Chemistry 111 Dr. Kevin Moore

Transcription

1 Chemistry 111 Dr. Kevin Moore

2 Black Body Radiation Heated objects emit radiation based on its temperature Higher temperatures produce higher frequencies PhotoElectric Effect Light on a clean metal surface will eject electrons Atomic Emission Spectra Electrically excited atoms emit light of different frequencies

3 Electromagnetic Energy (Radiation) X-Rays Visible Light Infrared Microwaves Increasing ν & E Gamma X-Ray UV Visible Infrared Microwaves Increasing λ Radio/TV nm 1 mm 1 m

4 Frequency(ν or ƒ) - # of wave peaks that pass a point per unit time (typically seconds) Wavelength(λ) distance between two peaks on a wave λ Speed of Light (c) 3.00 x 10 8 m/sec

5 Black Body Absorbs all incident radiation Emits energy Frequency increases with temperature Ultraviolet Catastrophe Max Planck Energy emitted is an integer multiple of a tiny fraction Planck s constant h J sec

6 E h c E hc h=6.626 x J sec

7 The blue light given off by a mercury streetlamp has a wavelength of 436 nm, what is its frequency? c m/sec m sec

8 A laser emits light at a frequency of 4.69 x sec -1. What is the energy of a pulse of light containing 5.0 x photons? E ( J sec)( sec 1 ) J 16 E J 17 tot (5.010 photons) 166 photon J

9 Compare the blue glow of a Hg Lamp to the yellow glow of a Na Lamp. Which has the highest Energy? Mercury Which has the longest wavelength? Sodium

10 White Light can be broken into its component colors by a prism

11 White Light is composed of all frequencies Excited atoms give off discreet frequencies of light Sodium Yellow Hydrogen Blue Neon Red Mercury Blue Potassium Purple Lithium - Violet The lines of specific frequencies is called the Line Spectra

12 Glowing Hydrogen Lamp

13 Atoms absorb energy Released at some random point later Specific frequencies of light released We see Light Energy Released Excited State Energy Energy Ground State New Excited State

14 Excited atoms of any element give off specific frequencies of light Discovered in the mid-1800 s Scientists began categorizing all elements Johann Balmer Determined a pattern in the wavelengths of Hydrogen (m=2) R m n R nm

15 Calculate the wavelength of a Hydrogen emission in the Balmer Series when n=3. R nm 1 1 R nm nm

16 All excitations do not result in visible light Determine the wavelength of light emitted when n=4 and m=3. R nm 1 1 R nm nm

17 Shine light on a clean metal surface and electrons are ejected Frequency of light must be above a minimum level regardless of the intensity High Intensity Low Frequency Low Intensity High Frequency Metal Surface

18 Einstein said that light was like a particle of energy ( photon ) Planck s quantum concept The energy of the photon depends on the frequency, not the intensity The intensity is a measure of the # of photons which have that frequency # of electrons ejected is proportional to intensity

19 Photon carries a quantum of energy The energy needed to free the electron from metal must be matched by the photon Led Niels Bohr (with Rutherford) to propose a solar system model for a Hydrogen atom

20 Electrons lie in orbits which are energy levels (n=1, 2 ) Emission spectra is the result of an excited electron relaxing to a lower energy level Only specific energy levels (orbits) are possible hν

21 Bohr calculated the individual energy of every orbit using classical Physics E 18 ( J) n 1 2 n Quantum #

22 Difference of Energies gave Rydberg Equation Proved that the spectral lines are orbit jumps by electrons Failed to predict any other element Fundamentally flawed because it does not correlate the movement of electrons Classical Physics cannot be used to describe the atom

23 Predicted the emission spectra of Hydrogen What good is the Bohr Model? Said that the energy states were quantized Energy produced is caused by electron movement between energy levels (not really traditional orbits) 1 1 E R h H n n 2 2 i f R H =2.18x10-18 J

