Quantum Theory

Front 1

Atoms and molecules are governed by same or different laws?

Back 1

Atoms and molecules are not governed by the same physical laws as larger objects

Front 2

Max Planck discovered?

Back 2

Energy is continous, and Atoms and molecules emit energy only in certain discrete quantities or quanta.

Front 3

What is a wave?

Back 3

vibrating disturbance by which energy is transmitted

Front 4

How do water waves travel?

Back 4

Wave repeats itself at regular intervals.

Front 5

Waves can be characterized by?

Back 5

length, height, and number of waves per second

Front 6

What is wavelength?

Back 6

distance between identical points on successive waves

Front 7

What is frequency?

Back 7

number of waves that pass through a particular point in 1 second

Front 8

What is amplitude?

Back 8

vertical distance from the midline of a wave to the peak or trough

Front 9

What is speed?

Back 9

Depends on the type of wave and nature of the medium. Speed is a product of wavelength and frequency
u = yV

Front 10

What is wavelenght and frequency measured in?

Back 10

m, cm, nm.
1 Hz = 1 cycle/second

Front 11

Visbile light is made up of?

Back 11

Electromagnetic waves

Front 12

What are Electromagnetic waves?

Back 12

Waves with an electrical field component and a magnetic field component

Front 13

The two components of electromagnetic waves differ in?

Back 13

Same speed, but travel in perpendicular planes

Front 14

What is electromagnetic radiation?

Back 14

Emission and transmission of energy in the form of electromagnetic waves

Front 15

What speed do electromagnetic waves travel at?

Back 15

3.00 * 10^8 m/s

Front 16

Long waves

Back 16

Are emitted from large antennas (radio, cell phones). Radio waves lowest frequency

Front 17

Short waves

Back 17

Have higher energy radiation. Gamma rays have the shortest wavelenght and highest frequency

Front 18

What is quantum?

Back 18

smallest quantity of energy that can be emitted or absorbed in the form of electromagnetic radiation

Front 19

What is the wavelength (in m) of an electromagnetic wave whose frequency is 3.64 * 10^7 Hz?

Back 19

ANSWER

Front 20

What is the frequency (in Hz) of a wave whose speed is 713 m/s and wavelength is 1.14 m?

Back 20

ANSWER

Front 21

What is Planck's constant?

Back 21

E=hv
H= 6.63 * 10^-34

Front 22

What is the photoelectric effect?

Back 22

Electrons are ejected from the surface of certain metals exposed to light of at least a certain minimum frequency called the
Threshold frequency

Front 23

Number of electrons depends on?

Back 23

intensity of light but energy of electrons does not

Front 24

Einstein theorized light is made up of?

Back 24

A stream of particles called photons.

Front 25

What do you need to break electrons free from a metal?

Back 25

Requires light of sufficiently high frequency

Front 26

What is KE and BE?

Back 26

E = hn = KE + BE
KE = kinetic energy of electron
BE = binding energy of the electron in the metal
The higher the frequency the greater the KE

Front 27

KE is dependent on?

Back 27

Frequency of the light. Ejected

Front 28

Light behaves as?

Back 28

Both as a particle and wave depending on the property being measured. All matter actually exhibits this dual nature.

Front 29

The energy of a photon is 5.87 * 10-20 J. What is the wavelength in nm? h = 6.63 * 10-34 J*s

A photon has a wavelength of 624 nm. Calculate the energy of the photon in J.

Back 29

ANSWER

Front 30

What is an emission spectra?

Back 30

Either continous or line spectra of radiation emitted by substances

Front 31

What is a line spectra?

Back 31

The light emmission only at specific wavelenghts. Produced by atoms

Front 32

Continous

Back 32

Emission spectra of the sun or heated solids

Front 33

Why is Bohr's model no accurate?

Back 33

Because it does not explain the spectral lines

Front 34

What is Rydberg's constant?

Back 34

Rydberg constant (RH) = 2.18 * 10^-18 J
Energy of electron

Front 35

Where are free electrons?

