Structure 1.3 - Electron Configurations PDF

Guiding question: How can we model the energy states of electrons in atoms? Structure 1.3 Time: SL and HL - 3 h Electron configurations AHL - 3 h...

Guiding question: How can we model the energy states of electrons in atoms? Structure 1.3 Time: SL and HL - 3 h Electron configurations AHL - 3 h 1  Prior knowledge 1. Which types of electromagnetic radiation are relevant in these contexts? 2. What is different about different types of electromagnetic radiation? 2 SL and HL content 3  Part A The electromagnetic spectrum S1.3.1 4 What is the relationship between wavelength and the energy of electromagnetic radiation? 5 What is electromagnetic radiation? Radiation that has both electric and magnetic fields can be modelled as both: Waves Particles 6 Properties of electromagnetic radiation Wavelength Frequency Wavelength, 𝜆 (lamda), is the distance between Frequency, f, is how often a full oscillation is two peaks (m) made over time (1 / t) 7 How are wavelength, frequency and energy related? c=𝜆f E=hf C = speed of light Energy of a photon 𝜆 = wavelength E = energy f = frequency h = Planck's constant Shorter wavelength, = 6.62×10−34 J-s higher frequency f = frequency Therefore, energy is proportional to wavelength 8 Quick check: The visible 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? 9 Practice questions Identify the region of the electromagnetic spectrum that has: 1. Shortest wavelength 2. Lowest frequency 3. Lowest energy 10  Part B Emission spectra and the Bohr model of the atom S1.3.2 11  What is white light? Where might we find it? Can you explain what is happening in the image above? 12 Practical: Continuous and line 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? 13 Continuous and line emission spectra Continuous spectrum Line spectrum Shows all frequencies Shows specific frequencies 14 The Bohr model: What causes the release of energy? The one electron in H can absorb energy and move Energ y in from ground state to a higher excitation (E) state 15 The Bohr model: What causes the release of energy? The electron can then emit/release energy and drop emitted down to a lower energy level Energy (relaxation) 16 Hydrogen emission spectrum Transitions down to Transitions down to Transitions down to n = 1 are large. (UV) n = 2 are in visible n = 3 are in IR light spectrum spectrum 17 The Bohr model and the hydrogen emission spectrum 18 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? 19 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. 20 Hydrogen emission spectrum: Excitation to n=∞ As we get further out, energy levels converge. “Edge” of the atom is known as n = ∞ If electrons absorb enough energy to reach n = ∞, they leave the atom and form an ion 21 Practice questions 1. Sketch an 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. 22  Part C Energy levels, sublevels and orbitals S1.3.3 S1.3.4 23 When is an electron no longer an electron? When it’s a probability. “We cannot know both the momentum and position of an Heisenberg’s Uncertainty Principle electron at the same time.” It’s simply a probability of finding an electron. Schrӧdinger’s Wave Equation 24  Maximum number of Energy level electrons n=1 2 n=2 8 n=3 18 n=4 ? 1. Can you identify 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. 25 Main energy levels Maximum number Energy level of electrons n=1 n=2 n=3 n=4 26 Main energy levels → sublevels Which sublevels are Main energy level Number of sublevels used? n=1 n=2 n=3 n=4 27 What is an orbital? Orbital: a region of space where there is a high probability of finding an electron. - Each orbital can contain 2 electrons of opposite spins. - Orbitals in the same sublevel have identical energy levels (degenerate orbitals) 28 Sublevels → s orbitals Spherical Energy level Subtype Number of Max number of orbitals electrons n=1 s 1 2 29 Sublevels → p orbitals Dumbbell shaped Energy level Subtype Number of Max number of orbitals electrons n=2 p 3 4 p-1 p0 p1 30 Sublevels → d orbitals Energy level Subtype Number of Max number of orbitals electrons n=3 d 5 10 31 Sublevels → f orbitals Energy level Subtype Number of Max number of orbitals electrons n=4 f 7 14 32 Sublevels → orbitals Number of orbitals in Sublevel Shapes of orbitals sublevel s 1 p 3 d 5 f 7 33 Summary Energy level (n) Type of sublevel Number of orbitals per type Total number of orbitals per energy Max. number of electrons within level (n2) energy level (2n2) 1 s 1 1 1 x 2e- = 2 s 1 2 1+3=4 4 x 2e- = 8 p 3 s 1 3 p 3 1+3+5=9 9 x 2e- = 18 d 5 s 1 p 3 *Remember: 4 1 + 3 + 5 + 7 = 16 16 x 2e- = 32 each orbital d 5 holds 2 e- f 7 34 Practice questions 1. State the sublevels found in the n=1, n=2, n=3 and n=4 main energy levels. 