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CHEM1110 > SECTION 1 - QUANTUM MECHANICAL MODEL OF ATOM > Flashcards

SECTION 1 - QUANTUM MECHANICAL MODEL OF ATOM Flashcards

1
Q

Quantum mechanics

A

Model that describes how electrons exist in atoms and helps explain chemical and physical properties of the elements

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2
Q

Parts of an atom:

A
  • Proton (p)
  • Neutron (n)
  • Electron (e-)
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3
Q

Electromagnetic radiation

A

Tool to understand electronic structure in atom and consists of a wave of energy

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4
Q

Frequency

A

Number of cycles per second (s-1) (v)

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5
Q

Wavelength

A

Distance the waves travels during one cycle (nm, m , um) (lamda)

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6
Q

Amplitude

A

Height of crest of each wave

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7
Q

Relationship between frequency and wavelength

A

Inversely proportional. Greater frequency = small wavelength, small frequency = big wavelength)

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8
Q

What type of ER has the longest wavelength (smallest frequency)

A

Radio

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9
Q

What type of visible light has the shortest frequency

A

Red

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10
Q

What type of ER has the shortest wavelength (biggest frequency)

A

Gamma ray (produced by sun, stars, etc)

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11
Q

Speed of light

A

2.998 x 10^8 ms-1

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12
Q

Relationship between visible light and brightness/color

A

Amplitude determines the brightness and frequency/wavelength determines color

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13
Q

Interference

A

Phenomenon in which 2 waves combine by adding or subtracting waves

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14
Q

In phase (constructive interference)

A

Waves are in line and combine together = higher amplitude, higher intensity (bright)

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15
Q

Out phase (destructive interference)

A

Waves are not in line and subtract one another = no amplitude, no intensity (dark)

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16
Q

Wave-particle duality of light

A

Waves can exhibit particle like properties and particles can exhibit wave-like properties (photons)

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17
Q

Photoelectric effect

A

Albert Einstein demonstrated that light has properties of a particle. The experiment consisted of the emission of electrons when electromagnetic radiation hits a metal

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18
Q

2 reasons why light is not proven as a wave (photoelectric effect) (2)

A
  • Light had to be a certain minimum frequency or no electron was ejected (if light was a wave, any frequency of light would work)
  • There was no lag time and the current flowed as soon as light was hit at the metal (if light was a wave, dim light would have taken a while to have enough energy to kick out the electron)
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19
Q

Photon

A

Term by Einstein of a particle of light and viewing it as a “packet” of electromagnetic radiation and has characteristics of Plank’s constant

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20
Q

Planck’s constant

A

6.626 x 10^-34

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21
Q

Blackbody radiation

A

Conducted by Max Planck. Blackbody is an idealized object that absorbs all electromagnetic radiation it comes into contact with. It then emits thermal radiation in continuous spectrum according to temperature. Planck theorized that energy emitted by hot object is in the form of packets (photon)

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22
Q

Ephoton formula

A

h = c/lamda

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23
Q

What is the energy of a packet dependant on

A

Dependant on frequency of light. Photon of the right frequency hits an electron on the surface and supplies the right amount of kinetic energy to kick electron

