1. Electrons

2. The Modern Atom Not along after Rutherford established that the positive charge in the atom resided in the nucleus. Scientists began putting together a new model of the atom using other discoveries that had previously not been explainable. The main discovery that got them rethinking the atom was the realization that the energy of light is quantized.

3. Bohr’s Model developed the planetary model of the atom the paths the electrons “traveled” were referred as orbitals energies were determined by the order from the nucleus

4. Bohr’s Model When energy is added to an atom it causes electrons to move from an orbital close to the nucleus to orbitals that are farther away. Starting Orbital = Ground State Higher Energy = Excited State

5. Quantum: of energy is the amount of energy required to move an electron from one energy level to another.of energy is the amount of energy required to move an electron from one energy level to another. the energy of the electron is quantized the energy of the electron is quantized a quantum leap is an abrupt change a quantum leap is an abrupt change The quantum of electromagnetic radiation is the photon.

6. Light is Energy Light is Energy

7. Neils Bohr Studied the Hydrogen line spectrum Energy is closely related to its color

8. What is Light? Light is how we refer to Electromagnetic RadiationLight is how we refer to Electromagnetic Radiation Electromagnetic Radiation applies to infrared light, UV light, microwaves, radio waves, X-Rays etc.Electromagnetic Radiation applies to infrared light, UV light, microwaves, radio waves, X-Rays etc. Light is a waveLight is a wave

9. Bozeman Waves Bozeman Waves

10. speed

11. Waves An Analogy: An Analogy: The waves of the ocean can transmit huge amounts of energy, even though they aren’t concrete objects. The waves of the ocean can transmit huge amounts of energy, even though they aren’t concrete objects.

12. Wavelength and Frequency λ = wavelength ν = frequency c = speed of light (3.00 e 8 m/s) c = λ ν

13. Bozeman Light Bozeman Light

14. Spectrum of White Light

15. Electromagnetic Spectrum In increasing energy, ROY G BIV

17. An excited lithium atom emitting a photon of red light to drop to a lower energy state.

18. Flame Test Flame Test

20. Energy of Light is Related to Frequency E = hν E = energy in JoulesE = energy in Joules h = Plank’s constant = 6.626 e -34 Jsh = Plank’s constant = 6.626 e -34 Js ν = frequency (in Hz or /sec)ν = frequency (in Hz or /sec) Ex: Light with a frequency of 450 GHz has an energy of 3.0 e -22 J

21. Diffraction of Electrons Diffraction of Electrons

22. Electrons as Waves Louis de Broglie (1924) Applied wave-particle theory to e - e - exhibit wave properties QUANTIZED WAVELENGTHS

23. Electrons as Waves EVIDENCE: DIFFRACTION PATTERNS ELECTRONS VISIBLE LIGHT

24. Heisenberg Uncertainty Principle It is impossible to know exactly both the velocity (speed and direction) and position of a particle at the same time. It is impossible to know exactly both the velocity (speed and direction) and position of a particle at the same time.

26. Erwin Schrodinger (1887-1961) In 1926 Schrodinger gave us our modern mathematical description of the atom: In 1926 Schrodinger gave us our modern mathematical description of the atom: The Quantum Mechanical Model The Quantum Mechanical Model

27. Quantum Mechanics Quantum Mechanics

28. Quantum Mechanical Model of the Atom Orbitals: Regions where statistically electrons are likely to be found. Orbitals: Regions where statistically electrons are likely to be found. Electrons are 3 dimensional waves Electrons are 3 dimensional waves Radial Distribution Curve Orbital

29. The Quantum Mechanical Model Electrons have specific energy levels (like Bohr predicted.) Electrons have specific energy levels (like Bohr predicted.) BUT electrons are not in specific ORBITS (unlike what Bohr predicted). BUT electrons are not in specific ORBITS (unlike what Bohr predicted).

30. The Quantum Mechanical Model Determines: Determines:  Energy level-the allowed energies an electron can have  Atomic orbital- A region of space around the nucleus in which there is a high probability of finding an electron.

31. Arrangement of Electrons in Atoms Electrons in atoms are arranged as Electrons in atoms are arranged as LEVELS (n) LEVELS (n) SUBLEVELS (l) SUBLEVELS (l) ORBITALS (m l ) ORBITALS (m l )

32. Types of Orbitals The most probable area to find these electrons takes on a shape The most probable area to find these electrons takes on a shape So far, we have 4 shapes. They are named s, p, d, and f. So far, we have 4 shapes. They are named s, p, d, and f. No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise

33. s orbital—1 spherical -2 electrons

34. Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

35. p-orbital-3 types-6 electrons

36. p Orbitals this is a p sublevel with 3 orbitals this is a p sublevel with 3 orbitals These are called x, y, and z These are called x, y, and z this is a p sublevel with 3 orbitals this is a p sublevel with 3 orbitals These are called x, y, and z These are called x, y, and z There is a thru the nucleus, which is an area of zero probability of finding an electron There is a PLANAR NODE thru the nucleus, which is an area of zero probability of finding an electron 3p y orbital

37. p Orbitals The three p orbitals lie 90 o apart in space The three p orbitals lie 90 o apart in space

38. d-orbital-5 types-10 electrons

39. d Orbitals d sublevel has 5 orbitals

40. The shapes and labels of the five 3d orbitals.

41. f-orbitals-14-electrons

42. f Orbitals For l = 3 ---> f sublevel with 7 orbitals For l = 3 ---> f sublevel with 7 orbitals

44. Orbitals and the Periodic Table grouped in s, p, d, and f orbitals grouped in s, p, d, and f orbitals s orbitals p orbitals d orbitals f orbitals

