Early History of the Periodic Table

Early Attempts to Organize Elements

By the year 1700, only a handful of elements had been identified and isolated. Several of these, such as copper and lead, had been known since ancient times. As scientific methods improved, the rate of discovery dramatically increased. With the ever-increasing number of elements, chemists recognized that there may be some kind of systematic way to organize the elements. The question was: how?

A logical way to begin grouping elements together was by their chemical properties. In other words, putting elements in separate groups based on how they reacted with other elements. In 1829, a German chemist, Johann Dobereiner (1780-1849), placed various groups of three elements into groups called triads . One such triad was lithium, sodium, and potassium. Triads were based on both physical as well as chemical properties. Dobereiner found that the atomic masses of these three elements, as well as other triads, formed a pattern. When the atomic masses of lithium and potassium were averaged together , it was approximately equal to the atomic mass of sodium (22.99). These three elements also displayed similar chemical reactions, such as vigorously reacting with the members of another triad: chlorine, bromine, and iodine. While Dobereiner’s system would pave the way for future ideas, a limitation of the triad system was that not all of the known elements could be classified in this way.

English chemist John Newlands (1838-1898) ordered the elements in increasing order of atomic mass and noticed that every eighth element exhibited similar properties. He called this relationship the “Law of Octaves .” Unfortunately, there were some elements that were missing and the law did not seem to hold for elements that were heavier than calcium. Newlands’s work was largely ignored and even ridiculed by the scientific community in his day. It was not until years later that another more extensive periodic table effort would gain much greater acceptance and the pioneering work of John Newlands would be appreciated.

Mendeleev’s Periodic Table

The periodic table was first built using a set of cards. With this strategy, Mendeleev could organize and rearrange material until patterns emerged.

In 1869, Russian chemist and teacher Dmitri Mendeleev (1836–1907) published a periodic table of the elements. The following year, German chemist Lothar Meyer independently published a very similar table. Mendeleev is generally given more credit than Meyer because his table was published first and because of several key insights that he made regarding the table.

Mendeleev was writing a chemistry textbook for his students and wanted to organize all of the known elements at that time according to their chemical properties. He famously organized the information for each element on to separate note cards that were then easy to rearrange as needed. He discovered that when he placed them in order of increasing atomic mass, certain similarities in chemical behavior repeated at regular intervals. This type of a repeating pattern is called “periodic.” A pendulum that swings back and forth in a given time interval is periodic, as is the movement of the moon around the Earth.

In Figure 2, atomic mass increases from top to bottom of vertical columns, with successive columns going left to right. As a result, elements that are in the same horizontal row are groups of elements that were known to exhibit similar chemical properties. One of Mendeleev’s insights is illustrated by the elements tellurium (Te) and iodine (I). Notice that tellurium is listed before iodine even though its atomic mass is higher. Mendeleev reversed the order because he knew that the properties of iodine were much more similar to those of fluorine (F), chlorine (Cl), and bromine (Br) than they were to oxygen (O), sulfur (S), and selenium (Se). He simply assumed that there was an error in the determination of one or both of the atomic masses. As we will see shortly, this turned out not to be the case, but Mendeleev was indeed correct to group these two elements as he did.

Notice that there are several places in the table that have no chemical symbol, but are instead labeled with a question mark. Between zinc (Zn) and arsenic (As) are two such missing elements. Mendeleev believed that elements with atomic masses of 68 and 70 would eventually be discovered and that they would fit chemically into each of those spaces. Listed in the Table below are other properties that Mendeleev predicted for the first of these two missing elements, which he called “eka-aluminum,” compared with the element gallium.

	Eka-Aluminum (Ea)	Gallium (Ga)
Atomic mass	68 amu	69.9 amu
Melting point	Low	30.15°C
Density	5.9 g/cm 3	5.94 g/cm 3
Formula of oxide	Ea 2 O 3	Ga 2 O 3

The element gallium was discovered four years after the publication of Mendeleev’s table, and its properties matched up remarkably well with eka-aluminum, fitting into the table exactly where he had predicted. This was also the case with the element that followed gallium, which was named eventually named germanium.

Mendeleev’s periodic table gained wide acceptance with the scientific community and earned him credit as the discoverer of the periodic law. Element number 101, synthesized in 1955, is named mendelevium after the founder of the periodic table. It would, however, be several years after Mendeleev died before the several discrepancies with the atomic masses could be explained and before the reasons behind the repetition of chemical properties could be fully explained.

Periodic Law

When Mendeleev put his periodic table together, nobody knew about the existence of the nucleus. It was not until 1911 that Rutherford conducted his gold foil experiment that demonstrated the presence of the nucleus in the atom. Just two years later, in 1913, English physicist Henry Moseley (1887-1915) examined x-ray spectra of a number of chemical elements. He would shoot X-rays through crystals of the element and study the wavelengths of the radiation he detected. Moseley found that there was a relationship between wavelength and atomic number. His results led to the definition of atomic number as the number of protons contained in the nucleus of each atom. He then realized that the elements of the periodic table should be arranged in order of increasing atomic number rather than increasing atomic mass.

