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## Periodic Properties

Copper and gold – the two colored metals In periodic table many metals are colorless(shiny white) and some metals have yellowish or bluish tinge. But particularly two metals, copper and gold, have distinct color. Copper is reddish brown and gold is yellow orange in color. The color of metals copper and gold is mainly due to absorption of light when a d electron jumps to an s orbital. Copper has a strong absorption at a slightly lower energy, with orange being most strongly absorbed and re‐emitted. Down the group there must be a increase in the gap between d and s orbitals. As usual the silver is white in color since an increase in energy absorption moves to UV region. But in case of gold due to relativistic correction the energy gap decreases and the electron emits in yellow‐orange region.

## After completing the topic, the student will be able to:

• Define the periodic properties such as atomic radius, ionization energy, electron affinity and electronegativity.
• Distinguish van der Waals, covalent and atomic radii.
• Explain the reason for the high magnitudes of ionization energy and electron affinity in some elements.
• Recognize the periodic trends in properties of elements.
• Explain the reasons for trends in the reactivity and behavior of elements in the periodic table.
• Compare the properties of isoelectronic species.
• Explain the relationship between ionization enthalpy and metallic character.
• Understand the effect of electronegativity on chemical reactivity of elements.
Periodic table Elements were discovered over the centuries. It took nearly 3 centuries to give the present look of the periodic table. This is a very well constructed table which can give a panoramic view of the properties of elements.
Periodic Properties

The periodic table is a table of the chemical elements in which the elements are arranged in order according to atomic number in such a way that the periodic properties (chemical periodicity) of the elements are made clear.

The standard form of the table includes periods (usually horizontal in the periodic table) and groups (usually vertical). Elements in groups have some similar properties to each other. There is no one single or best structure for the periodic table but by whatever consensus there is, the form used here is very useful.

The periodic table is a masterpiece of organized chemical information. The evolution of chemistry's periodic table into the current form is an astonishing achievement with major contributions from many famous chemists and other eminent scientists. Of the 116 elements in periodic table, chemists have grouped them as alkali metals, alkaline earth metals, transition elements, inner transition elements, p‐block elements (excluding halogens), halogens and noble gases. The location of these seven groups is in an inter‐linked manner based on their physical and chemical properties. For example, the 1st group of elements in periodic table is named for the alkaline (or basic) nature of their oxides and for the basic solutions the elements form in water.

General outer electronic configuration Elements present in groups will have a general electronic configuration which would be followed by the elements present in the periodic table. Given above is the outer electronic configuration of the groups present in the table.
Periodic trends

The properties of the elements in a periodic table exhibit trends. These trends can be predicted using the periodic table and can be explained and understood by analyzing the electronic configurations of the elements.

Elements tend to gain or lose valence electrons to achieve stable octet configuration. Stable octets are seen in the inert gases or noble gases, of Group VIII of the periodic table. In addition to this activity, there are two other important trends. First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it.

Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels (which shield the outermost electrons from attraction to the nucleus) increases downward within each group. These trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity and electronegativity.

By observing the trends in the atomic radii of the elements in a period and those within a group of the periodic table we can predict the relative sizes of atomic and ionic radii within an isoelectronic series. We can also estimate certain physical properties of a particular element if the properties of elements above and below the given element in a group of elements are known.

Let us explore the variations in atomic radius, ionization energies, electron affinities in a periodic table.

Trends of atomic radius in periodic table Atomic radius decreases from left to right in the periodic table. Alkali metals will have highest possible atomic radius in the period. It increases down the group as the extra shells are added each time. Cesium (Cs‐265) is the metal with highest electronic configuration.

In 1920, shortly after it had become possible to determine the sizes of atoms using X‐ray crystallography, it was suggested that all atoms of the same element have the same radii. However, in 1923, when more crystal data had become available, it was found that the approximation of an atom as a sphere does not necessarily hold when comparing the same atom in different crystal structures. The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean or typical distance from the nucleus to the boundary of the surrounding cloud of electrons.

Atomic radius decreases, in general, as the atomic number increases across any period of the periodic table. Thus, the alkali metals have the largest atoms and the noble gases the smallest. Furthermore, within a group, the atomic radius increases as the atomic number increases. Another trend in chemical behavior is the diagonal relationship. This refers to similarities in some properties between a second row element, and the third row element that is one space to the right in the table. Two such elements are considered to be on a diagonal.

