Chapter 5 Electrons in Atoms

1. Chapter 5Electrons in Atoms Different colors of light are associated with the movement of electrons

2. 5.1 Models of the Atom • Key Concepts • What was inadequate about Rutherford’s atomic model? • What was the new assumption in the Bohr model of the atom? • What does the quantum mechanical model determine about the electrons in an atom? • How do sublevels of principal energy levels differ?

3. Rutherford (1897) • Electrons move around a nucleus made of protons and neutrons • Could not explain the _______________ properties of elements • Ex: why do metals change colors when heated? chemical

4. Bohr (1913) proposed that electrons are found in specific paths, or _________, around a nucleus • Each orbit has a __________ energy level • Ex: rungs of a ladder The higher an electron is on the ladder, the farther it is away from the nucleus orbitals fixed

5. Quantum  the amount of energy needed to move an electron from one energy level to another • Moving up energy levels = ____________ energy • Moving down energy levels = _______ energy (light) • Energy levels are not equally spaced • It takes less energy to move between higher rungs or energy levels than lower ones absorb emit

6. Bohr’s model only worked for hydrogen • De Broglie (1923) proposed that small particles, like electrons, have _____________ properties • Schrodinger (1926) devised a mathematical calculation to describe the wave-like movement of electrons • _________________________ • Determines the allowed energies an electron can have and how _________ it is to find the electron at various locations around the nucleus wave-like Quantum Mechanical Model likely

7. Electron may be found anywhere within the electron cloud Movement of electron is similar to movement of a propeller blade Blade may be anywhere in the blurry region

8. principle quantum numbers • The energy levels in the quantum mechanical model are labeled by _____________________________ (n) • n = 1, 2, 3, 4, 5, 6… • Each principle energy level has sublevels that correspond to different cloud shapes, called _______________________ Name: Principle Quantum Number Symbol: n Represents: energy levels Values: any integer ≥ 1 atomic orbitals

9. spherical 4 Different Atomic Orbitals • S orbital = _____________ (1) • P orbitals = __________________ (3) 3. D orbitals = ____________________ (5) 4. F orbitals = ___________________ (7) dumbbells four leaf clovers Complex shapes

10. Orbitals get bigger because they can hold more electrons

11. Notice that n = # of sublevels and # of different shapes • The order of the shapes moves from lowest energy to highest energy (s, p, d, f…) • Example: n = 2; sublevel shapes = s & p • Name the shape(s) be for the following: • n = 1 • n = 3 • n = 4 s s, p, d s, p , d, f

12. Notice the number of orbitals per shape • s = 1 - d = 5 • p = 3 - f = 7 • If n = 3, then we will have 3 different shapes (s, p & d) • The s shape has 1 orbital available • The p shape has 3 orbitals available • And, the d shape has 5 orbitals available • How many orbitals does the 3rd energy level have? It has 9 orbitals

13. What is the relationship between n = 3 and 9 available orbitals? • 9 = 32 • So, we can say that n2 = # of orbitals available • We represent each orbital by listing the energy level followed by the orbital shape 1s n = 1, orbital shape = sphere 3d n = 3, orbital shape = clover

14. How many orbitals are in the 3s sublevel? • Well, we know that n = 3 and that the orbital shape = s • The 3rd energy level can have 9 total orbitals between s, p, & d…but we are only looking the s orbital  There is only 1 orbital in the 3s sublevel • Each orbital available can hold 2 electrons • s = 1 orbital = 2 electrons • p = 3 orbitals = 6 electrons • d = 5 orbitals = 10 electrons • f = 7 orbitals = 14 electrons

15. So, when n = 4, we can have… • s = 1 orbital = 2 electrons • p = 3 orbitals = 6 electrons • d = 5 orbitals = 10 electrons • f = 7 orbitals = 14 electrons 32 total electrons = 2n2

16. # of electrons at each main level = 2n2 n = 1 _____ electrons n = 2 _____ electrons n = 3 _____ electrons n = 4 _____ electrons 2 8 18 32

17. Key Answer #1 • What was inadequate about Rutherford’s atomic model? It did not explain the chemical properties of elements

18. Key Answer #2 • What was the new assumption in the Bohr model of the atom? An electron is found only in specific orbits

19. Key Answer #3 • What does the quantum mechanical model determine about the electrons in an atom? The likelihood or probability of finding electrons in various locations around the nucleus

20. Key Answer #4 • How do sublevels of principal energy levels differ? The sublevels are the different shapes of the electron clouds

21. 5.2 Electron Arrangement in Atoms • Key Concepts • What are the three rules for writing the electron configurations of the elements? • Why do actual electron configurations for some elements differ from those assigned using the aufbau principle?

