Chapter 3: Electronic Structure and the Periodic Law

1. Chapter 3: Electronic Structure and the Periodic Law Chemistry 140 HCC/TCHS Charles Lee-Instructor

2. LEARNING OBJECTIVES/ASSESSMENT • When you have completed your study of this chapter, you should be able to: • 1. Locate elements in the periodic table on the basis of group and period designations. • 2. Determine the number of electrons in designated atomic orbitals, subshells, or shells. • 3. Determine the number of valence shell electrons and the electronic structure for atoms, and relate this information to the location of elements in the periodic table. • 4. Determine the following for elements: the electronic configuration of atoms, the number of unpaired electrons in atoms, and the identity of atoms based on provided electronic configurations. • 5. Determine the shell and subshell locations of the distinguishing electrons in elements, and based on their location in the periodic table, classify elements into the categories (representative element, transition element, inner‐transition element, noble gas) and (metal, metalloid, nonmetal). • 6. Recognize property trends of elements within the periodic table, and use the trends to predict selected properties of the elements.

3. Locate elements in the periodic table on the basis of group and period designations. • Periodic means “repeated in a pattern.” • In the late 1800s, Dmitri Mendeleev, a Russian chemist, searched for a way to organize the elements. • When he arranged all the elements known at that time in order of increasing atomic masses, he discovered a pattern.

4. Locate elements in the periodic table on the basis of group and period designations. • Because the pattern repeated, it was considered to be periodic. Today, this arrangement is called a periodic table of elements. • In the periodic table, the elements are arranged by increasing atomic number and by changes in physical and chemical properties.

5. Mendeleev had to leave blank spaces in his periodic table to keep the elements properly lined up according to their chemical properties. • He looked at the properties and atomic masses of the elements surrounding these blank spaces.

6. Mendeleev’s Predictins • From this information, he was able to predict the properties and the mass numbers of new elements that had not yet been discovered.

7. Mendeleev’s Predictions • This table shows Mendeleev’s predicted properties for germanium, which he called ekasilicon. His predictions proved to be accurate.

8. Improving the Periodic Table • On Mendeleev’s table, the atomic mass gradually increased from left to right. If you look at the modern periodic table, you will see several examples, such as cobalt and nickel, where the mass decreases from left to right.

9. Improving the Periodic Table • In 1913, the work of Henry G.J. Moseley, a young English scientist, led to the arrangement of elements based on their increasing atomic numbers instead of an arrangement based on atomic masses. • The current periodic table uses Moseley’s arrangement of the elements.

10. The Atom and the Periodic Table • The vertical columns in the periodic table are called groups, or families, and are numbered 1 through 18. • Elements in each group have similar properties.

11. Electron Cloud Structure • In a neutral atom, the number of electrons is equal to the number of protons. • Therefore, a carbon atom, with an atomic number of six, has six protons and six electrons.

12. Electron Cloud Structure • Scientists have found that electrons within the electron cloud have different amounts of energy.

13. Electron Cloud Structure • Scientists model the energy differences of the electrons by placing the electrons in energy levels.

14. Electron Cloud Structure • Energy levels nearer the nucleus have lower energy than those levels that are farther away. • Electrons fill these energy levels from the inner levels (closer to the nucleus) to the outer levels (farther from the nucleus).

15. Electron Cloud Structure • Elements that are in the same group have the same number of electrons in their outer energy level. • It is the number of electrons in the outer energy level that determines the chemical properties of the element.

16. Energy Levels • The maximum number of electrons that can be contained in each of the first four levels is shown.

17. Energy Levels • For example, energy level one can contain a maximum of two electrons. • A complete and stable outer energy level will contain eight electrons.

18. Rows on the Table • Remember that the atomic number found on the periodic table is equal to the number of electrons in an atom.

19. Rows on the Table • The first row has hydrogen with one electron and helium with two electrons both in energy level one. • Energy level one can hold only two electrons. Therefore, helium has a full or complete outer energy level.

20. Rows on the Table • The second row begins with lithium, which has three electrons—two in energy level one and one in energy level two. • Lithium is followed by beryllium with two outer electrons, boron with three, and so on until you reach neon with eight outer electrons.

21. Rows on the Table • Do you notice how the row in the periodic table ends when an outer level is filled? • In the third row of elements, the electrons begin filling energy level three. • The row ends with argon, which has a full outer energy level of eight electrons.

22. Electromagnetic radiation – energy that travels through space in the form of a wave Photon – a unit, quanta, of electromagnetic radiation Frequency – the number of waves that pass a point in one second What is the basis of electron theory?

23. Order of high frequency to low frequency, short wavelength to long wavelength Cosmic rays Gamma rays X rays Ultraviolet Visible light Infrared light Microwaves Radio waves Electrical power What property do the above waves have in common? Examples of Electromagnetic Radiation

24. Radiation from excited atoms is analyzed and the frequency is measured. The equation, E = hf is used to determine the energy of the outer shell electrons. This information is gathered for all elements and a model of electron arrangement is developed. The Answer!