For example, Li and Mg are on a diagonal, and they have similar atomic radii. At Lithium to Beryllium the atomic radius decreases. Next, on moving down to Mg, the radius will increase again. If the decrease and increase values are about the same, then Li and Mg will have approximately the same radius (155 pm for Li versus 160 pm for Mg). This relationship does not hold for all properties of two diagonal elements.

Van der Waals radius The half of the distance between the nuclei of the two non bonded atoms of adjacently placed molecules is the van der Waal radius.

Since the boundary is not a well‐defined physical entity, there are various non‐equivalent definitions of atomic radius. Depending on the definition, the term may apply only to isolated atoms, or also to atoms in condensed matter, covalently bound in molecules, or in ionized and excited states, and its value may be obtained through experimental measurements, or computed from theoretical models. Under some definitions, the value of the radius may depend on the atom's state and context. Widely used definitions of atomic radius include:

Van der Waals radius: It is half the minimum distance between the nuclei of two atoms of the element that are not bound to the same molecule.

Ionic radius Ionic radius is the measure of inter atomic distance between the two adjacent ions in the crystal lattice.

It is the nominal radius of the ions of an element in a specific ionization state, deduced from the spacing of atomic nuclei in crystalline salts that include that ion. In principle, the spacing between two adjacent oppositely charged ions (the length of the ionic bond between them) should equal the sum of their ionic radii.

It is the nominal radius of the atoms of an element when covalently bound to other atoms, as deduced the separation between the atomic nuclei in molecules. In principle, the distance between two atoms that are bound to each other in a molecule (the length of that covalent bond) should equal the sum of their covalent radii.

Metallic radius Metallic radius is defined as one‐half of the distance between neighboring atoms in the metal.

It is the nominal radius of atoms of an element when joined to other atoms by metallic bonds.

Bohr radius: It is the radius of the lowest‐energy electron orbit predicted by Bohr model of the atom (1913). It is only applicable to atoms and ions with a single electron, such as hydrogen, singly ionized helium, and positronium. Although the model itself is now obsolete, the Bohr radius for the hydrogen atom is still regarded as an important physical constant.

Atomic trend in a periodic table: Size increases down a column. Size decreases across (to the right of) a row.

The number of energy levels increases down a group as the number of electrons increases. Each subsequent energy level is further from the nucleus than the last. Therefore, the atomic radius increases as the group and energy levels increase. Across a period, electrons are added to the same energy level. At the same time, protons are being added to the nucleus. The concentration of more protons in the nucleus creates a "higher effective nuclear charge". In other words, there is a stronger force of attraction pulling the electrons closer to the nucleus resulting in a smaller atomic radius.

Atomic radii of elements in picometers
Atomic radius trends among main-group elements

Atomic size greatly influences other atomic properties and is critical to understand element behavior. The adjacent figure shows the atomic radii of the main‐group elements and most of the transition elements. Among the main‐group elements, note that atomic size varies within both a group and period. Atomic radius generally increases from top-to-bottom and decreases from left-to-right. Across a period, as the electrons are added to the same shell effective nuclear charge increases leading to the decrease in size due to pull of electrons by nucleus.

These variations in atomic size are the result of two opposing influences:

• Changes in ‘n’: As the principal quantum number (n) increases, the probability that the outer electrons will spend more time farther from the nucleus increases as well, thus, the atoms are larger.
• Changes in Zeff: As the effective nuclear charge (Zeff ‐ the positive charge "felt" by an electron) increases, outer electrons are pulled closer to the nucleus, thus the atoms are smaller.
Atomic radius of alkali metals and noble gases Alkali metals will have highest possible atomic radius values whereas the noble gases will have the least possible values.
Effect of principal quantum number and Z eff

The net effect of changes in the principal quantum number (n) and change in effective nuclear charge (Zeff) influences depends on shielding of the increasing nuclear charge by inner electrons:

• Down the group, ‘n’ dominates. As we move down a main group, each member has one more level of inner electrons that shield the outer electrons very effectively. Even though calculations show Zeff on the outer electrons rising moderately for each element in the group, the atoms get larger as a result of the increasing ‘n’ value. Atomic radius generally increases in a group from top to bottom.
• Across a period, Zeff dominates. As we move across a period of main‐group elements, electrons are added to the same outer level, so the shielding by inner electrons does not change. Because outer electrons shield each other poorly, Zeff on the outer electrons rises significantly, and so they are pulled closer to the nucleus. Atomic radius generally decreases in a period from left to right.

