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In this lesson, we will be reviewing the properties of the periodic table and the patterns of the elements displayed therein. The material will concentrate on empirical characteristics of the elements that are reflected by their valance electrons.
We have covered an extensive (but not exhaustive) review of the main disciplines of the sciences: scientific investigation, physics, chemistry, Earth science, and biology. The last lessons concentrated on the chemical and physical aspects of biology.
The first periodic table that ordered the elements based on their atomic weight was developed by Dmitri Mendeleyev, a Russian chemist, in the 1860s. Mendeleyev predicted the existence of previously unknown elements and paved the way for modern chemistry. In the 1860s he knew:
Mendeleyev probably wrote the properties of elements known at the time on placards and arranged them in similar groups to discover that they were repeating, or periodic, thus the name of the table today, the periodic table. Where there were elements missing, he left room for elements that were yet to be discovered. One of the most striking patterns reveals an octet arrangement, or that certain characteristics of the elements repeat every eight elements. We know today that the observed properties of the elements are determined by their outermost electron cloud or shell, which has been described as the rule of octet rule.
Today, the periodic law is based on the number of protons in the nucleus of an element or the atomic number, not the atomic mass. The atomic number of the elements is the number of the protons in the nucleus of that element and is the method by which the periodic table is constructed. The atomic mass is the number of protons and neutrons in the nucleus. There are only a couple of places on the table where there is a discrepancy between the atomic mass and the atomic number. These stem from the nature of radioactivity when elements are unstable and have a short half-life. Remember that sub-atomic particles are too small to be seen — even with the most powerful microscopes in use today — and finding evidence for these particles requires sophisticated measurement techniques and equipment like particle accelerators. Nevertheless, the table organizes elements in a very systematic way that is invaluable to chemists and the basic arrangement first suggested by Mendeleyev is still valid today.
The periodic table is an important tool for all scientific endeavors. Understanding the significance of Groups IA through VIIIA is the key to the rest of the periodic table. The increasing Group A number corresponds to the increasing number of valence (outer) electrons. Group VIIIA elements all (except He) have 8 valence electrons, hence the rule of octet. Helium (He) is such a small atom that only two electrons have room to fit so near the nucleus and fill the outer shell. The elements in Groups IA and IIA easily lose electrons to become positive ions. The tendency to lose electrons easily identifies elements that are commonly called metals. Group IIIA elements tend to lose electrons but not as readily as the Group IA and IIA elements. They often combine with other atoms by sharing electrons. The elements in Group VIA and VIIA easily gain electrons to become negative ions. The tendency to gain electrons easily identifies elements that are commonly called non-metals. The Group V elements tend to gain electrons, but not as readily as Groups VIA and VIIA. They too often combine with other atoms by sharing electrons. The Group IVA elements invariably combine with other atoms by sharing electrons. They do not readily transfer electrons. The Group B elements have more complicated electron structures and will be discussed shortly.
Many of the elements in the table were named according to characteristic properties or sources for where they were found. For example, gold comes from the Latin term aurum meaning “shining dawn” which references its appearance. Phosphorus comes from Greek meaning “light bearer,” which is appropriate considering it is the main ingredient in match heads that spontaneously ignite with enough friction. Helium came from the Greek word helios that means “the sun” which is composed of helium and hydrogen. Moreover, several elements have been named after where they were first discovered or synthesized (e.g., Ytterbium, Yttrium, Berkelium, Lutetium, Americium, Dubnium, Germanium, and Scandium — just to name a few).
The periodic table organizes the elements in such a way that even someone unfamiliar with the substances can predict reactions (with both reactivity and composition) of the resulting compounds from the reaction, a significant advantage to having to be familiar with each individual substance. Now would be an excellent time to reacquaint yourself with the periodic table.
An elemental group is simply the vertical column of that element in the periodic table. There are 18 groups in the standard periodic table and they each correspond to the shared characteristics of those elements, based mainly on their shared number of valence electrons.
