1. Electrons in Atoms Chapter 13

2. Connection to Chapter 5 From the atomists to Rutherford, we discussed the evolution of subatomic particles. The discussion continues here…. *Chemical properties of atoms, ions, and molecules are related to the arrangement of electrons In this chapter (and the next few) we take a look at electrons in atoms!

3. Review John Dalton: Atoms are hard, indivisible spheres. JJ Thomson: discovered negatively charged subatomic particles called electrons –Assumed a positive charge was present to cause the atom to have no overall charge –Plum Pudding Model (or cookie dough model) https://reich-chemistry.wikispaces.com/Fall.2008.MMA.Cushman.Hutchinson.Timeline

4. Thompson’s model does not explain…  How many protons and electrons?  Where are the protons and electrons located?  What happens when electrons are gained/lost?

5. Ernest Rutherford Gold foil experiment: (a review) –Atom is mostly empty space –Small, positively charged nucleus with most of the atom’s mass –Electrons “float” around nucleus

6. Neils Bohr Proposed that electrons travel in circular paths around the nucleus. Electrons in a path have a fixed energy which prevents them from falling into the nucleus. Energy level: The region around the nucleus where the electron is likely to be moving.

7. Think of energy levels like a ladder: What is the difference between the two ladders?

8. Lowest Energy Level Highest Energy Level Electrons cannot exist between energy levels. Highest energy is farthest away from the nucleus Lowest energy level is closest to the nucleus

9. The quantum To move from rung (or orbit) to the next electrons must gain or lose the right amount of energy. –This amount is called a quantum of energy.

10. Erwin Schrödinger Came up with a mathematical equation to describe the location of an electron. This modern description is called the Quantum Mechanical (QM) Model of the atom. –Does not define a path for the electron, but a probability of where you will find an electron.

11. Schrödinger Equation Darker areas where electron is likely to be found

12. The Quantum Mechanical Model of Electron Behavior in Atoms According to the quantum mechanical model of electron behavior, the precise paths of electrons moving around the nucleus cannot be determined accurately. Instead of circular orbits, the location and energy of electrons moving around the nucleus is specified using the three terms principle energy level, sublevel and orbital.

13. Principal Energy level The location of electrons is indicated by assigning a number n to the level and all electrons located in that level The value of n can be 1, 2, 3, 4, etc. Must be a whole number The higher the n value, the higher is the energy of the shell and the contained electrons.

14. Sublevels Each principal energy level is made up of one or more sublevels that are designated by a letter from the group s, p, d, or f. The number of sublevels = the energy level (n) –EXAMPLE: »Energy level 3 (n = 3) »3 sublevels: designated 3s, 3p, and 3d.

15. Atomic Orbitals These are the types of sublevels…. Each sublevel consists of one or more atomic orbitals Four types of orbitals with different shapes: –s orbitals: Spherical in shape –p orbitals: shaped like a dumbell Notice the three different orientations of the p orbitals

16. Atomic Orbitals continued… d orbitals: 5 different orientations f orbitals: 7 different orientations

17. Atomic Orbitals Summary All s sublevels consist of a single s orbital. All p sublevels consist of three p orbitals. All d sublevels consist of five d orbitals. All f sublevels consist of seven f orbitals.

18. Quantum Mechanical Model Says… A maximum of two electrons per atomic orbital 2e - 2e - 2e - Maximum total of 6 electons for p orbitals

19. How Many in the d orbitals?

20. That is a lot to remember… A Summary… The principle energy level (n) is an integer. Low energies – lower n values. # of sublevels = n –Example: How many sublevels in n=3? Ans: 3 Types of sublevels in n=3: –3s, 3p, 3d How many orbitals in n=3? –3s (1) + 3p(3) + 3d(5) = 9 –Or # of orbitals = n 2

21. Now, each orbital can hold 2 electrons. How many electrons in the n=3? –Has 9 orbitals = 9x2= 18 electrons

22. Let’s do another example: Look at n=4 (the 4 th principal energy level) # sublevels? – # sublevel = n = 4 Types of sublevels? –= 4s, 4p, 4d, 4f Number of orbitals? –4s (1), 4p (3), 4d (5), 4f (7) = 16 Number of electrons? –16x2= 32 electrons fit in the 4 th principal energy level

23. Types of problems you might encounter How many electrons fit in the 3d sublevel? –There are 5 3d orbitals, each holding 2 electrons for 10 total electrons in this sublevel.

24. 13.2 Electron Configurations Electron Configuration: the arrangement of electrons around the nucleus Specifics about the principal energy levels, sublevels, orbitals and the number of electrons

25. How electrons fill orbitals (writing electron configurations) 1. Aufbau Principle Electrons enter orbitals of lowest energy first. Look at figure 13.6. The lowest energy level is the 1s. Electrons go here first

26. How electrons fill orbitals (writing electron configurations) 2. The Pauli Exclusion Principle –Only 2 electrons per orbital –Don’t squish more than 2 electrons in a box! –Each of the electrons must have opposite spin –One up arrow and one down arrow

27. How electrons fill orbitals (writing electron configurations) 3. Hund‘s Rule When electrons occupy orbitals of equal energy (ex. 2p) one electron enters each orbital until all orbitals contain one electron with parallel (the same) spin.

28. Writing Electron Configurations for the first 10 elements. H He Li Be B C N O F Ne

30. How do I remember that figure?

31. Exceptions to the rule… Aufbau diagrams work for the most part. They don’t work for Cr and Cu. It ends up that things are more stable when the d orbitals are half filled so Cr and Cu move things around to make this happen.