These trends hold well for the main‐group elements but not as consistently for the transition elements.

Atomic radii of transition elements Transition elements do not exhibit the trend consistently. They vary irregularly. This is due to the counter effect of shielding due to the inner d‐electrons to the increase in effective nuclear charge which increases with addition of electrons. But this trend is not followed by the f‐block elements. The size decreases from ‘Ce’ to ‘Lu’ as a result of lanthanide contraction. The effect results from poor shielding of nuclear charge by 4f electrons. The 6s electrons are drawn towards the nucleus, thus resulting in a smaller atomic radius.
Atomic radius trends among the transition elements

As we move from left to right, size shrinks through the first two or three transition elements because of the increasing nuclear charge. But, from then on, the size remains relatively constant because shielding by the inner 'd' electrons counteracts the usual increase in Zeff. For instance, vanadium (V, Z = 23), the third period 4th transition metal, has the same atomic radius as zinc (Zn, Z = 30), the last period 4th transition metal. This pattern of atomic size shrinking also appears in periods 5 and 6 in the d‐block transition series and in both series of inner transition elements. The lack of a vertical size increase from the period 5 to 6 transition metal is especially obvious.

This shielding by 'd' electrons causes a major size decreases from Group IIA (2) to Group IIIA (13), the two main groups that flank the transition series. The size decrease in periods 4, 5, and 6 (with a transition series) is much greater than in period 3 (without a transition series). Because electrons in the np orbitals penetrate more than those in the (n−1) 'd' orbitals, the first np electron [Group IIIA (13)] "feels" a Zeff that has been increased by the protons added to all the intervening transition elements. The greatest change in size occurs in period 4, in which calcium (Ca, Z = 20) is nearly 50% larger than gallium (Ga, Z = 31). In fact, shielding by the 'd' orbitals in the transition series causes such a major size contraction that gallium is slightly smaller than aluminum (Al, Z = 13), even though ‘Ga’ is below ‘Al’ in the same group!

Atomic and ionic radii The atomic radii and ionic radii of some main group elements are arranged in periodic table format. Metal atoms are represented by blue color, they decrease in size when they lose the electron in the outermost shell. Non metal atoms are represented by red color which increases in size when an electron is added to their outermost shell.

The size of an atom or ion depends on the size of the nucleus and the number of electrons. Generally atoms with higher number of electrons have larger radii than those with smaller number of electrons. Thus, ions will have radii different from the atoms because ions will have either gained or lost electrons.

The number of positive charges in the nucleus determines both the number of electrons that surround an atom and the number of electrons that can be lost or gained to form ions. The relative size of atoms can also be studied by measuring the radii of their ions. There are two types of ions: cations and anions. Cations are atoms that have lost one or more electrons and carry a positive charge; anions are the atoms that have gained one or more electrons and carry a negative charge. Cations are always smaller than neutral atoms of their same element. Cations in which the entire shell has been lost are only half the size of the neutral atoms. This further decrease in size is due to the fact that cations have more protons than electrons.

On the other hand, anions are always larger than the neutral atoms. Many are almost twice the size. Since the added electrons go into the same shell, get shielded and because of which the electron cloud expands, resulting in the large anion radius.

• Thus as the charge on the ion becomes more positive, there will be less electrons and the ion will have a smaller radius.
• As the charge on the ion becomes more negative, there will be more electrons and the ion will have a larger radius.
• As the atomic number increases in any given column of the periodic table, the number of protons and electrons increases and thus the size of the atom or ion increases.
Arrangement of LiI in a crystal The cation radius (rLi+) and the anion (rI) radius each make up a portion of the total distance between the nuclei of adjacent ions a crystalline ionic compound. The interatomic distance between the two anions can also be found as 2rI.

