Periodic Table   Chapter 14
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Periodic Table Chapter 14

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Periodic Table   Chapter 14 Periodic Table Chapter 14 Presentation Transcript

  • Chemical Periodicity Or How I Learned to Love Blowing Things Up!
  • Development of the Periodic Table About 70 elements had been discovered by the mid-1800’s, but no one had found a way to relate the elements in a systematic, logical way . Dmitri Mendeleev listed the elements in columns in order of increasing atomic mass. He then arranged the columns so that the elements with the most similar properties were side by side. Mendeleev left blank spaces in the table because there were no known elements with the appropriate properties and masses.
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  • Mendeleev and others were able to predict the physical and chemical properties of the missing elements. Eventually these elements were discovered and were found to have properties similar to those predicted. In 1913, Henry Moseley, a British physicist, determined the atomic number of the atoms of the elements. Moseley arranged the elements in a table by order of atomic number instead of atomic mass. That is the way the periodic table is arranged today.
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  • The Elements Song Neat-o Animation
  • The Modern Periodic Table Again, notice that the elements are listed in order of increasing atomic numbers, from left to right and from top to bottom. The horizontal rows of the periodic table are called periods . There are seven periods. The properties of the elements change as you move from left to right from element to element. The pattern of properties within a period repeats, however, when you move from one period to the next.
  • This gives rise to the periodic law : When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. Elements that have similar chemical and physical properties end up in the same column in the periodic table. Each vertical column of elements in the periodic table is called a group , or family. The elements in any group of the periodic table have similar physical and chemical properties. Each group is identified by a number and the letter A or B.
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  • The Group A elements are made up of Group 1A through Group 7A and Group 0 (or 8A or 18) (the group at the far right). Group A elements are called the representative elements because they exhibit a wide range of both physical and chemical properties . The representative elements can be divided into three broad classes . The first are metals , which have a high electrical conductivity and a high luster when clean. They are ductile (able to be drawn into wires) and malleable (beaten into sheets). Except for hydrogen, the representative elements on the left side of the periodic table are metals.
  • The Group 1A elements are called the alkali metals , and the Group 2A elements are called the alkaline earth metals . Most of the remaining elements that are not Group A elements are also metals . These include the transition metals and inner transition metals , which together make up the Group B elements. Copper, gold, and silver are familiar transition metals. The inner transition metals, which appear below the main body of the periodic table, are also called the rare-earth elements. Approximately 80% of all the elements are metals.
  • Metals
  • Inner Transition Metals
  • The nonmetals occupy the upper-right corner of the periodic table. Nonmetals are elements that are generally nonlustrous and that are generally poor conductors of electricity. Two groups on nonmetals are given special names. The nonmetals of Group 7A are called the halogens , which include chlorine and bromine. The nonmetals of Group 0 (or 18 or 8A) are known as the noble gases , which are sometimes called the inert gases because they undergo few chemical reactions.
  • Nonmetals
  • On your periodic table you will notice a heavy stair-step line. This line divides the metals from the nonmetals. Most of the elements that border this line are metalloids , elements with properties that are intermediate between those of metals and nonmetals. Silicon and germanium are two important metalloids that are used in the manufacture of computer chips and solar cells.
  • Metalloids
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  • Classification Of The Elements The periodic table is probably the most important tool in chemistry. Among other things, it is very useful for understanding and predicting the properties of the elements. For example, if you know the physical and chemical properties of one element in a group or family, you can predict the physical and chemical properties of the other elements in the same group - and perhaps even the properties of the elements in neighboring groups.
  • Classifying Elements by Electron Configuration Of the three major subatomic particles, the electron plays the most significant role in determining the physical and chemical properties of an element. The arrangement of elements in the periodic table depends on these properties. Thus there should be some relationship between the electron configurations of the elements and their placement in the table.
    • Elements can be classified into four categories according to their electron configurations.
    • The noble gases . These are elements in which the outermost s and p sublevels are filled.
    • The noble gases belong to Group 0. The elements in this group are sometimes called the inert gases because they do not participate in many chemical reactions. The electron configurations for the first four noble-gas elements are listed below. Notice that these elements have filled outermost s and p sublevels.
    Helium 1 s 2 Neon 1 s 2 2 s 2 2 p 6 Argon 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 Krypton 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 10 4 s 2 4 p 6 Notice that all of these elements have filled outermost s and p sublevels
  • 2. The representative elements . In these elements, the outermost s and p sublevel is only partially filled. The representative elements are usually called the Group A elements. For any representative element, the group number equals the number of electrons in the outermost energy level. For example, the elements in Group 1A (lithium, sodium, etc.) have one electron in the outermost energy level. Lithium 1 s 2 2 s 1 Sodium 1 s 2 2 s 2 2 p 6 3 s 1 Potassium 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 1 Notice the one electron in the outermost energy level
  • Carbon, silicon, and germanium, in Group 4A, have four electrons in the outermost energy level. Carbon 1 s 2 2 s 2 2 p 2 Silicon 1 s 2 2 s 2 2 p 6 3 s 2 3 p 2 Germanium 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 10 4 s 2 4 p 2 3. The transition metals . These are metallic elements in which the outermost s sublevel and nearby d sublevel contains electrons. The transition elements, called Group B elements, are characterized by addition of electron to the d orbitals. Scandium 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 1 Yttrium 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 10 4 p 6 5 s 2 4 d 1 Notice the four electrons in the outermost energy level Notice the outermost “s” and “d” electrons are filling.