45. Energy Levels Each energy level has a number called the PRINCIPAL QUANTUM NUMBER, n Each energy level has a number called the PRINCIPAL QUANTUM NUMBER, n Currently n can be 1 thru 7, because there are 7 periods on the periodic table Currently n can be 1 thru 7, because there are 7 periods on the periodic table

46. Energy Levels n = 1 n = 2 n = 3 n = 4

47. Why are d and f orbitals always in lower energy levels? d and f orbitals require LARGE amounts of energy d and f orbitals require LARGE amounts of energy It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy

48. Principal energy level Number of sublevels Type of sublevel n = 1 1 1 s (1 orbital) n = 2 2 2s (1 orbital) 2 p (3 orbital) n = 3 3 3s (1 orbital) 3p (3 orbitals) 3d (5 orbitals) n = 4 4 4s (1 orbital) 4p (3 orbitals) 4d (5 orbitals) 4f (7 orbitals)

49. The maximum number of electrons that can occupy and energy level = 2n 2 Energy level (n) Maximum number of electrons 12 28 318 432

50. s orbitals d orbitals Number of orbitals Number of electrons p orbitals f orbitals How many electrons can be in a sublevel? Remember: A maximum of two electrons can be placed in an orbital.

51. Electron Configurations A list of all the electrons in an atom (or ion) A list of all the electrons in an atom (or ion) Must go in order (Aufbau principle) Must go in order (Aufbau principle) 2 electrons per orbital, maximum 2 electrons per orbital, maximum We need electron configurations so that we can determine the number of electrons in the outermost energy level. We need electron configurations so that we can determine the number of electrons in the outermost energy level. The number of valence electrons determines how many and what this atom (or ion) can bond to in order to make a molecule The number of valence electrons determines how many and what this atom (or ion) can bond to in order to make a molecule 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 … etc.

52. Electron Configurations 2p 4 2p 4 Energy Level Sublevel Number of electrons in the sublevel 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 … etc.

53. Three rules tell you how to find the electron configurations of atoms. The Aufbau Principle—electrons occupy the orbitals of lowest energy first. The Aufbau Principle—electrons occupy the orbitals of lowest energy first. Pauli Exclusion Principle—an atomic orbital may hold at most two electrons. Pauli Exclusion Principle—an atomic orbital may hold at most two electrons. Hund’s Rule—one electron enters orbitals of equal energy until all orbitals contain one electron with the same spin direction. Hund’s Rule—one electron enters orbitals of equal energy until all orbitals contain one electron with the same spin direction.

54. Electrons fill the lowest energy orbitals first. Electrons fill the lowest energy orbitals first. Aufbau Principle

55. Each orbital can hold TWO electrons with opposite spins. Each orbital can hold TWO electrons with opposite spins. Pauli Exclusion Principle

56. RIGHT WRONG Hund’s Rule Hund’s Rule Within a sublevel, place one e - per orbital before pairing them. Within a sublevel, place one e - per orbital before pairing them. “Empty Bus Seat Rule” “Empty Bus Seat Rule”

57. Shorthand Configuration Shorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 Notation Longhand Configuration Longhand Configuration

58. Orbital Diagrams Graphical representation of an electron configuration Graphical representation of an electron configuration One arrow represents one electron One arrow represents one electron Shows spin and which orbital within a sublevel Shows spin and which orbital within a sublevel Same rules as before (Aufbau principle, d 4 and d 9 exceptions, two electrons in each orbital, etc. etc.) Same rules as before (Aufbau principle, d 4 and d 9 exceptions, two electrons in each orbital, etc. etc.)

59. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

60. O 8e - Orbital Diagram Orbital Diagram Electron Configuration 1s 2 2s 2 2p 4 Notation of Oxygen 1s 2s 2p

61. Shorthand Notation A way of abbreviating long electron configurations A way of abbreviating long electron configurations Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration

62. Shorthand Notation Step 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ]. Step 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ]. Step 2: Find where to resume by finding the next energy level. Step 2: Find where to resume by finding the next energy level. Step 3: Resume the configuration until it’s finished. Step 3: Resume the configuration until it’s finished.

63. Shorthand Notation Chlorine Chlorine Longhand is Longhand is 1s 2 2s 2 2p 6 3s 2 3p 5 You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s 2 2s 2 2p 6 The next energy level after Neon is 3 The next energy level after Neon is 3 So you start at level 3 (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17 So you start at level 3 (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17 [Ne] 3s 2 3p 5 [Ne] 3s 2 3p 5

64. Exceptions to the Aufbau Principle d 4 is one electron short of being HALF full d 4 is one electron short of being HALF full In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d 5 instead of d 4. In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d 5 instead of d 4. For example: Cr would be [Ar] 4s 2 3d 4, but since this ends exactly with a d4 it is an exception to the rule. Thus, Cr should be [Ar] 4s 1 3d 5. For example: Cr would be [Ar] 4s 2 3d 4, but since this ends exactly with a d4 it is an exception to the rule. Thus, Cr should be [Ar] 4s 1 3d 5. Procedure: Find the closest s orbital. Steal one electron from it, and add it to the d. Procedure: Find the closest s orbital. Steal one electron from it, and add it to the d.

65. Exceptions to the Aufbau Principle OK, so this helps the d, but what about the poor s orbital that loses an electron? OK, so this helps the d, but what about the poor s orbital that loses an electron? Remember, half full is good… and when an s loses 1, it too becomes half full! Remember, half full is good… and when an s loses 1, it too becomes half full! So… having the s half full and the d half full is usually lower in energy than having the s full and the d to have one empty orbital. So… having the s half full and the d half full is usually lower in energy than having the s full and the d to have one empty orbital.