Department of Physics and Geology The Elements and the Periodic Table

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1 Department of Physics and Geology The Elements and the Periodic Table Physical Science 1422 Equipment Needed Qty Periodic Table 1 Part 1: Background In 1869 a Russian chemistry professor named Dmitri Mendeleev and Lothar Meyer from Germany independently established an arrangement of elements similar to our present day periodic table. Mendeleev produced a chart summarizing the known properties of the elements for his students. When the elements were arranged in order of increasing atomic weight, Mendeleev noticed that the chemical properties of the elements repeated. Lithium and sodium, for example, are similar not only in their general physical properties (appearance, the conduction of electricity, etc.), but also in how they combine with other elements (chemical properties). Any element that combines with lithium will combine in exactly the same proportion with sodium. In order to properly align elements down a column, Mendeleev occasionally had to leave a blank space that was not filled by any known element. Instead of looking upon these blank spaces as a defect with his chart, Mendeleev predicted that new elements were still waiting to be discovered! By studying the properties of the elements around the gap and the variation of these properties across the rows and down the columns, he went on to predict the properties of these unknown elements. These predictions helped lead to the eventual discovery of those elements. Figure 1. Mendeleev s Periodic Table. (Source: web.fccj.org/~ethall/2045/ch5/mendelev.htm) Mendeleev's periodic table helped confirm that all matter is made of atoms. At the time though he had no idea why the properties of the elements repeated themselves. Today we know that this is because of the structure of the atom. The periodic tables designed by Mendeleev and Meyer were virtually identical but since Mendeleev drafted his table earlier in 1869 and because his table included blanks for yet to be discovered elements to fit, Mendeleev is considered the father of the modern periodic table.

2 The structure of the atom and the modern periodic table The atom is composed of a small, heavy nucleus that is surrounded by probability clouds of negatively charged particles called electrons. The nucleus is composed of protons (a heavy, positivelycharged particle) and neutrons (a heavy particle with no charge). The number of protons in an atom is known as its atomic number. This number defines each element. For example, hydrogen is the element that has one proton and lithium is the element that has three protons. Because all atoms under normal conditions are neutral, the number of electrons in an atom is equal to the number of protons. Unlike the number of protons, the number of neutrons in the nucleus can vary for a given element. For example, hydrogen can have no neutrons (most common hydrogen), one neutron (a hydrogen isotope called deuterium) or two neutrons (a hydrogen isotope called tritium). These different forms of an element are called isotopes. Although different isotopes of a given chemical element will have slightly different weights, the isotopes of the given element will all have the same chemical behavior. Each element is normally a mixture of several isotopes. This is why the atomic mass of most elements is not exactly an integer. Each element in the periodic table has a name and a symbol. In most cases the symbol is just an abbreviation. For example the symbol for aluminum is Al. Those elements, which have been known for a very long time, usually have a symbol that is related to their Latin name. For example the Latin word for gold is aurus (Au), and iron is ferrus (Fe). In the modern periodic table, the rows are called periods and the columns are called groups. Elements are arranged strictly in order of increasing atomic number. Atomic structure As indicated above, the number of protons determines how many electrons an element will have. In turn, the number of electrons determines the chemical behavior of an element. The possible states of each electron in an atom are described by a set of four quantum numbers: n, l, ml, and ms. Two different electrons cannot share the same state. Thus, two different electrons in an atom have to be represented by two different sets of values of quantum numbers. The principle quantum number n (n = 1, 2, 3,...) corresponds approximately to the average distance of an electron from the nucleus. Electrons with the same value of n have about the same energy and are said to occupy the same shell. The orbital angular momentum quantum number l (l = 0, 1, 2,..., n-1) corresponds approximately to the shape of the orbital (probability cloud distribution) of an electron around the nucleus. Electrons with the same values of n and l are said to be in the same subshell. Standard notation represents the values of the quantum number l with letters: l designation s p d f g h i p. 2

