Chapter 5 The Periodic Law

Transcription

1 Chapter 5 The Periodic Law 5-1 History of the Periodic Table A. Mendeleev and Chemical Periodicity 1. Dmitri Mendeleev went about organizing the elements according to their as you might organize information for a. He placed the name of each known element on a, together with the atomic of the element and a list of its observed and properties. He then arranged them according to various properties and looked for or. 2. Mendeleev noticed that when the elements were arranged in order of atomic, certain similarities in their chemical properties appeared at intervals. Such a repeating pattern is referred to as. 3. His first periodic table was published in. a. He placed, (atomic mass ), after, (atomic mass ). It allowed him to place in a group of elements with which it shares similar. b. Mendeleev s procedure left several in his periodic table. He boldly the existence and properties of the elements that would fill of the spaces. Today these elements are known as,, and. Their properties are to those predicted by Mendeleev. B. Moselely and the Periodic Law 1. In 1911, Henry Moseley examined the of 38 different metals. He discovered a previously unrecognized - The elements in the periodic table were arranged in increasing order according to, or the number of in the. 2. Moseley s work led to both the modern definition of and the recognition that, not, is the basis for the organization of the periodic table 3. Today, Mendeleev s principle of chemical is correctly stated in what is known as the periodic law: The physical and chemical properties of the elements are functions of their C. The Modern Periodic Table The Noble Gases 1. The periodic table is an arrangement of the elements in order of their so that elements with similar properties fall in the same, or. 2. In 1894 scientists discovered,, a gas in the atmosphere that had previously because of its total lack of chemical. In 1868 another noble gas,,, had been discovered as a component of the, based on the emission of. 3. Ramsay proposed a new group which he placed between the groups now known as Group (the family) and Group (the family). 4. In 1898 the noble gases,, and,, were added to the new group. The final noble gas,,, was discovered in The Lanthanides 5. In the early 1900 s the chemistry of the lanthanides was finally understood. 6. The lanthanides are the elements with atomic numbers from (, ) to (, ). 7. Because these elements were so in chemical and physical properties, the process of and them was a tedious task. 1

2 The Actinides 8. The actinides are the elements with atomic numbers (, ) to (, ). 9. The lanthanides and actinides belong in Periods and, respectively, between Groups and. To save, they are usually set off the main portion of the periodic table. Periodicity 10. Periodicity with respect to atomic can be observed in any of elements in the periodic table. 11. The differences in atomic number between Group metals follows the same as the differences in atomic number between the. 12. The reason for is explained by the arrangement of around the nucleus. 5-2 Electron Configuration and the Periodic Table A. Intro 1. The Group elements of the periodic table (the ) undergo few. This stability results from the gases special electron. 2. Generally, the electron configuration of an atom s energy level governs the atom s. B. - Periods and Blocks of the Periodic Table 1. Elements are arranged vertically in the periodic table in that share similar chemical. They are also organized in rows, or periods. There are a total of periods in the modern periodic table. Instructor s Note: The book spends a lot of time in this section on the relation between electron configuration and the blocks in the periodic table. Since this concept was used to teach you electron configuration in Chapter 4, it will be assumed that you understand the concept and it will be skipped in this reading guide. The s-block Elements: Groups 1 and 2 2. The elements of the s-block are chemically. 3. The elements of Group 1 of the periodic table are called the. a. They are so reactive that they are not found in as elements b. They combine vigorously with most, and they react strongly with to produce gas. 4. The elements of Group 2 of the periodic table are called the - metals. a. They are reactive than the alkali metals, and are also too reactive to be found in as elements. Hydrogen and Helium 5. Hydrogen does not share the same as the elements of Group. It is a element, with properties that do not closely those of any group. 6. Because helium s highest occupied energy levelis by electrons, it possesses special chemical, exhibiting the unreactive nature of a Group element. The d-block Elements: Groups The d-block elements are with typical properties and are often referred to as elements. a. They are typically than alkali metals and alkaline-earth metals. b. Palladium, and are among the least reactive of all elements. 2