Atomic and ionic radii also depend on the type of bonding that takes place between the constituents. Thus, atomic and ionic radii will vary somewhat as a function of the environment in which the atoms or ions are found. The first ionic radii were obtained by studying the structure of LiI (Lithium iodide) which contains a relatively small positive ion and a relatively large negative ion. The analysis of the structure of LiI was based on the following assumptions. The relatively small Li+ ions pack in the holes between the much larger I ions, the relatively large I ions touch one another. The Li+ ions touch the I ions.

If these assumptions are valid, the radius of the I ion can be estimated by measuring the distance between the nuclei of adjacent iodide ions. The radius of the Li+ ion can then be estimated by subtracting the radius of the I ion from the distance between the nuclei of adjacent Li+ and I ions.

Unfortunately, only two of the three assumptions that were made for LiI are correct. The Li+ ions in this crystal do not quite touch the I ions. As a result, this experiment overestimated the size of the Li+ ion. Repeating this analysis with a large number of ionic compounds, however, has made it possible to obtain a set of more accurate ionic radii. The negative ion is much larger than the atom from which it was formed. In fact, the negative ion can be more than twice as large as the neutral atom. For example, the covalent radius of neutral F, Cl, Br, and I atoms with the radii of their F, Cl Br, and I ions was much larger than the atom from which it was formed.

Cations are smaller than their parent atoms
Li 0.123 0.068
Na 0.157 0.095
K 0.2025 0.133
Rb 0.216 0.148
Cs 0.235 0.169
When a cation forms, electrons are removed from the outermost shell. The decrease of electron repulsions allows the nuclear charge to pull the remaining electrons closer.
Why anion radius is more than cation ?

The only difference between an atom and its ions is the number of electrons that surround the nucleus. E.g.: A neutral chlorine atom contains 17 electrons, while a Cl ion contains 18 electrons.

Cl : [Ne] 3s2 3p5
Cl : [Ne] 3s2 3p6

Because the nucleus can't hold the 18 electrons in the Cl ion as tightly as the 17 electrons in the neutral atom, the negative ion is significantly larger than the atom from which it forms. For the same reason, positive ions should be smaller than the atoms from which they are formed. For example, the 11 protons in the nucleus of Na+ ion, should be able to hold the 10 electrons on this ion more tightly than the 11 electrons on a neutral sodium atom.

The adjacent table provide data to test this hypothesis. They compare the covalent radii for neutral atoms of the Group IA elements with the ionic radii for the corresponding positive ions. In each case, the positive ion is much smaller than the atom from which it forms.

In an isoelectronic species (which have same number of electrons), ionic size decreases with increasing positive charge, and increases with increase in negative charge.

Atoms become larger as we go down a column of the periodic table. We can examine trends in ionic radii across a row of the periodic table by comparing data for atoms and ions that are isoelectronic atoms (ions that have the same number of electrons). The adjacent table summarizes data on the radii of a series of isoelectronic ions and atoms of second and third‐row elements.

Radii for second‐row and third‐row atoms or ions:

The data in the adjacent table is easy to explain if we note that these atoms or ions all have 10 electrons but the number of protons in the nucleus increases from 6 in the C4− ion to 13 in the Al3+ ion. As the charge on the nucleus becomes larger, the nucleus can hold a constant number of electrons more tightly. As a result, the atoms or ions become significantly smaller.

The relative size of positive and negative ions has important implications for the structure of ionic compounds. The positive ions are often so small they pack in the holes between planes of adjacent negative ions. In NaCl, for example, the Na+ ions are so small that the Cl ions almost touch.

First I.E of main group elements vs atomic number Across a period, as the size decreases higher energy is required to remove the electro. Hence ionization energy increases along the period. Down the group as the size increases it is easier to remove the electron. Hence ionization energy decreases. Noble gases have highest ionization energy values whereas alkali metals will have the lowest.
Ionization energies

The ease with which electrons can be removed from an atom is an important indicator of the atom's chemical behavior. Therefore, the ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion.

First ionization energy: The energy required to remove the outermost (highest energy) electron from a neutral atom in its ground state. In a periodic table down a group, first ionization energy decreases As electrons are far from the nucleus it would be easier to remove the outermost one. The shielding effect of inner electrons also decreases the effective nuclear charge so the first ionization energy decreases down the group.

