Download to read offline
![1. Has the e-
configuration 1s2
2s2
2p6
3s2
.
2. Has 2 electrons in the second energy level.
3. Has the e-
configuration 1s2
.
4. Contains [Ar] in its condensed e-
configuration.
5. Finishes with p2
in its e configuration.
6. Contains an unpaired electron in an s-orbital.
7. Contains a half filled set of d-orbitals.
8. Contains a complete p sub-level.
9. Has 28e-
in its 2+ ion.
Retrieval practice: State the name of an element that...
69](https://image.slidesharecdn.com/structure1-250226130252-80bb84b2/85/Structure-1-3-Electron-configurations-pptx-69-320.jpg)
Structure 1.3 - Electron configurations.pptx
- 1. Structure 1.3 Electron configurations Guidingquestion: How can we model the energy states of electrons in atoms? Time: SL and HL - 3 h AHL - 3 h 1
- 2. Prior knowledge 1. Whichtypes of electromagnetic radiation are relevant in these contexts? 2. What is different about different types of electromagnetic radiation? 2
- 3. SL and HLcontent 3
- 4.
- 5. What is therelationship between wavelength and the energy of electromagnetic radiation? 5
- 6. What is electromagneticradiation? Radiation that has both electric and magnetic fields can be modelled as both: 6
- 7. Properties of electromagneticradiation: Wavelength 7
- 8. Properties of electromagneticradiation: Frequency 8
- 9. How are wavelength,frequency and energy related? c = 𝜆 𝜈 E = h 𝜈 9
- 10. Quick check: Thevisible region of electromagnetic radiation 1. What are the different colours of light in the visible region of the electromagnetic spectrum? 2. Which colour of visible light has: a. the longest wavelength? b. the highest frequency? c. the greatest energy? 10
- 11. Practice questions Identify theregion of the electromagnetic spectrum that has: 1. Shortest wavelength 2. Lowest frequency 3. Lowest energy 11
- 12. Part B Emission spectraand the Bohr model of the atom 12
- 13. What is whitelight? Where might we find it? Can you explain what is happening in the image above? 13
- 14. Practical: Continuous andline emission spectra A diffraction grating separates the different components of electromagnetic radiation. With our eyes, we will be see any colours present in the visible region. Record your observations when: 1. Looking at white light 2. Looking at a hydrogen gas lamp What differences do you notice? 14
- 15. Continuous and lineemission spectra 15
- 16. The Bohr model:What causes the release of specific frequencies of energy? 16
- 17. The Bohr modeland the hydrogen emission spectrum 17
- 18. An energy leveldiagram for the hydrogen emission spectrum 18
- 19. Hydrogen emission spectrum:Drops to n=1 and n=3 seen in the visible region of the EM spectrum 19
- 20. Quick check 1. (MCQ)What do the lines on the hydrogen emission spectrum represent? a. The energy levels in a hydrogen atom b. The energy gap between energy levels in a hydrogen atom c. Excited electrons in hydrogen atoms d. Relaxed electrons in hydrogen atoms 2. The hydrogen emission spectrum is caused by electrons dropping to which energy level? 3. Why do the lines on an emission spectrum converge at higher energies? 20
- 21. Practical: Flame tests 1.Hold wooden splints soaked in different metal ion solutions in a Bunsen burner flame. 2. Record your observations. 3. Identify the likely metal ion present in the two unknown solutions. 21
- 22. Hydrogen emission spectrum:Excitation to n=∞ 22
- 23. Practice questions 1. Sketchan energy level diagram with 6 energy levels. Add and label arrows showing an electron transition… a. Caused by the absorption of energy. b. Representing a line in the hydrogen emission spectrum. c. That would represent the largest possible energy emission. 2. Explain why we do not see lines on a hydrogen emission spectrum caused by drops to n=1. 23
- 24. Part C Energy levels,sublevels and orbitals 24