Back 35

infinitely far from the nucleus

Front 36

Why is there a negative sign in equation of Rydberg's constant?

Back 36

Negative sign – is to assign a lower energy of electron in an atom than the energy of the free electron (arbitrarily assigned a value of zero)

Front 37

What happens as the electron gets closer to the nucleus?

Back 37

Becomes more stable and E becomes more negative, which corresponds to the most stable state.

Front 38

What is the ground state?

Back 38

The lowest energy state of a system or most stable

Front 39

What is the excited state?

Back 39

Higher energy than the ground state

Front 40

When is radiant energy emitted?

Back 40

When electrons drop from a higher energy orbital to a lower energy orbital

Front 41

The quantity of energy produced is dependent only on what?

Back 41

Initial and final states
If an electron starts at ni and drops to a lower energy state of nf the change in energy is given by Equation=

Front 42

What happen in the states when energy is given off.

Back 42

Ni>Nf change in energy is negative

Front 43

Each line on the emission spectrum corresponds to?

Back 43

To a transition in the H atom
When a large number of H atoms are examined all the lines of the spectrum are visible

Front 44

Electrons bound to a nucleus behave like?

Back 44

Electrons bound to a nucleus behave like a standing wave
Waves that can be generated by plucking a string

Front 45

Some points on a standing wave are

Back 45

Nodes
Do not move at all
The amplitude at this point is zero
Nodes are located at the end of the string and maybe in the middle

Front 46

De Broglie says that if an electron does not behave like a wave then?

Back 46

De Broglie said that if an electron does behave like as wave then the wave must perfectly fit the circumference of the orbit
The circumference of the orbit is related to the wavelength by the equation
=

Front 47

When can a particle be a wave and a wave be a particle?

Back 47

A particle in motion can be treated as a wave
A wave can also exhibit properties of a particle

Front 48

Protons can be accelerated to speeds near the speed of light in particle accelerators. Estimate the wavelength (in nm) of such a proton moving at 2.90 * 108 m/s. (mp = 1.673 * 10-27 kg)
A baseball has a mass of about 255 g. Calculate the wavelength of the baseball if it is thrown 100. mph.

Back 48

Answer=

Front 49

What is Heisenberg uncertainity principle?

Back 49

It is impossible to know simultaneously both the momentum p and the position of a particle

Front 50

Applying the Heisenberg uncertainty principle to the H atom?

Back 50

We see that the electron can not orbit the nucleus in a circular orbital.
If this were the case then we could know both the position and momentum of the electron at the same time.

Front 51

Erwin schrodinger formulated what?

Back 51

Formulated an equation to describe the behavior and energies of submicroscopic objects
This equation is very complicated and requires Calculus to solve

Front 52

The equation incorporated?

Back 52

The Schrödinger equation incorporates particle behavior of electrons in the form of mass and wave behavior in the form of wave function (y)

Front 53

Why is the wave funciton significant?

Back 53

Is that the square of the wave function (y2) is proportional to the probability of where the electron is located

Front 54

Where is an electron most likely to be?

Back 54

The most likely place an electron will be is where y2 is greatest

Front 55

What tells Schrodinger equation tell us?

Back 55

The Schrödinger equation gives possible energy states and identifies the wave function of the electrons.
These are characterized by quantum numbers. Quantum mechanics gives probability of an electron in a particular region of the atom (electron density)

Front 56

In quantum mechanics the orbits are called?

Back 56

Atomic orbitals

Front 57

What is an atomic orbital?

Back 57

The wave function of an electron in an atom.
This is to differentiate from the orbits in Bohr’s model.
Each atomic orbital has a characteristic energy and therefore a characteristic distribution of electron density.

Front 58

An assumptions must be made

Back 58

The difference between hydrogen and atoms with more than one electron is not that large.

Front 59

What are quantum numbers?

Back 59

Describe the distribution of electrons in hydrogen and other atoms.

