1. 1 How are an atom’s electrons configured? Section 3.3

2. 2 Electrons Most important part of an atom to chemists. Most important part of an atom to chemists. Light emission Light emission Electromagnetic spectrum Electromagnetic spectrum

3. 3 Electromagnetic Spectrum

4. 4 Light is an Electromagnetic Wave Prism Prism Waves Waves Waves can be described using speed, wavelength, and frequency. Waves can be described using speed, wavelength, and frequency.

5. 5 Light Travel Light travels at a constant speed. Light travels at a constant speed. The speed of light equals 2.998 x 10 8 m/s The speed of light equals 2.998 x 10 8 m/s 500 seconds to travel 150 million km from sun to Earth 500 seconds to travel 150 million km from sun to Earth

6. 6 Wavelength Wavelength – distance between 2 consecutive peaks or troughs of a wave (measured in meters) Wavelength – distance between 2 consecutive peaks or troughs of a wave (measured in meters) Range from 10 -13 for gamma rays to over 10 5 for radio waves Range from 10 -13 for gamma rays to over 10 5 for radio waves

7. 7 Frequency Frequency – the number of waves that pass thru a stationary point in one second Frequency – the number of waves that pass thru a stationary point in one second One wave per second is a hertz, Hz, (the unit for frequency) One wave per second is a hertz, Hz, (the unit for frequency) Ranges from less than 1000 Hz to more than 10 22 Hz Ranges from less than 1000 Hz to more than 10 22 Hz

8. 8 Light Waves High frequency/ low wavelength High frequency/ low wavelength Low frequency/ high wavelength Low frequency/ high wavelength

9. 9 Relationship Frequency Frequency x wavelength ____________ Speed of light

10. 10 Calculations The wavelength of red light is 7.6 x 10 -7 m, calculate the frequency. The wavelength of red light is 7.6 x 10 -7 m, calculate the frequency. F = speed = 2.998 x 10 8 m/s F = speed = 2.998 x 10 8 m/s wavelength 7.6 x 10 -7 m wavelength 7.6 x 10 -7 m

11. 11 Calculations Frequency = 3.9 x 10 14 /s = 3.9 x 10 14 Hz Frequency = 3.9 x 10 14 /s = 3.9 x 10 14 Hz

12. 12 Line Emission Spectra Purple hydrogen gas Purple hydrogen gas Splits and sent to prism Splits and sent to prism Line-emission spectra Line-emission spectra Lines of colored light produced when the light from excited atoms of an element is passed thru a prism Lines of colored light produced when the light from excited atoms of an element is passed thru a prism

13. 13 Excited electrons emit light Bohr theory – an electron can exist in different energy states Bohr theory – an electron can exist in different energy states Electrons can move to different energy levels. Electrons can move to different energy levels. Gain energy Gain energy Lose energy Lose energy

14. 14 States Ground state – lowest possible energy state an e- can be in Ground state – lowest possible energy state an e- can be in Excited state – e- can move here when it gains energy Excited state – e- can move here when it gains energy

15. 15 Excited State Unstable State Unstable State Higher energy than ground state Higher energy than ground state Electrons can “fall” from excited state to ground state and give off energy in the form of light Electrons can “fall” from excited state to ground state and give off energy in the form of light

16. 16 Excited State When the electron gives off energy, it is associated with a frequency that reflects the energy released. When the electron gives off energy, it is associated with a frequency that reflects the energy released. We see a color which produces the line emission spectra. We see a color which produces the line emission spectra.

17. 17 Energy of an electron Energy of an electron = Energy of an electron = -2.179x10 -18 J -2.179x10 -18 J n 2 n 2 n is any positive whole number n is any positive whole number

18. 18 Quantum Number If n has only certain values it can be, it is called a quantum number If n has only certain values it can be, it is called a quantum number The energy is said to be quantized The energy is said to be quantized The effects are only apparent at the atomic level The effects are only apparent at the atomic level

19. 19 Concept Check 1. What is the line emission spectra? 2. What does it mean to say that energy is quantized. 3. Look at fig 3-20. What wavelengths and frequencies define light in the visible spectrum?

20. 20 Orbitals Replace Orbits Bohr’s atomic model describes electrons in terms of E states. Bohr’s atomic model describes electrons in terms of E states. Present quantum model says e- have properties of particles and waves. Present quantum model says e- have properties of particles and waves.

21. 21 Quantum Numbers Each e- in an atom is assigned 3 quantum numbers: n, l, m Each e- in an atom is assigned 3 quantum numbers: n, l, m Q #’s are like seats at a concert – section, row, and seat number Q #’s are like seats at a concert – section, row, and seat number

22. 22 Orbitals e- are located in orbitals e- are located in orbitals Orbitals are regions in space that you are likely to find an e- Orbitals are regions in space that you are likely to find an e- They are like e- clouds They are like e- clouds Only certain combinations are acceptable. Only certain combinations are acceptable.