- 25. 1. Can youidentify a mathematical formula that relates the energy level (n) and its maximum number of electrons? 2. Predict the maximum number of electrons held in the n=4 energy level. Energy level Maximum number of electrons n=1 2 n=2 8 n=3 18 n=4 ? 25
- 26. Main energy levels Energylevel Maximum number of electrons n=1 n=2 n=3 n=4 26 2 8 18 32
- 27. Main energy levels→ sublevels Main energy level Number of sublevels Which sublevels are used? n=1 s n=2 s, p n=3 3 s, p, d n=4 4 s, p, d, f 27 1 2
- 28. Sublevel Number of orbitalsin sublevel Shapes of orbitals s 1 p 3 px, pppppp d 5 f 7 Sublevels → orbitals 28
- 29. Summary Main energy levelSublevels Total number of orbitals Maximum no. electrons n=1 s 1 2 n=2 p 3 6 n=3 d 5 10 n=4 f 7 14 29
- 30. Practice questions 1. Statethe sublevels found in the n=1, n=2, n=3 and n=4 main energy levels. 2. Define ‘orbital’. The area where you have high probability of finding electron 3. State the number of electrons found in a single orbital. 2 4. Sketch the shape of an s orbital. 5. Sketch a pz orbital on these axes → 6. State the number of orbitals are found in the s, p, d and f sublevels. 1, 3, 5, 7 7. Explain, in terms of sublevels and orbitals, why the n=4 energy level can contain a maximum of 32 electrons. It has 4s, 4p, 4d, and 4f sublevels, the s orbital has 1 orbital, p 3, d 5, f 7. and one orbital contains 2 electron. We have 16 orbital in total, so there are 32 electrons 30
- 31.
- 32. Which of theseorbitals might you be expected to sketch in IB Chemistry? 32
- 33. How do werepresent orbitals more simply? 33
- 34. Orbital box diagramfor n=1, n=2 and n=3 energy levels 34 Reminder: Main energy level Sublevels Total number of orbitals n=1 s 1 n=2 s, p 1 + 3 n=3 s, p, d 1 + 3 + 5
- 35. Orbital box diagrams Howmight we draw the first 8 electrons in an atom? 35
- 36. Relative energy levels of electrons ingaseous atoms of the first twenty elements Increasing energy s p d f 1s 2s 2p 3s 3d 3p 4s 4p Electrons fill the lowest available energy level 4s fills before 3d Electrons remain unpaired as far as possible Cr an electron is promoted from 4s to 3d to give a half-filled 3d subshell Cu an electron is promoted from 4s to 3d to give a full 3d subshell Click to add electrons 36
- 37. Relative energy levels of electrons ingaseous atoms of the first twenty elements Increasing energy s p d f 1s 2s 2p 3s 3d 3p 4s 4p Electronic configuration in shorthand nomenclature Click to add electrons H 1s1 He 1s2 Li 1s2 2s1 Be 1s2 2s2 B 1s2 2s2 2p1 C 1s2 2s2 2p2 N 1s2 2s2 2p3 O 1s2 2s2 2p4 F 1s2 2s2 2p5 Ne 1s2 2s2 2p6 Na 1s2 2s2 2p6 3s1 Mg 1s2 2s2 2p6 3s2 Al 1s2 2s2 2p6 3s2 3p1 Si 1s2 2s2 2p6 3s2 3p2 P 1s2 2s2 2p6 3s2 3p3 S 1s2 2s2 2p6 3s2 3p4 Cl 1s2 2s2 2p6 3s2 3p5 Ar 1s2 2s2 2p6 3s2 3p6 K 1s2 2s2 2p6 3s2 3p6 4s1 Ca 1s2 2s2 2p6 3s2 3p6 4s2 Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1 Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2 V 1s2 2s2 2p6 3s2 3p6 4s2 3d3 Cr 1s2 2s2 2p6 3s2 3p6 4s1 3d5 Mn 1s2 2s2 2p6 3s2 3p6 4s2 3d5 Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Co 1s2 2s2 2p6 3s2 3p6 4s2 3d7 Ni 1s2 2s2 2p6 3s2 3p6 4s2 3d8 Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10 Zn 1s2 2s2 2p6 3s2 3p6 4s2 3d10 Ga 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1 Ge 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 As 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 Se 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 Kr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 37
- 38. Quick check Symbol Fulle- configuration Condensed e- configuration Li C Cu Fe Cr 38
- 39. How do wewrite e- configurations for ions? Atom Full e- configuration of atom Ion Full e- configuration of ion Li Li+ O O2- Mg Mg2+ Fe Fe3+ Cu Cu2+ 39
- 40. Using the periodictable to write e- configurations 40
- 41. Practice Symbol Full e- configuration Condensede- configuration Be N Si Cu Ti4+ Mn2+ Cu+ Se2- 41
- 42.