25. Periodic Table - Encarta

26. 1. Principal quantum number - indicates shell or energy level - 1,2,3,... K,L,M,... 2. Suborbital quantum number - s,p,d,f,g,... NUMBER SHAPE ORBITALS/SHELL MAXIMUM # e- s sphere 1 2 p figure 8 3 6 d “ 5 10 f “ 7 14 Quantum Numbersnumbers used to describe electrons

27. 3. Magnetic quantum numbers - indicate state of magnetic fields around the electron 4. Spin quantum number - indicates direction of spin of the electron on its axis electron pair - two electrons occupying the same space orbital spinning in opposite directions Only two electrons can occupy the same path. Quantum Numbers (continued)

28. Shapes and Orientations of Orbitals

29. the quantum theory helps to explain the structure of the periodic table. n - 1 indicates that the d subshell in period 4 actually starts at 3 (4 - 1 = 3). Periodic table arrangement

30. Note that electron configurations are true whether we are speaking of an atom or ion: 1s22s22p6 describes both Ne and Na+ Q – based the shorthand electron configurations for Br–, Sn, Sn2+, Pb? A – [Ar]4s23d104p6, [Kr]5s24d105p2, [Kr]5s24d10, [Xe]6s24f145d106p2 or [Xe] 4f145d106s26p2 Periodic table and quantum theory

31. Look at your value for Cu ([Ar]4s23d9). The actual value for Cu is [Ar]4s13d10… why? The explanation is that there is some sort of added stability provided by a filled (or half-filled subshell). The only exceptions that you need to remember are Cr, Cu, Ag, and Au. The inner transition elements also do not follow expected patterns. Unusual electron configurations

32. Electrons are difficult to visualize. As a simplification we will picture them as tiny wave/particles around a nucleus. Heisenberg’s uncertainty principle The location of electrons is described by: n, l, ml n = size, l = shape, ml = orientation • Heisenberg showed it is impossible to know both the position and velocity of an electron. • Think of measuring speed & position for a car. Slow Fast

33. The distance between 2+ returning signals gives information on position and velocity. A car is massive. The energy from the radar waves will not affect its path. However, because electrons are so small, anything that hits them will alter their course. The first wave will knock the electron out of its normal path. Thus, we cannot know both position and velocity because we cannot get 2 accurate signals to return. Heisenberg’s uncertainty principle

34. Although we cannot know how the electron travels around the nucleus we can know where it spends the majority of its time (thus, we can know position but not trajectory). The “probability” of finding an electron around a nucleus can be calculated. Relative probability is indicated by a series of dots, indicating the “electron cloud”. Electron clouds • 90% electron probability/cloud for 1s orbital (notice higher probability toward the centre)

35. p orbitals look like a dumbell with 3 orientations: px, py, pz(“p sub z”). Summary: p orbitals and d orbitals Four of the d orbitals resemble two dumbells in a clover shape. The last d orbital resembles a p orbital with a donut wrapped around the middle.

36. Each subshell (1s, 3p, 2d, 5f, 1g, etc.) has a specific shape derived from mathematics. As we move to higher energy level, the shapes get stranger You need to know 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s Q -How many shells are shown in 3s? Q- Explain why a p sub-shell has the different orientations it does (refer to quantum numbers). Q- Why does s have only one orientation? Q- How far do the probabilities extend from the nucleus (for 1s for example)? Q- Why do we represent the electron’s position as a probability?

37. n l ml ms 3d 4s 3p 3s 2p 2s 1s 1 0(s) 0 ENERGY 2 0(s) 0 1(p) -1, 0, 1 3 0(s) 0 1(p) -1, 0, 1 2(d) -1, 0, 1, 2 -2, 4 0(s) 0 Movie: periodic table of the elements: t10-20

38. Configuration notation is expressed by the principal quantum number written first, the suborbital quantum number second, and the number of electrons in the suborbital written as a power. Example: N - 1s22s22p3 In Class Assignment Show the electron configuration for the elements H - Kr. Configuration Notation

39. 1. A short line, ____, will represent an electron path or suborbital. 2. An arrow pointing up,  will represent an electron spinning in a particular direction. 3. An arrow pointing down,  represents an electron spinning in the opposite direction. 4. The principal quantum numbers, 1,2,3..., will represent the shells or energy levels. Orbital Diagrams

40. symbol 1s 2s 2p 3s 3p He-4 __ __ __ __ __ __ __ __ __ N-14 __ __ __ __ __ __ __ __ __ Mg-24 __ __ __ __ __ __ __ __ __ S-32 __ __ __ __ __ __ __ __ __ Ar-40 __ __ __ __ __ __ __ __ __ Format for Orbital Diagrams

41. The following order must be followed when filling out the orbital notation for the elements. 1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p, 7s,5f, 6d,7p Notice, 4s is filled before 3d because 4s electrons have lower energy than 3d electrons, and the same is true for the rest. Order for Filling Space Orbitals