24 Louis de Broglie Showed that an electron has wavelike properties Used Einstein s equation for Energy and mass relationship 2 E mc Using the equation for wavelength hc hc h 2 E mc mc h, for matter, c v mc h mv

25 What is the wavelength of a golf ball (45.69 g) if it is traveling m/sec? h mv kgm sec kg( m /sec) 34 m Infinitely small

26 Uncertainty Principle Always an uncertainty in determining the position and velocity of an electron The attempt to determine the electron s position will modify its velocity ( x)( mv) h 4

27 Proposed a quantum mechanical model of the atom Supposes that the electron is a wave Moves away from the solar system concept of the atom Produces wave functions for the electron (ψ) Also called orbitals ψ 2 is the probability of finding the electron in space

28 Four Quantum numbers define each electron The address of the electron Principal Quantum # (n) Help define the size and energy of the orbital Shell Always positive integers (1, 2, 3 ) Angular Momentum Quantum # (l) Define the shape of the orbitals in a subshell Always integers from 0 to n-1

29 Magnetic Quantum # (m l ) defines the orientation (3D) of the orbital within a subshell From l to 0 to +l Spin Quantum # (m s ) defines the spin of the electron in the orbital +½ or ½ Spin up or Spin down

30 Each electron in an atom must have a unique set n=1 l=0 m l =0 m s =+½, ½ 2 possible electrons n=2 l=0 m l =0 m s =+½, ½ 2 possible electrons n=2 l=1 m l =-1, 0, 1 m s =+½, ½ 6 possible electrons

31 l=0 One possible orientation Spherical Contains 2 possible electrons, each having opposite spin s orbital l=1 3 Possible Orientations x, y, z orientations in space Each contains 2 electrons 6 total p orbital

32 l=2 5 possible orientations (m l = -2, -1, 0, 1, 2) Shaped like a 4 leaf clover 2 electrons per orbital 10 total electrons l=3 7 possible orientations (m l = -3, -2, -1, 0, 1, 2, 3) Shaped like two flowers in opposite directions 2 electrons per orbital 14 total electrons

33 n is the principal quantum number and determines the size of the subshells Energy increases with n l is the angular momentum quantum number and determines the shape of the orbital Energy increases with l s<p<d<f Two competing sources of energy for the electron Distance to the nucleus Complexity of the orbital

34 Orbital is describing the wave function of the electron Nodes exist Areas of zero probability The center of a p orbital is zero probability An s orbital has areas of zero probability inside of the sphere Zero probability is the zero amplitude position of the wave Wave has a +/- region (above the line and below the line)

35 Pauli Exclusion Principle No two electrons in an atom can have the same 4 quantum # s All orbitals in a subshell (same l value) have the same energy (degenerate) Hund s Rule Electrons occupy orbitals of equal energy with parallel spins

36 5s 4d Energy 4s 4p 3p 3d 3s 2p 2s 1s

37 Represent the configuration of electrons with the n & l values only 1s 1subshell (only subshell of the 1 st shell) 2s, 2p 2 subshells (d cannot exist in 2 nd shell) 3s, 3p, 3d 3 subshells (f cannot exist in 3 rd shell) Each shell contains s, p, d & f after the 3 rd shell Electrons completely fill orbitals in the subshell having the lowest energy 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d

38 H: 1s 1 He: 1s 2 Li: 1s 2 2s 1 O: Mg: Ga: 1s 2 2s 2 2p 4 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1

39 Every element has a core configuration which is equivalent to its closest Noble Gas 2 nd Row: [He] 3 rd Row: [Ne] 4 th Row: [Ar] Replace the core with the Noble Gas in brackets

40 Configuration initiates with the s group on the row of the Periodic Table on which the element lies Number of electrons is Z Z Noble Gas O: [He]2s 2 2p 4 Mg: [Ne]3s 2 Fe: [Ar]4s 2 3d 6 Pb: [Xe]6s 2 4f 14 5d 10 6p 2

