2A2 Atomic Spectra and Electron Energy Flashcards
Describe the relationship between atomic spectra, electron energy levels, and electron energy transitions. (49 cards)
Define:
Ground state
The lowest energy configuration of its electrons.
Atoms are most stable in the ground state.
Define:
Electron energy level
The specific energy state an electron can occupy in an atom, determined by its distance from the nucleus.
Each electron naturally occupies the lowest available energy level (ground state, n = 1) unless it absorbs energy and transitions to an excited state (n > 1).
What is the principal quantum number?
It indicates the main energy level of an electron in an atom, determining its average distance from the nucleus.
The principal quantum number (n) starts at 1, with higher values (n = 2, 3, etc.) indicating electrons in higher energy levels.
Who proposed the Bohr model of the atom?
Niels Bohr
in 1913
The Bohr model introduced the concept of quantized energy levels, explaining why electrons occupy specific orbits instead of moving freely.
Fill in the blank:
When an electron moves to a higher energy level, it is in an _______ state.
excited
An electron absorbs energy to move to an excited state, temporarily occupying a higher energy level before returning to its ground state.
What distinguishes the valence shell in an atom?
It is the outermost electron shell of an atom, containing electrons with the highest energy that participate in chemical bonding.
Electrons in the valence shell are called valence electrons.
True or false:
Electrons can transition to any energy level without restriction.
False
Electrons can only transition between quantized energy levels by absorbing or emitting specific amounts of energy, as described by quantum mechanics.
What happens when an electron returns to its ground state?
It emits energy as light or another form of radiation.
The ground state is the lowest energy level an electron can occupy, where it is most stable.
What factor determines the wavelength of light emitted when an electron transitions between energy levels?
The energy difference between the two initial and final energy levels.
A larger energy difference results in a shorter wavelength (higher frequency), such as ultraviolet (UV) light, while a smaller energy difference produces longer wavelengths like infrared (IR) light.
Explain why electrons do not stay in excited states for long.
They lose energy quickly and return to lower energy levels.
This process stabilizes the atom.
Why are emission spectra also called bright-line spectra?
Because they appear as distinct bright lines against a dark background due to specific electron transitions releasing quantized energy.
Each bright line corresponds to a specific electron transition, with its wavelength determined by the energy difference between levels.
How do electron transitions produce the bright-line spectra of elements?
Each bright line corresponds to a specific electron transition.
These lines are unique for each element, like a spectral fingerprint, and occur when an electron moves from a higher to a lower energy level, emitting light at a precise wavelength.
True or false:
Electrons can move between energy levels without absorbing or emitting energy.
False
Energy must be absorbed or released for a transition to occur.
What happens when an electron absorbs a photon?
It can be excited to a higher energy level.
The photon must have the right energy for this to occur.
Fill in the blank:
When an electron transitions from a higher to a lower energy level, it releases a _______.
photon
The energy of the emitted photon exactly equals the energy difference between the two levels, following the equation E = hƒ (Planck’s equation)
What type of energy is typically involved in electron transitions?
Electromagnetic energy, such as light or photons.
The energy of emitted or absorbed photons corresponds to specific wavelengths in the atomic emission or absorption spectrum, forming unique spectral lines for each element.
True or false:
An electron can skip intermediate energy levels during a transition.
True
The transition depends on the energy absorbed or emitted, not the path.
True or false:
Electron transitions can only occur in visible light wavelengths.
False
Transitions can occur across the electromagnetic spectrum, including UV and infrared.
What is the significance of the Rydberg formula?
It calculates the wavelengths of light emitted or absorbed during electron transitions in hydrogen atoms.
This formula applies to transitions in hydrogen’s energy levels.
What equation relates a photon’s energy (E) to its wavelength (λ)?
E=hc/λ
Where:
λ is the wavelength,
h is the Planck constant, and
c is the speed of light.
Fill in the blank:
Electrons closer to the nucleus have _______ energy compared to electrons farther away.
lower
The attraction to the positively charged nucleus reduces the energy of inner electrons.
How do energy transitions relate to chemical reactions?
They influence bond formation, bond breaking, and the release or absorption of energy.
Electron transitions are key to understanding reactivity and energy changes.
Explain how ionization energy relates to electron transitions.
It’s the amount of energy needed to remove an electron from an atom.
This energy is sufficient to force the electron to transition beyond the atom’s highest energy level and escape its attraction to the nucleus.
What are atomic spectra?
The fingerprints of different elements, represented as an array of lines produced by light emitted or absorbed due to their chemical composition.
They are used to determine the composition of various materials, including stars.