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Electrons
Model of atoms

2. The Wave Nature of Light
Light was first recognized as manifestation of electromagnetic energy and it was called electromagnetic radiation
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4. Produced by motion of electrically charged particles
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Produced by motion of electrically charged particles
Move through vacuum (at x 108 m/s or 186,282 mi/hr), air and other substances
Have characteristic wavelengths/frequencies
Visible radiation has wavelengths between 400 nm (violet) and 750 nm (red)
As electromagnetic wave, light has some characteristics in common with all forms of electromagnetic energy

5. Substance or material that carries wave
Wave: disturbance of medium which transports energy without permanently transporting matter
Medium
Substance or material that carries wave
Merely carries wave from source to other location
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6. Light is a repeating waveform in motion
Rest Position: no energy present
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(amt. energy found in wave)

7. If frequency ↑, wavelength ↓
Frequency, wavelength, and velocity are inversely proportional to each other
If frequency ↑, wavelength ↓
Ex. Purple light has a frequency of x Hz.  What is its wavelength?
c = 
3.00 x 108 m/s = 7.42 x 1014 Hz ()
 = 4.04 x 10-7
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9. Photons: packets of energy that make up light
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Photons: packets of energy that make up light
Each carries specific energy related to its wavelength
Photons of short wavelength (blue light) carry more energy than long wavelength (red light) photons
Einstein successful explained photo-electric effect within context of quantum physics

10. Particulate Theory of Light
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Light is series of energy packets passing through space
Size of energy packets vary/change color of light
Quantized: electron limited to specific quantities of energy, not random value of energy
Distance between energy packets = wavelength
# photons passing point in period of time = frequency
Particulate Theory of Light

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12. One of simplest working models of atom developed by Niels Bohr
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Suggested H’s electron moves around nucleus in only certain allowed orbits
Smaller orbit, lower energy level
Larger orbit, higher energy level
Electron can have different energy levels
Ground state: lowest level
Excited state: atom gains energy
One of simplest working models of atom developed by Niels Bohr

13. Assigned quantum number, n, to each orbit
Electron in ground state (1st energy level, n = 1)
Does not radiate energy
Quantum jump: electron moves from one energy level to another by gaining energy (excited state) or losing energy (ground state) in continuously changing amounts
Electron drops from higher to lower energy orbit
Photon with specific energy emitted as light
Shown as different colored line spectrums (atomic spectrum)
Every element has its own
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14.
Atomic emission spectrum (amount of electromagnetic radiation of each frequency gas emits when heated/excited)
Photon hits metal, is absorbed as electron takes up energy
Einstein deduced each photon possesses energy
Different metals require different minimum frequencies for electrons to exhibit photoelectric effect
Above threshold frequency, # electrons ejected depend on intensity of light
If photon’s frequency below minimum, electron remains bound to metal surface
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16. 3/25/2017 Fe

