Electron Configuration! Chapter 5

Transcription

1 Electron Configuration! Chapter 5

2 DO NOW - Finish coloring your periodic tables! (5 min)

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4 State at Room Temperature Appearance Conductivity Malleability and Ductility Metals - solid except for mercury (a liquid) - shiny lustre - good conductors of heat and electricity - malleable - ductile Non-Metals - some gases - some solids - only bromine is a liquid - not very shiny - poor conductors of heat and electricity - brittle - not ductile Metalloids - solids - can be shiny OR dull - may conduct electricity - poor conductors of heat -brittle - not ductile

5 Helpful Videos

6 Quantum Mechanics Better than any previous model, quantum mechanics does explain how the atom behaves. Quantum mechanics treats electrons as particles that act like waves (like light waves) which can gain or lose energy. But they can t gain or lose just any amount of energy. They gain or lose a quantum of energy. A quantum is just an amount of energy that the electron needs to gain (or lose) to move to the next energy level. In this case it is losing the energy and dropping a level.

7 Atomic Orbitals Much like the Bohr model, the energy levels in quantum mechanics describe locations where you are likely to find an electron. Remember that orbitals are geometric shapes around the nucleus where electrons are found. Quantum mechanics calculates the probabilities where you are likely to find electrons.

8 Atomic Orbitals Of course, you could find an electron anywhere if you looked hard enough. So scientists agreed to limit these calculations to locations where there was at least a 90% chance of finding an electron. Think of orbitals as sort of a "border for spaces around the nucleus inside which electrons are allowed. No more than 2 electrons can ever be in 1 orbital. The orbital just defines an area where you can find an electron.

9 Energy Levels Quantum mechanics has a principal quantum number. It is represented by a little n. It represents the energy level n=1 describes the first energy level n=2 describes the second energy level etc. Red Orange Yellow Green Blue Indigo Violet n=1 n=2 n=3 n=4 n=5 n=6 n=7 Each energy level represents a period (row) on the periodic table.

10 Sub-levels = Specific Atomic Orbitals Each energy level has 1 or more sub-levels which describe the specific atomic orbitals for that level. Blue = s block Yellow = p block Red = d block Green = f block n = 1 has 1 sub-level (the s orbital) n = 2 has 2 sub-levels (s & p) n = 3 has 3 sub-levels (s, p, & d) n = 4 has 4 sub-levels (s, p, d & f) There are 4 types of atomic orbitals: s, p, d and f Each of these sub-levels represent the blocks on the periodic table.

11 Orbitals d s p In the s block, electrons are going into s orbitals. In the p block, the s orbitals are full. New electrons are going into the p orbitals. In the d block, the s and p orbitals are full. New electrons are going into the d orbitals.

12 Energy Level Sub-levels Total Orbitals Total Electrons Total Electrons per Level n=1 s 1 (1s orbital) 2 2 n=2 s p 1 (2s orbital) 3 (2p orbitals) n=3 s p d 1 (3s orbital) 3 (3p orbitals) 5 (3d orbitals) Complete the chart in your notes as we discuss this. n=4 s 1 (4s orbital) 2 32 Thep first level (n=1) has an s orbital. It has only 1. 3 (4p orbitals) 6 There orbitals in the first d are no5other (4d orbitals) 10 energy level. f 7 (4f orbitals) 14 We call this orbital the 1s orbital.

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14 Where are these Orbitals? 1s 2s 3s 4s 5s 6s 7s 2p 3p 3d 4d 5d 6d 4f 5f 4p 5p 6p 7p

15 Electron Configurations What do I mean by electron configuration? The electron configuration is the specific way in which the atomic orbitals are filled. Think of it as being similar to your address. The electron configuration tells me where all the electrons live.

16 Rules for Electron Configurations 3 rules govern electron configurations: Aufbau Principle Pauli Exclusion Principle Hund s Rule Using the orbital filling diagram at the right will help you figure out HOW to write them Start with the 1s orbital. Fill each orbital completely and then go to the next one, until all of the elements have been accounted for.

17 Fill Lower Energy Orbitals FIRST Each line represents an orbital. 1 (s), 3 (p), 5 (d), 7 (f) High Energy The Aufbau Principle states that electrons enter the lowest energy orbitals first. The lower the principal quantum number (n) the lower the energy. Within an energy level, s orbitals are the lowest energy, followed by p, d and then f. F orbitals are the highest energy for that level. Low Energy

18 Write the electron config. below for each of these elements: Sodium: Iron: Bromine: Barium:

19 Write the electron config. below for each of these elements: Sodium: 1s22s22p63s1 Iron: 1s22s22p63s23p64s23d6 Bromine: 1s22s22p63s23p64s23d104p5 Barium: 1s22s22p63s23p64s23d104p65s24d105p66s2

20 No more than 2 Electrons in Any Orbital ever The next rule is the Pauli Exclusion Principal. The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full. The spins have to be paired. We usually represent this with an up arrow and a down arrow. Since there is only 1 s orbital per energy level, only 2 electrons fill that orbital. Quantum numbers describe an electrons position, and no 2 electrons can have the exact same quantum numbers. Because of that, electrons must have opposite spins from each other in order to share the same orbital.

21 Hund s Rule Hund s Rule states that when you get to degenerate orbitals, you fill them all half way first, and then you start pairing up the electrons. What are degenerate orbitals? Degenerate means they have the same energy. Don t pair up the 2p electrons until all 3 orbitals are half full.

22 DO NOW 1. Please take your packet out for me to check 2. Go over packet 3. Start lab activity

23 Noble Gas Notation

24 Let s try it out! NOW that we know the rules, we can try to write some electron configurations. Lets write some electron configurations for the first few elements, and let s start with hydrogen.