Beyond the 18 different groups, there are also classifications that are sometimes more meaningful and carry broader characteristics across more elemental groups. For example, there are nine basic classes of elements using this method:
The alkali metals are the elements in Group 1A on the periodic table (all the elements in the first column excluding hydrogen H): lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). With a single electron in their outermost or valence shell, these elements are some of the most reactive with other elements. Hence, they are rarely found in their pure elemental forms in nature. If or when they are in their pure forms, they tend to be extremely volatile and potentially dangerous. They are all low-density silver gray metals that are relatively soft. They react with water to form very basic hydroxides and also with halogens to form ionic salts. They lose their valence electrons most easily of all the elements. This lowest energy-state in losing an electron causes them to form positively charged ions. Thus they form ionic substances with those elements that readily gain electrons.
Hydrogen is not an alkali metal because it occurs naturally as a diatomic gaseous molecule and removal of its single electron requires considerably more energy than that for the alkali metals.
With two electrons in their valence orbit, the alkaline Earth metals are the elements in Group 2A (the second row from the left) of the periodic table: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). The alkaline Earth metals were named for their placement between the alkali metal oxides and the rare Earth metal oxides. These elements are also silver in color, low-density, soft metals, but they do not react as readily with halogens and water as the alkali metals. In this class, beryllium is the exception since it does not react with water and its halides are covalent compounds. Otherwise, this is a cohesive class of elements.
The rare Earths (lanthanides and actinides) are often the elements usually put in a separate section below the periodic table and share many properties that make them difficult to tell apart. The common properties for the lanthanides and actinides include similar coloration (silver, white, or gray), high luster, and good electrical conductivity. As a group, they also have very few differences in solubility and reactivity. The elements of the lanthanide series are the 15 rare Earth elements from lanthanum to lutetium (atomic numbers 57 through 71) on the periodic table. The actinide series encompasses the 15 chemical elements that lie between actinium and lawrencium (atomic numbers 89 – 103) on the periodic table. Although actinides are chemically similar to the lanthanides, the actinides are radioactive. Some actinides with higher atomic numbers are not found in nature and have extremely short half-lives.
These are the 39 elements in the very middle of the periodic table (in rows 3 through 7) that form at least one ion with a partially filled d electron orbit (except for zinc and scandium). The properties of the transition elements are the result of their ability to contribute valence electrons in a variety of ways the other elements do not.
The post-transition metals are the metallic elements on the right-hand side of the periodic table, occurring between the metalloids and the transition metals in rows 3-5: aluminum (Al), gallium (Ga), indium (In), tin (Sn), thallium (Tl), lead (Pb), and bismuth (Bi). These elements are usually soft at room temperature and they have relatively low melting and boiling temperatures. They are also more electropositive than the transition metals, but less than alkali metals and alkaline Earth metals.
The metalloids are the elements sandwiched between the post-transition metals and the non-metals in rows 3-7 on the periodic table: boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te), and polonium (Po). They share common features of ionization and reactivity and generally have intermediate properties between metals and non-metals.
Even though the non-metals are fewer in number than other groups, they play a major part of life on Earth. Non-metals are highly electronegative and gain electrons more readily than they give them up. The non-metals are the halogens, the noble gases, and hydrogen (H), carbon (C), nitrogen (N), oxygen (O), phosphorus (P), sulfur (S), chlorine (Cl), and selenium (Se). Most non-metals are found on the upper right-hand side of the periodic table (except for hydrogen in the uppermost left). Non-metals are generally either insulators or semiconductors and usually form ionic bonds with metals by gaining electrons. Many non-metals (hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine) are diatomic, and most of the rest are otherwise polyatomic.