22. Unstable arrangements become more stable by ____________ energy • Does the rock formation in the picture look stable? _______ • What will it eventually do to become more stable? losing No It will fall or rearrange

23. In an atom, _________ and the ___________ interact to make the most stable arrangement possible • The ways in which electrons are arranged around the nuclei of atoms are called __________ ____________________ electrons nucleus electron configurations

24. Three rules tell you how to find the electron configurations of atoms 1. ______________________ 2. ______________________ 3. ______________________ Aufbau Principle Pauli Exclusion Principle Hund’s Rule

25. lowest • Aufbau Principle – electrons occupy the orbitals of __________ energy first Each box represents an orbital Will electrons occupy the 2s orbital or the 2p orbital first? __________ 2s

26. What do the three boxes by 2p represent? • Which is higher, 3d or 4s? The three dumbbell orbitals = px, py, pz 3d

27. hold • Pauli Exclusion Principle – an atomic orbital may ________ at most two electrons that have ____________ spins opposite Counterclockwise Clockwise = ________________________ = ___________________

28. same • Hund’s Rule – electrons occupy orbitals of the same energy in a way that makes the number of electrons with the ______ spin direction as large as possible • Orbitals in the same sublevel have ________ energy levels • They all must have electrons with the same spin first • Practice – Put three electrons into the 2p sublevel equal 2p x y z

29. Practice – put four electrons into the 2p sublevel • Put five electrons into the 2p sublevel 2p x y z 2p x y z

30. unpaired • If an orbital only has one electron in its box, it is called ______________ Ex: How many unpaired electrons are in the following sublevel? 4 3d xy xz yz x2-y2 z2

31. We use all three rules to write a shorthand method for showing the electron configurations of an atom • It consists of… • The _____________ of the main energy levels • The ____________ of the atomic orbital • _____________ indicating the # of electrons Number (n) letter Superscript

32. 1 • Practice: Hydrogen • How many electrons does hydrogen have? ___ 1s1 _______________

33. 2 • Practice: Helium • How many electrons does helium have? ____ 1s2 _______________

34. 3 • Practice: Lithium • How many electrons does lithium have? ____ 1s22s1

35. 6 • Practice: Carbon • How many electrons does carbon have? ____ 1s22s22p2

36. 16 • Practice: Sulfur • How many electrons does sulfur have? ____ 1s22s22p63s23p4

37. 22 • Practice: Titanium • How many electrons does titanium have? ____ 1s22s22p63s23p64s23d2

38. LETS COLOR!!!

39. --- s block

40. --- s block --- p block

41. --- s block --- p block --- d block

42. Exceptional Electron Configurations • Some atoms do not follow the aufbau principle Ex: Copper – 29 electrons _________________ Actually is __________________ Ex: Chromium – 24 electrons_______________ Actually is ________________ 1s22s22p63s23p64s23d9 1s22s22p63s23p64s13d10 1s22s22p63s23p64s23d4 1s22s22p63s23p64s13d5

43. These atoms are trying to become more ______________ by having sublevels that are least ______ full. • Filled and half-filled energy sublevels are more stable than partially-filled energy sublevels. stable half

44. Key Answer #1 • What are the three rules for writing the electron configuration of elements? Aufbau Principle Pauli Exclusion Principle Hund’s Rule

45. Key Answer #2 • Why do the actual electron configurations for some elements differ from those assigned using the aufbau principle? They are trying to become more stable

46. What is the basis for exceptions to the aufbau diagram? Filled and half-filled orbitals are more stable than partially filled orbitals

47. 5.3 Physics and the Quantum Mechanical Model • Key Concepts • How does quantum mechanics differ from classical mechanics? • How are the wavelengths and frequency of light related? • What causes atomic emission spectra?

48. light waves • The quantum mechanical model grew out of the study of ___________, which also moves in ___________ • If all objects have a wavelike motion, why can’t we see this for ordinary objects like baseballs or trains? • Classical Mechanics – describes the motion of __________ objects • Quantum Mechanics – describes the motion of ________________ particles and __________ Mass has to be very small in order to detect wavelength large subatomic atoms

49. Complete Wave Cycle • Starts at _________ • Moves to its ___________ point and __________ point • Returns to ___________ zero highest lowest zero

50. zero height crest • Amplitude • a wave’s _________ from ______ to the __________ • Wavelength (λ) • Distances between the _____________ crests