2. Define ‘orbital’. 3. State the max number of electrons found in a single orbital. 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. 7. Explain, in terms of sublevels and orbitals, why the n=4 energy level can contain a maximum of 32 electrons. 35  Part D Orbital box diagrams and electron configurations S1.3.5 36 Which of these orbitals might you be expected to sketch in IB Chemistry? 37 How do we represent orbitals more simply? With Orbital Box Diagrams! 38 Orbital box diagram for n=1, n=2 and n=3 energy levels 3d degenerate Reminder: Main Total degenerate 3p energy Sublevels number of level orbitals 3s n=1 s 1 2p degenerate 2s n=2 s, p 1+3 1s n=3 s, p, d 1+3+5 1 2 3 n 39 How do we tell the difference between a 1s and 2s orbital? Nodes: Regions where there is zero probability of finding an electron 40 Orbital box diagram rules Aufbau Principle: electrons fill from lowest energy to highest. Pauli Exclusion Principle: electrons have spin states (it’s angular momentum, 𝑝). Two electrons in the same orbital must have opposite spins. Hund’s Rule: electrons in the same sub-level will occupy orbitals individually before sharing. 41 Orbital filling order n=1 1s2 n= 2 2s2 2p6 n=3 3s2 3p6 3d10 n=4 4s2 4p6 4d10 4f14 n=5 5s2 5p6 5d10 5f14 n=6 6s2 6p6 6d10 n=7 7s2 7p6 n=8 8s2 42 Relative energy levels s p d f of electrons in gaseous Electrons fill the lowest available atoms of the energy level first twenty Click to add electrons elements 4p 3d 4s 3p Cr Cu anelectron an 4s fills electron beforeis3d is promoted promoted from 4sfrom to 3d4s to to 3s giveto 3d a full give3d asubshell half-filled 3d subshell Electrons remain unpaired as far 2p as possible 2s Increasing energy 1s 43 Relative energy levels s p d f of electrons in gaseous Electronic configuration in atoms of the shorthand nomenclature first twenty Click to add electrons elements 4p 3d 4s 3p 122222222222 1222222222 561 234666666 222222 12 3462 6 1 566666 1 2221 221 22 32810 3s H 1s He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 1s 1s 1s 1s 1s 1s 1s 2s2s 2s 2s 2s 2s 2s 2s2p 2p 2p 2p 2p 2p 2p 3s3s 3s 3s 3s 3s3p 3p 3p 3p 3p 3p 4s 4s 4s 4s 4s 4s 3d 3d3d15610 3d 3d 3d 10 710 10 5 4p653412 4p 4p 2p 2s Increasing energy 1s 44 Writing Electron Configurations Eg. Lithium e- = 3 Full e- configuration: Ionic e- configuration: Li: 1s2 2s1 Closest Noble Gas: Helium Li+: 1s2 [He] Condensed e- configuration: Closest Noble Gas: Helium Li: [He] 2s1 45 Quick check Symbol Full e- configuration Condensed e- configuration Li 1s2 2s1 [He] 2s1 C Ar Fe Cr 46 Using the periodic table to write e- configurations 47 How do we write e- configurations for ions? Atom Full e- configuration of atom Ion Full e- configuration of ion Be Be2+ O O2- Mg Mg2+ Fe Fe3+ Cu Cu2+ 48 Practice - Condensed e- Symbol Full e configuration configuration B N Si Cu Ti4+ Mn2+ Cu+ Se2- 49 AHL content 50  Part E First ionization energy S1.3.6 51 What might these graphs look like? 52 Which factors affect the difficulty of removing an e- from an atom? Distance from nucleus: the further away an electron is from the nucleus, the easier it is to remove. Number of protons: more protons means a greater positive charge. 53 In what sense are outer e- ‘shielded’ by inner electrons? Electron shielding: the influence that inner electrons have on how strongly outer electrons feel the attraction of the nucleus. - Electrons repel each other. The negative charge of inner electrons subtracts from the positive charge felt by outer electrons. - Electrons between valence electrons and the nucleus are called shielding electrons (SE). 54 What is ‘first ionization energy’? Definition: The minimum energy required to eject an electron out of a neutral atom in its ground state. (ie. energy required to reach n = ∞) General formula: X(g) + energy → X+(g) + e- 55 How does 1st ionization energy change down a group? Valence electrons further from the nucleus experience less electrostatic attraction → less energy to remove. + + Li: Na: 3 K: 87 Rb: Cs: Fr: p 11 19 p 37 55 , 2+, SE 10 18 SE 36 54 86 56 Quick check Which best explains why it requires less energy to remove an electron from barium than calcium? a. There are more protons in the barium nucleus than the calcium nucleus. b. There are more electrons in barium than in calcium. c. The outer electrons are in a higher energy level in barium than calcium. d. Barium is in a different group to calcium. 