The halogens are the highly reactive elements in Group 7A of the periodic table: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). These elements usually form diatomic molecules in their natural form and require one electron to fill their outer electron orbits. Hence, they have a strong tendency to form negative ions (referred to as halide ions) and salts (known as halides). These elements have increasing melting and boiling points and decreasing electronegativity as you move down the row. Because halogens are highly reactive, they can be quite dangerous. For example, chlorine and iodine are commonly used as disinfectants in drinking water, swimming pools, and many surfaces we want to be sterile. We primarily utilize chlorine both internally (as an ion, chloride is involved in many biological processes and important molecules) and externally as a disinfectant.
The noble gases are the chemical elements in group 8A, on the right side of the periodic table: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xn), and radon (Rn). The noble gases are also known as inert gases since they are not usually re-active during most chemical reactions. Their full electron shells make the inert gases highly unreactive except under extreme heat and pressure. They have very low melting and boiling points.
Generally, the arrangement in the periodic table provides a clear delineation of metals and non-metals, which can be explained by the atomic structures of the elements. Although we just went over the major classes of elements, most of the elements on the periodic table are metals and they are characterized by:
Elements on the far right side of the periodic table are non-metals and have the following characteristics:
Elements in between metals and non-metals on the table have intermediate characteristics. For example:
Elements in each group tend to have a common stable oxidation number that can be used to predict reactions. For example, Group 1 elements can be thought of as having an oxidation number of +1 because these elements commonly lose one electron to become more stable. Group 2 can be thought of as having an oxidation number of +2 because these elements commonly lose two electrons to become more stable. Group 3 can be thought of as having an oxidation number of +3 because these elements commonly give away three electrons to become more stable. Group 4 can be thought of as having an oxidation number of + or -4 because these elements commonly gain or lose four electrons to become more stable, depending on what other element it reacts with. Group 5 can be thought of as having an oxidation number of -3 because these elements commonly gain three electrons to become more stable. Group 6 can be thought of as having an oxidation number of -2 because these elements commonly gain two electrons to become more stable. Elements in group 7 gain one electron to become more stable and are said to have an oxidation number of -1.
To predict how elements will react, simply determine the whole number ratio required to make an algebraic sum of zero. Since negative one and positive one have an algebraic sum of zero, these elements tend to bond with each other to form stable compounds, such as sodium chloride NA+1 + Cl-1 → NaCl. Sodium chloride is much more stable than sodium or chlorine are individually, and the combination is the common substance known as table salt. Any compound that has an algebraic sum of zero based on the sum of the oxidation numbers of its elements has the potential for reacting and can be predicted based on its location on the periodic table.
Barium chloride, BaCl 2 is an example where barium has the oxidation number of +2 and chloride -1 meaning that it requires two -1 ions or (two chloride ions) to balance a single +2 ion, barium, for an algebraic sum of zero. Barium chloride is a common salt used to defrost roads in the wintertime. Carbon dioxide, CO2 is an example where carbon has an oxidation number of +4 and oxygen has an oxidation number of -2. Carbon dioxide is the gas that is produced as a waste product in the process of cellular respiration in all living things. Carbon will even bond with itself very tightly to form the hardest pure substance known, which is a diamond. Chemists’ first theories were based on simple observable proportions such as these to determine the relationships on the periodic table without a complete understanding of why these proportions were observed.
Today, we understand that the periodic table summarizes both the chemical and physical properties that are based on atomic structure. As the distance from the center of the nucleus increases, the complexity of the arrangement of the electrons around the nucleus increases. This can also be identified by the distance from the top of the periodic table or row. For example, the difference between lithium (element number 3) and sodium (element number 11), is that sodium has one more shell or layer of electrons than lithium. They both have one electron in their outermost shell (because they are in group number one making them very reactive), but because the second shell of electrons has more volume to occupy as a function of distance from the center of the atom (V = π · r3), there is the potential for a more complex arrangement in space to find an electron within that space. Sodium is more stable than lithium but reacts in the same proportion as lithium with other elements.
The difference between helium (element number 2) and neon (element number 10), is that neon has one more shell of electrons than helium. The properties of the two elements are similar — colorless, odorless gases — but the additional shell of electrons means neon also has more protons in the nucleus which more strongly hold onto the electrons.