The Periodic Table; Chapter 5: Section 1 - History of the Periodic Table Objectives: Explain the roles of Mendeleev and Moseley in the development of

Transcription

1 The Periodic Table; Chapter 5: Section 1 - History of the Periodic Table Objectives: Explain the roles of Mendeleev and Moseley in the development of the periodic table. Describe the modern periodic table. Explain how the periodic law can be used to predict the physical and chemical properties of elements. - Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic number. Warm-Up: - List three things you can find out about an element by looking at a periodic table. - List two questions you have about the periodic table - what do you still want to know about it? - What does the term "periodic" mean? How does it apply to the periodic table?

2 Mendeleev's Periodic Table: Compare this table to our modern periodic table. - Similarities? - Differences? Where are the spaces for the elements Mendeleev predicted?

3 The Modern Periodic Table An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group. -Arranged by atomic number - Additions: - Noble gases - Lanthanides - Actinides - Periodicity - similar properties among members of a chemical group or family

4 Section 2-Electron Configuration and the Periodic Table Objectives: - Describe the relationship between electrons in sublevels and the length of each period of the periodic table. - Locate and name the four blocks of the periodic table. Explain the reasons for these names. - Discuss the relationship between group configurations and group numbers. - Describe the locations in the periodic table and the general properties of the alkali metals, the alkaline-earth metals, the halogens, and the noble gases. **************************************************************** Warm-up- 1. Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Kr]5s 1 is located. 2. a. Without looking at the periodic table, write the group configuration for Group 2 elements b. Without looking a the periodic table, write the complete electron configuration for the Group 2 elements in the fourth period.

5 Niels Bohr proposed, in 1913, what is now called the Bohr model of the atom. He suggested that electrons could only have certain classical motions: 1. Electrons in atoms orbit the nucleus 2. The electrons can only orbit at a certain discrete set of distances from the nucleus called orbitals 3. Electrons can only gain and lose energy by jumping from one allowed orbit to another according to Planks relationship: E=h

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7 Section 3-Electron Configuration and Periodic Properties Objectives: - Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity. - Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations. - Define valence electrons, and state how many are present in atoms of each maingroup element. - Compare the atomic radii, ionization energies, and electronegativities of the d- block elements with those of the main-group elements. **************************************************************** Warm-Up: 1. Identify the block, period, group, group name (where appropriate), and element name for the elements with the following electron configurations. a) [Ne]3s23p1 b) [Ar]3d104s24p6 c) [Kr] 4d105s1 2. Show the trend for electronegativity

8 Electron Configuration and Periodic Properties Atomic Radius - one-half the distance between the nuclei of identical atoms that are bonded together. Atomic radii decrease from left to right across a period and increase down a group.

9 The First Ionization Energy The process by which the first ionization energy of hydrogen is measured would be represented by the following equation. H(g) H+(g) + e- Ho = kj/mol

10 Ionic Radii - Ca+ions vs. Anions The ionic radius is different from the atomic radius of an element. Positive ions are smaller than their uncharged atoms. Negative ions are larger than their atoms. ca+ion - a positive ion formed by the loss of one or more electrons from a neutral atom (metal elements form ca+ions) anion - a negative ion formed by the gain of one or more electrons by a neutral atom (nonmetal elements form anions)

11 Valence Electrons Valence electrons are the electrons available to be lost, gained, or shared in the formation of chemical compounds. for main group elements, the valence electrons are those in the outermost s and p sublevels Examples: A Mg atom has 2 valence electrons. [Ne]3s2 An Al atom has 3 valence electrons. [Ne]3s23p1 A Br atom has 7 valence electrons. [Ar]3d104s24p5

12 Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. Fluorine is the most electronegative element! In general, electronegativities increase across a period and decrease down a group (when there is a trend). The most electronegative elements are in the upper right corner of the periodic table, while the least electronegative elements are located in the lower left corner of the periodic table.

13 Practice -ch 4 1. What is the wavelength of the radiation whose frequency is 5.00 X sec -1? c ; c x m / sec X 10 /sec 8 6.0x10 m(60 nm)

14 6. For each of the first five principal energy levels (n=1 through n=5), state the number of sublevels, the number of orbitals, and the maximum number of electrons that can occupy the energy level. What are the mathematical expressions for each of these? 6. n=1: 1 sublevel, 1 orbital, 2 electrons n=2: 2 sublevels, 4 orbitals, 8 electrons n=3: 3 sublevels, 9 orbitals, 18 electrons n=4: 4 sublevels, 16 orbitals, 32 electrons n=5: 5 sublevels, 25 orbitals, 50 electrons

15 7. State the three major rules that govern the filling of electrons in orbitals. Be sure you know the names of the rules and the way each affects the filling of orbitals and the writing of electron configurations. 7. Pauli Exclusion Principle Aufbau Principle Hund s Rule