Shielding − Inner electrons at lower energy levels essentially block or shield the proton's force of attraction towards the nucleus. It therefore becomes easier to remove the outer electron.

Across a period, first ionization energy increases. Across a period, the atomic radius decreases as the differentiating electron enters into same period. The outer electrons are closer to the nucleus and more strongly attracted to the center. Therefore, it becomes more difficult to remove the outermost electron.

Due to symmetry in shape, s-electrons are closer to the nucleus when compared to the electrons in other orbitals of the same shell. Hence s-electrons experience more attraction from the nucleus than p, d and f-electrons. Thus energy to remove an electron from a given energy level decrease in order of s > p > d > f.

Exceptions in trend of ionization energies across 2nd period IE increases from Li to Be as Be has fully filled outer shell. It decreases for boron followed by increase at ‘N’ due to stable half‐filled configuration. It is followed by decrease for oxygen. ‘Ne’ has highest IE value as it is difficult to remove the electron from its shell.
Exceptions to first ionization energy trends
• Xs1 > Xp1
Example: 3Li (2s1) > 5B(2p1)
The energy of an electron in an Xp orbital is greater than the energy of an electron in its respective Xs orbital. Therefore, it requires less energy to remove the first electron in a 'p' orbital than it is to remove one from a filled 's' orbital.
• Xp3 > Xp4
Example:
7N(3p3) > 8O(3p4)
It is difficult to remove electron from half‐filled configuration)
After the separate degenerate orbitals have been filled with single electrons, the fourth electron must be paired. The electron‐electron repulsion makes it easier to remove the outermost, paired electron.
Second ionization energy is more for Na. It attains the stable configuration of ‘Ne’ in its first ionization. Second ionization energy increases as the nucleus binds to the outer electrons more firmly in an ion than in a neutral atom.
Second ionization energies

Second ionization energy is the energy required to remove a second outermost electron from a ground state atom. Subsequent ionization energies increase greatly once an ion has reached the state like that of a noble gas. In other words, it becomes extremely difficult to remove an electron from an atom once it loses enough electrons to lose an entire energy level so that its valence shell is filled.

Alkali metals and hydrogen: First ionization energy is very low and second ionization is much higher.

Alkaline earth metals: First ionization energy is low and second ionization energy is also low.

Electron affinities of the main‐group elements Negative values indicate that the energy is released when the anion forms. Positive values, which occur in group 8A (18), indicate that energy is absorbed to form the anion; in fact, these anions are unstable and the values are estimated theoretically.
Electron affinities

Electron affinity is defined as the energy change that accompanies the addition of an electron to an atom. Some atoms readily attract electrons, and the electron affinity has a negative value, meaning that energy is released. Most atoms, however, do not accept additional electron readily and the electron affinity is a positive value, indicating that energy must be used to add the electrons. For example, the addition of electron to a chlorine atom is accompanied by an energy change of −349 kJ/mol. The negative sign indicates that energy is released during the process.

Cl (g) + eCl(g)       ΔE = −349 kJ/mol

Chlorine has the highest affinity for electrons and francium has the lowest. Electron affinity varies diagonally across the periodic table. The atom close to fluorine tends to accept electrons readily, and those close to francium do not. Therefore, in a periodic table electron affinity value decreases down the group, and increases across a period. Down the group, electron affinity decreases because of larger radii.

Periodic trend of electron affinity Periodic trend of EA is depicted as gradations in shading with arrows indicating the direction of general increase in a group. EA values of 8th group are not to be taken into account as the anions are difficult to form and are unstable if formed.
Exceptions in trends of electron affinity values

Among non‐metals, however, the elements in the first period have lower electron affinities than the elements below them in their respective groups. Elements with electronic configurations of ns2, ns2 np3 and ns2 np6 have electron affinities less than zero because they are unusually stable. In other words, instead of energy being given off, these elements actually require an input of energy in order to gain electrons. Electron affinity (EA) values are much smaller than ionization energies.

Electron affinity values of element with electronic configurations of ns2 are less than zero because it is a stable, dia‐magnetic atom with no unpaired electrons. EA value of element with electron configuration ns2 np3 is less than zero because it is a stable atom with 3 unpaired p‐orbital electrons each occupying its own sub‐shell. EA value of element with electron configuration ns2  np6 is also less than zero because it is a stable atom with filled valence (outermost) shell.

