Chem 1 > Topic 8 Atomic Spectroscopy > Flashcards

Topic 8 Atomic Spectroscopy Flashcards

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1
Q

What is the ground state of an atom?

A

The lowest energy state of the atom

For nitrogen, this includes 2 electrons in the n=1 shell and 5 in the n=2 shell.

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2
Q

True or false: All atoms have energy levels that are occupied by electrons in the ground state.

A

FALSE

Atoms have other energy levels that exist but aren’t occupied by electrons.

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3
Q

What is the energy of bound electrons?

A

Negative energy

Bound electrons are attracted to the nucleus and have negative energy.

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4
Q

What happens to an electron that is very far from the nucleus?

A

It has no binding energy and is considered unbound

The energy of such an electron is set to 0.

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5
Q

What is the Rydberg equation used for?

A

To calculate the energy levels for one-electron species

In the equation, Z represents the number of protons and n is the principal quantum number.

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6
Q

In the Rydberg equation, what happens when n is very large?

A

Energy approaches zero

Mathematically, 1 over infinity equals zero, indicating the electron is not bound.

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7
Q

As the value of n increases, what happens to the energy of the electron?

A

It becomes less negative

The electron becomes more weakly bound.

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8
Q

What is the relationship between isoelectronic species and the number of protons?

A

Electrons are more tightly bound to the species with more protons

This results in more negative orbital energies.

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9
Q

What is the energy level of lithium ion (LI2+) compared to hydrogen in the n=1 shell?

A

Lower in energy for lithium

This indicates that lithium’s electrons are more tightly bound.

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10
Q

What happens when an electron absorbs just the right amount of energy?

A

The electron is said to be excited

This process involves the electron transitioning to a higher energy level.

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11
Q

What is the term used when an electron emits just the right amount of energy?

A

The electron is said to relax

This process involves the electron transitioning to a lower energy level.

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12
Q

What are photons scientifically known as?

A

light

Light is a form of electromagnetic radiation.

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13
Q

What is electromagnetic radiation?

A

Energy travelling through space

It encompasses various forms of light, including visible light.

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14
Q

Define wavelength.

A

Distance between hills or valleys on a wave

These are also referred to as peaks and troughs.

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15
Q

What is frequency in the context of waves?

A

Number of cycles in a given time

It indicates how often a wave oscillates.

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16
Q

Define amplitude.

A

Height of the wave measured from middle point to a peak or a trough

It reflects the energy of the wave.

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17
Q

True or false: Two waves can have the same frequency and wavelength but different amplitudes.

A

TRUE

This indicates that they can carry different amounts of energy.

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18
Q

What is the relationship between intensity of waves and amplitude?

A

Intensity is classically related to amplitude

Higher amplitude generally means greater intensity or brightness.

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19
Q

What surprising discovery was made about the speed of light in the 1880s?

A

Light has the same speed, regardless of wavelength or frequency

This challenged previous assumptions about light behavior.

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20
Q

What happens when we hit the surface of a metal with red light?

A

Nothing happens; no electrons come off

Even with high intensity red light, no electrons are ejected.

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21
Q

What occurs when we hit the surface of a metal with blue light?

A

Electrons come off

Higher intensity blue light results in more electrons being ejected, but they have the same kinetic energy as those from low intensity blue light.

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22
Q

According to Einstein, the energy of light is proportional to its frequency. What is the equation that represents this relationship?

A

E = hf

Combining with the equation relating frequency to wavelength gives E = hc/λ.

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23
Q

Why does red light not have enough energy to eject electrons from metal?

A

Red light has the lowest frequency of the visible spectrum

Higher intensity red light has greater amplitude but the same energy due to unchanged frequency.

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24
Q

How does blue light differ from red light in terms of energy?

A

Blue light has a higher frequency and thus a higher energy

This allows blue light to kick out electrons from the metal.

