UNIT 5 THE PERIODIC TABLE THE PERIODIC TABLE EARLY ATTEMPTS OF CLASSIFICATION Many chemists started to organize and classify the elements according to their properties. In the 1790s, Antoine LaVoisier compiled the list of 23 elements at the time. He is the first person acknowledging that an organization of elements might soon be necessary. In 1829, Johann Dobereiner noticed that elements grouped in threes seemed to share properties. He called these triads. He noticed that the average mass of largest and the smallest of the 3 elements was equal to the mass of the middle element in the triad. Then in 1863, John Newlands, an English chemist, decided to arrange the elements by increasing atomic mass. He found that by putting the elements into a table format of columns and rows, that the elements exhibited similar properties within each column. Specifically, Newlands set up the table with 7 columns. He then devised the Law of Octaves which states that the properties of elements repeat every 8 elements in order of atomic mass. In 1869, Dmitri Mendeleev, a Russian chemist, argued that Newlands table was not complete. He believed that the properties of the elements were a function of the atomic mass. He also argued that Newland s table did not allow spaces for any unknown elements. He reorganized the table to have 8 columns and left spaces for elements that may exist but have not yet been discovered. His table was accurate enough that he was able to predict the properties of those unknown elements. He then devised the Periodic Law, which states that the properties of elements are a periodic function of their atomic masses. MODERN PERIODIC LAW Although Mendeleev s table was very accurate and functional, there were a few elements that did not fit well into it. In 1913, Henry Mosely made a slight, but vital revision. Instead of arranging the elements in order of increasing atomic mass, he decided to arrange the table in order of increasing atomic number. This change did not affect most of the elements on the table. But the few that it did affect made all the difference. The few elements that didn t quite fit into Mendeleev s table now fit perfectly into Mosely s. He then changed Mendeleev s Period Law to reflect that. The MODERN PERIODIC LAW states that the properties of elements are a periodic function of their atomic number. 1
The table is arranged by Groups and Periods. A Group is a vertical column and there are 18 of them. The f-sublevel elements are not part of a Group. A Period is a horizontal row on the Table. There are 7 Periods. Be careful, the f-sublevel is actually part of Period 6 and Period 7 on the Table. USING THE PERIODIC TABLE Atoms are stable (less reactive) when they have a full Valence Energy Level. Since the d- and f- sublevels are never part of the Valence Energy Levels of atoms, the most Valence electrons possible is 8. The Octet Rule is a GENERAL rule of thumb that states that an atom is stable when it has 8 Valence electrons. Atoms will gain, lose, or share these Valence electrons in order to become stable. An important property to remember is that metals must lose electrons and nonmetals tend to gain electrons. This rule allows you to determine how many electrons need to be gained or lost in order to become stable. The true rule for stability is that the Valence Energy Level of the atom be full. So atoms are most stable if their VEL is full, but this is not always possible. Some atoms will instead lose all the Valence electrons and empty out the VEL. This actually can give them a full VEL it is just now one level lower than it was. And yet still there are some elements that can t fill the current Valence level and emptying it out does not result in a filled lower level. These elements can become partially stable. A partially stable configuration can be obtained when all the orbitals in a sublevel are half-full. Remember that electrons are always moving and will locate themselves in the most energy efficient arrangement. So a Chromium atom which should have a 4s 2, 3d 4 arrangement will actually rearrange to have partial stability by moving one electron and arranging as 4s 1, 3d 5. Now the 4s and the 3d are all half full. TYPES OF ELEMENTS & GENERAL PROPERTIES 1) METALS hard, opaque, malleable, ductile, shiny, good conductors of heat and electricity, and tend to have high MPs. 2) NONMETALS brittle solids, transparent or translucent, dull, poor conductors of heat and electricity, and tend to have low MPs. 3) METALLOIDS Properties of both metals and nonmetals, depends on the individual element. FAMILIES A family is a collection of elements that have very similar properties. You should know some properties of each family. Elements are most similar when they are in the same group. 1) ALKALI METALS Group 1 (not H) *The most reactive metals.* 2) ALKALINE EARTH METALS Group 2 Very reactive metals. 3) TRANSITION METALS Groups 3 12 Form colored compounds. 4) INNER TRANSITION METALS Also called the Rare Earth Elements, f-sublevel elements a. LANTHANIDES Period 6 of the Inner Transition Metals Bright, silvery elements that are very reactive. b. ACTINIDES Period 7 of the Inner Transition Metals Radioactive. 5) HALOGENS Group 17 *The most reactive nonmetals.* 6) NOBLE GASES Also called the INERT GASES, Group 18 *UNREACTIVE elements* 2
PERIODIC PROPERTIES The elements are located on the Periodic Table in such a way that shows that the properties and the position of elements are specifically related to their electron configuration. Elements in the same Group have the same Valence electron configuration. This means that their properties will be similar. So, in general, elements in the same column exhibit similar properties. If you know that Fluorine is a very reactive gas, then you can safely say that Chlorine is also a very reactive gas. The intensity of the property may vary a bit, but that is why they properties are similar as opposed to the same. Not only do groups link the properties, but there are connections across a period as well. The idea that properties are similar both vertically and horizontally on the table is called a trend. By knowing the trend, you can predict properties of elements that you may be unfamiliar with. ATOMIC RADIUS Atomic Radius is