2. Do now: answer these questions In most of its compounds, this element exists as a monatomic cation: F, S, N or Ca Which of the following groups has the species correctly listed in order of decreasing radius? Cu 2+, Cu +, Cu V, V 2+, V 3+ F -, Br -, I – B, Be, Li Which of the elements may occur in the greatest number of different oxidation states? C, F, O or Ca? Why?

3. Back titration: A back titration, or indirect titration, is generally a two-stage analytical technique: Reactant A of unknown concentration is reacted with excess reactant B of known concentration. A titration is then performed to determine the amount of reactant B in excess.

4. Back titrations are used when: one of the reactants is volatile, for example ammonia. an acid or a base is an insoluble salt, for example calcium carbonate a particular reaction is too slow direct titration would involve a weak acid - weak base titration (the end-point of this type of direct titration is very difficult to observe)

5. Example: A student was asked to determine the concentration of ammonia, a volatile substance, in a commercially available cloudy ammonia solution used for cleaning. First the student pipetted 25.00 mL of the cloudy ammonia solution into a 250.0 mL conical flask. 50.00 mL of 0.100 M HCl (aq) was immediately added to the conical flask which reacted with the ammonia in solution. The excess (unreacted) HCl was then titrated with 0.050 M Na 2 CO 3(aq). 21.50 mL of Na 2 CO 3(aq) was required. Calculate the concentration of the ammonia in the cloudy ammonia solution.

6. Periodicity and Atomic Structure By: Thy Vuong-Schmick

7. Do now: What are the products when lithium carbonate is heated? LiOH and CO 2 Li 2 Oand CO 2 LiO and CO 2 LiC and O 2 LiO and CO Which elements react most readily with water? K Ca S O Mg Which metal will not react with aqueous hydrochloric acid? Fe Al Cu K Ni

8. Know your reactivity series Know solubility chart

9. Recap: titration Discuss with someone sits next to you: Joe thinks his drink tastes funny, should he do a titration to see what’s in it? Joe’s glass of water taste salty, should he do a titration to find out the concentration of salt in his water?

10. Also report: uncertainty, deviation, percent yield when possible.

11. Flame test: NaCl SrCl 2 LiCl KCl CuCl 2 CaCl 2 H 3 BO 3 - Dissolve in methanol - Doesn’t have to be exact, a couple spoonful is good - Sticks/wood splints - Bunsen burner - Dip the stick in solutions and light up to see color

12. Do Now: What element does this Graph represent? - Boron - Nitrogen - Aluminum - Phophorous IONIZATION ENERGY ELECTRONRELECTRONR

13. Photoemission spectroscopy (PES), How to analyze this spectra: Shoot xray at substance How does Xray work? Electron gun Electron emitted with specific energy Measure E Xray E – E emitted = IE Energy shot out

14. The higher the peak the more electrons there are in a shell So look at the graph, how many shells are there? At 100 and 104, electrons are closest to the nucleus so hard to remove them, so the IE is high Due to proton attract strongly to the electrons: Columbic’s law Now look at the peaks’ height to determine which shell and how many electrons are in each shell: 1s 2 …

15. Contents: Development periodic table (know the scientists: Bohr…) Light and electromagnetic spectrum Radiation and atomic spectra Particlelike properties of electromagnetic radiation: Planck Equation (derive in PChem) Wavelike properties of matter: de Broglie equation (derive PChem) Quantum mechanics and Heisenberg Uncertainty principle Wave functions and quantum numbers Shapes of orbitals Electron spin and Pauli exclusion principle Orbital energy levels in multielectron atoms Electron configuration Atomic radii…

16. Homework Problems Read Sections 6.1 - 6.4 & 21.6 Read Sections 6.5 – 6.9 Chapter 6 2, 4, 56, 7, 8, 50, 53, 54 12, 16, 1860, 62, 64, 66 24, 26, 28, 3068, 70, 71 32, 34, 3872, 73, 74 42, 43, 46 Chapter 21 45, 48, 50, 51 For the pervious chapter, I don’t think I gave enough practice problems  If you need me to reteach solutions/titration let me know

17. Review: physic Wavelength = ? Speed = ? units for speed, wavelength and frequency

18. Why study visible light and electromagnetic radiation? Atomic structure: Atoms give off light when heated (excited) Pass  narrow slit  prism Light emitted by excited atom, release electron Give a few wavelength (not a full rainbow) Give series discrete lines separated by blank areas  “line spectrum” Example: Excited Na = heat NaCl inflame of Bunsen Burner Flame test lab Give yellow light Excited H = bluish light Remember how metals in firework?