When ordered by atomic number, the discrepancies within Mendeleev’s table disappeared. Tellurium has an atomic number of 52, while iodine has an atomic number of 53. So even though tellurium does indeed have a greater atomic mass than iodine, it is properly placed before iodine in the periodic table. Mendeleev and Moseley are credited with being most responsible for the modern periodic law : When elements are arranged in order of increasing atomic number, there is a periodic repetition of their chemical and physical properties. The result is the periodic table as we know it today. Each new horizontal row of the periodic table corresponds to the beginning of a new period because a new principal energy level is being filled with electrons. Elements with similar chemical properties appear at regular intervals, within the vertical columns called groups .

Modern Periodic Table: Periods and Groups

The Modern Periodic Table

The periodic table has undergone extensive changes in the time since it was originally developed by Mendeleev and Moseley. Many new elements have been discovered, while others have been artificially synthesized. Each fits properly into a group of elements with similar properties. The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties appear in the same vertical column or group.

The figure below shows the most commonly used form of the periodic table. Each square shows the chemical symbol of the element along with its name. Notice that several of the symbols seem to be unrelated to the name of the element: Fe for iron, Pb for lead, etc. Most of these are the elements that have been known since ancient times and have symbols based on their Latin names. The atomic number of each element is written above the symbol.

A period is a horizontal row of the periodic table. There are seven periods in the periodic table, with each one beginning at the far left. A new period begins when a new principal energy level begins filling with electrons. Period 1 has only two elements (hydrogen and helium), while periods 2 and 3 have 8 elements. Periods 4 and 5 have 18 elements. Periods 6 and 7 have 32 elements because the two bottom rows that are separated from the rest of the table belong to those periods. They are pulled out in order to make the table itself fit more easily onto a single page.

A group is a vertical column of the periodic table, based on the organization of the outer shell electrons. There are a total of 18 groups. There are two different numbering systems that are commonly used to designate groups and you should be familiar with both. The traditional system used in the United States involves the use of the letters A and B. The first two groups are 1A and 2A, while the last six groups are 3A through 8A. The middle groups use B in their titles. Unfortunately, there was a slightly different system in place in Europe. To eliminate confusion the International Union of Pure and Applied Chemistry (IUPAC) decided that the official system for numbering groups would be a simple 1 through 18 from left to right. Many periodic tables show both systems simultaneously.

Metals

Chemists classify materials in many ways. We can sort elements on the basis of their electron arrangements. The way the electrons are distributed determines the chemical properties of the element. Another way is to classify elements based on physical properties. Some common physical properties are color, volume, and density. Other properties that allow us to sort on the basis of behavior are conduction of heat and electricity, malleability (the ability to be hammered into very thin sheets), ductility (the ability to be pulled into then wires), melting point, and boiling point. Three broad classes of elements based on physical properties are metals, nonmetals, and metalloids.

A metal is an element that is a good conductor of heat and electricity. Metals are also malleable, which means that they can be hammered into very thin sheets without breaking. They are ductile, which means that they can be drawn into wires. When a fresh surface of any metal is exposed, it will be very shiny because it reflects light well. This is called luster. All metals are solid at room temperature with the exception of mercury (Hg), which is a liquid. Melting points of metals display a very wide variance. The melting point of mercury is -39°C, while the highest melting metal is tungsten (W), with a melting point of 3422°C. The elements in blue in the periodic table below are metals. About 80 percent of the elements are metals.

Gold has been used by many civilizations for making jewelry (see Figure 3). This metal is soft and easily shaped into a variety of items. Since gold is very valuable and often used as currency, gold jewelry has also often represented wealth.

Copper is a good conductor of electricity and is very flexible and ductile. This metal is widely used to conduct electric current in a variety of appliances, from lamps to stereo systems to complex electronic devices (see Figure 4).

Mercury is the only metal to exist as a liquid at room temperature (see Figure 5). This metal was extensively used in thermometers for decades until information about its toxicity became known. Mercury switches were once common, but are no longer used. However, new federally-mandated energy-efficient light bulbs that are now used contain trace amounts of mercury and represent a hazardous waste.

Figure 3. Gold jewelry.

Figure 4. Copper wire exposed.

Figure 5. Pouring Mercury.

Nonmetals

When we sort parts in our shop or garage, we often classify them in terms of common properties. One container might hold all the screws (possibly sub-divided by size and type). Another container would be for nails. Maybe there is a set of drawers for plumbing parts.

When you get finished, you could also have a collection of things that don’t nicely fit a category. You define them in terms of what they are not. They are not electrical components, or sprinkler heads for the yard, or parts for the car. These parts may have some common properties, but are a variety of items.

