1. 12.1 Electron Configuration HL

2. Starter Element Bingo

3. References Moodle Powerpoints Workbook

4. Assessment Objective 12.1.3 State the Relative energies of the s,p,d and f orbitals in a single energy level. 12.1.4 state the maximum number of orbitals in a given energy level. 12.1.5 Draw the shape of an s – orbital and the shapes of the p x. P y and p z orbitals. 12.1.6 Apply the Aufbau principle to electron configurations, Hind Rule and the Pauli exclusion Principle to write the electron configuration for atoms up to Z =54

5. Why care about where electrons are located? Allows us to understand how atoms bond together to form chemicals and compounds. These lessons will be very helpful when we look at bonding…………………

6. Introduction Electrons are found in orbitals Orbital: A region in an atom where there is a high probability of finding electrons. (Electrons clouds) Orbital’s are filled from low to high energy or starting closest to the nucleus ORBITS ARE NOT LIKE A PLANETS!!!!!

7. Introduction The quantum mechanical model is the current model of the atom that describes the probability of finding an electron in a given region of space around the nucleus. It is called the electron cloud model.

8. Quantum numbers are the 4 numbers that specify the properties of atomic orbitals and the electrons that reside in them. 1. The principle quantum number (shell): electrons occupy the specific energy levels. 2. The angular momentum quantum number (orbital shape): specifies the shape of the orbital. 3. The magnetic quantum number (orbital orientation): specifies how this shape is arranged in three dimensions around the nucleus. 4. The spin quantum numbers: specifies in which direction the electrons are spinning.

9. 1. The Principle Quantum Number Electrons can occupy only specific energy levels. These levels are numbered 1 – 7. The higher energy levels indicate a higher energy state for that electron and a location further away from the nucleus. In other words, electrons near the nucleus are low energy electrons. Increasing Energy

10. 1. The Principle Quantum Number The energy level indicates the size of the electron cloud. The formula 2n 2 indicates the total possible electrons in an energy level.

11. 1. The Principle Quantum Number Example: Calculate the total possible electrons in the 1 st through 4 th energy levels. Remember: 2n 2 Energy Level # of Electrons 1 2 3 4

12. 2. Orbital Shape An energy level is actually made of many energy states called orbitals (subshells or sublevels).

13. An orbital is a three dimensional region around the nucleus that indicates the probable location of one pair of electrons. These orbitals are s, p, d and f. The second quantum number indicates the shape of the orbital. s orbital p orbital d orbitals

14. Assessment Objective 12.1.3 State the Relative energies of the s,p,d and f orbitals in a single energy level.

15. As move up each level and sub level the relative energy increases

16. 7 Different Main Energy Levels Increasing energy Sub levels in each level e.g level 2 has s and p sub levels Increasing energy

17. Assessment Objective 12.1.5 Draw the shape of an s – orbital and the shapes of the p x. P y and p z orbitals.

18. The s-orbital is sphere shaped. Need to Draw these!

19. The p-orbital is dumb bell shaped Need to Draw these!

20. d-orbitals look like leaves of clover

21. f-orbitals look like flowers

22. 3 Orientation The magnetic number indicates the orientation of an orbital around the nucleus. Since an ‘s’ orbital is spherical, it can have only 1 orientation.

23. 3 Possible Orientations for p-Orbitals

24. d-Orbitals have 5 Possible Orientations

25. f-Orbitals have 7 Possible Orientations

26. Orbital Orientation Orbital Type MAX # Electrons per Orbital Possible Orientations for Each Type of Orbital Max # of Electrons if Each Orbital Orientation is Filled s212 p236 d2510 f2714

27. 4 Spin The spin quantum number indicates that the 2 electrons occupying a single orbital must have opposite spin.