19. Chemist categorize elements into line spectra Monochromatic = ____________________ (laser) Compare to light bulb and stars_______________________ Spectrum: ________________________________ Continuous spectrum__________________________ Example: Rainbow Each element has its own spectral signature Line spectrum Use this info to identify elements in minerals and other substances

20. Johann Balmer 1885: Work with hydrogen atom Wavelength of 4 lines in H spectrum Came up with his equation: R is a constant called: Rydberg constant 1.097 x 10 -2 nm -1 *** due to quantum number “n” that on the pd table we do not see 2d, 1p… (talk about this later)

21. N and m: These deal with the quantum numbers: With n = the outer shell And m = the inner shell (some book use ni and nf, or n and n’) Remember the Bohr model? (talk about spiral electron, disobey law of physic, lead to Bohr’s 3 postulates…) Let’s draw the energy diagram to show where electrons go

22. Bohr: Electron orbits certain radii, corresponding to certain definite energies, are permitted for the electron in a hydrogen atom Electron in permitted orbit has specific energy and is in an allowed energy state. Electron in allowed energy state will not radiate energy so it doesn’t spiral into nucleus Energy emitted or absorbed by electron only as it changes from one allowed energy state to another (as a photon) E = hv

23. Lead to “principal quantum number” Each orbit corresponds to a different value of n, as radius of orbit gets larger as n increases Therefore first orbit allowed is n =1, we do not have n =0 Next is n = 2 and so forth… Talk about Node Talk about “ground state” n =1 Higher one climbs ladder, higher energy Excited state The point where electron is completely separated from the nucleus What is n = ? And Energy = ? When electron absorb or emit energy…

24. ∆E = E f – E i = E photon = hv Electron absorb or emit energy Initial to final state

25. Example of electron move from n 3 to n 1 :

26. If n f is smaller than n i, electron moves closer to the nucleus Release energy ∆E = - #

27. Use what we know for E to calculate wavelength or frequency: Since change in energy is negative, release energy, the 1.03 x10 -7 is wavelength “emitted”

28. Derive and lead to: Rydberg Equation

29. Practice: Using figure on the right, which of the following electronic transitions produces the spectral line having longest wavelength: n 2 to n1, n3 to n2 or n 4 to n3 Hint: E = hv ʎ = c/v

30. Practice: Which of these electronic transition emits or absorb energy? n from 3 to 1 n from 2 to 4

31. Problem with Bohr model? Which lead to _____________________

32. n is the principal quantum number Start with 1 because it is the first shell You can not have n = 0 It is also the initial energy level of the electron M is the final energy level of electron

33. Example: What are 2 longest wavelength lines in the Lyman series of the hydrogen spectrum? How do we solve for this?

34. Remember: Lyman series of the hydrogen spectral line m = 1 Balmer: m = 2 which means the electron jump up from or drop to the 2 nd level Paschen: m = 3 Brackett: m = 4 pFund: m = 5 And so on…

35. So to solve the problem: In the Lyman series: m = 1 always (m is the inner shell) Where n > 1 (n is the outer shell of the Bohr model) what does this mean??? Let’s look at the energy diagram (We will also look at the quantum number later where they also use the m to represent something else, so don’t get confuse)

36. So wavelength is greatest when “n” is smallest So n = 2 and 3 Now plug in to the equation we get wavelength is 121.5 nm and 102.6nm We can also test n = another number to make sure we will not have wavelength bigger than 121.5 nm

37. Another example: What is the shortest wavelength line in the Lyman series of the hydrogen spectrum? Recall: m is always 1, n is > 1 In this case, we want the shortest wavelength so the quantum number “n” has to be? Why?