In the chemical world, these “spare parts” would be considered non-metals, loosely defined as not having the properties of metals. A nonmetal is an element that is generally a poor conductor of heat and electricity. Most properties of nonmetals are the opposite of metals. There is a wider variation in properties among the nonmetals than among the metals.  Nonmetals exist in all three states of matter. The majority are gases, such as nitrogen and oxygen. Bromine is a liquid. A few are solids, such as carbon and sulfur. In the solid state, nonmetals are brittle , meaning that they will shatter if struck with a hammer. The solids are not lustrous. Melting points are generally much lower than those of metals.   The green elements in the table below are non-metals.

Non-metals have a wide variety of uses. Sulfur can be employed in gunpowder, fireworks, and matches to facilitate ignition (see Figure 6). This element is also widely used as an insecticide, a fumigant, or a means of eliminating certain types of fungus. An important role for sulfur is the manufacture of rubber for tires and other materials. First discovered in 1839 by Charles Goodyear, the process of vulcanization makes the rubber more flexible and elastic as well as being more resistant to changes in temperature. A major use of sulfur is for the preparation of sulfur-containing compounds such as sulfuric acid.

Bromine is a versatile compound, used mainly in manufacture of flame-retardant materials, especially important for children’s clothing (see Figure 7).

For treatment of water in swimming pools and hot tubs, bromine is beginning to replace chlorine as a disinfectant because of its higher effectiveness. When incorporated into compounds, bromine atoms play important roles in pharmaceuticals for treatment of pain, cancer, and Alzheimer’s disease.

Helium is one of the many non-metals that is a gas. Other non-metal gases include hydrogen, fluorine, chlorine, and all the group eighteen noble (or inert) gases. Helium is chemically non-reactive, so it is useful for applications such as balloons (see Figure 8) and lasers, where non-flammability is extremely important. Liquid helium exists at an extremely low temperature and can be used to cool superconducting magnets for imaging studies (MRI, magnetic resonance imaging). Leaks in vessels and many types of high-vacuum apparatus can be detected using helium. Inhaling helium changes the speed of sound, producing a higher pitch in your voice. This is definitely an unsafe practice and can lead to physical harm and death.

Metalloids

Some elements are “none of the above.” They don’t fit neatly into the categories of metal or non-metal because of their characteristics. A metalloid is an element that has properties that are intermediate between those of metals and nonmetals.  Metalloids can also be called semimetals. On the periodic table, the elements colored yellow, which generally border the stair-step line, are considered to be metalloids. Notice that aluminum borders the line, but it is considered to be a metal since all of its properties are like those of metals.

Examples of Metalloids

Silicon is a typical metalloid (see Figure 9). It has luster like a metal, but is brittle like a nonmetal. Silicon is used extensively in computer chips and other electronics because its electrical conductivity is in between that of a metal and a nonmetal.

Boron is a versatile element that can be incorporated into a number of compounds (see Figure 10). Borosilicate glass is extremely resistance to thermal shock. Extreme changes in the temperature of objects containing borosilicates will not create any damage to the material, unlike other glass compositions, which would crack or shatter. Because of their strength, boron filaments are used as light, high-strength materials for airplanes, golf clubs, and fishing rods. Sodium tetraborate is widely used in fiberglass as insulation and also is employed in many detergents and cleaners.

Arsenic has long played a role in murder mysteries, being used to commit the foul deed (see Figure 11). This use of the material is not very smart since arsenic can be easily detected on autopsy. We find arsenic in pesticides, herbicides, and insecticides, but the use of arsenic for these applications is decreasing due to the toxicity of the metal. Its effectiveness as an insecticide has led arsenic to be used as a wood preservative.

Antimony is a brittle, bluish-white metallic material that is a poor conductor of electricity (see Figure 12). Used with lead, antimony increases the hardness and strength of the mixture. This material plays an important role in the fabrication of electronic and semiconductor devices. About half of the antimony used industrially is employed in the production of batteries, bullets, and alloys.

Figure 10. Boron.

Figure 11. Arsenic.

Figure 12. Antimony.

Blocks of the Periodic Table

The elements in the periodic table could be considered to be similar to types of music. Each set of elements has its unique set of properties, with different sets of elements having some common characteristics in terms of electron arrangements. We can see patterns of electronic structure and reactivity in the periodic table that allow us to understand better the behavior of individual elements.

Periods and Blocks

There are seven horizontal rows of the periodic table, called periods. The length of each period is determined by the number of electrons that are capable of occupying the sublevels that fill during that period, as seen in the Table below .