28. Assessment Objective 12.1.4 state the maximum number of orbitals in a given energy level.

29. Principle Quantum # Main Energy Level (n) Types of Orbitals in Main Energy Level (n) # of Orientations per Orbital Type # of Orbitals per Main Energy Level (n 2 ) Max # Electrons In Filled Orbitals Max # of Electrons per Main Energy Level (2n 2 ) 1 s 1 1 2 2 2 s p 1 3 4 2 6 8 3 s p d 1 3 5 9 2 6 10 18 4 s p d f 1 3 5 7 162 6 10 14 32

30. Principle Quantum Number: Main Energy Level (n) Types of Orbitals in Main Energy Level (n) # Of Orientations per Orbital Type # of Orbitals Per Main Energy Level Max # of Electron per Filled Orbital Max # of electrons per Main Energy Level 5 spdfspdf 13571357 (n - 1) 2 16 2 6 10 14 2(n -1) 2 32 6 spdspd 135135 (n - 3) 2 9 2 6 10 2(n – 3) 2 18 7 spsp 1313 (n – 5) 2 4 2626 2(n – 5) 2 8

31. The row numbers on the periodic table are the same as the principle quantum numbers or energy levels (n).

32. Writing Electron Configurations

33. Assessment Objective 12.1.6 Apply the Aufbau Principle to electron configurations, Hind Rule and the Pauli Exclusion Principle to write the electron configuration for atoms up to Z =54

34. The Aufbau Principle Energy levels overlap so the diagram shows the order of the sublevels. The Aufbau principle states that an electron occupies the lowest energy orbital that can receive it.

35. Writing Electron Configurations Using the atomic number as the total number of electrons you can write the electron configuration for all of the elements. The number of electrons in each sublevel is written as a exponent.

36. Remember……………. Orbitals in each level Level 1 s only, level 2 s and p, level 3 s. p and d Level 4 and 5 s, p, d and f Level 6 s, p and d Level 7 s and p Number of electrons in each type of orbital. s =2, p = 6, d = 10, f = 14 Electrons fill from lowest energy levels first (Aufbau Principle) (but energy levels overlap)

37. Example Boron (in level 2 so s and p orbitals) Number of protons = 5 Therefore number of electrons = 5 Answer B 1s 2 2s 2 2p 1

38. Phosphorous ( in level 3 so s, p and d orbitals) Number of protons = number of electrons = 15 Answer P 1s 2 2s 2 2p 6 3s 2 3p 5

39. Lead ( level 6 so has s, p, d and f ) Atomic number = 82 = 82 protons = 82 electrons Pb 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 2

40. Using the periodic table

41. Write the electron configuration for magnesium #12 gallium #31, element #35.

42. Periodic Table Orbital Filling Method

43. Noble – Gas Notation The Group VIII elements, helium, argon, krypton, xenon, and radon are called the noble gases. The configurations of the noble gases are often used as a shorthand method for writing longer electron configurations. For Example: Sodium – Na has 11 electrons Electron Configuration 1s 2 2s 2 2p 6 3s 1 Noble – Gas Configuration [Ne]3s 1

44. Example 2: Arsenic, As, 33 electrons Electron Configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 Noble – Gas Configuration [Ar]4s 2 3d 10 4p 3 Example 3: Barium, Ba, 56 electrons Example 4: Rubidium, Rb

45. Use the periodic table to write the electron configurations for the following atoms. Example 1: Nitrogen, 7 electrons Example 2: Phosphorus, 15 electrons Example 3: Cerium, 58 electrons

46. Assessment Objective 12.1.6 Apply the Aufbau principle to electron configurations, Hind Rule and the Pauli exclusion Principle to write the electron configuration for atoms up to Z =54

47. Pauli and Hind Although the following do not effect the writing of electronic configurations it is important to note how these electrons are arranged within the orbitals

48. Pauli Exclusion Principle The maximum of two electrons can occupy a single atomic orbital Electrons move in 2 direction (spin) (no two electrons in a given atom can have the same 4 quantum numbers. i.e have same energy level, shape, orientation and spin)

49. Hind Rule Single electrons with same spin must occupy equal energy orbital before additional electrons can be added in the opposite spin.

50. In practice let us consider to element oxygen N 1s 2 2s 2 2p 3 There are 3 orientations for the p orbital (x, y and z) According to Pauli, only two electrons in each orientation (so max of six electrons for p sub-orbital)

51. According to Hind these electron MUST put one electron in each orientation in the same spin direction first. If we were to expand out the electron configuration for nitrogen N 1s 2 2s 2 2p x 1p y 1p z 1 (all p orbitals in same spin direction)

52. For oxygen, we start filling up the sub orbital with electrons in the other spin direction O 1s 2 2s 2 2p 4 If we expand this out we get: O 1s 2 2s 2 2p x 2p y 1p z 1 With the 2px orbital containing electrons with opposite spins.