23. 23 Rules for Assigning Quantum #’s Principal quantum number – n Principal quantum number – n n has to be a whole number and typically is not higher than 7 n has to be a whole number and typically is not higher than 7 The larger n is, the further from the nucleus the e- is and the higher the energy is The larger n is, the further from the nucleus the e- is and the higher the energy is

24. 24 Rules for Quantum Numbers The l quantum number can be a whole number between 0 and n-1 The l quantum number can be a whole number between 0 and n-1 If n = 3, l can be 0, 1, or 2 If n = 3, l can be 0, 1, or 2

25. 25 l Quantum Number When l = 0, it is called the s orbital. l = 1, is the p orbital, l =2 is the d orbital, and l=3 is f orbital When l = 0, it is called the s orbital. l = 1, is the p orbital, l =2 is the d orbital, and l=3 is f orbital If n = 3 and l = 1, it is the 3p orbital. If n = 3 and l = 1, it is the 3p orbital. 3p electron 3p electron

26. 26 m Quantum Number Can be a whole number Can be a whole number Depends on l quantum number Depends on l quantum number If l = 1, m can be -1, 0, or 1 If l = 1, m can be -1, 0, or 1 There can be 3 different p orbitals There can be 3 different p orbitals

27. 27 l and m quantum numbers They indicate shapes and orientations of the orbitals. They indicate shapes and orientations of the orbitals. Quantum theory tells us the exact energy, but not the exact location of the e-. It only tells the probability of the location of the e- Quantum theory tells us the exact energy, but not the exact location of the e-. It only tells the probability of the location of the e-

28. 28 Orbitals

29. Quantum Numbers Quantum # MeaningLimits of Number n Main E Level Only up to 7 l Angular Momentum Up to n-1 l=0=s l=1=p l=2=d l=3=f m Magnetic Quantum # s=1 p=3 d=5 f=7 Spin+1/2 or -1/2 29

30. 30 Pauli Exclusion Principle No more than two e- can occupy a single orbital at a time No more than two e- can occupy a single orbital at a time e- spin in opposite directions e- spin in opposite directions Spin quantum number, m Spin quantum number, m m can be +½ or -½ m can be +½ or -½

31. 31 Electron Configuration A description of the electron orbitals in an atom A description of the electron orbitals in an atom Aufbau principle – electrons in an atom will occupy the lowest energy orbitals available. Aufbau principle – electrons in an atom will occupy the lowest energy orbitals available.

32. 32 Electron Configurations Remember, the smaller the principal quantum number, the lower the energy, and the smaller the l quantum number, the lower the energy Remember, the smaller the principal quantum number, the lower the energy, and the smaller the l quantum number, the lower the energy

33. 33 Electron Configurations The order in which energy levels fill is … The order in which energy levels fill is … 1s<2s<2p<3s<3p 1s<2s<2p<3s<3p

34. 34 Electron Configurations After this, the energy levels are less straightforward After this, the energy levels are less straightforward The E levels of the 3d orbitals are slightly higher than those of the 4s orbitals. The E levels of the 3d orbitals are slightly higher than those of the 4s orbitals. 1s<2s<2p<3s<3p<4s≈3d

35. 35 Electron Configurations The next irregularity is 5s and 4d are close in energy The next irregularity is 5s and 4d are close in energy 1s<2s<2p<3s<3p<4s≈3d<4p<5s≈4d Still more irregularities exist with higher energy orbitals. Still more irregularities exist with higher energy orbitals.

36. 36 Electron Configurations Tells us how the 16 e- of S are configured Tells us how the 16 e- of S are configured The electron configuration for S is The electron configuration for S is S=1s 2 2s 2 2p 6 3s 2 3p 4 S=1s 2 2s 2 2p 6 3s 2 3p 4 Each s orbital has 2 e-, each p orbital can have 6 e- (2 per orbital) Each s orbital has 2 e-, each p orbital can have 6 e- (2 per orbital)

37. 37 Electron Configurations To save space, some configurations are written like To save space, some configurations are written like S = [Ne]3s 2 3p 4 S = [Ne]3s 2 3p 4 This means take neon’s configuration and add 3s 2 3p 4 to the end of it. This means take neon’s configuration and add 3s 2 3p 4 to the end of it.

38. 38 Electron Configurations There are still some irregularities with higher energy orbitals. There are still some irregularities with higher energy orbitals. Chromium is an example of this Chromium is an example of this Cr = [Ar] 3d 5 4s 1 Cr = [Ar] 3d 5 4s 1 There is one unfilled d orbital and a filled s orbital There is one unfilled d orbital and a filled s orbital

39. 39 Electron Configurations There are e- configs listed in the p.t. in the back of your book. There are e- configs listed in the p.t. in the back of your book. These are the ground state configs of the isolated atoms in the gas phase. These are the ground state configs of the isolated atoms in the gas phase. Under other conditions, the configs could be different Under other conditions, the configs could be different

40. 40 Rules for Writing e- Configs 1. Determine the # of e- the atom has (atomic #) Fluorine as an example has 9 electrons Fluorine as an example has 9 electrons

41. 41 Rules for Writing e- Configs 2. Fill orbitals in order of increasing energy S orbital – 2 e- P orbitals – 6 e- D orbitals – 10 e- F orbitals – 14 e-

42. 42 Rules for Writing e- Configs

43. 43 Rules for Writing e- Configs The configuration for F is 1s 2 2s 2 2p 5 The configuration for F is 1s 2 2s 2 2p 5 3. Make sure the total number of e- in the config match the atomic number

44. 44 Orbital Diagrams Use boxes to show e- location Use boxes to show e- location Boron = 1s 2 2s 2 2p 1 Boron = 1s 2 2s 2 2p 1 The orbital diagram looks like… The orbital diagram looks like… 1s 2s { 2p   

45. 45 Orbital Diagrams Hund’s Rule says the maximum stability for e- is when you have the maximum number of unpaired electrons when they have the same quantum number. Hund’s Rule says the maximum stability for e- is when you have the maximum number of unpaired electrons when they have the same quantum number.

46. 46 Orbital Diagrams What is the Carbon orbital diagram. What is the Carbon orbital diagram. The config is 1s 2 2s 2 2p 2 The config is 1s 2 2s 2 2p 2