SAT Chemistry Atomic Structure and the Periodic Table of the Elements - Transition Elements

SAT Chemistry Atomic Structure and the Periodic Table of the Elements - Transition Elements

The elements involved with the filling of a d sublevel with electrons after two electrons are in the s sublevel of the next principal energy level are often referred to as the transition ele­ments. The first examples of these are the elements between calcium, atomic number 20, and gallium, atomic number 31. Their electron configurations are the same in the Is, 2s, 2p, 3s, and 3p sublevels. It is the filling of the 3d and changes in the 4s sublevels that are of interest, as shown in the following table.



The asterisk (*) shows where a 4s electron is promoted into the 3d sublevel. This is because the 3d and 4s sublevels are very close in energy and that there is a state of greater stability in half-filled and filled sublevels. Therefore, chromium gains stability by the movement of an electron from the 4s sublevel into the 3d sublevel to give a half-filled 3d sublevel. It then has one electron in each of the five orbitals of the 3d sublevel. In copper, the movement of one 4s electron into the 3d sublevel gives the 3d sublevel a completely filled configuration.
The fact that the electrons in the 3d and 4s sublevels are so close in energy levels leads to the possibility of some or all the 3d electrons being involved in chemical bonding. With the variable number of electrons available for bonding, it is not surprising that transition ele­ments can exhibit variable oxidation numbers. An example is manganese with possible oxida­tion numbers of +2, +3, +4, +6, and +7, which correspond, respectively, to the use of no, one, two, four, and five electrons from the 3d sublevel.
The transition elements in the other periods of the table show this same type of anomaly, as they have d sublevels filling in the same manner.

Transition elements have several common characteristic properties.

• They often form colored compounds.
• They can have a variety of oxidation states.
• At least one of their compounds has an incomplete d electron subshell.
• They are often good catalysts.
• They are silvery blue at room temperature (except copper and gold).
• They are solids at room temperature (except mercury).
• They form complex ions.
• They are often paramagnetic due to unpaired electrons.

PERIODIC TABLE OF THE ELEMENTS

History
The history of the development of a systematic pattern for the elements includes the work of a number of scientists such as John Newlands, who, in 1863, proposed the idea of repeating octaves of properties.
Dimitri I. Mendeleev in 1869 proposed a table containing 17 columns and is usually given credit for the first periodic table since he arranged elements in groups according to their atomic weights and properties. It is interesting to note that Lothar Meyer proposed a similar arrangement about the same time. In 1871 Mendeleev rearranged some elements and pro­posed a table of eight columns, obtained by splitting each of the long periods across into a period of seven elements, an eighth group containing the three central elements (such as Fe, Co, Ni), and a second period of seven elements. The first and second periods of seven across were later distinguished by use of the letters A and B attached to the group symbols, which were Roman numerals. This nomenclature of periods (IA, IIA, etc.) has been revised in the present Periodic Table, even in the extended form of assigning Arabic numbers from 1-18 as shown in Table 4.




Mendeleev’s table had the elements arranged by atomic weights with recurring properties in a periodic manner. Where atomic weight placement disagreed with the properties that should occur in a particular spot in the table, Mendeleev gave preference to the element with the correct properties. He even predicted elements for places that were not yet occupied in the table. These predictions proved to be amazingly accurate and led to wide acceptance of his table.

Periodic Law
Henry Moseley stated, after his work with X-ray spectra in the early 1900s, that the properties of elements are a periodic function of their atomic numbers, thus changing the basis of the periodic law from atomic weight to atomic number. This is the present statement of the periodic law.