57 Why do atoms get smaller across a period? Valence electrons remain in the same energy level (and therefore distance from the nucleus), but the positive electrostatic attraction experienced increases. 58 General trend: 1st ionization energy across a period General trend increases towards Noble Gases 59 Quick check 60 Data for 1st ionization energy across a period 61 Practice questions 1. Predict and 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 → 62  Part F Successive ionization energies S1.3.7 63 64  1st ionization energy: X (g) → X+ (g) + e- Write a general equation for the 2nd and 3rd ionization energy. 65 Successive ionization energies of Na 1. Na: 1s2 2s2 2p6 3s1 2. Na+: 1s2 2s2 2p6 3. Na2+: 1s2 2s2 2p5 4. Na3+: 1s2 2s2 2p4 5. Na4+: 1s2 2s2 2p3 6. Na5+: 1s2 2s2 2p2 7. Na6+: 1s2 2s2 2p1 8. Na7+: 1s2 2s2 9. Na8+: 1s2 2s1 10. Na9+: 1s2 11. Na10+: 1s1 66 Successive ionization energies of Na Each time an e- is removed, repulsion between all e- decreases (ie. greater electrostatic attraction) Greater electrostatic attraction → more energy required to remove the next e- 67 Successive ionization energies of Na 1st e- removed: 3rd shell, furthest from nucleus, easy to remove. 2nd-7th e- removed: 2rd shell, closer from nucleus, harder to remove from 2p. 8th-9th: removed from 2s 10-11th e- removed: 1st shell, closest to nucleus, very hard to remove. 68 Quick check 69 Practice question Sketch the successive ionization energies for a nitrogen atom. 70  Part G The limit of convergence and ionization S1.3.8 71  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. 72 The limit of convergence As energy levels converge, we reach the edge of the atom at the convergence limit. 73 How do we calculate ionization energy from spectral data? The convergence limit for a hydrogen atom Data booklet: occurs at a frequency of 3.28 x 1015 s-1. E=hf Calculate the: h = 6.63 x 10-34 J s 1. First ionization energy for a hydrogen atom in J. 1 mol = 6.02 x 1023 atoms 2. First ionization energy for a hydrogen atom in kJ. 3. First ionization energy for hydrogen in kJ mol-1. 74 Practice question 1 The electron in a hydrogen atom reaches the Data booklet: convergence limit when it absorbs radiation E=hf with a wavelength of 9.15 x 10-8 m. Calculate the ionization energy in kJ mol-1. h = 6.63 x 10-34 J s 1 mol = 6.02 x 1023 atoms c=f𝜆 c = 3.00 x 108 m s-1 75 Practice question 2 A beam of electromagnetic radiation has an energy of Data booklet: 3.65 x 10-20 J per photon. E=hf a. Calculate the frequency of the radiation. b. Calculate the wavelength of the radiation. h = 6.63 x 10-34 J s c. Identify the type of electromagnetic radiation c=f𝜆 using the electromagnetic spectrum below. c = 3.00 x 108 m s-1 76 Guiding question and review 77 How can we model the energy states of electrons in atoms? Concept map: Make and annotate links between these concepts and the core Particles concepts on the right. Charge Electron Hydrogen Electrostatic forces Hydrogen emission spectrum Photon Potential energy Energy level Conservation of mass/energy Sublevel Orbital 78 Key terminology 1. Wavelength 2. Frequency 3. Continuous spectrum 4. Line spectrum 5. Photon 6. Hydrogen emission spectrum 7. Orbital 8. AHL: Ionization energy 79 Past-paper questions 80 Retrieval practice 81  Gamma rays Infrared Microwaves Visible light X-rays Radio waves Ultraviolet Retrieval practice: Place the following regions of the electromagnetic spectrum in order of increasing frequency. 82  Unknown star 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? 83  Total number of Maximum no. Main energy level Sublevels orbitals electrons n=1 n=2 n=3 n=4 Retrieval practice: Complete this table. 84 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... 85 1. Inner electrons shield 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? 86 The first four successive ionization energies are: 420, 3600, 4400 and 5900 kJ mol-1 Which group of the periodic would this element be found? 87 NOS and TOK 88 NOS/TOK: Evidence, models and theories 1. Use examples from Structure 1.3 to explain the relationship between these: Evidence Models Theories 2. Can you do the same using examples from another IB subject? 89 NOS: Logarithmic scales Why do we often use logarithmic scales in science? 90 Extension 91 Extension: Doppler shift How can emission and absorption spectra tell us about movement? 92 Extension: Atomic absorption spectra What are the difference between an emission line spectrum and an absorption line spectrum? What causes these differences? 93 Extension: Electron configurations and quantum numbers 94 Probabilities and quantum wavefunctions 95