38. Answer: N has to be approaching infinity. Plug it in and we should get wavelength = 91.2 nm

39. electromagnetic waves

40. Try: Rabbits Mate In Very Unusual eXpensive Gardens meaning: Radio Microwaves Infra-Red Visible light Ultra-violet X-rays Gamma rays

41. Example: Which wave has the higher frequency Which represent a more intense beam of light? Which is blue light and which is red light?

44. Answer: Blue light has more frequency than red because the wavelength is shorter and E = hv where h is a constant and v is frequency

45. Electronic structure of atom example problems 1: You can do number 1 to 5 on the worksheet

46. Do Now: this is Phosphorous - What peak represents the 2s subshell? - At what peak do electrons have the most Coulombic potential energy? - How many valence electrons does this atom have? IONIZATION ENERGY ELECTRONRELECTRONR

47. Do now: Why does an ion of phosphorus (P -3 ) have a larger radius than a neutral atom of phosphorus? A) there is a greater Coulombic attraction between the nucleus and the electrons in P -3 B) the core electrons in the ion exert a weaker shielding force than those of a neutral atom C) the nuclear charge is weaker in the ion than in the atom D) The electrons in ion have a greater Coulombic repulsion than those in neutral atom Binary compound contain only copper and iodine is found to be 16.5% copper by mass. What is the empirical formula of this compound? Tutoring after school for AP exam in May or other materials

49. Matt’s question: MRI is similar to NMR Scan Hydrogen proton in our body Why? Use magnetic properties in human body to produce images Element have charge (draw vector) Spinning on its axis Behave like a small bar magnet Usually spin randomly In MRI, they line up Create a vector MRI san this Switched off machine, vector returns to resting state Emit energy in term of radio wave Produce image

50. NMR: organic chemistry Nuclear Magnetic Resonance Can’t look at beaker to see structure because it is microscopic Different nucleus absorb different magnetic field So you start with low magnetic field, and increase it higher and higher until nucleus absorb that field Draw a picture of NMR scan with peaks Peak tells the absorption (from low to high) (down field, up field) (or chemical shift ppm) determining the structure of organic compounds Calculate how many hydrogen bond to carbon, double bonds, single bonds…… Elements have a characteristic spin, we call it “I” Organic: H, C, F, P… Have I = ½ (others have I = 1, 2,3… ) A spinning charge generates a magnetic field Draw this to show +1/2 and -1/2

51. More NMR In the presence of an external magnetic field 2 spin states exist, +1/2 and -1/2. The magnetic moment of the lower energy +1/2 state is aligned with the external field but that of the higher energy -1/2 spin state is opposed to the external field Then: calculate the difference in energy between the two spin states (ΔE) = MHz ranging from 20 to 900 Mz, radio and television broadcast spectrum.

52. Do now: Which neutral atom of the following elements would have the most unpaired electrons? Titanium Manganese Nickel Zinc Which element will have a higher electronegativity value: Chlorine or Bromine? Why?

54. Do now: 1.which of the following elements has its highest energy subshell completely full? 1.Sodium? 2.Aluminum 3.Chlorine 4.Zinc 2.Which of the following isoelectric species has the smallest radius? 1.Sulfur ion 2.Chloride ion 3.Argon 4.Potassium ion

55. Next lesson: Particle like properties of electromagnetic radiation Planck equation

56. Recall: Max Planck 1858 – 1947: Black body radiation: Visible glow that solid objects give off when heated Example: Electric stove Iightbulb Energy or intensity of blackbody radiation varies with wavelength E = hv Since we know the speed of light equation we can sub for v So E = ? H = 6.626 x 10 -34 J.s 1 J = 1 Kg. m 2 / s 2

57. Higher frequency: Shorter wavelengths Higher energy Higher intensity We just prove the homework problem from yesterday

58. Lower frequency: Longer wavelengths Lower energy

59. Electromagnetic energy = quantized Quantum = smallest bit of electromagnetic radiation that can be emitted Quantum = photon Light has properties of particles as well as waves As frequency increases, E increases Energy is not continuous, it is quantized (which means only certain energies are allowed) Electrons excited to certain shell It cannot stay in 1.5 shell or.5 shell Must be whole number shell Stairs example or bubbles