  • 4 . The inner transition metals . These are metallic elements in which the outermost s sublevel and nearby f sublevel generally contain electrons. The inner transition metals are characterized by the filling of f orbitals. Cerium [Xe]6 s 2 5 d 1 4 f 1 Thorium [Rn]7 s 2 6 d 1 5 f 1 If you consider both the electron configurations and the positions of the elements in the periodic table, another pattern emerges. The periodic table can be divided into sections, or blocks, that correspond to the sublevels that are filled with electrons. Notice the outermost “s” and “f” electrons are filling.
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  • The s block is the part of the periodic table that contains the elements with s 1 and s 2 outer electron configurations. It is composed of the elements in Groups 1A and 2A and the noble gas helium. The p block is composed of the elements in Groups 3A, 4A, 5A, 6A, 7A, and 0 with the exception of helium. The transition metals belong to the d block, and the inner transition metals belong to the f block.
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  • Periodic Trends You know from the quantum mechanical model that an atom does not have a sharply defined boundary that sets the limit of its size. Therefore, the radius of an atom cannot be measured directly. There are, however, several ways to estimate the relative size of atoms. The atomic radius is one-half of the distance between the nuclei of two like atoms in a diatomic molecule.
  • Group Trends Atomic size generally increases as you move down a group of the periodic table. Why? As you descend, electrons are added to successively higher principal energy levels and the nuclear charge increases. The outermost orbital is larger as you move downward. The shielding of the nucleus by electrons also increases with the additional occupied orbitals between the outermost orbital and the nucleus.
  • Periodic Trends Atomic size generally decreases as you move from left to right across a period. Why? As you go across a period, the principal energy level remains the same. Each element has one more proton and one more electron than the preceding element. The electrons are added to the same energy level, causing the increasing nuclear charge to pull them in closer.
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  • Trends in Ionization Energy When an atom gains or loses an electron, it becomes an ion. The energy required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom is called the ionization energy . Removing one electron results in the formation of a positive ion with a 1+ charge. Na( g ) Na + ( g ) + e -
  • The energy required to remove this first outermost electron is called the first ionization energy . To remove the outermost electron from the gaseous 1+ ion requires an amount of energy called the second ionization energy, and so forth. Nifty swell animation
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  • Notice the large increase in energy between the first ionization energy and the second ionization energy for sodium. This means it is relatively easy to remove one electron from a Group 1A metal to form an ion with a 1+ charge. It is difficult, however, to remove an additional electron. For the Group 2A metals, the large increase in ionization energy occurs between the second and third ionization energies. What does this tell you about the ease of removing one electron from these metals?
  • Group Trends The first ionization energy generally decreases as you move down a group of the periodic table. This is because the size of the atoms increases as you descend, so the outermost electron is farther from the nucleus. The outermost electron should be more easily removed, and the element should have a lower ionization energy.
  • Periodic Trends For the representative elements, the first ionization energy generally increases as you move from left to right across a period. The nuclear charge increases and the shielding effect is constant as you move across. A greater attraction of the nucleus for the electron leads to the increase in ionization energy.
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  • Trends in Ionic Size The atoms of metallic elements have low ionization energies. They form positive ions easily, because they have a tendency to lose electrons easily. By contrast, the atoms of nonmetallic elements readily form negative ions (they attract electrons). Why? And how do those differences affect ion sizes? Peachy keen animation
  • Group Trends Positive ions (cations) are always smaller than the neutral atoms from which they form. This is because the loss of outer-shell electrons results in increased attraction by the nucleus for the fewer remaining electrons. In contrast, negative ions (anions) are always larger than the neutral atoms from which they form. This is because the effective nuclear attraction is less for an increased number of electrons.
  • Periodic Trends Going from left to right across a row, there is a gradual decrease in the size of positive ions. Then beginning with Group 5A, the negative ions, which are much larger, gradually decrease in size as you continue to move right. The ionic radii of both anions and cations increase as you go down each group .
  • Cation vs. Anion
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  • Trends in Electronegativity The electronegativity of an element is the tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element. Electronegativity generally decreases as you move down a group. As you go across a period from left to right, the electronegativity of the representative elements increases.
  • The metallic elements at the far left of the periodic table have low electronegativities. By contrast, the nonmetallic elements at the far right (excluding the noble gases), have high electronegativities. The electronegativity of cesium, a metal, the least electronegative element, is 0.7; the electronegativity of fluorine, a nonmetal, the most electronegative element, is 4.0. Because fluorine has such a strong tendency to attract electrons, when it is chemically combined to any other element it either attracts the shared electrons or forms a negative ion. In contrast, cesium has the least tendency to attract electrons. Groovy animation
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  • Trends in Electron Affinity The electron affinity ( E.A. ) is the energy released upon adding an electron to a neutral atom in the gaseous state to form a negative ion (anion). 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6 F(g) Generally, the energy that results from this process (the electron affinity) is negative or close to zero. The more negative this energy the more this process is favored + e- F- .
  • Note that the noble gases, alkali metals and alkali earth metals have E.A. close to zero - indicating that these groups of elements do not particularly like to become anions. However, the nonmetals and especially the halogens are highly negative and thus readily become anions. A periodic trend is evident, as was the case for the ionization energy. Metals have lower (more positive) electron affinities and nonmetals have the higher (more negative) electron affinities . Phat animation
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