3 The p-block Elements: Groups The p-block elements together with the s-block elements are called the elements. 9. The properties of elements of the p-block. At its right hand end, it includes all of the except hydrogen and helium. All of the are also in the p-block. There are p-block metals. 10. The elements of Group are known as the halogens. They are the most nonmetals. The f-block Elements: Lanthanides and Actinides 11. There are a total of f-block elements between lanthanum and hafnium. There are also f-block elements, the, between actinium and element. 12. The actinides are all. The first actinides have been found on Earth. The remaining actinides are known only as -made elements. 5-3 Electron Configuration and Periodic Properties A. Atomic Radii 1. Atomic radius may be defined as - the distance between the of identical atoms that are bonded together. 2. The trend to atoms across a period is caused by the increasing charge of the nucleus. 3. As electrons occupy sublevels in successively higher main energy levels located from the nucleus, the sizes of the atoms. In general, the atomic radii of the -group elements down in a group. B. Ionization Energy 1. An electron can be from an atom if enough is supplied. Write the equation for ionization: The represents an of element A with a positive charge. a. An ion is an atom or group of atoms that has a or charge. 2. Any process that results in the of an is referred to as ionization. 3. Ionization energy (or ionization energy) is the energy required to one electron from a atom. Measurements of ionization energies are made on atoms in the phase. Write the equation for the first ionization of sodium (at #11), using the ionization energy from the table on page 143: Period Trends 4. Group metals have the first ionization energies in their respective periods. This is a major reason for the of this group. 5. Group elements, the, have the ionization energies. They do not lose electrons. 6. In general, ionization energies of the -group elements across each period a. The increase is caused by nuclear charge (and smaller radius) 3

4 Group Trends 7. Among the main-group elements, ionization energies generally down the groups. a. Electrons removed from atoms of each succeeding element in a group are in higher energy levels from the nucleus b. As atomic number increases going down a group, more lie between the nucleus and electrons in the highest occupied energy levels. This partially the electrons from the effect of the nuclear. Removing Electrons from Positive Ions 8. The energies for removal of electrons from an atom are referred to as the ionization energy, ionization energy, and so on. 9. Each electron removed from an ion feels an increasingly effective nuclear (the nuclear minus the electron ). 10. Large jumps in ionization energy occur when an ion assumes a -gas configuration. Write the equations for the first, second and third ionizations of Magnesium (at# 12), using ionization energies from Table 5-3 on page 145. Explain why there is such a great increase in ionization energy between the removal of the second and third electron in magnesium: C. Electron Affinity 1. Electron affinity is the change that occurs when an electron is by a neutral atom. 2. Most atoms energy when they acquire an electron: + The quantity of energy released is represented by a number. 3. Some atoms must be to gain an electron by the addition of energy: + The quantity of energy absorbed is represented by a number. An ion produced in this way will be and will the added electron spontaneously. 4. The (Group ) gain electrons most. 5. Ignore the text at this point, and look at the table on page 147 to answer the following two questions: a. IN GENERAL, as you go across a period, electron affinity tends to, except for irregularities due to the stability of -filled and -filled sublevels b. IN GENERAL, as you go down in a group of main-group element, electron affinity tends to 6. It is always more to add a second electron to an already charged ion. Therefore second electron affinities are all. 4

5 D. Ionic Radii 1. A cation is a ion. The formation of a cation by the of one or more always leads to a in atomic radius. 2. An anion is a ion. The formation of an anion by the addition of one or more always leads to an in atomic radius. 3. Within each period of the periodic table, the at the left tend to form and the at the upper right tend to form. Period Trends 4. Cationic radii across a period because the electron cloud due to the increasing charge acting on the electrons. 5. Anionic radii across each period for the elements in Groups -. Group Trends 6. Just as there is a gradual of atomic radii down a group, there is also a gradual of ionic radii. E. - Valence Electrons 1. Chemical compounds form because electrons are,, or between atoms. The electrons that interact in this manner are those in the energy levels. 2. Valence electrons are the electrons available to be,, or in the formation of chemical compounds. They are often located in -filled main energy levels. For main-group elements, the valence electrons are the electrons in the outermost and sublevels. F. Electronegativity 1. Electronegativity is a measure of the ability of an atom in a chemical compound to electrons. 2. The most electronegative element is, and it is arbitrarily assigned an electronegativity value of. 3. Electronegativities tend to across each period, although there are. a. The and -earth metals are the least electronegative elements b. Nitrogen,, and the are the most electronegative elements 4. Electronegativities tend to either down a group or remain the. G. Periodic Properties of the d- and f-block Elements Skip it *** On a separate piece of paper, answer Chapter Review Problems #36, 41 and 46 from pages 157 and 158. Attach your answers to THIS PAGE! 5