defined as the radius of an atom. As you move down a Group, energy levels are being added to the atom, which obviously makes the radius of the atom larger. So you can say that the trend of atomic radius of elements increase as you move down a Group. The trend of atomic radius also increases as you move from right to left across a Period. This is because as you move from left to right across the Period, you are adding one additional proton into the nucleus. This results in the nucleus having a stronger pull on the electrons. This trend holds true for all the elements on the table except for the Inert Gases. IONIC RADIUS Ionic Radius is the radius of the ions that are formed from atoms. We will address the trend of cations separately from anions. When an atom forms a cation, it has lost electrons to become positively charged. In most situations, the atom tends to lose enough electrons to empty out its Valence Energy Level to become stable. Therefore, the original Valence Energy Level is no longer present causing the resulting cation to always be smaller than the original atom. As far as the trend goes, as you move down a Group, there may be one less Energy Level than originally, but there is still one more than the Period above it. So ionic radius increases as you move down a Group. The ionic radius also increases from right to left. This is because the ions will most likely have lost all their Valence electrons, but not any protons. And since the number of protons increases from right to left, the pull on the electrons increases right to left, holding the electrons in closer. When an atom forms an anion, it gains electrons. Atoms are trying to fill their Valence Energy Level. Since the number of protons in the atom has not changed, the same number of protons is now trying to hold onto more electrons. The resulting anion is always larger than the original atom. 3
The trend of ionic radius of anions is similar to that of the cations. As you move from left to right across a Period, the atom needs to gain fewer electrons to become stable. And the ratio of protons in the nucleus to electrons in the atom will get closer to 1:1. The more electrons gained causes the ion to be larger since the pull of the protons on the electrons is weaker. So the ionic radius increases from right to left. As you move down a Period, again you are dealing with an additional energy level making the radius larger. IONIZATION ENERGY Ionization Energy is the amount of energy required to remove an electron from an atom. This measurement is expressed in kilojoules per mole, kl/mol. Ionization Energy is often used to identify whether an element is a metal or not. Metals have low ionization energies and nonmetals tend to have very high ionization energies. Since in some cases, more than one electron is removed from the atom, the amount of energy needed to remove each electron changes. First Ionization Energy refers to amount of energy needed to remove the most loosely held electron. The amount of energy needed to remove additional electrons (Second Ionization Energy, Third Ionization Energy, and so on) is always greater than the First Ionization Energy. This is because the pull on the electrons is now greater in comparison. The trend of increasing First Ionization Energy is to the right and up. The trend of ionization energy is determined by obvious values, but can be explained by four factors; Nuclear Charge, Shielding Effect, Radius, & Sublevel. Nuclear Charge The larger the positive charge in the nucleus, the harder it will be to remove the electron, causing a greater IE. Shielding Effect The shielding effect is when electrons in lower energy levels partially block the nucleus attraction to the outer electrons. The greater the shielding effect, the less energy it takes to remove those outer electrons lower IE. This is a main idea in understanding how electrons take part in reactions. On the right is an illustration to help you visualize this complex idea. Radius The radius of the atom affects the amount of energy needed to remove the electron. The greater the radius, the weaker the attraction between the nucleus and the valence electrons. This also increases the possibility of the shielding effect. So, the larger the atom, the lower the IE. Sublevel Electrons in a full or a half-full sublevel are more difficult to remove since there is some amount of stability in the atom. So they would require more energy higher IE. This explains why the trend increases, drops slightly, and then increases again as you move across a sublevel. 4
ELECTRONEGATIVITY Electronegativity is the ability of an atom to attract an electron in a covalent bond. This term is very similar to electron affinity. Electronegativity, EN, refers to an atom attracting a pair of electrons in a bonding situation. Electronegativity has a number value as a comparative scale. Electron affinity is a measure of the atom to attract one electron on its own. Metals tend to have low EN values, and nonmetals tend to have high EN values. The trend for Electronegativity increases as you move right along the Periodic Table and up. The same 4 factors that affect IE; nuclear charge, shielding effect, radius, and sublevel, also affect EN. Nuclear Charge The larger the positive charge in the nucleus, the easier it will be to attract the electron, causing a greater EN. Shielding Effect The greater the shielding effect, the less likely the atom will be able to attract electrons lower EN. Radius The radius of the atom affects the ability to attract the electron. The greater the radius, the weaker the attraction between the nucleus and the electron will be. This also increases the possibility of the shielding effect. So, the larger the atom, the lower the EN value will be. Sublevel Atoms that are stable or partially stable will not attract additional electrons that will throw off the stability they already have. This explains why the trend increases, drops slightly, and then increases again as you move across a sublevel. *The trend for Electronegativity DOES NOT apply to the Inert Gases. These elements are naturally stable. Therefore, there is no reason for them to gain or lose electrons. All changes in the atom s electron configuration would make it unstable. 5