60. Albert Einstein He support “energy quantized” not continuous Electron jump from one shell to another, not half shell… Explanation of “photoelectric effect:” Irradiating a metal surface with light cause electrons to be ejected from the metal Frequency of light used for this must be above some threshold value Different value for every metals Example: blue light cause metallic sodium to emit electrons Red light has no effect on sodium You can calculate the Energy of blue and red light base on wavelength and frequency

61. Photoelectric effect summary: A beam of light behaves as if it were composed of a stream of small particles Photons E = hv If v or E of photon strike metal is below a minimum value No electron ejected If above the value, which mean sufficient energy is transferred from photon for electron to overcome attractive forces holding it to the metal then the electron will be ejected Matter and energy occur only in discrete unit What does this mean?

62. Remember: Energy of photon depends on frequency Not intensity (intensity= number of photons in the beam) It doesn’t matter how many photon you shine on the metal, if the E is not sufficient enough electrons won’t be ejected Low intensity beam of high E might knock a few electrons loose from a metal High intensity beam of low E photons might not be able to knock loose a single electron Think of a golf ball and a bunch of ping pong balls hitting a window

63. New lesson: wavelike properties of matter De Broglie equation He said: Matter and light are wavelike as well as particlelike

64. He use Einstein equation: E = mc 2 Solve for m Combine with Planck equation: E = hv Sub E to derive a relationship between mass and wavelength. Combine with the speed of light equation to get: M = h/(ʎc) Replace speed of light by speed of electron c  v To calculate wavelength of electron or any other particle/object of mass m moving at velocity v VELOCITY IS NOT THE SAME AS SPEED !!! 

65. Let’s look at E = mc 2 mass and energy of an object are proportional. If a system loses mass, it loses energy (exothermic) if it gains mass, it gains energy (endothermic) Because the proportionality constant in the equation, c 2, is such a large number, even small changes in mass are accompanied by large changes in energy.

67. Example: What is the de Broglie wavelength of a pitched baseball with mass of 120 g and speed of 44.7 m/s?

68. Example: As electron in a hydrogen atom jumps from the n 3 orbit to n 7 orbit, does it absorb energy or emit energy? Also, can I solve for energy? If so, what is it?

69. Flame test question from gen chem: What happens when you mix two or more metal ion solutions in a flame test? Do the colors they produce mix?

70. Answer Nik’s question: Segre chart Equation for n turn to p and e and gama Spontaneous Exothermic Loss of mass Problem from yesterday with the Co  Ni Assume: left with a nuclei Mass defect

71. Do now: What is the most likely electron configuration for a sodium ion in its ground state? 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 1 Which of the following statements is true regarding sodium and chlorine? Sodium has greater EN and a larger first IE Sodium has a larger first IE and larger atomic radius Chlorine has a larger atomic radius and greater EN Chlorine has greater EN and larger first IE

72. How to calculate photons: Every photon has a characteristic energy associated with it. The energy of a photon is dependent on its frequency. After we know the energy of a single photon, we can find out how many it takes to get the total energy of the pulse. Divided the energy of the whole pulse by the energy per photon to get the number of photons.

73. New lesson: quantum mechanics and Heisenberg Uncertainty principle Bohr proposed a model (1185 to 1962): hydrogen atom as a nucleus with an electron circling around it Certain specific orbits correspond to certain specific energy level for the electron Nik’s question Bohr only knew about the one electron, fails for atoms with more than one electron

74. Schrodinger: quantum mechanical model He said: don’t think of electrons as planet orbit the sun Think of electron travels like wave Heisenberg uncertainty principle: It is impossible to know the exact location of the electron Because electron is always in motion Just like a car moving You can’t tell

75. So if you really want to… To see electron, light photon interact with and bounce off electron Problem: Transfer energy from photon to electron Increase energy of electron making it move faster So you just change it’s (electron) position

76. Heisenberg Uncertainty principle: (∆x) (∆mv) ≥ h/(4∏) This equation: we can never know both the position and the velocity of an electron beyond a certain level of precision. -If we know velocity (certain) then look at the equation, the other term is large (uncertain) -If we know the position of electron (certain) then the velocity is large (uncertain) -Remember ∆ is the change in something so you don’t know the final and initial state

77. Let’s prove it Hydrogen: simplest model Plug in the equation: Mass of electron is 9.11 x 10 -31 kg Velocity of electron in H atom 2.2 x10 6 m/s Let’s assume we know the velocity to be within 10% then v =.2 x10 6 m/s Now plug in our equation we get??? The diameter of H atom is 240 pm so what do we conclude?

