1. Writing Electron Configurations

2. Ok...let’s simplify this. Every atom has a nucleus. In that nucleus we have protons (positive charge) and neutrons (no charge) Surrounding the nucleus we have electrons. Electrons have a negative charge, however they do not “fall” into the nucleus (despite being attracted to the positive proton), why?

3. Ok...let’s simplify this Because electrons exist in “orbitals” An orbital is the region of space in which you will find an electron. (For our purposes) This is where it gets confusing... There are different levels of orbitals. We represent these different levels of orbitals with the letter “n”.

4. So if an electron is in the 1 st orbital surrounding the nucleus it is n=1. If an electron is in the 3 rd orbital surrounding the nucleus it is n=3. Easy enough, right?

5. There’s more. We also have the letter “l” this represents what type of orbital the electron is in. The different types of orbitals are as follows... S, p, d, f, and g These letters refer to the shape of the orbital.

6. So now, we have two pieces of information to help us determine where an electron is. We know the number of orbital it is in and we know the shape of the orbital.

7. Review: Quantum Numbers Principle Quantum Number (n) represents the energy orbital number (how far the orbital is from the nucleus). Orbital-Shape Quantum Number Orbital ( l) represents the shape of the orbital sublevel (s is round, p is like 2 balloons, 2 d orbitals put together) Magnetic Quantum Number represents the direction of the orbital sublevel ( l )

8. n can be any number from 1 on. l can be any number from 0 to (n-1) m l is –l to +l

9. Ok...back to new stuff The Fourth Quantum Number (ms) is used to represent the direction of the spin on the electron. There are only two possible values for this number, + ½ or – ½. Electrons spin like a top.

10. Electron Spin

11. Pauli Exclusion Principle We have a new principle to learn! The Pauli Exclusion Principle states that no two electrons in any atom will have the same 4 quantum numbers. For example you would not find two electrons with n=4, l=3, ml=-3 and ms= - ½ You could find n=4, l=3, ml=-3, ms= - ½ and n=4, l=3, ml= -3, ms= + ½

12. Example

13. Electron Configurations

14. An electron configuration can be written for each atom. That is, we can write out where we can find each electron in an atom. And remember, no two electrons will have the same set of 4 quantum numbers in the same atom.

15. Generally when we write electron configurations we write the configuration of the atom in it’s ground state. Ground state means that the electron is neutral and is not an isotope.

16. Example Hydrogen, in it’s ground state, has one electron (atomic number 1, mass of 1). The electron configuration for Hydrogen is 1s 1. Remember, n=1, l=0 and the small, raised 1 symbolizes that there is one electron in the s sublevel.

17. Helium is He 1s 2 symbolizing there are 2 electrons in the s sublevel. Lithium is 1s 2 2s 1 symbolizing that there are 2 electrons in the 1 s sublevel and 1 electron in the 2 s level. Fluorine is written as 1s 2 2s 2 2p 5 Why do you think that there are only 2 electrons in the s sublevel, but there can be 5 electrons in the p sublevel?

18. Remember the Pauli Exclusion Principle? What did the Pauli Exclusion Principle tell us? That no two electrons can have the same 4 quantum numbers in an atom. Let’s write out the possible values of the 4 quantum numbers in the s sublevel. How many electrons do you think is possible in the p sublevel? How about the d sublevel?

19. There is a maximum of 2 electrons in the s sublevel. There is a maximum of 6 electrons in the p sublevel. There is a maximum of 10 electrons in the d sublevel.

20. Aufbau Principle Each electron occupies the lowest energy orbital available Atom’s are “built-up”

22. Do you think that the direction of the arrow symbolizes anything? Why do you think that in the Carbon electron configuration the 2 p level has two electrons in the same direction?

23. We can tell the electron configuration of an element just by looking at the periodic table! The “long form periodic table” shows us which sublevel contains the atom’s valence electrons.

25. A helpful guide to what goes where

26. Let’s try it. How would we write the configuration for: – Hydrogen – Helium – Boron – Carbon – Nitrogen

28. Noble Gas Notation The last column of the periodic table is called the Noble Gases. The Noble Gases have full valence electron shells. So we when we out electron configurations for our elements we can shorten it like this...

29. Shorten Electron Configuration Notation Ne – 1s 2 2s 2 2p 6 Na – 1s 2 2s 2 2p 6 3s 1 – [Ne]3s 1

30. Exceptions! There are almost always exceptions to the rules! In this case, Chromium and Copper do not fill their energy shells as we would expect them to. How so and why? Evidence shows that these two elements achieve a ground state in ways that do not follow the rules.

32. Hidden Information Recall that in the periodic table a group runs vertically (up and down) and a period rums horizontally (left to right). Recall that the last group number is also the number of valence electrons in that atom.

33. Hidden Information The n value of the highest energy level shell (where the valence electrons are) is also the period number where you find the element.

34. s-block The s-block includes elements in group 1 and 2 (and hydrogen and helium) In the s-block all elements have their valence electrons in an s block (be it 3s, 4s, etc.)

35. p-block The p-block includes elements in groups 13-18 The elements in this block have their valence electrons in the p level.

36. d-block The d-block includes all of the transition elements (groups 3-12) These elements have valence electrons in the d shell

37. f-block We (in our class) have not looked at any electron configurations in the f-block, but we know that it exists. The f-block consists of the “inner transition” metals.

39. Recall We said that the last number of the group number (i.e. 15....5) is the number of valence electrons in an atom. We also said that the last (highest) n value is the period number. So, using the periodic table what element is represented by [He]2s 2 ?

40. How about this one? [Kr]5s 2 4d 10 5p 5 ?

41. Questions What would the valence shell be for Arsenic (As)? 4p Let’s see: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 OR [Ar] 4s 2 3d 10 4p 3

42. How about Cesium (Cs)? 6s Let’s see: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 6 6s 1 OR [Xe] what? How can we determine what comes next? Xe valence shell is 5p How many more electrons does Cs have more than Xe? What sell do we fill after 5p? [Xe]6s 1

43. Atomic Radius Atomic radius is a measurable property of an atom. What is radius?

45. Atomic Radii decreases across a period (left to right) Increases down a group (up and down) There are two factors that affect atomic radii

46. When n increases, that means we are moving farther from the nucleus and thus giving an larger area of finding an electron and thus a larger atomic radii. So it makes sense that as you increase n (remember n corresponds to the period number) that the atomic radii would increase down a group.

47. Another factor that affects atomic radii is called Z eff, this stands for the attraction between the nucleus and the electrons. The further away from the nucleus a valence shell is, the more it will be shielded from the positive attraction of the nucleus.

48. HOWEVER, as you travel across a period, the n value stays the same (same distance away) but each time an electron is added and the attraction to the nucleus gets stronger.

50. Ionization Energy Ionization Energy is the energy needed to completely remove an electron from an atom. Hint: an ion is an atom with a charge, meaning it has lost an electron or has gained an electron. Remember that ionization energy has to do with removing an electron, thus creating an ion.

51. The trends of ionization energy are the exact opposite as atomic radii. Ionization energy decreases down a group and increases across a period.

52. Trends in Ionization Energy and Chemical Reactivity Summarize the table on page 155 and create a chart in your notebook.

53. Electron Affinity Change in energy that occurs when an electron is added to a gaseous atom. What does gaseous atom mean? Read and summarize the rules of electron affinity on page 156.

54. Homework Ionization Energy and chemical reactivity chart Electron affinity summary Section summary (pg. 157) questions #1, 3 and 4.