41 Actual electron distribution within a sublevel Up arrow for spin up, Down arrow for spin down Remember Hund s Rule Fe: [Ar]4s 2 3d 6 4s 3d Ne: [He]2s 2 2p 6 2s 2 p

42 Elements prefer electrons configured in completed or half completed subshells Lowest energy point 2s 1, 2s 2, 2p 3 or 2p 6 Energy Difference between s and p is significant Energy Difference between s and d is extremely small Chromium [Ar]4s 2 3d 4 [Ar]4s 1 3d 5 4s 3d Copper [Ar]4s 2 3d 9 [Ar]4s 1 3d 10 4s 3d

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44 ns and np n=row # (n-1)d - d is always the # of the previous row (n-2)f f is always 2 rows behind f will fill before d Sn 5 th Row Element 5s, 5p 4d W 6 th Row Element 6s 4f & 5d

45 Electrons in outer shells are shielded by electrons in inner shells Less pull from the nucleus Effective Z Group electrons by shells and subshells Treat s & p together Treat d in a separate group Treat f in a separate group Example Fe: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 [1s: 2] [2s,p: 8] [3s, p: 8] [3d: 6] [4s: 2]

46 Slater s Rules Do not count the e - you are looking at s & p electrons 35% of s & p in the same group (ignore d or f) 85% of all e - in the previous shell (n-1) 100% of all e - in the shell (n-2) d & f electrons 35% of all e - in the same subshell (d ignores f) 100% of all other e -

47 Calculate the shielding (s) Z eff = Z s Which electron would be easier to remove in Fe? The 4s or the 3d? Fe: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 [1s: 2][2s,p: 8][3s,p: 8][3d: 6][4s: 2] S (3d) = 0.35*(5) *(18) = Z eff = = 6.25 S (4s) = 0.35*(1) *(14) *(10) = Z eff = = d electron is held 2x stronger than 4s

48 Shielding (Video) Electrons are pulled toward the nucleus Electrons shield one another from nucleus Electrons in the same shell are less effective (~35%) Beryllium Oxygen +4 +8

49 Going across Periodic Table Protons increase Electrons increase Poorly shielded because in the same Shell Protons are +1, while electrons are only -0.35! Z eff increases across the table! Compare Z eff of Chlorine and Sodium (outer e - ) S (Cl) = 0.35* * *2 = 10.9 Z eff = = 6.1 S (Na) = 0.85* *2 = 8.8 Z eff = = 2.2 The pull of the nucleus on the outer electrons is greater in Chlorine than in Sodium

50 Atomic Radii decrease across the Periodic Table Chlorine is almost ½ size of Sodium Cl: 99 pm Na: 186 pm Bromine is almost ½ size of Potassium Br: 114 pm K: 227 pm Radii increases down the Periodic Table A new shell is present causing a natural enlargement of the atom

51 Distance from nucleus to outermost electron Decreases across the Periodic Table Increases down the Periodic Table

52 Transition Metals d & f electrons are the most ineffective at shielding Tend to sit at the very edge of shell farthest from nucleus Size decreases until d block is close to complete 6 th row has no significant increase in size Virtually identical to 5 th row f electrons are furthest out from nucleus Poor shielding

53 Atoms lose the outermost electron first Metals ions lose outer s e - before outer d (Z eff ) Cations are smaller than the normal element Lose an e -, but not a proton Z eff increases Anions are larger than the normal element Gain an e -, but not a proton Z eff decreases

54 Electron Configuration of Ions Cations: Lose outermost shell first Regardless of the order of original filling Fe: [Ar]4s 2 3d 6 jj 4s 3d Fe +2 : [Ar]3d 6 jj 4s 3d Fe +3 : [Ar]3d 5 j 4s 3djj

55 p subshell typically gains electrons O: [He]2s 2 2p 4 j 2s 2p O -2 : [He]2s 2 2p 6 j 2s 2p C: [He]2s 2 2p 2 j 2s 2p C +2 : [He]2s 2 j 2s 2p