The Table
The horizontal rows of the periodic table are called periods or rows. There are seven periods, each of which begins with an atom having only one valence electron and ends with a com­plete outer shell structure of an inert gas. The first three periods are short, consisting of 2, 8, and 8 elements, respectively. Periods 4 and 5 are longer, with 18 each, while period 6 has 32 elements, and period 7 is incomplete with 22 elements, most of which are radioactive and do not occur in nature.
In Table 4, you should note the relationship of the length of the periods to the orbital structure of the elements. In the first period, the Is² orbital is filled with the noble gas helium, He. The second period begins with the 2s¹ orbital and ends with the filling of the 2p⁶ orbital, again with a noble gas, neon, Ne. The same pattern is repeated in period three, going from 3s¹ to 3p⁶. The eight elements from sodium, Na, to argon, Ar, complete the filling of the n- 3 energy level with 3s² and 3p^s. In the fourth period, the first two elements fill the 4s² orbital. Beyond calcium, Ca, the pattern becomes more complicated. As discussed in the section “Order of Filling and Notation,” the next orbitals to be filled are the five 3d orbitals whose elements represent transition elements. Then the three 4p orbitals are filled, ending with the noble gas krypton, Kr. The fifth period is similar to the fourth period. The 5s² orbital filling is represented by rubidium, Rb, and strontium, Sr, both of which resemble the elements directly above them on the table. Next come the transition elements that fill the five 4d orbitals before the next group of elements, from indium, In, to xenon, Xe, complete the three 5p orbitals. (Table 3 should be consulted for the irregularities that occur as the d orbitals fill.) The sixth period follows much the same pattern and has the filling order 6s²,4ƒ¹⁴,5d¹⁰,6p⁶. Here, again, irregularities occur and can best be followed by using Table 3.
The vertical columns of the Periodic Table are called groups or families. The elements in a group exhibit similar or related properties. In 1984 the IUPAC agreed that the groups would be numbered 1 through 18.

PROPERTIES RELATED TO THE PERIODIC TABLE
Metals are found on the left of the chart (see Table 4) with the most active metal in the lower left comer. Nonmetals are found on the right side with the most active nonmetal in the upper right-hand corner. The noble or inert gases are on the far right. Since the most active metals react with water to form bases, the Group 1 metals are called alkali metals. As you proceed to the right, the base-forming property decreases and the acid-forming properties increase. The metals in the first two groups are the light metals, and those toward the center are heavy metals.
The elements found along the dark line in the Periodic Table (Table 4) are called metal­loids. These elements have certain characteristics of metals and other characteristics of nonmetals. Some examples of metalloids are boron, silicon, arsenic, and tellurium.
Here are some important general summary statements about the Periodic Table:
■ Acid-forming properties increase from left to right on the table.
■ Base-forming properties are high on the left side and decrease to the right.
■ The atomic radii of elements decrease from left to right across a period.
■ First ionization energies increase from left to right across a period.
■ Metallic properties are greatest on the left side of the table and decrease to the right.
■ Nonmetallic properties are greatest on the right side of the table and decrease to the left.

Study Table 4 carefully because it summarizes many of these properties. For a more detailed description of metals, alloys, and metalloids see pages 274-276 in Chapter 13.
Radii of Atoms
The size of an atom is difficult to describe because atoms have no definite outer boundary. Unlike a volleyball, an atom does not have a definite circumference. To overcome this problem, the size of an atom is estimated by describing its radius. In metals, this is done by measuring the distance between two nuclei in the solid state and dividing this distance by 2. Such measurements can be made with X-ray diffraction. For a nonmetallic element that exists in pure form as a molecule, such as chlorine, measurements can be made of the distance between nuclei for two atoms covalendy bonded together. Half  of this distance is referred to as the covalent radius. The method for finding the covalent radius of the chlorine atom is illustrated in the following diagram.



Figure 10 shows the relative atomic and ionic radii for some elements. As you review this chart, you should note two trends:
1. Atomic radii decrease from left to right across a period in the Periodic Table (until the noble gases).
2. Atomic radii increase from top to bottom in a group or family.
The reason for these trends will become clear in the following discussions.

Atomic Radii in Periods
Since the number of electrons in the outer principal energy level increases as you go from left to right in each period, the corresponding increase in the nuclear charge because of the additional protons pulls the electrons more tightly around the nucleus. This attraction more than balances the repulsion between the added electrons and the other electrons, and the radius is generally reduced. The inert gas at the end of the period has a slight increase in radius because of the electron repulsion in the filled outer principal energy level. For example, lithium’s atomic radius in Figure 10 is 0.152 nm at the one end of period 2 whereas fluorine has a radius of only 0.064 nm at the far end of the period. This trend can be seen in Figure 10 across every period.

