of having a full valence (outermost) shell of electrons. In such a state, the resulting charged atom has the electron configuration of a noble gas.
Addition of an electron will disrupt the proton-electron balance and leave the atom negatively charged. Removal of an electron will, conversely, leave the atom positively charged. These charged atoms are known as ions.
Formation of Monatomic Ions
Monatomic ions are formed by the addition or removal of electrons from an atom’s valence shell. The inner shells of an atom are filled with electrons that are tightly bound to the positively charged atomic nucleus and so do not participate in this kind of chemical interaction, but the valence shell can be very reactive depending on the atom and its electron configuration. The process of gaining or losing electrons from a neutral atom or molecule is called ionization.
Atoms can be ionized by bombardment with radiation, but the more purely chemical process of ionization is the transfer of electrons between atoms or molecules. This transfer is driven by the stabilization that comes by obtaining stable (full shell) electronic configurations. Atoms will gain or lose electrons depending on which action takes the least energy.
For example, Group 1 element sodium (Na) has a single electron in its valence shell, with full shells of 2 and 8 electrons beneath. Removal of this one electron leaves sodium stable: Its outermost shell now contains eight electrons, giving sodium the electron configuration of neon. Having gained a positive charge, the sodium ion is called a cation. The ionization of sodium can be chemically illustrated as follows:
Na → Na+ + e−
Sodium could gain electrons, but it would require seven more to achieve a full valence shell. Removing one electron is much easier than gaining seven, and thus sodium will in every chemical scenario achieve its octet by becoming a cation.
On the other hand, a chlorine atom (Cl) has seven electrons in its valence shell, which is one short of a stable, full shell with 8 electrons. Thus, a chlorine atom tends to gain an extra electron and attain a stable 8-electron configuration (the same as that of argon), becoming a negative chloride anion in the process:
Cl + e− → Cl−
Combining the propensity of sodium to lose an electron and of chloride to gain an electron, we observe complimentary reactivity. When combined, the uncharged atoms can exchange electrons and in doing so, achieve complete valence shells. The resulting ions stick together due to ionic bonds (opposite charges attract), leaving a crystal lattice structure of NaCl, more commonly known as rock salt. The reaction is as follows:
Na+ + Cl− → NaCl
Polyatomic and Molecular Ions
Ionization is not limited to individual atoms; polyatomic ions can also be formed. Polyatomic and molecular ions are often created by the addition or removal of elemental ions such as H+ in neutral molecules. For example, when ammonia, NH3, accepts a proton, H+, it forms the ammonium ion, NH4+. Ammonia and ammonium have the same number of electrons in essentially the same electronic configuration, but ammonium has an extra proton (the H+) that gives it a net positive charge.
When writing the chemical formula for an ion, its net charge is written in superscript immediately after the chemical structure for the molecule or atom. The net charge is written with the magnitude before the sign, that is, a doubly charged cation is indicated as 2+ instead of +2. However, the magnitude of the charge is omitted for singly charged molecules or atoms; for example, the sodium cation is indicated as Na+ and not Na1+.
An alternative way of showing a molecule or atom with multiple charges is by drawing out the signs multiple times; this is often seen with transition metals. Chemists sometimes circle the sign; this is merely ornamental and does not alter the chemical meaning. A twice-positively charged iron atom can also be expressed as Fe2+ or Fe++.
In the case of transition metals, oxidation states can be specified with Roman numerals; for example, Fe2+ is occasionally referred to as Fe(II) or FeII. The Roman numeral designates the formal oxidation state of an element, whereas the superscripted numerals denotes the net charge. The two notations are therefore exchangeable for monatomic ions, but the Roman numerals cannot be applied to polyatomic ions. However, it is possible to mix the notations for the individual metal center with a polyatomic complex, as demonstrated using the uranyl ion (UO2) as an example.
It should be noted that it is possible to remove many electrons from an atom. The energy required to do so may be recorded in a successive ionization energy diagram.