Period Length and Sublevels in the Periodic Table

Period	Number of elements in period	Sublevels in order of fill
1	2	1 s
2	8	2 s 2 p
3	8	3 s 3 p
4	18	4 s 3 d 4 p
5	18	5 s 4 d 5 p
6	32	6 s 4 f 5 d 6 p
7	32	7 s 5 f 6 d 7 p

Recall that the four different sublevels each consist of a different number of orbitals. The s sublevel has one orbital, the p sublevel has three orbitals, the d sublevel has five orbitals, and the f sublevel has seven orbitals. In the first period, only the 1 s sublevel is being filled. Since all orbitals can hold two electrons, the entire first period consists of just two elements. In the second period, the 2 s sublevel, with two electrons, and the 2 p sublevel with six electrons, are being filled. Consequently, the second period contains eight elements. The third period is similar to the second, filling the 3 s and 3 p sublevels. Notice that the 3 d sublevel does not actually fill until after the 4 s sublevel. This results in the fourth period containing 18 elements due to the additional 10 electrons that are contributed by the d sublevel. The fifth period is similar to the fourth. After the 6 s sublevel fills, the 4 f sublevel with its 14 electrons fills. This is followed by the 5 d and the 6 p . The total number of elements in the sixth period is 32. The later elements in the seventh period are still being created. So while there are a possible of 32 elements in the period, the current number is slightly less.

The period to which a given element belongs can easily be determined from its electron configuration. For example, consider the element nickel (Ni). Its electron configuration is [Ar]3 d ⁸ 4 s ² . The highest occupied principal energy level is the fourth, indicated by the 4 in the 4 s ² portion of the configuration. Therefore, nickel can be found in the fourth period of the periodic table.

Based on electron configurations, the periodic table can be divided into blocks denoting which sublevel is in the process of being filled. The s , p , d , and f blocks are illustrated below.

The figure also illustrates how the d sublevel is always one principal level behind the period in which that sublevel occurs. In other words, the 3 d sublevel fills during the fourth period. The f sublevel is always two levels behind. The 4 f sublevel belongs to the sixth period.

Hydrogen and Alkali Metals

One value of the periodic table is the ability to make predictions about the behavior of individual elements. By knowing which group an element is in, we can determine the number of reactive electrons and say something about how that element will behave.

The periodic table is arranged on the basis of atomic numbers (number of protons in the nucleus). One of the valuable consequences of this arrangement is that we can learn a lot about the electron distribution in these atoms. The colors in the table below indicate the different groupings of atoms based on the location and number of electrons in the atom.

If we look at Group I (red column), we see that it is labeled alkali metals . Also note the green H above the alkali metals. All of these elements have a similar configuration of outer-shell electrons (see Table below ).  In each case, there is one electron in the outer orbital and that is an s -orbital electron. Hydrogen is not an alkali metal itself, but has some similar properties due to its simple one proton (loctated in the nucleus), one electron arrangement. The lone electron exists in a s -orbital around the nucleus. For lithium, there are two 1 s electrons in an inner orbit and one 2 s electron in the outer orbit. The same pattern holds for sodium and potassium.

Element	Symbol	Electron Configuration
hydrogen	H	1 s 1
lithium	Li	[He]2 s 1
sodium	Na	[Ne]3 s 1
potassium	K	[Ar]4 s 1
rubidium	Rb	[Kr]5 s 1
cesium	Cs	[Xe]6 s 1
francium	Fr	[Rn]7 s 1

Even an atom with a very complex electron composition such as cesium still has the single s electron in its outer orbital (see Figure 13).

This one electron is very easily removed during chemical reactions.  The group I elements react rapidly with oxygen to produce metal oxides. They are very soft metals, which become liquid just above room temperature.

The alkali metals also react readily with water to produce hydrogen gas and metal hydroxides in the following video: Alkali Metals: Explosive reactions.

Li reacts with water to produce hydrogen gas. Sodium also reacts the same way, just more rapidly. Potassium reacts rapidly with water producing hydrogen gas and heat which ignites the hydrogen gas. Rubidium and cesium react yet more vigorously and explode on contact with water.

Alkaline Earth Metals

Group 2 elements are referred to as “alkaline earth” metals (tan column below). The name “alkaline” comes from the fact that compounds of these elements form basic (pH greater than 7) or alkaline solutions when dissolved in water. If the Group 1 elements all have one s electron in their outer orbital, we can predict that the Group 2 elements will have two electrons in that outer shell.

The beryllium atom, the first element of Group 2, has an atomic number of four. The atom has the 1 s shell filled as well as the 2 s shell, giving a total of four electrons (1 s ² 2 s ² ). Note that there are two s electrons in the outer shell, a structure that is characteristic of the Group 2 elements. Barium (atomic number 56) has the same outer shell structure of two electrons in the s orbital, even though the internal electron structure for barium is quite complicated.

Radium (atomic number 88) has similar properties to barium and is also in the Group 2 category. However, radium is a radioactive element and is generally under the category of radioisotopes in addition to being an alkaline earth metal, because it is not a stable element.

The Group 2 elements tend to be less reactive than their Group 1 counterparts. The need to remove two electrons in order for the material to react means more energy is needed for electron removal. However, these elements are reactive enough that they do not exist in their elemental forms in nature, but are present as compounds.

Uses of Alkaline Earth Compounds

Since magnesium burns brightly, it is used in flares and fireworks. Magnesium alloys with aluminum provide light weight and sturdy materials for airplanes, missiles, and rockets. Several antacids use magnesium hydroxide to neutralize excess stomach acid.