Electron affinity trend of some elements EA values show large variation than the expected because of half‐filled and fully filled outer shells. Fluorine due to its small size and high electron density is not favorable for addition of electron.
Electron affinity of halogens and noble gases

In the adjacent plot of EA versus atomic number, highest peaks are for the halogens and lowest are for noble gases. Local minima occur for filled sub‐shells and half‐filled 'p' sub‐shells. The halogens which are one electron shy of a filled 'p' sub‐shell have the most negative electron affinities. By gaining an electron, a halogen atom forms a stable negative ion that has a noble‐gas configuration (8 electrons). Chlorine has the highest electron affinity among all elements.

The addition of an electron to a noble gas however would require that the electron reside in a new, higher energy sub‐shell. This is energetically unfavorable. So the electron affinity is positive, meaning that ion will not form. E.g: Be, N, Ne. Alkaline earth elements (Group IIA) and noble gases (Group VIII A) do not form stable negative ions.

EN values on Pauling scale Fluorine has the highest electronegativity value of 4 on Pauling scale, while ‘Cs’ has the lowest value of 0.7 which is not shown in the picture. Location of hydrogen is placed as such due to its ability to act as metallic element (H+) which loses electron easily, and also as it forms hydride (H) as the halogens.
Electronegativity

Electronegativity is a chemical property that describes an atom's ability of attraction on bonded electrons. The concept of electronegativity was developed by Linus Pauling to describe the attraction of electrons by individual atoms. Since electronegativity describes a qualitative property, there is no standardized method of calculating electronegativity. However, the scale that most chemists use in quantifying electronegativity is the Pauling scale. The numbers assigned by the Pauling scale are dimensionless due to electronegativity being largely qualitative. Electronegativity is a combination of ionization energy, electron affinity and other factors.

Electronegativities show the same diagonal trend as do ionization energies and electron affinities. Fluorine has the highest electronegativity of 3.98 Pauling units and francium the lowest. The electronegativity concept is used in determining how electrons are distributed in molecules.

Electronegativity decreases with increase of atomic radius General trend of EN values across the periodic table: increases along a period from left to right and decreases along a group from top to bottom.
Electronegativity trends in a periodic table

In a periodic table, the electronegativity of elements shows an increasing trend from the lower left corner to the upper right corner. This is because, the elements on the left side of the periodic table have less than a half‐filled valence shell. The energy required to gain electrons is significantly higher compared to the energy required to lose electrons. As a result, the elements on the left side of the periodic table generally lose electrons in forming bonds. Conversely, elements on the right side of the periodic table are more energy efficient in gaining electrons to create a complete valence shell of 8 electrons. Down the group, electronegativity decreases. This is because the atomic number increases down the group and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius.

Exceptions: Important exceptions include the noble gases, lanthanides, and actinides. The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess a more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have standard electronegativity values. As for the transition metals, while they have values, there is little variance across the period and up and down a group. This is because of their metallic properties that affect their ability to attract electrons as easily as the other elements.

Trends in metallic behavior The gradation in metallic behavior among the elements is depicted (as gradation in shading) from bottom-left to top-right, with arrows showing the direction of increase.
Metals, non-metals and metalloids

Metals are located in the left and lower three‐quarters of the periodic table. They are typically shiny solids with moderate to high melting points, are good thermal and electrical conductors, can be drawn into wires and rolled into sheets, and tend to lose electrons to non‐metals.

Non‐metals are located in the upper right quarter the table. They are typically not shiny, have relatively low melting points, are poor thermal and electrical conductors, are mostly crumbly solids or gases, and tend to gain electrons from metals. Metalloids are located in the region between the other two classes and have properties between them as well. Thus, metallic behavior decreases left-to-right and increases top-to-bottom in the periodic table.

It's important to realize, however, that an element's properties may not fall neatly into above described categories. For instance, the non‐metal carbon in the form of graphite is a good electrical conductor. Iodine, another non‐metal, is a shiny solid. Gallium and Cesium are metals that melt at temperatures below body temperature, and mercury is a liquid at room temperature and iron is quite brittle. Despite such exceptions, we can make several generalizations about metallic behavior.