17. h = Planck's constant, 6.626 x 10-34 J·s ν= frequency of radiation
Ephoton = nhν
n = # photons
h = Planck's constant, x J·s
ν= frequency of radiation
Convert wavelength from nanometers to meters:
1 x 10-9 meters = 1 nm
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18. Groups of lines observed in emission spectrum of hydrogen atoms
UV visible IR IR IR
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19. Calculate the energy of a photon of yellow light with a frequency of 5
Calculate the energy of a photon of yellow light with a frequency of 5.09 x 1014 s-1.
E= nhn = (1)(6.626 x J.s)(5.09 x 1014 s-1)= 3.37 x J
Calculate the energy of a photon of wavelength x 104 nm (infrared).
E = nhn = nhc/l = (1)(6.626 x J·s)(3.00 x 108 m/s)
(5.00 x 10-5 m)
= 3.98 x J
Calculate energy of mole of photons of yellow light with a frequency of 5.09 x 1014 s-1.
E = nhn = (6.022 x 1023)(6.626 x J·s) (5.09 x 1014 s-1)
= 2.03 x 105 J
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20. What is the frequency in hertz of blue light having a wavelength of 425 nm?
7.06 X 1014 Hz
A certain substance strongly absorbs infrared light having a wavelength of 6,500 nm.  What is the frequency in hertz of this light?
4.62 X 1013 Hz
Yellow light has a wavelength of 600 nm. What is its frequency in hertz?
5.00 X 1014 Hz
Green light has a wavelength of 550 nm. What is its frequency in hertz?
5.45 X 1014 Hz
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21. Intense microwaves have a frequency of 9. 5 X 1011 Hz
Intense microwaves have a frequency of 9.5 X 1011 Hz. What is the wavelength of these particular microwaves?
3.16 X 10-4 m  =  mm = 316 micrometers = 3.16 X 105 nm
Infrared waves can be seen if you look down the railroad tracks or a road on a hot day. They heat the air as they go past causing the air to refract or bend the light.  If infrared rays of 9.75 X 1013 Hz are being reflected off the tracks or road what will be the size of  the wavelengths in micrometers?
3.08 X 10-6 m  = 3.08 micrometers
A sunbather forgot their sunblock. On the beach they get a unheathy dose of UV radiation of 5.66 X 1016 Hz.  What is the wavelength of these particular UV waves?
5.30 X 10-9 m  = 5.00 nm
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22. Sodium vapor lamps are used to sometimes light streets
Sodium vapor lamps are used to sometimes light streets. If the frequency of the light coming from them is 5.09 X 1014 Hz what is the energy in each photon?
3.37 X J/photon
What is the energy of each photon of red light that has a frequency of 4.0 X 1014 Hz?
2.65 X J/photon
Calculate the energy in joules/photon for green light having a wavelength of 550 nm.
3.62 X J/photon
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23. Microwaves are used to heat food in microwave ovens
Microwaves are used to heat food in microwave ovens. The microwave radiation is absorbed by moisture in the food. This heats the water, and as water becomes hot, so does the food.   How many photons having a wavelength of 3.00 mm would have to be absorbed by 1.00 g of water to raise its temperature by 1oC?
6.63 X J/photon; 6.31 X 1022 photons
The wavelengths of X-rays are much shorter than those of ultraviolet or visible light. Show quantitatively why continued exposure to X-rays is more damaging than exposure to sunlight.
X-rays: 6.63 X J/photon, UV rays: X J/photon, X-rays are 100 times more powerful than UV rays.
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24. Flame Tests
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Prepare 0.5 M solutions of barium/calcium/potassium lithium/sodium/ and strontium chloride (nitrates can be used).
Fold the end of a nichrome or platinum wire into a ball and tap the straight end to a wooden stick.
Dip the end into dilute hydrochloric acid, hold it in the burner until no color shows.
Dip the end into a test tube of one of the solutions, place it in the flame, record color on chart.

25. Homework: Read 5.1, pp. 116-126 Q pg. 126, #8-10
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Read 5.1, pp
Q pg. 126, #8-10
Q pp , #33, 36, 37, 65, 66, 70, 71, 74, 76
Homework:

26. de Broglie (de-broy-lee) 1924
By the mid-1920s, scientists convinced Bohr atomic model was incorrect, formulated new explanations of how electrons arranged in atoms
de Broglie (de-broy-lee) 1924
If light could act as both particles and waves, so could electrons
Since energy E of photon equals Planck’s constant times frequency f, or E = hf, momentum p of electron would equal Planck’s constant divided by wavelength
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29. Heisenberg’s Uncertainty Principle applied de Broglie’s hypothesis
Impossible to determine with perfect accuracy both position and momentum of particle simultaneously
Making measurements on object alters location/ momentum enough to disturb accuracy of reading location/momentum
More certain we are about particle's position, less certain we are about its velocity, and vice versa
Bohr ran into trouble because he tried to predict electron’s movement too precisely
Restricting electron to certain locations and having it move in orbits violated Heisenberg Uncertainty Principle
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Heisenberg’s Uncertainty Principle applied de Broglie’s hypothesis

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31. Is light a wave or a particle?
Electromagnetic radiation has dual "personality“
Acts like waves/photons with no mass
Displays behaviors characteristic of any wave (reflection, refraction, diffraction, interference, exhibits Doppler effect) that would be difficult to explain with pure particle-view
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Is light a wave or a particle?