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24
Q

Kinetic energy formula

A

KE = energy of photon - binding energy

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25
Atomic spectra
The spectrum of the electromagnetic radiation emitted or absorbed by an electron during transitions between different energy levels within an atom. Viewable via prism
26
Niels Bohr
Developed a model for the atom to explain the atomic line spectra for the H atom
27
What did Bohr popstulate (3)
- Electrons can move in a circular orbit (wrong) - Electrons have a fixed set of allowed orbits (energy states) at fixed distances from nucleus with principle quantum number n (correct) - Electrons can jump from one allowed orbit to another with a certain energy
28
Energy of electron in an orbit formula
En = -2.179 x 10^-18J (z^2/n^2)
29
Excitation
Electrons absorbs energy and gains enough energy to move to a higher orbit
30
Relaxation
Electron releases energy (emitted) to go to a lower orbit
31
Ground state
Lowest energy level (n=1)
32
Excited state
Other energy levels (n = or 2 or greater)
33
Difference in energy (delta E = Ef - Ei) formula
(-2.179 x 10^-18 J)(z^2/n^2f - z^2/n^2i)
34
Ionization energy
Minimum amount of energy needed to remove 1 electron from a gaseous ground state atom to an infinite distance
35
Louis De Broglie
Suggested that small particles like electrons have wave length properties (contrary to Einstein)
36
What is the wavelength dependant on
Mass and velocity
37
De Broglie relation formula
lamda = h/m x v
38
Mass of electrons
9.1094 x 10^-31 Kg
39
Heisenburg uncertainty principle
States that we cannot know both the position and speed of a particle (photon/energy) with perfect accuracy
40
Heisenburg uncertainty principle formula
Delta x = m delta v -> h/4pi
41
Relationship between position/velocity and uncertainty
Inversely proportional (the more uncertain in position, the smaller uncertainty in velocity)
42
Schroiders equation
Describes the probability of finding an electron in an atom
43
Dual nature of electrons
Electrons in hydrogen atoms are seen as standing waves due to their dual nature (similar to light)
44
Consequences of Schroiders equation (3)
- Electron energy is quantized (specific allowed values) - Each energy level has 1 or more associated wave functions - Squaring wave function gives high probability map of electron position in 3 dimensional space (orbital)
45
3 quantum numbers:
- Principle quantum number (n) - Angular momentum quantum number (l) - Magnetic quantum number (mI)
46
Principle quantum number (4)
- Indicates size and orbital energy - Integer - Larger n = larger orbital - As n increases, so does the distance from nucleus and energy increase
47
Angular momentum quantum number (3)
- Indicates orbital space - Integer - Every number from 0 up to n-1
48
Magnetic quantum number (3)
- Determines orientation of orbital in space - Integer from -l through 0 to +l - The number of ml = number of orbitals
49
How are shapes of orbitals calculated
Based on the probability of finding an electron in a specific location of space outside nucleus 90% of the time
50
S orbital
Spherical shape and size increases as n increase
51
Radial/spherical nodes
Circular regions of 0 probability of finding electrons in 2s and 3s orbitals (# of radial nodes = n-l-1)
52
P orbital
First occurs in level n=2 and has 3 different orientations that are the same size, shape, energy
53
What is found at the nucleus of every p orbital
Angular node
54
Spin orientation quantum number (ms)
Electron spin orientation and is +1/2 if arrow is up and -1/2 if arrow is down
55
Degenerate
Orbitals with the same energy
56
Effective nuclear charge Zeff:
Actual amount of positive charge experienced by an electron in a multi electron atom (Zeff = Z - S). Z is atomic number, S is core electrons
57
3 rules for electron configuration:
- Aufbau principle - Pauli exclusion principle - Hunds rule
58
Aufbau principle
Orbitals of lowest energy filled first
59
Pauli exclusion principle
Each electron in atom has a unique set of 4 quantum numbers meaning only 2 electrons placed in orbital must have opposing spins
60
Hunds rule
Electrons singally filled in each orbital of same energy before pairing (easier to decrease repulsion)
61
Valence electrons
Electrons in the largest and least stable occupied orbitals (outer electrons)
62
Core electrons
lectrons in the smaller and more stable occupied orbitals (inner electrons). Core electrons have same configuration as preceding noble gas in table
63
Short hand notation
Using noble gas to represent core electron configuration with valence electrons (eg. [He])
64
Important exceptions for transition metals
Transition metals have general configuration of 4s2 3d# but electrons can be removed from 4s to satisfy full shell of d
65
Periodic table blocks (4)
- S block: alkali metal + alkali earth - P block: basic metal + metalloid + nonmetal + halogen + noble gas - D block: transition metal - F block: lanthanoids + actinoids
66
Diamagnetic atoms
Atom/ion with all electrons paired and has no magnetic properties
67
Paramagnetic atoms
Atom/ion with unpaired electrons and shows net magnetic properties. The amount of paramagnetism = number of unpaired electrons
68
Main group metals tend to ___ enough electrons to adopt the previous noble gas configuration
Lose
69
Main group nonmetals tend to ___ enough electrons to adopt the next noble gas configuration
Gain
70
Transitional metal ion configurations
Transition metals first lose ns electron before d electrons to form cations
71
How do we perceive color
Through the presence of a variety of wavelengths in white line is how we perceive color in objects
72
X ray
Passes through substances, then blocks visible light and used to image bones and internal organs. Can cause cancer
73
UV
In between x ray and visible light and is the component of sunlight that produces burns or tans. Can cause cancer
74
Visible light
Ranges from violet (shortest wavelength, higher energy) to red (longest wavelength, lower energy). Impacts certain molecules in our eyes to change shape and sends signal to our brains that results in our ability to see
75
Infrared radiation
The heat that is felt when hand is placed near hot object
76
Microwave
Efficiently absorbed by water and can heat substances that contain water
77
Radio waves
Transmits signals responsible for AM and FM radio, telephone, tv
78

CHEM1110 (18 decks)

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