2. Chemistry 140 HCC/TCHS Charles Lee-Instructor

3. 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. LEARNING OBJECTIVES/ASSESSMENT

4. 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. Locate elements in the periodic table on the basis of group and period designations.

5. 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. Locate elements in the periodic table on the basis of group and period designations.

6. 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.

7. 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.

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

9. 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.

10. 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.

11. 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.

12. 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.

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

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

15. 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).

16. 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.

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

18. 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.

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

20. 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.

21. 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.

22. 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.

23. Shapes and Orientations of Orbitals

24. Periodic table arrangement 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).

25. Periodic table and quantum theory Note that electron configurations are true whether we are speaking of an atom or ion: 1s 2 2s 2 2p 6 describes both Ne and Na + Q – based the shorthand electron configurations for Br –, Sn, Sn 2+, Pb? A – [Ar]4s 2 3d 10 4p 6, [Kr]5s 2 4d 10 5p 2, [Kr]5s 2 4d 10, [Xe]6s 2 4f 14 5d 10 6p 2 or [Xe] 4f 14 5d 10 6s 2 6p 2

26. Unusual electron configurations Look at your value for Cu ([Ar]4s 2 3d 9 ). The actual value for Cu is [Ar]4s 1 3d 10 … 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.

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

28. Heisenberg’s uncertainty principle 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.

29. Electron clouds 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”. 90% electron probability/cloud for 1s orbital (notice higher probability toward the centre)

30. Summary: p orbitals and d orbitals p orbitals look like a dumbell with 3 orientations: p x, p y, p z (“p sub z”). 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.

31. 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?

32. 4s 3s 2s 1s 2p 3p 3d ENERGYENERGYENERGYENERGY n l mlml msms 10(s) 2 1(p) 0 0 -1,10, 30(s) 1(p) 0 -1,10, 2(d) -1,1,0, -2, 2 40(s)0 Movie: periodic table of the elements: t10-20

33. Testing concepts Q- How many shells are shown in Fig 6.24 ‘3s’ A- Just one (the 3s). In an atom containing a 3s subshell both of the other s orbitals would also be present (superimposed on 3s). Q- Which orbitals do not contain nodes? A- Just the 1s subshell/orbital. Q- Explain with reference to quantum numbers why it makes sense that a p subshell has the different orientations it does. A- For p (l=1), m l can be -1,0,1. These three orbitals correspond to the three possible orientations of p. Recall that ml = orientation.

34. Testing concepts Q- Why does s have only one orientation? A- Because it’s spherical (or because it has only one value for m l ). Q- How far do probabilities extend from the nucleus (for 1s for example)? A-Theoretically, infinitely. Orbital shapes show where the electron will be 90% of the time. Q- Why do we represent the electron’s position as a probability? A- Heisenberg’s uncertainty principle shows we cannot know both position and velocity.