Front 60

What are three quantum numbers that describe the distributioni of electrons?

Back 60

Principal quantum number
Angular momentum quantum number
Magnetic quantum number

Front 61

What is the fourth quantum number that describes the behavior of a specific electron?

Back 61

Spin quantum number

Front 62

Principal quantum numbers

Back 62

Represented by n
Can have values of integers
Relates the average distance from the electron to the nucleus in a particular orbital.

Front 63

Larger the n what happens?

Back 63

The greater distance of an electron in the orbital from the nucleus and therefor the larger the orbital.

Front 64

Angular Momentum Quantum Number

Back 64

Represent by l
Tells the shape of the orbital
Dependent on n
For any n, l = any integer from 0 to (n-1)
For n = 1, l = 0
For n = 3, l = 0, 1, or 2
l is normally designated by the letter that symbolizes the different atomic orbitals

Front 65

Magnetic quantum number

Back 65

Represented by ml
Depends on l
For any value l, there are (2l +1) values for ml
If l = o then ml = 0
If l = 1 then there are three possible values of ml, (-1, 0, 1)
The value of ml indicates the number of orbitals in the subshell with value l

Front 66

Spin quantum numbers

Back 66

Represented by ms
It was noticed that the application of a magnetic field could split the lines in an emission spectra
The only way this could be explain is if electron behave as tiny magnets
If the electrons are thought of as spinning on their own then the magnetic field can be explained
The spinning charge generates a magnetic field

Front 67

Spin quantum numbers are always?

Back 67

–½ or +½

Front 68

Atomic orbital related to angular momentum

Back 68

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Front 69

What is a shell?

Back 69

A collection of orbitals with the same value of n

Front 70

What are subshells?

Back 70

Orbitals with the same n and l value are called subshells
For example n = 2
Two subshells: l = 0 and l = 1
Called the 2s and 2p subshells

Front 71

What is the shape of the orbitals?

Back 71

Shapes are not well defined
Wave function defining orbital extends from the nucleus to infinity
It is convenient to think of the orbitals as having a shape
It is especially helpful when talking about chemical bonds

Front 72

What shape and size are s orbitals?

Back 72

All s orbitals are spherical
The s orbitals do change in size however
The increase in size is the reason for the increase in the principal quantum number

Front 73

Electron density

Back 73

Falls off rapidly as the electron get farther from the nucleus.

Front 74

p orbitals start with the principal quantum number? What happen when n=2, and l=1?

Back 74

2
three possible orbital so there are 3 2p orbitals

Front 75

What are the shape, and size of orbitals?

Back 75

They are oriented along the axes of a 3-d plot
The three orbitals differ only in the orientation
They are identical in shape, size and energy
p orbital are thought of as two lobes on either side of the nucleus

Front 76

d orbitals

Back 76

n must equal at least 3 and l must equal 2 for the d orbitals to exist

Front 77

How many d orbitals are there?

Back 77

There are 5 d orbitals
The orbitals differ in orientation as well as one having a different shape.
All 3d orbitals have the same energy
d orbitals which have a larger n value have similarly shaped orbitals that are larger

Front 78

Why are f orbitals important?

Back 78

are important for accounting for the behavior element with atomic number > 57
In this class we are not concerned with orbitals having l > 3

Front 79

Give the quantum numbers associated with the following orbitals:
2p
3s
5d

Back 79

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Front 80

Energy of orbitals increase as?

Back 80

n increases
1s<2s=2p<3s=3p=3d<4s=4p=4d=4f<…

Front 81

Does electron density change for 2s and 2p orbitals?

Back 81

Although the electron density is different for 2s and 2p orbitals the energy remains the same

Front 82

Which is orbital is the most stable?

Back 82

1s

Front 83

The total energy depends on what?