Calcium compounds are widely found in limestone, marble, and chalk. Calcium is an important constituent of cement. Other uses include calcium chloride as a deicer and limestone as a white pigment in paints and toothpaste.

Strontium is widely used in fireworks and magnets. Barium compounds can be used in paints, filler for rubber, plastic, and resins, and as a contrast medium for X-rays. Many beryllium compounds are toxic, but these materials have been employed in metal alloys.

Noble Gases

The reactivity of an element can give us important clues as to the electron configuration of that material. If an element is extremely unreactive, this suggests that the electron configuration is such that adding or removing electrons is very unlikely. There must be a stable electron configuration that resists further reaction.

The Group VIII (new group 18) elements are essentially chemically inert (light blue column on the right). All these elements exist as monatomic gases at room temperature. If we look at the electron configurations, we see that helium (atomic number 2) has a full shell of two s electrons. Since there are no electrons shielding this shell from the nucleus, these two electrons will be very difficult to remove, making helium unreactive.

The remaining elements in the group have full outer shells consisting of two s electrons and six p electrons for an outer shell content of eight electrons. This particular arrangement renders the atoms fairly unreactive. This group has been referred to as the “inert” gases, indicating that they are chemically inert, or unreactive. Another popular term is “ noble gases ,” suggesting that these gases do not like to have much to do with the other, more common materials (or that they don’t do a lot of work).

Noble Gas Compounds

In more recent years, a number of reactions using the noble gas elements have been discovered. Although the conventional wisdom was that the complete outer shells of these elements would not allow them to react, some scientists believed that the outer electrons of the larger elements (such as Rn, Xe, and Kr) were far enough away from the nucleus that they should be able to be displaced under the right set of conditions. The first compound (XePtF₆ ) was made with xenon in 1962. Since then, several compounds have been made with radon, xenon, krypton, and argon. Only helium and neon have not formed compounds at this time.

Colors of Noble Gases

The different gases glow when an electric current is passed through them. Many of these gases are used in displays because of their chemical inertness. They are stable and will not react with other materials in the system. Radon also will give a reddish glow, but is not used because it is radioactive and will not retain its structure as radon for any significant length of time.

Halogens

How do you study a gas that does not exist as such in nature? It’s not as easy as you think. Fluorine is so reactive that we cannot find it free in nature. None of the halogens exist free in nature (unlike some of the metals suchas gold and silver) because they are very reactive. The video below shows how violently elemental fluorine reacts with other materials.

Some elements are much more reactive than others. The Group I (red)  and Group II (tan) elements can easily lose electrons during a reaction. Elements of other groups are much more likely to accept electrons as they react.

The elements of Group VIIA (new Group 17 – fluorine, chlorine, bromine, iodine, and astatine) are called the halogens (tan column). The term “halogen” means “salt-former” because these elements will readily react with alkali metal and alkaline earth metals to form halide salts. The halogens all have the general electron configuration ns ² np ⁵ , giving them seven valence electrons. They are one electron short of having the full outer s and p sublevel, which makes them very reactive.

Physical Properties of Halogens

As elements, chlorine and fluorine are gases at room temperature, bromine is a dark orange liquid, and iodine is a dark purple-gray solid. Astatine is so rare that its properties are mostly unknown. In the picture below we see chlorine gas on the left (green), bromine solid and vapor in the middle (orange), and solid iodine (grey) on the right. Fluorine is not shown in the picture below because it is too corrosive and will destroy the glass container.

None of these elements are found free in nature because of their reactivity. Fluorine is found in combination with cations in several minerals. Chlorine is found in table salt, in the oceans (which are about 2% chloride ion by weight) and in lakes such as the Great Salt Lake in Utah. Small amounts of bromide and iodide salts can be found in the oceans and in brine wells in several states.

Watch the following two video experiments of p block elements:

This first video is of bromine reacting with aluminum:

This second video shows halogen reactions:

http://youtu.be/cbFCWFksYoM

Transition Elements

Many electron configurations of elements are simple and straightforward. We can look at the outer shell and easily understand how that set of elements will react in terms of electron gain or loss. However, there are sets of elements that are more complex in their behavior. One such group is called the transition elements.

Transition elements are the elements that are found in Groups 3-12 (old groups IIA-IIB) on the periodic table (salmon-colored block in the middle of the table). The term refers to the fact that the d sublevel , which is in the process of being filled, is in a lower principal energy level than the sublevel filled before it. For example, the electron configuration of scandium, the first transition element, is [Ar]3d¹4s².

Remember that the configuration is reversed from the fill order—the 4 s filled before the 3 d begins. Because they are all metals, the transition elements are often called the transition metals. As a group, they display typical metallic properties and are less reactive than the metals in Groups 1 and 2. Some of the more familiar ones are so unreactive that they can be found in nature in their free, or uncombined state. These include platinum, gold, and silver. Because of this unique filling order, the transition elements are often referred to as “ d -block” elements.