78. The uncertainty in the electron’s position is similar in size to the atom itself! Whaaaaa?

79. Example: You are traveling at a speed of 90 km/h in a car with a mass of 1250 kg. If the uncertainty in the velocity is 1%, what is the uncertainty (in meters) in the position of the car? How does this compare with the uncertainty in the position of an electron in a hydrogen atom?

80. New lesson: Wave functions and quantum numbers Schrodinger’s quantum mechanical model of atomic structure: Wave function (think of it as adding waves to get orbitals) Solution to the wave equation = wave function = orbital = (symbol) Think of wave function as expression whose square is psi 2 Defines the probability of finding electrons with in a given volume of space around the nucleus Recall the proof we did on the probability of finding electrons in a H atom???

81. Summary: Wave equation (solve)  wave function or orbital (psi)  probability of finding electron in a region of space (psi 2 ) Extra credits: why do we square the wave function to find probability???

82. Answer: Review from physic (Resnick- Halliday law) Squared amplitude of the Electric field in an electromagnetic wave gives the intensity of the wave at that point measure of concentration of the substance In quantum physic it means the probability density of finding the photon at that point Or… Born Rule: People rejected this idea at first… Read the article

83. Wave function: There are 3 variables We call them the “quantum numbers” N, l, and m l These describe the energy level of orbital Also describe the 3D shape which is the region in space occupied by a given electron

84. Let’s start with n: N is the principal quantum number: Positive n = 1, 2, 3, 4… must be whole number because energy is QUANTIZE Determine the size and energy level of orbitals Think of “n” as the shell, or layers around the atom For n = 3 means 3 rd shell Draw the Bohr model Example: Hydrogen has 1 electron so the Energy of the orbital depends only on “n” But other atoms with more than 1 electrons and the Energy depends on “n” and “l”

85. What happens when “n” increases?

86. As “n” increases: Draw the Bohr model The number of orbitals increase Size of atom larger Electrons are farther from nucleus Takes E to separate – from + So this increase distance between electron and nucleus means Energy of electron in orbital increases as quantum number “n” increases Think about the last statement, does it make sense???

87. Explanation: Energy of electron is already there (waiting to be used) so the energy is ready to overcome the threshold value. So all it takes is a little bit of external energy from the outside source to remove the outer electron Since that electron already has too much energy because it is farther from the nucleus WHAT KIND OF ENERGY ARE WE TAKING ABOUT HERE HINT: ENERGY IN AN OBJECT

88. Angular momentum quantum number “l” This tells you the 3D shape of orbitals For orbital whose principal quantum number is “n” The “l” can have integral value from 0 to n- 1 With in each shell = n different shapes orbital Basically: For example: When n = 1 then l = 0 n = 2 then l = 1, 0 n = 3 then l = 2, 1, 0

89. So… If n = shell Then “l” = subshell Which stands for those letter s, p, d, f… (sharp, principal, diffuse and fundamental) When l = 0, 1, 2, 3, 4 Sub shell notation is s, p, d, f, g So at: 0 we have s 1 we have p 2 we have d 3 we have f 4 we have g Remember this!