- 43. Part E First ionizationenergy 43
- 44. In what senseare the outer electrons in an atom ‘shielded’ by inner electrons? 44
- 45. What is ‘firstionization energy’? Definition: General formula: 45
- 46. Data for 1stionization energy across a period 46
- 47.
- 48. How does 1stionization energy change down a group? Predict the general trend in 1st ionization energies down a group. 48
- 49. Practice questions 1. Predictand explain the general trend in 1st ionization energy down group 2. 2. Explain the general trend in 1st ionization energy across a period. 3. Explain the discontinuities (exceptions) to the general trend in the 1st ionization energies of period 3 elements → 49
- 50.
- 51. 1st ionization energy:X (g) → X+ (g) + e- Write a general equation for the 2nd and 3rd ionization energy. 51
- 52. Successive ionization energiesof Na 52
- 53. Quick check Sketch thesuccessive ionization energies for a nitrogen atom. 53
- 54.
- 55. Part G The limitof convergence and ionization 55
- 56. E = h𝜈 h = 6.63 x 10-34 J s A radio station transmits at a frequency of 1.089 x 106 s-1 . Use the data provided to calculate the energy of the photon of the transmission waves in kJ. 56
- 57. The limit ofconvergence 57
- 58. How do wecalculate ionization energy from spectral data? The convergence limit for a hydrogen atom occurs at a frequency of 3.28 x 1015 Hz. Calculate the: 1. First ionization energy for a hydrogen atom in J. 2. First ionization energy for a hydrogen atom in kJ. 3. First ionization energy for hydrogen in kJ mol-1 . Data booklet: E = h 𝜈 h = 6.63 x 10-34 J s 1 mol = 6.02 x 1023 particles c = 𝜈 𝜆 c = 3.00 x 108 m s-1 58
- 59. Practice question 1. Abeam of electromagnetic radiation has an energy of 3.65 x 10-20 J per photon. a. Calculate the frequency of the radiation. b. Calculate the wavelength of the radiation. c. Identify the type of electromagnetic radiation using the electromagnetic spectrum below. 59
- 60. Practice question 2. Theelectron in a hydrogen atom reaches the convergence limit when it absorbs radiation with a wavelength of 9.15 x 10-8 m. Calculate the ionisation energy in kJ mol-1 . 60
- 61.
- 62. How can wemodel the energy states of electrons in atoms? 62 Where do we find electrons? How do we represent the position of electrons? How do electrons change energy states? AHL - How can we ionize an atom?
- 63. Key terminology 1. Wavelength 2.Frequency 3. Continuous spectrum 4. Line spectrum 5. Photon 6. Hydrogen emission spectrum 7. Orbital 8. AHL: Ionization energy 63
- 64.
- 65.
- 66. Gamma rays Infrared Microwaves Visible light X-rays Radiowaves Ultraviolet Retrieval practice: Place the following regions of the electromagnetic spectrum in order of increasing frequency. 66
- 67. Nature of science: 1.Which elements might be found in the unknown star? 2. How do emission spectra provide evidence for the existence of different elements? Unknown star 67
- 68. Retrieval practice: Completethis table. Main energy level Sublevels Total number of orbitals Maximum no. electrons n=1 n=2 n=3 n=4 68
- 69. 1. Has thee- configuration 1s2 2s2 2p6 3s2 . 2. Has 2 electrons in the second energy level. 3. Has the e- configuration 1s2 . 4. Contains [Ar] in its condensed e- configuration. 5. Finishes with p2 in its e configuration. 6. Contains an unpaired electron in an s-orbital. 7. Contains a half filled set of d-orbitals. 8. Contains a complete p sub-level. 9. Has 28e- in its 2+ ion. Retrieval practice: State the name of an element that... 69
- 70. 1. Inner electronsshield valence electrons from the positive charge in the nucleus. 2. Elements in the same group have the same number of shielding electrons. 3. Elements in the same period have the same number of shielding electrons. 4. Down a group, 1st ionization energy increases as valence electrons are found further from the nucleus. 5. Across a period, the general trend in 1st ionization energy in increasing. Retrieval practice: True or false? 70
- 71. The first foursuccessive ionization energies are: 420, 3600, 4400 and 5900 kJ mol-1 Which group of the periodic would this element be found? 71
- 72.