Compounds of many transition elements are distinctive for being widely and vividly colored. As visible light passes through a transition metal compound dissolved in water, the d-orbitals absorb light of various energies. The visible light of a given energy level which is not absorbed produces a distinctly colored solution.

Lanthanides and Actinides

We see some hidden “layers” in chemistry. As we look at the periodic table below, we see two pink boxes – one between Ba (element 56) and Hf (element 72) and the other between Ra (88) and Rf (104). These elements all have unfilled f -sublevels. Because of the uniqueness of the electron configurations, these elements fit into the two boxes in the larger periodic table.

As the number of elecrons in an atom increases, we begin to see some strange behaviors. Due to the way the electron energy levels work, some inner levels fill after one or more outer levels do. We see this in two similar groups of elements – the lanthanides and the actinides .

The f-Block

The first of the f sublevels to begin filling is the 4 f sublevel. It fills after the 6 s sublevel, meaning that f sublevels are two principal energy levels behind. The general electron configuration for elements in the f block is ( n – 2 )f ^1-14 ns ² . The seven orbitals of the f sublevel accommodate 14 electrons, so the f block is 14 elements in length. It is pulled out of the main body of the period table and is shown at the very bottom. Because of that, the elements of the f block do not belong to a group, being wedged in between Groups 3 and 4. The lanthanides are the 14 elements from cerium (atomic number 58) to lutetium (atomic number 71). The word comes from the Greek “lanthanein” meaning “to be hidden.” The name probably arose because these elements all hide behind one another in the periodic table. The 4 f sublevel is in the process of being filled for the lanthanides. They are all metals and are similar in reactivity to the Group 2 alkaline earth metals.

The actinides are the 14 elements from thorium (atomic number 90) to lawrencium (atomic number 103). The 5 f sublevel is in the process of being filled. The actinides are all radioactive elements and only the first four have been found naturally on Earth. All of the others have only been artificially made in the laboratory. The lanthanides and actinides together are sometimes called the inner transition elements.

Uses of Lanthanides

Lanthanides have been widely used as alloys to impart strength and hardness to metals. The main lanthanide used for this purpose is cerium, mixed with small amounts of lanthanum, neodymium, and praseodymium. These metals are also widely used in the petroleum industry for refining of crude oil into gasoline products.

Erbium and other lanthanides are widely used in some optical devices, such as night vision goggles, laser beams, and phosphorescent materials.

Figure 16. Oil refinery.

Figure 17. Night vision goggles.

Uses of Actinides

The actinides are valuable primarily because they are radioactive. These elements can be used as energy sources for applications as varied as cardiac pacemakers and generation of electrical energy for instruments on the moon. Uranium and plutonium have been employed in nuclear weapons and in nuclear power plants.

Periodic Trends: Atomic Radius

The size of atoms is important when trying to explain the behavior of atoms or compounds. One of the ways we can express the size of atoms is with the atomic radius. This data helps us understand why some molecules fit together and why other molecules have parts that get too crowded under certain conditions.

The size of an atom is defined by the edge of its orbital. However, orbital boundaries are fuzzy and in fact are variable under different conditions. In order to standardize the measurement of atomic radii, the distance between the nuclei of two identical atoms bonded together is measured. The atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together.

Atomic radii have been measured for elements. The units for atomic radii are picometers, equal to 10 ^-12 meters. As an example, the internuclear distance between the two hydrogen atoms in an H ₂ molecule is measured to be 74 pm. Therefore, the atomic radius of a hydrogen atom is .

Periodic Trend

The atomic radius of atoms generally decreases from left to right across a period. There are some small exceptions, such as the oxygen radius being slightly greater than the nitrogen radius. Within a period, protons are added to the nucleus as electrons are being added to the same principal energy level. These electrons are gradually pulled closer to the nucleus because of its increased positive charge. Since the force of attraction between nuclei and electrons increases, the size of the atoms decreases. The effect lessens as one moves further to the right in a period because of electron-electron repulsions that would otherwise cause the atom’s size to increase.

Group Trend

The atomic radius of atoms generally increases from top to bottom within a group. As the atomic number increases down a group, there is again an increase in the positive nuclear charge. However, there is also an increase in the number of occupied principle energy levels. Higher principal energy levels consist of orbitals which are larger in size than the orbitals from lower energy levels. The effect of the greater number of principal energy levels outweighs the increase in nuclear charge and so atomic radius increases down a group.

Periodic Trends: Ionization Energy

There is an on-going tension between the electrons and protons in an atom. Reactivity of the atom depends in part on how easily the electrons can be removed from the atom. We can measure this quantity and use it to make predictions about the behaviors of atoms.

Ionization energy is the energy required to remove an electron from a specific atom. It is measured in kJ/mol, which is an energy unit, much like calories. The ionization energies associated with some elements are described in the Table below. For any given atom, the outermost valence electrons will have lower ionization energies than the inner-shell kernel electrons. As more electrons are added to a nucleus, the outer electrons become shielded from the nucleus by the inner shell electrons. This is called electron shielding.

