Spectropy

Question 1

What does the atomic emission spectrum look like ?

Answer 1

All black w/ green (to the left) and yellow (right) stripe

Question 2

What does the atomic absorption spectrum look like ?

Answer 2

COLORFUL

Question 3

Was Rutherford’s model good ? Why ?

Answer 3

No, it wasn’t sustainable (atoms would eventually blow up)

Question 4

Who told Rutherford “your model is ass” and came up with what ?

Answer 4

• Bohr (model)

Question 5

Explain energy diagrams

Answer 5

Goes from n=1 up to n=6 (increasing energy)

Question 6

In the Bohr model, what is atomic excitation vs de-excitation ? what’s happening to the photons ?

Answer 6

• Excitation: jump to higher energy level; needs energy/photons

-De-excitation: jumps to lower energy level; releases photons

Question 7

In energy diagrams, what’s it called when it goes to a lower energy level ? When it jumps to a higher one ?

Answer 7

• to lower: emitting
• to higher: absorption

Question 8

Practice: Energy Diagrams:
(what color)
n=2 –> n=1

Answer 8

yellow

Question 9

Practice: Energy Diagrams:
(what color)
n=1 –> n=3

Answer 9

green

Question 10

How to calculate when electrons jump energy levels ?

Answer 10

2nd level - 1st level

Question 11

How to calculate wavelength of electron moving ?

Answer 11

Plancks’ constant divided by (mass x velocity)

Question 12

What is Plancks’ constant ?

Answer 12

6.626 x 10^-34

Question 13

What is the mass of an electron ?

Answer 13

9.1 x 10^31

Question 14

What is the evidence of electrons as waves ?

Answer 14

electromagnetic radiation = wave so electrons = waves

Question 15

What’s the Heisenberg Uncertainty Principle ? What did it prove ?

Answer 15

• One can’t measure both the exact position and momentum of a small particle (e-).
• Proved Bohr’s model wrong

Question 16

What are the 4 quantum numbers ? explain.
DO I INCLUDE THIS OR NAH

Answer 16

1. n = energy level (all real numbers)
2. l = subshell shape (0 -> n-1)
3. ml = orientation (-l or +l)
4. ms = spin of e-
   - aligned w/ field: 1/2
   (clockwise)
   - against: -1/2 (anti-
   clockwise)

Question 17

no answer
REMEMBER: go over periodic table orbitals (1s, 2s, etc)

Question 18

What is the Aufbau principle ?

Answer 18

e- occupy lowest level orbitals first

Question 19

What is Hund’s rule ?

Answer 19

e- occupy orbitals of the same energy one by one (don’t pair up till they have to)

Question 20

How to do electron configuration ?

Answer 20

1. Look at amount of electrons (top number)
2. Use up amount keeping in mind each level’s capacity
3. fill out the arrow using Hund’s rule and the Aufabau principle

Question 21

What is the capacity for each energy level ? s ? p ? d ?

Answer 21

s: 2
p: 6
d: 10

Question 22

How to do electron configuration for ions ?
ex. S^-2

Answer 22

1. atomic # - charge (16 - (-2) = 18
2. rest, the same

Question 23

Practice: Electron configuration

N^7 ?

Answer 23

1.) e- = 7
2.) 1s^2 2s^2 2p^3
3.) two arrows (going opposite) for both 1s and 2s and 3 separate arrows for 2p (one on e/a line)

Question 24

What are core electrons ?

Answer 24

The ones closest to the nucleus

Question 25

What are valence electrons ?

Answer 25

Outside (ONLY last shell); high on energy and used for reactions

Question 26

How to find short hand electron configuration ?

Answer 26

1. Look at last noble gas before element
2. do electron configuration from gas to element you trynna find

Question 27

How many electrons can each orbital hold in a Bohr Model ?

Answer 27

2 –> 8 –> 10

Question 28

The electron configuration of an element is 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^3. How many core and valence electrons ?

Answer 28

• core: 28
• valence: 5

Question 29

Green light has a wavelength of 445 nm. What is the energy of one photon of green light ?

Answer 29

29.5 x 10^-32 = 4.47 x 10^19 J

Question 30

How to calculate the energy of one photon of light ?

Answer 30

(Planks constant(h) x c) / wavelength