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34. Principles of Quantum Mechanics (Schrödinger)
Equation contains both wave and particle terms
Electrons do not have planetary orbit
Location of electron is probability, not certain position
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35. Quantum Theory
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Describes mathematically wave properties of electrons and other very small particles
Cloud shapes now called orbitals
3-D region around nucleus that indicates probable location of electron (“probability regions”)
Electrons not confined to fixed circular path

36. Quantum numbers
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Specify properties of atomic orbitals/electrons in orbitals
1st 3 from Schrödinger equation (main energy level, shape, and orientation of orbital)
4th is spin quantum number
Electrons have specific energy levels (1st, 2nd)
Different energy levels associated w/different orbits
Those nearer nucleus have lower energy than those farther away
Electrons cannot exist between energy levels

37. Quantization of energy
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Place a ball at the top of the stairs and roll it gently toward the flight of stairs.
Observe the motion and intermittent resting points of the ball as it moves down the stairs.
What is its final resting place on one step analogous to?
Toss a small ball toward the top of the stairs with as little spin as possible.
Where does it come to rest? What happens if you throw it harder (use more energy)?
What is the amount of energy you use analagous to?

38. Principal Quantum Number, n
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Main energy level occupied by electron/ size of orbital
As n becomes larger, atom becomes larger and electron is further away from nucleus
Principal Quantum Number, n

39. Cartesian coordinate system (x, y, and z axes) as frame of reference; nucleus located at origin
Boundary surface diagrams: volume of space that encloses 90% probability of finding electron within orbital’s boundary surfaces
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40. Azimuthal Quantum Number, l (angular momentum)
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Shape of cloud
Divides shells into subshells (sublevels) (l) in each principal energy level (l = n-1)
Azimuthal Quantum Number, l (angular momentum)
n = 1, 1 sublevel (s)
n = 2, 2 sublevels (p)
n = 3, 3 sublevels (d)
n = 4, 4 sublevels (f)

41. Divides subshell into orbitals which hold electrons
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Divides subshell into orbitals which hold electrons
Specifies 3-D orientation of each orbital around nucleus
Magnetic Quantum Number, ml (effect of different orientations of orbitals 1st observed in presence of magnetic field)

42. Each orbital has specific # sublevels
s has 1 sublevel
p has 3 sublevels (px, py, pz)
d has 5 sublevels (dxy, dyz, dxz, dx2-y2, dz2)
f has 7 sublevels
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43. Magnetic Quantum Number, ms (spin quantum number)
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Specifies orientation of spin axis of electron
Creates magnetic field because it spins, oriented in one of two directions
Pairs (diamagnetic) not attracted to magnets
Unpaired (paramagnetic) weakly attracted to magnets
Magnetic Quantum Number, ms (spin quantum number)

44. Each sublevel can contain maximum of two electrons
s has lowest energy (max 2 electrons)
p (max 6)
d (max 10)
f has highest energy (max 14)
Must have opposite spins
1st electron to fill orbital has a ↑/+ spin
2nd electron to fill the orbital has a ↓/- spin
You can use /, N/S, +/-
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46. Homework: Read 5.2, pp. 127-134 Q pg. 134, #13, 15, 16
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Read 5.2, pp
Q pg. 134, #13, 15, 16
Q pg. 146, #42, 45, 49, 52, 56,
Homework:

47. All orbitals related to energy sublevel are of equal energy (All three 2p orbitals are of equal energy)
Sublevels w/in principal energy level have diff. energies (Three 2p orbitals are of higher energy than 2s orbital)
In order of increasing energy, sequence of energy sublevels within principal energy level is s, p, d, and f
Orbitals within one principal energy level can overlap orbitals related to energy sublevels within another principal level (Orbital related to atom’s 4s sublevel has lower energy than five orbitals related to 3d sublevel)
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Aufbau (“building up” in German) principle: each electron occupies lowest energy orbital available

48. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
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49. Pauli exclusion principle -atomic orbital has at most 2 electrons
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No more than 2 electrons, each with opposing spin (↾⇂), can be located in energy level
No two electrons can have the same set of quantum numbers
If 1 energy level is available, then 2 electrons can be accommodated
Pauli exclusion principle -atomic orbital has at most 2 electrons

50. n l m
Subshell
notation
# orbitals in subshell
# electrons needed to fill subshell
Total # electrons in subshell 1 1s
2 (± ½) 2 2s
-1,0,1 2p 3 6 8 3s 3p
-2,-1,0,1,2 3d 5 10 18 4 4s 4p 4d
-3,-2,-1,0,1,2,3 4f 7 14 32
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Hund’s rule: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by 2nd electron, and all electrons in singly occupied orbitals must have same spin
WRONG
RIGHT

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56. Exceptions to electron configuration rules
Half filled 4s and 3d is more stable than expected electron configuration
expected 1
observed
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58. Electron Configurations of Ions
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1. Which of the following sets of atomic number and configuration represent the ground state electron configuration of an atom or ion? State which atom or ion it is.        a)   A = 8,  1s2 2s2 2p4        b)   A = 11,  1s2 2s2 2p6        c)   A = 14,  1s2 2s2 2p6 3s2        d)   A = 22,  1s2 2s2 2p6 3s2 3p6 4s2
2. Write the correct electron configurations for:        a)  Pb4+        b)  S2-        c)  Fe3+        d)  Zn2+
Electron Configurations of Ions

59. 3. Give the electron configurations for the following transition metal ions:        a)  Sc3+        b)  Cr2+        c)  Ag1+        d)  Ni3+
4. Of the following species (Sc0, Ca2+, Cl0, S2-, Ti3+), which are isoelectric?
5. Identify the group containing the element composed of atoms whose last electron:         a)  enters and fills and 's' subshell.         b)  enters but does not fill an 's' subshell.         c)  is the first to enter a 'p' subshell.         d)  is the next to the last in a given 'p' subshell.         e)  enters and fills a given 'p' subshell.         f)   is the first to enter a 's' subshell.         g)  half fills a 'd' subshell.
6. Write the electron configuration for argon. Name two positive and two negative ions that have this configuration.
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60. 6. 1s2 2s2 2p6 3s2 3p6 4s2 = Ar = S-2, Cl-1, K+1 and Ca+2
1. a)   oxygen as a neutral atom     b)   lithium as a +1 ion     c)   silicon as a +2 ion     d)   titanium as a +2 ion
2. a)  Pb4+   1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d8     b)  S2-      1s2 2s2 2p6 3s2 3p6     c)  Fe3+    1s2 2s2 2p6 3s2 3p6 4s2 3d3     d)  Zn2+    1s2 2s2 2p6 3s2 3p6 4s2 3d8
3. a)  Sc3+    1s2 2s2 2p6 3s2 3p6     b)  Cr2+    1s2 2s2 2p6 3s2 3p6 4s2 3d2     c)  Ag1+    1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s0 4d10     d)  Ni3+    1s2 2s2 2p6 3s2 3p6 4s2 3d5
4. Ca2+and  S2-have the same electronic configuration with 18 electrons each.
5.  a)  The alkali earth metals      b)  The alkali metals      c)  The boron group      d)  The halogens      e)  The noble gases      f)   The alkali metals      g)  The manganese group
6.  1s2 2s2 2p6 3s2 3p6 4s2   = Ar = S-2, Cl-1, K+1 and Ca+2
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Read 5.3, pp Q pg. 141, #25, 27 Q pp , #60, 64, 78 a/d, 79 a-d, 80 a/c/f Test practice, pg. 149, all questions Use link for quiz and submit as before. e/science.php?qi=520
Homework: