Atomic Orbital Energy (7.3.3) Flashcards
• When an electron is infinitely separated from a nucleus, the energy is defined to be zero.
• When an electron is infinitely separated from a nucleus, the energy is defined to be zero.
• For a hydrogen atom, the principal quantum number indicates the energy level of the electron.
• For a hydrogen atom, the principal quantum number indicates the energy level of the electron.
When an electron is infinitely separated from a
nucleus, the energy is defined to be zero.
The 1s orbital is the lowest energy state of an
electron.
As an electron moves from a lower energy level to a
higher energy level, its energy becomes closer to
zero. Energy is required to move an electron to a
higher energy level.
When an electron is infinitely separated from a
nucleus, the energy is defined to be zero.
The 1s orbital is the lowest energy state of an
electron.
As an electron moves from a lower energy level to a
higher energy level, its energy becomes closer to
zero. Energy is required to move an electron to a
higher energy level.
As the principal quantum number increases, the
average distance of an electron from a nucleus
increases. Therefore, the energy gets closer to
zero.
The energy of an electron in hydrogen (En) is equal
to negative one multiplied by the Rydberg
constant (R) multiplied by the nuclear charge (Z)
squared, divided by the principal quantum number
(n) squared. For hydrogen, the only variable is n.
Therefore, for the hydrogen atom, the principal
quantum number (n) indicates the energy level of
the electron.
The solutions to the Schrödinger equation
accurately predict the behavior of any atom that
contains only one electron (any hydrogen-like atom,
such as He+, Li2+, or Be3+). However, since these
solutions do not take into account the interactions
between electrons, they do not accurately predict
the behavior of atoms with more than one electron.
As the principal quantum number increases, the
average distance of an electron from a nucleus
increases. Therefore, the energy gets closer to
zero.
The energy of an electron in hydrogen (En) is equal
to negative one multiplied by the Rydberg
constant (R) multiplied by the nuclear charge (Z)
squared, divided by the principal quantum number
(n) squared. For hydrogen, the only variable is n.
Therefore, for the hydrogen atom, the principal
quantum number (n) indicates the energy level of
the electron.
The solutions to the Schrödinger equation
accurately predict the behavior of any atom that
contains only one electron (any hydrogen-like atom,
such as He+, Li2+, or Be3+). However, since these
solutions do not take into account the interactions
between electrons, they do not accurately predict
the behavior of atoms with more than one electron.
How does the nuclear charge relate to the size of the 1s orbital and the energy of the orbital?
The orbital size decreases and the energy of the orbital becomes more negative as the nuclear charge increases. (D)
Because the positive charge of the nucleus increases, the attraction to the electron increases, pulling the electron closer to the nucleus. This decreases the size of the orbital. The energy of an electron becomes more negative the closer it is to the nucleus. Therefore, when the orbital size decreases because of increased nuclear charge, the energy of the orbital becomes more negative.
Determine the energy of an electron that would be assigned for n = 4 for the Li 2+ ion.
−1.23 × 10^−18 J (B)
How much energy is required to move the electron of the hydrogen atom from the 1s to the 2s orbital?
1.64 × 10^−18 J (D)
What is the energy of an electron in the 2s orbital of a hydrogen atom?
−5.45 × 10^−19 J (A)
When does an electron have zero energy?
When it has been removed from the atom. (C)
When the electron has been removed from the atom we say that its energy is zero in relation to the nucleus. All of the orbitals of an atom have energies that become increasingly negative the closer the electron is to the nucleus. Only when an electron has been removed and is infinitely far from the nucleus does it have zero energy.
Determine the energy required to remove the electron from the 1s orbital of a Helium ion (He+ ).
8.72 × 10^−18 J (C)
For which of the orbitals below is the electron closest to the nucleus?
2s (A)
An electron in the n = 2 will be lower in energy than an electron in the n = 3 or the n = 4 level regardless of the type of orbital. The lowest energy level is closest to the nucleus.
Electrons in the 2s and 2p orbitals have ___ quantum numbers n and electrons in the 2s and 3s orbitals have ___ quantum numbers n.
the same; different (A)
Electrons in orbitals with the same principal quantum number have the same energy. Energies increase as the principal quantum number increases.
Determine the frequency of light required to move the electron of the Be3+ ion from the n = 1 to the n = 4 orbital. The atomic number of beryllium is 4.
4.94 × 10^16 s−1 (A)
What frequency of light would cause the electron of a Li2+ ion to be ejected from the 1s orbital? The energy needed to remove this electron is 1.96 × 10−17 J.
2.96 × 10^16 s−1 (B)
Divide the energy required to remove the electron by Plank’s constant to find the frequency of light needed to remove the electron from the ion.
As the number of protons increases…
The size of the ion will decrease, the energy will be more negative, and it will cost more energy to remove the electron from the atom.