- 73. NOS/TOK: Evidence, modelsand theories 1. Use examples from Structure 1.3 to explain the relationship between these: 2. Can you do the same using examples from another IB subject? Evidence Models Theories 73
- 74. NOS: Logarithmic scales Whydo we often use logarithmic scales in science? 74
- 75.
- 76. Extension: Doppler shift Howcan emission and absorption spectra tell us about movement? 76
- 77. Extension: Atomic absorptionspectra What are the difference between an emission line spectrum and an absorption line spectrum? What causes these differences? 77
- 78. Extension: Electron configurationsand quantum numbers 78
Editor's Notes
- #5 Data booklet section 3.
- #6 A wave and a particle (photon)
- #7 Wavelength, 𝜆 (lamda), is the distance between two peaks (m). Images: https://www.britannica.com/science/wavelength
- #8 Frequency, 𝜈 (nu), is the number of waves that pass a point per second (s-1 or Hz). Images: https://www.britannica.com/science/wavelength
- #9 Key points: Shorter wavelength, higher frequency. Energy is proportional to wavelength. The formulas are not required for SL.
- #14 Helpful simulation: https://phet.colorado.edu/sims/cheerpj/discharge-lamps/latest/discharge-lamps.html?simulation=discharge-lamps
- #15 Continuous shows all frequencies of visible light. Line shows only specific frequencies.
- #16 The one electron in a H atom can absorb energy and move from the ground state up to a higher E level (excited state). The electron can then emit/release energy and drop back down to a lower energy level (relaxation).
- #17 Lines on the spectrum converge because the energy levels converge at higher energy.
- #19 Transitions down to n=1 are larger in energy and would be seen in the UV region. Transitions down to n=3 are smaller in energy and would be seen in the IR region.
- #20 B A
- #22 As energy levels converge at higher energies, we will soon reach the ‘edge’ of the atom known as n=∞. If an electron absorb enough energy to reach n=∞, it will leave the atom (forming a positive ion).
- #25 Maximum number of electrons = 2n2
- #26 It can be helpful to mention ‘electron arrangements’ here. Although not required in the IB, many students have come across them in previous chemistry courses.
- #27 Note that each additional main energy level has 1 additional sublevel. In order of increasing energy: s < p < d < f Evidence for sublevels can be seen in ionization energy data. Explained by quantum mechanics. (not required) Names of sublevels: sharp, principal, diffuse, fundamental
- #28 An orbital is a region of space where there is a high probability of finding an electron. Each orbital can contain 2 electrons of opposite spin. Orbitals in the same sublevel have identical energy. Note that each set of axes shows one orbital (the light and dark colours technically shows the phases of the lobes which helps explain bonding and antibonding molecular orbitals). https://www.mathsisfun.com/physics/atom-orbitals.html
- #29 n=1 s 1 1 x 2e- = 2 n=2 s, p 1 + 3 = 4 4 x 2e- = 8 n=3 s, p, d 1 + 3 + 5 = 9 9 x 2e- = 18 n=4 s, p, d, f 1 + 3 + 5 + 7 = 16 16 x 2e- = 32
- #33 A box can be used to represent individual orbitals.
- #35 (specific names not required) Aufbau principle: Electrons fill orbitals from lowest energy to highest. Pauli exclusion principle: Quantum mechanics states that electrons possess a property called spin. Two e- in the same orbital must possess opposite spin. Hund’s rule: Electrons in the same sub-level will occupy orbitals individually before sharing.
