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1 Name Period PRE-AP 14-1 Development of the Periodic Table Ch. 14 The Periodic Table p Dmitri Mendeleev published the first periodic table in He organized the elements by atomic mass. He even left holes for elements he said were yet to be discovered. One of these holes was below silicon. He named this element Ekasilicon and predicted its properties. Later germanium was discovered and was determined to be Mendeleev s Ekasilicon. Compare the properties that Mendeleev predicted to the actual properties of Germanium with Figure 5-5 on p. 162 of your text. In 1913 H.G.C. Mosley while working in Rutherford s lab, discovered a way to determine the atomic number of each element. He then improved on Mendeleev s periodic table when he reorganized the table by atomic number instead of atomic mass. 1

2 The Periodic Law says that when elements are arranged in increasing atomic number, their chemical and physical properties show a periodic pattern. We have already seen one example of this law in the electron configuration Reading the Periodic Table Groups or families are vertical columns on the periodic table. Periods are horizontal rows on the periodic table. Figure 14-5 on p. 395 shows the long form of the periodic table. The elements of all the "A" groups are called representative elements. (S & P Block) The elements of all the "B" groups are called transition elements. The elements in the bottom two rows (f-block) are the inner transition elements. Figure 14.2 on p identifies some group and section names. 2

3 YOU NEED COLOR PENCILS AND A NEW PERIODIC TABLE FOR THIS SECTION!!!! Group 1A are called the alkali metals. (Note: Hydrogen is its own group.) Group 2A are called the alkaline earth metals. 3

4 Group 3A are called the Boron group. Group 4A are called the Carbon group. Group 5A are called the Nitrogen group. Group 6A are called the Oxygen group. Group 7A are called the halogens. (Halogen means salt forming.) Group 8A are called the Noble Gases. The number of the A group elements is the number of valence electrons in the atom Elements in a group have similar properties because they have the same number and location of valence electrons.. Figure 14.2 on p is a periodic table with phases at room temperature (20 o C) and metallic properties. 4

5 Metals are elements that have high luster (looks shiny), are malleable (can be hammered into sheets), ductile (can be pulled into wire), and are good electrical and heat conductors. Almost all metals are solids at room temperature. Metals are on the left of the zigzag line. Nonmetals are elements that have low luster (looks dull), are NOT malleable or ductile, and poor conductors of heat and electricity. Most nonmetals are gases at room temperature, but several are solids and one is a liquid. Nonmetals are on the right of the zigzag line plus hydrogen. Semimetals (a.k.a. metalloids) are elements that have some characteristics of both metals and nonmetals Semimetals are on the zigzag line except Aluminum and Polonium both of which are metals. 5

6 14-2 Periodic Trends YOU NEED COLOR PENCILS AND A NEW PERIODIC TABLE FOR THIS SECTION!!!! Periodic trends are the properties of the elements that cycle in predictable ways as you move through the periodic table. Atomic radius (a.k.a. atomic size) is the distance from the center of the atom s nucleus to the outermost electron. Atomic radius decreases as you move left to right in any period. Atomic radius increases as you move down in any group. Fig on p The downward trend occurs because as you move down a group you increase the principal quantum number, n. As n increases, so does the size of the orbital s. The horizontal trend occurs because as you move left to right in a period you are increasing the number of protons in the nucleus, which increases the positive charge of the nucleus, which pulls the negatively charged electrons closer to the nucleus, which decreasing the size of the atom. 6

7 Ionic Radii: There are often several different ions each element can make. However, there is typically an ion of a particular charge that is most common. The list of common ion charges is below. Group 1A ions are Group 5A ions are Group 2A ions are Group 6A ions are Group 3A ions are Group 7A ions are Group 4A ions are Group 8A is Please note that each atom gain or loses electrons so that it has the electron configuration of the closest noble gas.. This is because the noble gas configuration is very stable because it has a full s and p outer sublevel also know as a stable octet. The size of these common ions is also periodic. Look at Fig.. Ionic size increases as you move down a group and decreases as you move left to right with a jump up at 5A. The reasons for both the downward and horizontal trends are the same as those for atomic size. However, in ionic size you have a jump because of the increase in the number of energy levels the ion has. How many electrons does C 4+ have? What is its configuration? How many electrons does N 3- have? What is its configuration? Density, Melting point, and boiling point all increase as you go down a group. These do not have significant left to right trends. Metals increase in reactivity as you go down a group. Nonmetals increase in reactivity as you go up a group. Again, these have no significant left to right trends. Ionization Energy: Ionization energy is the energy needed to remove an electron from the atom. Remember the SI unit for energy is the joule (J). The unit for a large number of atoms is the mol, abbreviated mol. So ionization energy is measured in kj/mol. 7

8 Li + energy --> Li + + e - You can think of ionization energy as how strongly the atom holds its outer-most electron. Ionization energy increases as you move left to right in any period. Ionization energy decreases as you move down in any group. See Fig p Please note these trends are opposite of the atomic size trends. The larger the attraction is between the nucleus and the outer-most electron, the smaller the atom will be and the larger the ionization energy will be. The energy needed to remove the first electron is sometimes called the first ionization energy. The energy to remove the second electron is the second ionization energy. Li ---> Li + + e - First ionization energy = 521 kj/mol Li + ---> Li 2+ + e - Second ionization energy = 7,304 kj/mol Li > Li 3+ + e - Third ionization energy = 11, 752 kj/mol Look at Fig on p Notice the large increase in ionization energy for sodium from its first to second. For magnesium the large increase is from second to third. Why do you suppose this is? It is again the idea of a stable octet or a stable noble gas configuration. Electron affinity is the energy change that occurs when an electron is added to an atom. Ne + e - ---> Ne - Electron affinity = 29 kj/mol A positive electron affinity means that amount of energy had to be added to the atom to make it take the electron. A negative electron affinity means that amount of energy was given off by the atom when the electron was added. F + e - --> F - Electron affinity = -328 kj/mol You can think of electron affinity as the amount of attraction an atom has for an extra electron. Remember that a negative electron affinity means the atom has more attraction than a positive one. 8

9 In general, metals have positive (low) electron affinities and nonmetals have negative (high) electron affinities. Again this trend relates to the stable octet rule. Electronegativity is the attraction an atom has for electrons in a bond. Electronegativity increase as you move left to right in any period. Electronegativity decreases as you move down in any group. See Figure 14.2 on p

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