25 Electron Configurations Element Configuration Element Configuration H Z=1 1s1 He Z=2 1s2 Li Z=3 1s22s1 Be Z=4 1s22s2 B Z=5 1s22s22p1 C Z=6 1s22s22p2 N Z=7 1s22s22p3 O Z=8 1s22s22p4 F Z=9 1s22s22p5 Ne Z=10 1s22s22p6 (2p is now full) Na Z=11 1s22s22p63s1 Cl Z=17 1s22s22p63s23p5 K 1s22s22p63s23p64s1 Sc Z=21 1s22s22p63s23p64s23d1 1s22s22p63s23p64s23d6 Br Z=35 1s22s22p63s23p64s23d104p5 Z=19 Fe Z=26 Note that all the numbers in the electron configuration add up to the atomic number for that element. Ex: for Ne (Z=10), = 10

26 One last thing. Look at the previous slide and look at just hydrogen, lithium, sodium and potassium. Notice their electron configurations. Do you see any similarities? Since H and Li and Na and K are all in Group 1A, they all have a similar ending. (s1)

27 Electron Configurations Element Configuration H Z=1 1s1 Li Z=3 1s22s1 Na Z=11 1s22s22p63s1 K 1s22s22p63s23p64s1 Z=19 This similar configuration causes them to behave the same chemically. It s for that reason they are in the same family or group on the periodic table. Each group will have the same ending configuration, in this case something that ends in s1.

28 Quantized Energy

29 Electromagnetic Radiation Visible light is a type of electromagnetic radiation A form of energy that exhibits wavelike behavior through space. Electrons are particles that act like waves. Electrons that transition from excited to ground state has a different amount of energy and a different wavelength of light.

30 Electromagnetic Wave Relationship Amplitude is the wave s height from the origin to a crest

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32 Electromagnetic Spectrum The electromagnetic spectrum (EM spectrum) includes all forms of electromagnetic radiation, with the only difference in types of radiation being frequencies and wavelengths.

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35 Electromagnetic Wave Relationship Frequency and wavelength are inversely proportional c= c: speed of light (3.00 x 108 m/s) : wavelength (m, nm, etc.) x v v: frequency (Hz)

36 Lets try a problem... Calculate the wavelength of a wave that has the frequency of 3.44 x 109 Hz.

37 Calculate the wavelength of a wave that has the frequency of 3.44 x 109 Hz. c= x v = c/v 8 = 3.00 x 10 m/s x 10 Hz Hz = S -1

38 Calculate the wavelength of a wave that has the frequency of 3.44 x 109 Hz. 8 = 3.00 x 10 m/s x 10 s -2-1 = 8.72 x 10 m

39 Energy of Quantum A quantum is the minimum amount of energy that can be gained or lost by an atom. Max Planck studied light by heating objects. These different colors correspond to different frequencies and wavelengths. Equantum = h x v E: Energy (Joules or kj) h: Planck s constant x J x S v: frequency (Hz or s-1)

40 What is a photon? A massless particle that carries a quantum of energy. They can be absorbed or released by electrons!

41 Photoelectric Effect When a photon hits an electron on a metal surface, the electron can be emitted. The emitted electrons are called photoelectrons.

42 Lets try a problem... What is the energy of a photon from the violet portion of the Sun s light if it has a frequency of x 1014 s-1?

43 Lets try a problem... What is the energy of a photon from the violet portion of the Sun s light if it has a frequency of x 1014 s-1? Equantum = h x v *Unknown is E E = (6.626 x J.s) (7.230 x 1014 s-1) E = x J

44 Atomic Emission Spectrum

45 Flame Test

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47 Atomic Emission Spectra A set of frequencies of the electromagnetic waves emitted by atoms of the element. Each element is unique!

48 Periodic Trends

49 Periodic Trends More than 20 properties change in predictable way based location of elements on PT Some properties: Density Melting point/boiling point Atomic radius Ionization energy Electronegativity

50 Going down group 1 Increasing # of energy levels... Period Element Configuration 1 H 1 2 Li Na K Rb Cs Fr

51 Increasing number of energy levels

52 Atomic Radius How large the radius is of an atom is what do you think influences that?

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54 Increasing Atomic Radius Increasing number of energy levels

55 Cs has more energy levels, so it s bigger next Li: Group 1 Period 2 Cs: Group 1 Period 6

56 As we go across, elements gain electrons, but they are getting smaller! Family IA or 1 IIA or 2 IIIA or 13 IVA or 14 VA or 15 VIA or 16 VIIA or 17 VIIIA or 18 Element Li Be B C N O F Ne Configuration

57 Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius

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59 Why does this happen? As you go from left to right, you again more protons (the atomic number increases) You have greater proton pulling power Remember the nucleus is + and the electrons are - so they get pulled towards the nucleus The more protons your have, the more Proton Pulling Power

60 Ionization Energy Amount energy required to remove a valence electron from an atom in gas phase 1st ionization energy = energy required to remove the most loosely held valence electron (e- farthest from nucleus)

61 previous index next Cs valence electron lot farther away from nucleus than Li electrostatic attraction much weaker so easier to steal electron away from Cs THEREFORE, Li has a higher Ionization energy than Cs

62 Increased Ionization Energy (harder to remove an electron) Increasing number of energy levels Increasing Atomic Radius Increased Electron Shielding Decreasing Atomic Radius Decreased Ionization Energy (easier to remove an electron)

63 Electronegativity ability of atom to attract electrons in bond noble gases tend not to form bonds, so don t have electronegativity values Fluorine: most electronegative element = 4.0 Paulings

64 Increasing number of energy levels Increasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Increased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius Decreased Electronegativity

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66 Electronegativity ability of atom to attract electrons in bond noble gases tend not to form bonds, so don t have electronegativity values Fluorine: most electronegative element = 4.0 Paulings