3 That is, l=0 is known as the s-orbital, l=l is the p-orbital, l=2 is the d-orbital, l=3 is the f-orbital, and so on. When n=1, there is only one possible subshell (l=0); for n=2 there are two possible subshells (l=0 and l=l); and so on. That is, for a given value of n, there are n possible values for l. When denoting a subshell, the standard notation is to use a number (principal quantum number) accompanied with a letter (orbital angular momentum quantum number); for example the subshell 3p corresponds to subshell with quantum numbers n=3 and l=1. Quantum number ml (ml = -l, -l+1,, -2, -1, 0, 1, 2,, l-1, l ) corresponds approximately to the orientation in space of the orbital ( probability cloud distribution) of an electron around the nucleus. When l=0, there is only one possible value for ml (ml = 0). When l=1, there are three possible values for ml (ml = -1,0,1). In general, for a given value of l, there are 2l+1 possible values for ml. Quantum number spin ms corresponds to the spin orientation of the electron. ms has two possible values -1/2 and +1/2. Since two different electrons cannot be in the same state, for each subshell there is a maximum number of electrons. The maximum number of electrons that can be present in a given subshell is determined by the orbital quantum number l. For example, for an s-subshell (l=0, for a given n), the only possible value for ml is 0. Considering that ms has two possible values, the total number of electrons that can be in an s- subshell is 2. Principal quantum number: n Orbital Angular Momentum: 0 (s-orbital) Possible states: n=principal, l=0, ml = 0, ms = -1/2 n=principal, l=0, ml = 0, ms = +1/2 Total # states: 2 Similarly for a p-subshell (l=1, for a given n), one finds that the total number of electrons that can be there is 6: Principal quantum number: n Orbital Angular Momentum: 1 (p-orbital) Possible states: n=principal, l=1, ml = -1, ms = -1/2 n=principal, l=1, ml = -1, ms = +1/2 n=principal, l=1, ml = 0, ms = -1/2 n=principal, l=1, ml = 0, ms = +1/2 n=principal, l=1, ml = +1, ms = -1/2 n=principal, l=1, ml = +1, ms = +1/2 Total # states: 6 The larger the value of l, the more electrons a particular subshell can hold. In general, for a subshell with orbital quantum atomic number l, the total number of possible states in the subshell is given by 2(2l+1), equal to 4l +2. Therefore every time we increase l by one, the number of possible states increases by four. Thus a d-subshell (l=2) has 10 possible states (6+4); an f-subshell (l=3) has 14 possible states (10+4); and so on. p. 3

4 The sequence in which electron subshells are filled in an atom corresponds to the order of increasing electron energy: ls, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s Most periodic tables give the electron configuration of each element - this describes how the electrons are arranged around the nucleus. Each subshell is identified by its principle quantum number n followed by the letter corresponding to its orbital quantum number l (standard notation). A superscript after the letter indicates the number of electrons in that subshell. For example, nitrogen has seven electrons and its electron configuration is 1s 2 2s 2 2p 3. This tells us that this element contains two 1s electrons (n=1, l=0), two 2s electrons (n=2, l=0), and three 2p electrons (n=2, l=l). We can visually describe this arrangement by placing a straight line above a subshell for each possible value of quantum number ml, and by representing each electron with an up or down arrow indicating its spin (up indicating ms = +1/2, down indicating ms = -1/2). For example, beryllium has 4 electrons, and thus its electron configuration is 1s 2 2s 2, which can be represented as: 1 2 Another example would be the electronic configuration for sodium (atomic number = 11) which is 1s 2 2s 2 2p 6 3s 1 and can be represented as: A shell or a subshell that contains its full quota of electrons is said to be closed. In sodium the 1s, 2s and 2p subshells are closed, but 3s is considered open since it can still have one more electron. Chemical Activity To understand how electrons determine the chemical activity of an atom, you need to remember electrons in closed shells do not contribute to chemical reactions. Only the electrons in the outermost subshell contribute to chemical reactions. All metal atoms have one or more electrons outside a closed shell or subshells. Since these electrons are weakly bound, metal atoms combine chemically to nonmetal atoms by losing these electrons. All nonmetal atoms need one or more electrons to achieve a closed shell or a subshell and the maximum stability. In chemical reactions, these atoms pick up electrons from metal atoms or from sharing electrons with other nonmetal atoms. The inert gases have closed shells or subshells. This makes it hard for them to gain or lose electrons. As a result, they have almost no ability to react chemically. Answer the questions in the Lab Report section. p. 4

5 Name Class Date Part 3: Lab Report The Elements and the Periodic Table Questions 1. What is the name of the element with 6 protons? 2. What is the name of the element with 16 protons? 3. What is the name of the element with 36 electrons? 4. What element has the symbol Pb? 5. What element has the symbol Ar? 6. Elements in the same group have similar chemical properties. A) True B) False 7. Phosphorus and Sulfur belong to the same... A) Group B) Period 8. Two atoms have the same number of neutrons, but a different number of protons. Can they be the same element? A) Yes B) No Put the correct number of electrons in each orbital for: Boron: Aluminum: Oxygen: Argon: Sodium: p. 5

6 14. In Fluorine, which orbitals are closed? 15. Which orbital has quantum numbers n=2 and l=0? A) 1s B) 2s C) 2p 16. Which orbital has a higher energy than the 4s orbital? A) 2p B) 3d C) 1s 17. Give the electronic configuration of cobalt. 18. Give the electronic configuration of magnesium. 19. How many electrons in a silicon atom can contribute to chemical reactions (i.e., are not in closed shells)? 20. How many electrons in a neon atom can contribute to chemical reactions (i.e., are not in closed shells)? p. 6

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