90. Let’s prove it:

91. Another example: At the principal quantum number n = 3 and l = 2, describe this orbital

92. Answer: 3d orbital Which means 3 = 3 rd shell L is 2 means it is the d sublevel because -2, -1, 0, 1, 2

93. Magnetic quantum number: Represent by m l : Which tells you about the spatial orientation of orbital with respect to axes For every “l” the “m l ” is –l to + l Which means each subshell (orbital with the same shape or value l) there are 2l + 1 Different spatial orientation for those orbitals (we’ll talk more about this later) Let’s do an example to be clear

94. Example: What are all of the m l when l = 0, l = 1, and l = 2

95. Let’s summarize everything in a table We call this the “table of quantum numbers” nlMlMl Orbital notation # orbital in subshell # orbital in shell 1 2 3 4

96. Do now: An atom of silicon in its ground state is subjected to a frequency of light that is high enough to cause electron ejection. An electron from which subshell of silicon would have the highest kinetic energy after ejection? 1s 2p 3p 4s The wavelength range for infrared radiation is 10 -5 while ultraviolet is 10 -8 m, which type of radiation has more energy? Why? Ultraviolet due to higher frequency Ultraviolet due to longer wavelength Infrared due to lower frequency Infrared due to shorter wavelength

97. Do now: An atom of silicon in its ground state is subjected to a frequency of light that is high enough to cause electron ejection. An electron from which subshell of silicon would have the highest kinetic energy after ejection? 1s 2p 3p 4s The wavelength range for infrared radiation is 10 -5 while ultraviolet is 10 -8 m, which type of radiation has more energy? Why? Ultraviolet due to higher frequency Ultraviolet due to longer wavelength Infrared due to lower frequency Infrared due to shorter wavelength

98. F orbital:

99. Orbital energy diagram: KJ/mol This tells you the energy level of different orbital in hydrogen atom depend only on “n” Why? Tells you that energy levels of orbital in multielectron atoms depend on “n” and “l”

100. Electrons in box: After you write your electron configuration, draw the boxes and fill in arrows. Orbitals can be represented as boxes with the electrons in them. one up arrow, one down arrow Example: for 1s orbital: 1s 2

101. Think of the quantum numbers

102. How electrons fill orbitals: Electrons fill low energy orbitals before fill higher energy ones. Why? Because they are closer to the nucleus If there is a choice between orbitals of equal energy, they fill the orbitals singly as far as possible People on the bus example

104. Let’s practice for: Hydrogen carbon

105. Practice this: Identify the shell and subshell of an orbital with the quantum number: N = 3 L = 1 M l = 1

106. Answer: 3p

107. Practice problem: Give the possible combination of quantum number for 4p orbital?

108. Answer: When n = 4 And “P” means it is -1, 0, 1 So l = 1 when m l = -1, 0, 1 n = 4 at l = 1 and m l = -1 n = 4 at l = 1 and m l = 0 n = 4 at l = 1 and m l = 1

109. Shapes of orbital:

110. Recall: shape of orbital defined by “l” when “l” = 0 is s orbital, 1 is p orbital, 2 is d orbital and so forth

111. “s” orbital: Spherical Which mean probability of finding “s” electron depends only on distance from nucleus (not on direction) Only one possible orientation of sphere in space s orbital has m l = 0 and there is only one s orbital per shell Value of psi 2 for s orbital is greatest near nucleus and drops off as distance from nucleus increase

112. Draw the “s” orbitals for 1s, 2s, 3s Show nodes and talk about probability of finding electrons… Define Node =

113. This is the s orbitals:

114. Critical thinking: How does an electron get from one region of the orbital to another if it’s not allowed to be at the node?

115. P orbital: Dumbbell shaped/figure 8 Electron distribution concentrated in identical lobes on either side of nucleus Separated by planar node cut through nucleus Draw this to illustrate

116. Probability of finding electron near nucleus is zero The two lobes of p orbital have different phases Draw this to illustrate different shading Show px, py, and pz on axis with nodal plane Only lobes of same phase can interact in forming covalent chemical bonds

117. Continue… There are: 3 allowable values of m 1 because l = 1 So we have 3 p orbitals 90 degree angles to one another along three coordinate axes x, y, z Designated by 2p x, 2p y, 2p z P orbitals in higher shells are larger than those in smaller shell and farther from nucleus Draw this to illustrate

118. D and f orbitals: For the “d” you have 5 orbitals 4 Cloverleaf and 1 donut shaped Clover leaf: Have four lobes with max electron probability separated by two nodes through nucleus Draw this Donut shaped: Draw this