Back 83

The total energy depends on the sum of orbital energies as well as the repulsive forces
It turns ends up that the total energy is lower when the 4s orbital fills before the 3d

Front 84

Order of atomic orbitals filling

Back 84

1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p

Front 85

2s quantum numbers

Back 85

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Front 86

Ms

Back 86

has no affect on energy, size, shape, or orientation of the orbital but determines how electrons are arranged in an orbital

Front 87

What is electron configuration?

Back 87

How the electrons of an atom are distributed among the orbitals
This is how the electrons are distributed in a ground state atom

Front 88

The number of electrons in an atom is equal to?

Back 88

Atomic number

Front 89

Ground State H

Back 89

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Front 90

Pauli exclusion principle?

Back 90

No two electrons in the same atom can have the same four quantum numbers.
If they have the same n, l, and ml, then they must have different ms.
These two electrons would be in the same orbital but have opposite spins.
Thus each orbital can contain only two electrons

Front 91

Paramagnetism

Back 91

Contain net unpaired spins and are attracted by a magnet

Front 92

Diamagnetism

Back 92

Do not contain unpaired spins and are slightly repelled by a magnet. A He atom with opposite spins in the orbital

Front 93

What happens if the spins in an orbital do match up?

Back 93

The magnetic fields reinforce each other. This would make an atom paramagnetic.

Front 94

Odd and even numbered atoms have?

Back 94

Odd numbered atoms always have one or more unpaired electrons
Even numbered atoms may or may not have unpaired electrons

Front 95

Which orbital is filled first?

Back 95

The 1s orbital is filled before electrons are start to fill the 2s or 2p orbital

Front 96

2s and 2p orbitals

Back 96

Both the 2s and 2p orbitals have electrons that spend more time away from the nucleus than electrons in the 1s orbital
The electrons of the 2s and 2p orbitals are shielded from the attractive forces of the nucleus by the 1s electrons
This reduces the electrostatic interactions between the nucleus and the 2s and 2p electrons

Front 97

Shielding

Back 97

Experimentally the 2s orbital gives us a lower energy than the 2p
Although the 2s electron spend more time on average farther from the nucleus than a 2p electron; the denisty near the nucleus is greater for a 2s electron
So the 2s orbital is more penetrating
The 2s orbital is less shielded
For the same values of n, the penetrating power decreases as l increases

Front 98

How is the stability of the electron is determined?

Back 98

The stability of the electron is determined by the strength of attraction to the nucleus

Front 99

Shielding effect

Back 99

2s orbitals are less shielded than 2p orbitals so it follows that 2s orbitals have lower energy. Less energy is required to remove a 2p electron than a 2s electron

Front 100

Hund's Rule

Back 100

The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins

Front 101

Rules for assigning electrons to orbitals

Back 101

Each shell (or principal quantum number) n contains n subshells
Each subshell consists of quantum number l contain (2l + 1) orbitals
No more than 2 electrons can be placed in an orbital
The maximum number of electrons in principal level n is equal to 2n2

Front 102

How many electrons can be present in the prinicpal level n=4?
What are the principal quantum numbers for the last electron in boron (B)?

Back 102

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Front 103

Aufbau principle

Back 103

As protons are added one by one to the nucleus to build up elements, electrons are similarly added to the atomic orbitals

Front 104

Noble gas core

Back 104

a method of showing electron configurations where the noble gas most nearly preceding the element being considered is used.The noble gas is followed by the electron configuration of the most highly filled subshells

Front 105

Transition metals have either

Back 105

Have either incompletely filled d orbitals or give rise to cations that have incompletely filled d subshells

Front 106

Two irregularities in the fourth preiod

Back 106

Chromium – [Ar]4s13d5
Copper – [Ar]4s13d10
The reason for this is that there is actually more stability in a half-filled or filled d orbital

Front 107

Lanthanides and Actinide

Back 107

Lanthanides – have incompletely filled 4f orbitals or readily give rise to cations with incompletely filled 4f subshells
Actinide series – last row of elements, most are not found in nature but have been synthesized

Front 108

Write the ground-state electron configuration for Sr.
Write the ground-state electron configuration for Ga.

Back 108

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