Ionization Energies (kJ/mol) of the First 18 Elements

Element	IE 1	IE 2	IE 3	IE 4	IE 5	IE 6
H	1312
He	2373	5251
Li	520	7300	11,815
Be	899	1757	14,850	21,005
B	801	2430	3660	25,000	32,820
C	1086	2350	4620	6220	38,000	47,261
N	1400	2860	4580	7500	9400	53,000
O	1314	3390	5300	7470	11,000	13,000

If we plot the first ionization energies vs. atomic number for the main group elements, we would have the following trend

Moving from left to right across the periodic table, the ionization energy for an atom increases. We can explain this by considering the nuclear charge of the atom. The more protons in the nucleus, the stronger the attraction of the nucleus to electrons. This stronger attraction makes it more difficult to remove electrons.

Within a group, the ionization energy decreases as the size of the atom gets larger. On the graph, we see that the ionization energy increases as we go up the group to smaller atoms. In this situation, the first electron removed is farther from the nucleus as the atomic number (number of protons) increases. Being farther away from the positive attraction makes it easier for that electron to be pulled off.

Electron Shielding

The attraction between an electron and the nucleus of the atom is not a simple issue. Only with hydrogen is there a one-to-one relationship that can be discussed in terms of direct charge attraction. As the size of the atom increases, the number of protons and electrons also increase. These changes influence how the nucleus attracts electrons.

In general, the ionization energy of an atom will increase as we move from left to right across the periodic table. There are several exceptions to the general increase in ionization energy across a period. The elements of Group 13 (B, Al, etc.) have lower ionization energies than the elements of Group 2 (Be, Mg, etc.). This is an illustration of a concept called “electron shielding.” Outer electrons are partially shielded from the attractive force of the protons in the nucleus by inner electrons.

To explain how shielding works, consider a lithium atom. It has three protons and three electrons—two in the first principal energy level and its valence electron in the second. The valence electron is partially shielded from the attractive force of the nucleus by the two inner electrons. Removing that valence electron becomes easier because of the shielding effect.

There is also a shielding effect that occurs between sublevels within the same principal energy level. Specifically, an electron in the “s” sublevel is capable of shielding electrons in the “p” sublevel of the same principal energy level. This is because of the spherical shape of the “s” orbital. The reverse is not true – electrons in “p” orbitals do not shield electrons in “s” orbitals.

The electron being removed from an Al atom is a 3p electron, which is shielded by the two 3s electrons as well as all the inner core electrons. The electron being removed from a Mg atom is a 3s electron, which is only shielded by the inner core electrons. Since there is a greater degree of electron shielding in the Al atom, it is slightly easier to remove the valence electron and its ionization energy is less than that of Mg. This is despite the fact that the nucleus of the Al atom contains one more proton than the nucleus of the Mg atom .

There is another anomaly between Groups 15 and 16. Atoms of Group 16 (O, S, etc.) have lower ionization energies than atoms of Group 15 (N, P, etc.). Hund’s rule is behind the explanation. In a nitrogen atom, there are three electrons in the 2p sublevel and each is unpaired. In an oxygen atom, there are four electrons in the 2p sublevel, so one orbital contains a pair of electrons. It is that second electron in the orbital that is removed in the ionization of an oxygen atom. Since electrons repel each other, it is slightly easier to remove the electron from the paired set in the oxygen atom than it is to remove an unpaired electron from the nitrogen atom.

Electron Affinity

When electrons are removed from an atom, that process requires energy to pull the electron away from the nucleus. Addition of an electron releases energy from the process.

In most cases, the formation of an anion by the addition of an electron to a neutral atom releases energy. This can be shown for the chloride ion formation below:

Cl + e⁻  → Cl⁻ + energy

The energy change that occurs when a neutral atom gains an electron is called its electron affinity. When energy is released in a chemical reaction or process, that energy is expressed as a negative number. The figure below shows electron affinities in kJ/mole for the representative elements. Electron affinities are measured on atoms in the gaseous state and are very difficult to measure accurately.

The elements of the halogen group (Group 17) gain electrons most readily, as can be seen from their large negative electron affinities. This means that more energy is released in the formation of a halide ion than for the anions of any other elements. Considering electron configuration, it is easy to see why. The outer configuration of all halogens is ns ² np ⁵ . The addition of one more electron gives the halide ions the same electron configuration as a noble gas, which we have seen is particularly stable.

Period and group trends for electron affinities are not nearly as regular as for ionization energy. In general, electron affinities increase (become more negative) from left to right across a period and decrease (become less negative) from top to bottom down a group. However, there are many exceptions, owing in part to inherent difficulties in accurately measuring electron affinities.

Ionic Radii

Electrons and protons are strongly attracted to one another. The strength of that attraction and the relative numbers of the two particles in a given atom or ion have a significant influence on the size of that species. When an atom loses one or more electrons, the resulting ion becomes smaller. If electrons are added to the atom, the ion becomes larger.