Atomic Radii in Groups
For a group of elements, the atoms of each successive member have another outer principal energy level in the electron configuration, and the electrons there are held less tightly by the nucleus. This is so because of their increased distance from the nuclear positive charge and the shielding of this positive charge by all the core electrons. Therefore, the atomic radius increases down a group. For example, oxygen’s atomic radius in Figure 10 is 0.066 nm at the top of group 16, whereas polonium has a radius of 0.167 nm at the bottom of the same group. This trend can be seen in Figure 10 down every group.

Ionic Radius Compared with Atomic Radius
Metals tend to lose electrons in forming positive ions. With this loss of negative charge, the positive nuclear charge pulls in the remaining electrons closer and thus reduces the ionic radius below that of the atomic radius.
Nonmetals tend to gain electrons in forming negative ions. With this added negative charge, which increases the inner electron repulsion, the ionic radius is increased beyond the atomic radius. See Figure 10 for relative atomic and ionic radii values.

Electronegativity
The electronegativity of an element is a number that measures the relative strength with which the atoms of the element attract valence electrons in a chemical bond. This electro­negativity number is based on an arbitrary scale going from 0 to 4. In general, a value of less than 2 indicates a metal.



Notes: The atomic radius is usually given for metal atoms, which are shown in gray, and the covalent radius is usually given for atoms of nonmetals, which are shown in black.
Notice in Table 5 that the electronegativity decreases down a group and increases across a period. The inert gases can be ignored. The lower the electronegativity number, the more electropositive an element is said to be. The most electronegative element is in the upper right corner—F, fluorine. The most electropositive is in the lower left comer of the chart—Fr, francium.



Ionization Energy
Atoms hold their valence electrons, then, with different amounts of energy. If enough energy is supplied to one outer electron to remove it from its atom, this amount of energy is called the first ionization energy. With the first electron gone, the removal of succeeding electrons becomes more difficult because of the loss of repulsive effects that were present with a greater number of electrons. It should be noted that the lowest ionization energies are found with the least electronegative elements.
Ionization energies can be plotted against atomic numbers, as shown in the graph below. Follow this discussion on the graph to help you understand the peaks and valleys. Not sur­prisingly, the highest peaks on the graph occur for the ionization energy needed to remove the first electron from the outer energy level of the noble gases, He, Ne, Ar, Kr, Xe, and Rn, because of the stability of the filled p orbitals in the outer energy level. Notice that, even among these elements, the energy needed gradually declines. This can be explained by con­sidering the distance of the involved energy level from the positively charged nucleus. With each succeeding noble gas, a more distant p orbital is involved, therefore making it easier to remove an electron from the positive attraction of the nucleus. Besides this consideration, as more energy levels are added to the atomic structure as the atomic number increases, the additional negative fields associated with the additional electrons screen out some of the positive attraction of the nucleus. Within a period such as that from Li to Ne, the ionization energy generally increases. The lowest occurs when a lone electron occupies the outer s orbital, as in Li. As the s orbital fills with two electrons at atomic number 4, Be, the added stability of a filled 2s orbital explains the small peak at 4. At atomic number 5, B, a lone elec­tron occupies the 2p orbital. This electron can be removed with less energy, and therefore a dip occurs in the graph. With the 2p orbitals filling according to Hund’s Rule (refer to Table 2, page 72), with only one electron in each orbital before pairing occurs, again a slightly more stable situation and, therefore, another small peak occurs at atomic number 7. After this peak, a dip and continual increases occur until the 2p orbitals are completely filled with paired electrons at the noble gas Ne. As you continue to associate the atomic number with the line in the chart, you find peaks occurring in the same general pattern. These peaks are always related to the state of filling of the orbitals involved and the distance of these orbitals from the nucleus.