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My Packs: UWO (CHEM 1301) > Chapter 2: 2.1 Atomic Structure > Flashcards

Chapter 2: 2.1 Atomic Structure Flashcards

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1
Q

Define:
Wavelength

A

The distance between two closest equivalent points of the wave

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2
Q

Define:
Frequency (in waves)

A

The number of oscillation cycles that occur at a fixed point in one second

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3
Q

What does the amplitude determine?

A

Amplitude determines the intensity

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4
Q

Value of Planck’s Constant

A

6.6260696 * 10^-34 J*s

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5
Q

How is the wave’s intensity obtained?

A

By squaring the amplitude of the wave

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6
Q

What are the wavelengths between 400-750 nm known as?

A

The visible spectrum/visible region

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7
Q

State the wavelengths of the visible spectrum

A

400-750 nm

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8
Q

How does the energy of light (photons) relate to wavelength?

A

Inverse relationship
(Longer wavelength = Lower energy, vice versa)

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9
Q

Unit for frequency

A

Hertz

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10
Q

True or False:
Energy of light can be absorbed or emitted only as small discrete portions

A

True

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11
Q

When a photon is emitted, the system _____ the energy

A

Loses

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12
Q

When a photon is gained, the system _____ the energy

A

Gains

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13
Q

What is happening when hydrogen atoms emit light?

A

The electron is emitting energy while falling back to ground state

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14
Q

What are the 3 series discovered during the research of the emission spectrum of atomic hydrogen?

A

Balmer Series (Visible)
Lyman Series (IR)
Paschen Series (UV)

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15
Q

Value of Rydberg Constant

A

10973731.6 / m

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16
Q

Ground state of Lyman emission series

A

n=1

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17
Q

Ground state of Balmer emission series

A

n=2

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18
Q

Ground state of Paschen emission series

A

n=3

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19
Q

What does m and n stand for in the Rydberg Formula?

A

m = Excited state
n = Ground state

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20
Q

True or False:
The electron does not have to return to ground state orbit

A

False, the electron has to return to lower energy orbit

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21
Q

What were the drawbacks of Bohr model?

A
  1. Doesn’t explain how the electron doesn’t fall into the nucleus
  2. Each line when enlarged is actually comprised of two very closely spaced lines
  3. Why are only certain orbits allowed?
  4. Didn’t work for atoms with more than one electron
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22
Q

Define:
Diffraction

A

Waves expanding to fill all space available (e.x. on other side of wall with a slit)

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23
Q

What is it called when waves interact with each other?

A

Interference

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24
Q

What is it known as when both waves interfere “in phase”?

A

Constructive interference

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25
Q

What is it known as when the two interacting waves are in opposite phases?

A

Destructive interference

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26
Q

How do we know electrons behave as waves, not as particles?

A

If an electron behaves as a particle, there would be a two single slits on the other side of two slits
However, there are multiple bands of varying intensity on the other side showing the wave property

27
Q

In the double-band experiment, waves in phase creates a:

A

Bright band

28
Q

In the double-band experiment, waves that are out of phase creates a:

A

Dark band

29
Q

Define:
Wave

A

A disturbance that travels through space and time, usually accompanied by a transfer of energy

30
Q

What are the two types of waves?

A

Standing
Travelling

31
Q

Travelling waves vs. Standing waves

A

Travelling: Every point undergoes oscillatory motion with the same amplitude
Standing: Different points oscillate with different amplitudes

32
Q

True or False:
All points in Standing waves all points oscillate

A

False, in Standing waves some points won’t oscillate at all

33
Q

Define:
Nodes

A

Points of a standing wave that are never displaced

34
Q

Complete the equation:
Greater number of nodes =

A

Higher the energy of the wave (think how much effort it takes to shake a rope into a wave with more nodes)

35
Q

What does the psi symbol (trident-like symbol) mean?

A

Means wavefunction, describes the wave

36
Q

de Broglie waves of individual electrons are mathematically described by:

A

Electronic wavefunctions

37
Q

What does the Schrodinger Equation calculate?

A

Calculates shape of electronic wavefunctions and their associated energies

38
Q

What are wavefunctions of individual electrons?

A

Orbitals

39
Q

True or False:
Electrons in an atom orbit the nucleus

A

False, electrons in an atom do not “orbit” the nucleus

40
Q

Finish the equation:
Greater the value of the square of the orbital at a particular point =

A

Greater the probability of finding the electron near that point

41
Q

To characterize the state of an electron, what four quantum numbers are used?

A

n
l
m(l)
m(s)

42
Q

What are the quantum numbers n, l, and m(l) also known as? Where do they come from?

A

Orbital quantum numbers
They arise in process of solving Schrodinger equation for H atom

43
Q

What does the m(s) number describe?

A

M(s) describes an intrinsic property of each electron called spin

44
Q

What is n known as? What does it determine?

A

n is known as the Principal Quantum Number
Determines the energy and size of an orbital

45
Q

Fill in the sentences:
Principal quantum numbers can assume only ________ ______ ______
Orbitals with the same n values are said to form a _____

A

Positive integer values
Shell

46
Q

What number gives the total number of nodes for the orbital?

A

The principal quantum number, n

47
Q

As n increases…

A

Size and energy of orbital increases

48
Q

State the formula for number of nodes

A

n - 1

49
Q

What is l known as? What does it determine?

A

Azimuthal Quantum Number or Orbital Angular Momentum
Describes the shape of an orbital and the angular momentum of an electron in that orbital

50
Q

True or False:
l numbers can be any positive integers

A

False, only specific values of l are allowed, ranged from 0 to n - 1

51
Q

Define:
Sublevel

A

Orbitals with particular numbers n and l

52
Q

What l values correspond to 0, 1, 2, and 3? What do these values represent?

A

Sharp, principal, diffuse, and fine
Represents shapes, introduces the spdf notation

53
Q

What is m(l) known as? What does it determine?

A

Magnetic Quantum Number
Describes the orientation of an orbital with a given l relative to the xyz axes

54
Q

What m(l) values are possible for a given l?

A

From -l, through 0, to l

55
Q

What does it mean for orbitals to be “degenerate?

A

To have the same energy

56
Q

What does each unique combination of n, l, and m(l) represent?

A

An orbital

57
Q

What is m(s) known as? What does it determine?

A

Spin Quantum Number
Associated with an intrinsic angular momentum of each electron called spin

58
Q

Define:
Spin

A

A vector that - loosely speaking - can only have 2 orientations with respect to an identified direction:
Up or Down

59
Q

What are the possible values of m(s)?

A

+1/2 or -1/2

60
Q

Describe:
Pauli Exclusion Principle

A

No two electrons in an atom can have the same four quantum numbers
Thus, if two electrons are in the same orbital they must have opposite spins

61
Q

True or False:
Wavefunctions are one dimensional

A

False. We sometimes use the one -dimensional wavefunction for illustrative purposes. However, the actual wavefunction of the electron is in 3d

62
Q

True or False:
An orbital is like a balloon or a container where the electron lives

A

False

63
Q

What is an orbital?

A

An orbital is an function. The square of this function describes how likely the electron is to be found near various points in space

64
Q

What does the square of the orbital give us?

A

The electron density, gives us the probability of finding an electron at various points

My Packs: UWO (CHEM 1301) (17 decks)

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