The ionic radius for an atom is measured in a crystal lattice, requiring a solid form for the compound. These radii will differ somewhat depending upon the technique used. Usually X-ray crystallography is employed to determine the radius for an ion.

The removal of electrons always results in a cation that is considerably smaller than the parent atom. When the valence electron(s) are removed, the resulting ion has one fewer occupied principal energy level, so the electron cloud that remains is smaller. Another reason is that the remaining electrons are drawn closer to the nucleus because the protons now outnumber the electrons. One other factor is the number of electrons removed. The potassium atom has one electron removed to for the corresponding ion, while calcium loses two electrons.

The addition of electrons always results in an anion that is larger than the parent atom. When the electrons outnumber the protons, the overall attractive force that the protons have for the electrons is decreased. The electron cloud also spreads out because more electrons results in greater electron-electron repulsions. Notice that the group 16 ions are larger than the group 17 ions. The group 16 elements each add two electrons while the group 17 elements add one electron per atom for form the anions.

Periodic Trends: Electronegativity

Valence electrons of both atoms are always involved when those two atoms come together to form a chemical bond. Chemical bonds are the basis for how elements combine with one another to form compounds. When these chemical bonds form, atoms of some elements have a greater ability to attract the valence electrons involved in the bond than other elements.

Electronegativity

Electronegativity is a measure of the ability of an atom to attract the electrons when the atom is part of a compound. Electronegativity differs from electron affinity because electron affinity is the actual energy released when an atom gains an electron. Electronegativity is not measured in energy units, but is rather a relative scale. All elements are compared to one another, with the most electronegative element, fluorine, being assigned an electronegativity value of 3.98. Fluorine attracts electrons better than any other element. The table below shows the electronegativity values for the elements.

Since metals have few valence electrons, they tend to increase their stability by losing electrons to become cations. Consequently, the electronegativities of metals are generally low. Nonmetals have more valence electrons and increase their stability by gaining electrons to become anions. The electronegativities of nonmetals are generally high.

Trends

Electronegativities generally increase from left to right across a period. This is due to an increase in nuclear charge. Alkali metals have the lowest electronegativities, while halogens have the highest. Because most noble gases do not form compounds, they do not have electronegativities. Note that there is little variation among the transition metals. Electronegativities generally decrease from top to bottom within a group due to the larger atomic size.

Of the main group elements, fluorine has the highest electronegativity (EN = 4.0) and cesium the lowest (EN = 0.79). This indicates that fluorine has a high tendency to gain electrons from other elements with lower electronegativities. We can use these values to predict what happens when certain elements combine. The following video shows this.

http://www.youtube.com/watch?v=Kj3o0XvhVqQ

When the difference between electronegativities is greater than ~1.7, then a complete exchange of electrons occurs. Typically this exchange is between a metal and a nonmetal. For instance, sodium and chlorine will typically combine to form a new compound and each ion becomes isoelectronic with its nearest noble gas. When we compare the EN values, we see that the electronegativity for Na is 0.93 and the value for Cl is 3.2. The absolute difference between ENs is |0.93 – 3.2| = 2.27. This value is greater than 1.7, and therefore indicates a complete electron exchange occurs.

Metallic and Nonmetallic Character

Development of the periodic table has helped organize chemical information in many ways. We can now see trends among properties of different atoms and make predictions about the behavior of specific materials.

Metallic character refers to the level of reactivity of a metal. Metals tend to lose electrons in chemical reactions, as indicated by their low ionization energies. Within a compound, metal atoms have relatively low attraction for electrons, as indicated by their low electronegativities. By following the trend summary in the figure below, you can see that the most reactive metals would reside in the lower left portion of the periodic table. The most reactive metal is cesium, which is not found in nature as a free element. It reacts explosively with water and will ignite spontaneously in air. Francium is below cesium in the alkali metal group, but is so rare that most of its properties have never been observed.

Reactivity of metals is based on processes such as the formation of halide compounds with halogens and how easily they displace hydrogen from dilute acids.

The metallic character increases as you go down a group. Since the ionization energy decreases going down a group (or increases going up a group), the increased ability for metals lower in a group to lose electrons makes them more reactive. In addition, the atomic radius increases going down a group, placing the outer electrons further away from the nucleus and making that electron less attracted by the nucleus.

Nonmetals tend to gain electrons in chemical reactions and have a high attraction for electrons within a compound. The most reactive nonmetals reside in the upper right portion of the periodic table. Since the noble gases are a special group because of their lack of reactivity, the element fluorine is the most reactive nonmetal. It is not found in nature as a free element. Fluorine gas reacts explosively with many other elements and compounds and is considered to be one of the most dangerous known substances.

Note that there is no clear division between metallic and non-metallic character. As we move across the periodic table, there is an increasing tendency to accept electrons (non-metallic) and a decrease in the possibility that an atom would give up one or more electrons.