16 8Write the electron configurations for the following atoms: Mg P Br Xe

17 Identify the elements that have the following electron configurations: a) 1s 2 2s 2 2p 6 3s 2 3p 1 b) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 c) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 7 4s 2 9. a) Al b) Kr c) Co 11. a)fr b) Mn c) Ne d) C e)al,si,p,s,cl f)cd

18 Consider the elements neon, bromine, and phosphorus. Which has: a) three electrons in its 3p sublevel? b) its highest energy level completely filled? c) the highest occupied energy level? 10. a) P b) Ne c) Br

19 For each of the first five principal energy levels, a. list the types of sublevels (s, p, d, f, etc.) found within each level, b. the number of orbitals within each sublevel, and c. the number of electrons that can occupy each sublevel. d. Show that the sum of the number of electrons in all the sublevels within one principal energy level (n) equals 2n 2. s p d f

20 (Ch5)p.166 # 5,6,11-13,16,17,21,22-26, 36-38,41,45 (in class) 5. What determines the length of each period in the periodic table? 5. The length of a period is determined by the total number of electrons that can fill the outer sublevels of the elements of that period.

21 6. What is the relationship between the electron configuration of an element and the period in which that element appears in the periodic table? 6. An element s period corresponds to its highest occupied main energy level.

22 11. What name is sometimes used to refer to the entire set of d-block elements? the transition elements/metals

23 12. a. What types of elements make up the p block? a) The p-block consists of non-metals at the right, metalloids in the middle, and metals at the left. b. How do the properties of the p-block metals compare with those of the metals in the s and d blocks? b) The p-block metals are generally harder and more dense than the s-block metals, but softer and less dense than the d-block metals

24 13. a. Which elements are designated as the halogens? a) The halogens are the Group 17 elements. b. List three of their characteristic properties. b) The halogens are the most reactive non-metals; they react vigorously with most metals to form salts. They are also among the most electronegative elements.

25 16. a. What are the main-group elements?. a) The main-group elements are those in the s and p blocks (plus hydrogen and helium). b. What trends can be observed across the various periods within the main-group elements? b) Across the periods in the main-group elements: atomic size decreases, ionization energy increases, electron affinity increases, electronegativity increases, cation size decreases, anion size decreases

26 17. Write the noble-gas notation for the electron configuration of each of the following elements, and indicate the period in which each belongs. a. Li b. O c. Cu d. Br e. Sn

27 22. a. What is meant by atomic radius? b. What trend is observed among the atomic radii of main-group elements across a period? c. Explain this trend. 22. a) one-half the distance between the nuclei of two bonded identical atoms b) They decrease. c) As electrons are added to s and p sublevels in the same main energy level, the increasing positive charge of the nucleus pulls electrons closer to the nucleus, resulting in decreasing atomic radii.

28 25. a. How do the first ionization energies of main group elements vary across a period and down a group? b. Explain the basis for each trend. 25. a) They increase across a period and decrease down a group. b) Across a period, the increasing nuclear charge attracts electrons in the same energy level more strongly and makes them more difficult to remove. Down a group, the electrons to be removed from each successive element are farther from the nucleus in increasingly higher energy levels and are thus more easily removed.

29 45. For each element listed below, determine the charge of the ion that is most likely to be formed and the identity of the noble gas whose electron configuration is thus achieved. a. Li e. Mg i. Br b. Rb f. Al j. Ba c. O g. P d. F h. S

30 48. Identify which trends in this diagram below describe atomic radius, ionization energy, electron affinity, and electronegativity.

31 48. Identify which trends in this diagram below describe atomic radius, ionization energy, electron affinity, and electronegativity.

32 48. Identify which trends in this diagram below describe atomic radius, ionization energy, electron affinity, and electronegativity.

33 Chem I Honors: HW 5.3p. 164 #1 a. atomic radii a. generally decreases across periods and increases down groups b. first ionization energy generally increases across periods and decreases down groups c. electron affinity generally increases across periods among Groups and decreases down groups, but many exceptions are observed d. ionic radii generally decreases across periods and increases down groups e. electronegativity gradually increases across periods and either decreases or remains about the same down groups

34 Chem I Honors WH 5.3 p 167 # a. A ca+ion is a positive ion, and an anion is a negative ion. b. Ca+ions are always smaller than the atoms from which they are formed; anions are always larger. 28. a. Valence electrons are atomic electrons available to be lost, gained, or shared in the formation of chemical compounds. b. Valence electrons are located in an atom s outermost energy level. 29. a. lost, 1 b. lost, 2 c. lost, 3 d. gained, 2 e. gained, 1 f. neither lost nor gained, 0

35 Chem I Honors WH 5.3 p 167 # a. Electronegativity is the ability of an atom in a chemical compound to attract electrons from other atoms. Linus Pauling initiated concept b. Fluorine is the most electronegative element and is arbitrarily assigned an electronegativity of 4.0(Pauling scale). The values for all other elements are assigned in relation to this value. 31. Group 17, Group 1