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25
What analogy is used to explain the concept of **light intensity**?
Light comes in little packets, like balls ## Footnote More intense light means more packets of energy (photons) knocking off more electrons.
26
What is a **photon**?
A packet of light ## Footnote The energy of the photon is proportional to its frequency.
27
How are the energy and wavelength of a photon related?
They are inversely proportional ## Footnote Higher frequency radiation has higher energy and lower frequency radiation has lower energy.
28
What happens when **electrons** in a hydrogen atom relax to lower energy levels?
A **photon** is emitted ## Footnote The energy of the photon is equal to the magnitude of the change in the energy of the electron.
29
The energy of the **photon** is equal to what?
The absolute value of the **ΔE** of the electron ## Footnote This relationship is crucial for understanding how light is emitted from excited atoms.
30
How do **neon lights** work?
* Gas tubes are plugged into a socket * Electricity excites electrons to higher states * Electrons relax and emit photons ## Footnote Different gases emit different photons, resulting in various colored neon lights.
31
What is the **Balmer series**?
The spectrum corresponding to transitions to the **n=2** energy level from higher levels. The transitions to the n = 2 level give rise to photons in the visible region of the EM spectrum. This is called the VISIBLE LINE SPECTRUM. ## Footnote This series was analyzed by scientist **Balmer**.
32
What does the **Lyman series** correspond to?
Transitions from higher energy levels to **n=1** ## Footnote In this series, hydrogen emits light in the **UV** region of the EM spectrum.
33
What does the **Paschen series** correspond to?
Transitions to the **n=3** energy level ## Footnote In this series, hydrogen emits light in the **IR** region of the EM spectrum.
34
The **line spectrum** of hydrogen corresponds to transitions to which energy level?
**n=2** ## Footnote This spectrum is visible and results from transitions from higher energy levels.
35
What is the relationship between the energy of the photon and its **wavelength**?
The energy of the photon is related to energy changes of the **electron** ## Footnote This relationship allows for the calculation of the wavelength for each energy transition.
36
What equation is used to calculate the **energy of each energy level** for one-electron atoms or ions?
Rydberg equation ## Footnote In this equation, Z equals the nuclear charge of the species in question.
37
What do the **orbitals** in an atom have?
Well defined energies ## Footnote Each orbital corresponds to a specific energy level for electrons.
38
What is the term for when electrons are in their lowest energy levels?
GROUND STATE ## Footnote This state indicates that the atom has its electrons in the lowest available energy orbitals.
39
What does **n = infinity (oo)** represent in atomic energy levels?
The highest energy level ## Footnote At this level, the electron is no longer bound to the nucleus.
40
True or false: All atoms have **higher energy orbitals** that are unoccupied by electrons.
TRUE ## Footnote This means that even in their ground state, atoms have potential energy levels that can be occupied.
41
What equation is used to calculate the **energy levels** for 1 electron atomic species?
E = -(2.178 x 10^-18 J)Z²/n² ## Footnote This equation is known as the Rydberg Equation.
42
In the Rydberg Equation, what does **Z** represent?
Z: nuclear charge of species (atomic number) ## Footnote Z indicates the number of protons in the nucleus.
43
In the Rydberg Equation, what does **n** represent?
n: the shell number ## Footnote n indicates the principal quantum number of the electron shell.
44
What happens to the energy (E) as **n** approaches infinity?
E = 0 ## Footnote As n increases, the energy levels approach zero.
45
As you get closer to the nucleus, what happens to the **attractive force**?
The attractive force gets larger ## Footnote This is due to the increased proximity to the positively charged nucleus.
46
In isoelectronic species, how does the number of protons in the nucleus affect the **binding of electrons**?
Electrons are more tightly bound to the nucleus that has more protons ## Footnote This results in energies that are more negative.
47
In the context of energy levels, what does **n = ∞** represent?
The electron is free from the nucleus ## Footnote This indicates that the electron is no longer bound to the atom.
48
What does **n = 1** represent in the context of electron energy levels?
The ground state ## Footnote n = 1 is the lowest energy level for an electron in an atom.
49
What happens to an electron when it absorbs just the right amount of **energy**?
It is excited to a higher energy level ## Footnote This transition can occur between various energy levels, such as from n = 1 to n = 2.
50
What happens to an electron when it emits just the right amount of **energy**?
It decays to a lower energy level ## Footnote This transition can occur between various energy levels, such as from n = 2 to n = 1.
51
When does **energy release** occur in relation to electrons?
When an electron goes from a higher energy shell to a lower energy shell ## Footnote Energy is released in the form of a photon.
52
What is the process called that allows us to probe **atomic structure**?
Spectroscopy ## Footnote Spectroscopy involves measuring the energy of photons.
53
True or false: The energy of a photon can be measured.
TRUE ## Footnote This measurement is a key aspect of spectroscopy.
54
The intensity of a wave is proportional to the square of its **amplitude**. What is the formula for intensity?
Intensity = (amplitude)² ## Footnote This relationship indicates how energy is carried by the wave.
55
The equation for the **speed of a wave (c)** is __________.
C = λ ν ## Footnote In this equation, C represents the speed, λ represents the wavelength, and ν represents the frequency.
56
Who formalized the idea that light can act as a **particle**?
Albert Einstein ## Footnote This concept was a significant development in understanding the nature of light.
57
What experiment did Einstein explain in 1905 that demonstrated the particle nature of light?
Photoelectric Effect ## Footnote This experiment was crucial in supporting the particle theory of light.
58
True or false: The **wave model** of light is sufficient to explain all phenomena related to light.
FALSE ## Footnote The wave model does not account for the particle-like behavior of light.
59
Who is associated with the statement: **The wave model doesn’t work**?
Max Planck ## Footnote Max Planck is a key figure in the development of quantum theory.
60
In the **photoelectric effect**, what happens when light strikes the surface of a metal?
Electrons are emitted from the surface ## Footnote The energy of these emitted electrons can be measured.
61
What can be measured in the **photoelectric effect** after electrons are emitted?
The kinetic energy (KE) of the electrons ## Footnote This measurement supports the concept of light behaving as particles.
62
What is the principle of **energy conservation** in spectroscopy?
ДЕ = 0 ## Footnote This principle states that the total energy remains constant during the process.
63
What happens when an **electron** goes from a higher energy shell to a lower energy shell?
Energy is released in the form of a photon ## Footnote This process is fundamental in understanding how light is emitted from atoms.
64
What is the **visible line spectrum** for hydrogen?
* Transition 3 to 2 * Transition 4 to 2 * Transition 5 to 2 * Transition 6 to 2 ## Footnote The transitions to the n = 2 level give rise to photons in the visible region of the EM spectrum.
65
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