119. D orbitals:

120. Summary: Nodal planes through nucleus and overall geometry increase with “l” quantum numbers: S have 1 lobe and 0 node P has 2 lobe and 1 node D has 4 lobe and 2 nodes F has 8 lobe and 3 nodes

121. Practice: How many nodal planes through the nucleus do you think g orbital has?

122. Practice: Give possible combination of n and l quantum numbers for the following fourth shell orbital:

123. Answer: N = 4 L = 2 Why?

124. Quantum mechanics and atomic spectra: Revisit Balmer-Rydberg equation and talk about m 2 and n 2

125. Electron spin and Pauli Exclusion principle Recall: What do n, l and m l tell you? Some lines of multielectron atoms occur in pairs So to add to quantum mechanics, we need another symbol: m s Electron spin

126. Continue… Electrons behave as if they spin around axis Clock wise or counter clock wise Produce magnetic field Which give “spin quantum number” m s +1/2 or -1/2

127. Why study this? Electrons occupy specific orbitals in multielectron atoms Pauli Exclusion principle: No two electrons in an atom can have the same 4 quantum numbers The set of 4 quantum numbers associated with an electron acts as a unique address for that electron in an atom No 2 electrons can have the same address

128. Illustrate this: Look at n = 2, what is the l, m l and m s ?

129. To conclude: An orbital can hold only 2 electrons, which must have opposite spins

130. Orbital energy levels in multielectron atoms: Recall that for H atom we only depend on the quantum number n Why? Now we have a situation of multielectron atoms: We have a situation of electron-electron repulsion Electron shielding

131. Z eff = Z actual – Electron shielding Where Z eff is the net nuclear charge felt by an electron, which is lower than the actual nuclear charge Talk about s orbital, p, d and Z eff trend

132. Electron configuration: This tells you which orbitals are occupied by electrons Aufbau principle: Electron must fill lowest energy level first Ground state p orbital has 6 electrons (3 sub orbitals) so the energy is the same in those 3 = degenerate Explain ground state vs. excited state Draw picture of this

133. Rules of Aufbau principle: Lower energy orbitals fill before higher Draw Energy orbital diagram to show this Orbital hold 2 electrons Must have opposite spins If 2 or ore degenerate orbitals are available, one electron goes into each until all are half full = Hund’s rule Then the second electron fill one of the orbitals Same value spin in each of the singly occupied orbitals

134. Example: people on the bus Due to electron-electron repulsion

135. Practice: Nitrogen Carbon Oxygen Fluorine neon

136. Show how they came up with valence electrons:

137. Practice examples: Show electron configurations for ions

138. Special electron configurations Look at Chromium’s predicted and actual configuration Look at copper This happens when atomic numbers are greater than 40 Where energy differences between subshells are small Transfer electron from one subshell to another lowers total energy of atom because of a decrease in electron-electron repulsions

139. Last but not least… When we talk about radius: how we can talk about definite size for atom (remember electron clouds around atoms have no specific boundaries)

140. Atomic radius: Half distance between nuclei of two identical atoms when they are bonded together Example: Cl-Cl is 198/2 = 99 C-C in diamond is 154/2 = 77 To check this you can add distance between Cl and C nuclei when those two atoms bonded In CH 3 Cl = 178

141. Finally: Now can you explain why atom size/radius gets larger going down the periodic table, and smaller going from left to right? (in terms of quantum mechanics)

142. Summary: Rule for valence-shell electron: strongly shielded by electrons in inner shells, which are closer to nucleus Less strongly shielded by other electrons in same shell, according to s>p>d>f Weakly shielded by other electrons in same subshell which are at the same distance from nucleus

143. Do now: Which of the following nuclei has 3 more neutrons than protons? 11 B 37 Cl 24 Mg 70 Ga Which of the following ions has the smallest ionic radius? Oxygen ion Fluoride ion Magnesium ion Aluminum ion Which of the following is true of the halogens when comparing them to other elements in the same period? They have larger atomic radii than other elements within their period They have lower ionization energy than elements within their period They have less peaks on PES than other elements within their period Their electronegativity is higher than other elements within their period