- #36 Click in presentation mode to work through this slide. As the main energy levels converge at higher energy, we begin to see the sublevels overlap. 3p < 4s< 3d. This occurs more and more when including the n=5 and n=6 energy levels.
- #37 Click in presentation mode to work through this slide.
- #40 https://commons.wikimedia.org/wiki/File:Simple_Periodic_Table_Chart-blocks.svg
- #44 Repulsion between inner and outer electrons reduces the pull on outer electrons towards the nucleus.
- #45 IB definition:
- #46 Valence e- are all in the 2nd energy level so experience same shielding from 1st energy level electrons. Increasing number of protons in nucleus → increasing electrostatic attraction between nucleus and electrons → increasingly more energy required to remove a valence electron. General trend: Increasing across a period as valence e- are in same E level (so experience the same shielding) but experience greater attraction to the nucleus as the number of protons increases. General trend predicted by Bohr model is accurate → increasing 1st IE across a period. Exceptions at B and O: Valence e- in B is in 2p sublevel → slightly further from nucleus → less electrostatic attraction → less energy required to remove. Easiest valence electron to remove in O is found paired with another electron in a 2p orbital → electrostatic repulsion between paired electrons requires less energy to remove.
- #47 Answer: C. The last 2 values are significantly lower than the first 3 so are likely elements in the next period.
- #48 Valence electrons are further and further from the nucleus so experience less and less electrostatic attraction → less energy required to remove electrons. You might note that the number of protons also increases significantly, however the number of shielding electrons (SE) increases in proportion to this so distance from the nucleus becomes the primary factor to consider: Li has 3 p+ and 2 SE Na has 11 p+ and 10 SE K has 19 p+ and 18 SE Rb has 37 p+ and 36 SE Cs has 55 p+ and 54 SE Fr has 87 p+ and 86 SE
- #49 https://chemistry.stackexchange.com/questions/145745/why-does-the-ionization-enthalpy-of-elements-across-a-period-not-follow-a-regula
- #52 General trend is increasing: each time an electron is removed, repulsion between all electrons decreases and they move closer to nucleus → greater electrostatic attraction → more energy required to remove next electron. Explaining the jumps 1st electron removed is from 3rd energy level → farthest from nucleus → easiest to remove 2nd-9th electron removed are from 2nd energy level → closer to nucleus → more difficult to remove. Small jump between 7th and 8th electron: 2nd-7th electron removed are from 2p sublevel 8th and 9th electron removed are from 2s sublevel 10th and 11th electron removed are from 1st energy level → closest to nucleus so very difficult to remove.
- #53 General increase as more electrons are removed. Small jump between 3 and 4 as electrons removed move from 2p to 2s. Large jump between 5 and 6 as electrons removed from n=2 to n=1 energy level.
- #54 Answer: B. An element in group 13 will have a significant jump in IE for the 4th electron removed as it is an energy level closer to the nucleus.
- #56 Note: The units of Hz is the same as s-1.
- #57 As energy levels converge at higher energies, we soon reach the edge of the atom at the convergence limit.
- #58 E = hv = 2.17x10-18 J 2.17x10-18 J x 1/1000 = 2.17x10-21 kJ 2.17x10-21 kJ x 6.02x1023 mol-1 = 1310 kJ mol-1
- #59 v=E/h = 5.51 x 1013 s-1 (or units in Hz) λ=c/v = 5.45 x 10-6 m Radio waves
- #66 Radio waves < microwaves < infrared < visible light < ultraviolet < x-rays < gamma rays
- #71 Answer: Group 1. Largest increase between 1st and 2nd IE.
- #73 Possible relationships: Evidence is used to propose and support models and theories. Models contain the characteristics used in explanatory theories. Theories use models to explain the behaviour of matter. Models and theories must be supported by evidence.
- #74 A logarithmic scale condenses a large range of values into a smaller range with simpler numbers. This can help with the graphical presentation of data.
- #76 https://webbtelescope.org/contents/articles/spectroscopy-101--beyond-temperature-and-composition