Bohr proposed a model of the atom with a dense, positively charged nucleus surrounded by electrons revolving around the nucleus in defined pathways of distinct energy levels called orbits.
The energy of an electron is quantized, which is to say that there is not an infinite range of energy levels available to an electron. Electrons can exist only at certain energy levels, and the energy of an electron increases the farther it is from the nucleus. The energy difference between energy levels is called a quantum.
For an electron to jump from a lower energy level to a higher one, it must absorb an amount of energy precisely equal to the energy difference between the two levels. Every element has a characteristic atomic absorption spectrum. When electrons return from the excited state to the ground state, they emit an amount of energy that is exactly equal to the energy difference between the two levels. Every element has a characteristic atomic emission spectrum. Sometimes the electromagnetic energy emitted corresponds to a frequency in the visible light range.
The quantum mechanical model posits that electrons do not travel in defined orbits but rather in complex patterns called orbitals. An orbital is a region of space around the nucleus defined by the probabilities of finding an electron in that region of space. The Heisenberg uncertainty principle states that it is impossible to know at the same time both an electron’s position and its momentum.
There are four quantum numbers. These numbers completely describe any electron in an atom. The principal quantum number, n, describes the average energy of an orbital. The azimuthal quantum number, l, describes the subshells within a given principal energy level. The magnetic quantum number, ml , specifies the particular orbital within a subshell where an electron is likely to be found at a given moment in time. The spin quantum number, ms, indicates the spin orientation of an electron in an orbital.
The system of designating the placement of electrons into the principal energy levels, subshells, and orbitals is electron configuration. For example, 1s22s22p63s2 is the electron configuration for magnesium. A neutral magnesium atom has 12 electrons: two in the s-orbital of the first energy level, two in the s-orbital of the second energy level, 6 in the p-orbitals of the second energy level, and 2 in the s-orbital of the third energy level. The two electrons in the s-orbital of the third energy level are the valence electrons for the magnesium atom.
Electrons fill the principle energy levels and subshells according to increasing energy, which can be determined by the (n + l) rule. Electrons fill orbitals according to Hund’s rule, which states that electrons prefer to be unpaired with parallel spins.
Valence electrons are those electrons in the outermost shell and/or those available for interaction (bonding) with other atoms. For the representative elements, the valence electrons are found in s- and/or p-orbitals. For the transition elements, the valence electrons are found in s-, d-, and f-orbitals. Many atoms interact with other atoms to form bonds so as to complete the octet in the valence shell.
EQUATIONS TO REMEMBER
Practice Questions
1. Which of the following is the correct electron configuration for Zn2+?
A. 1s22s22p63s23p64s03d10
B. 1s22s22p63s23p64s23d8
C. 1s22s22p63s23p64s23d10
D. 1s22s22p63s23p64s03d8
2. Which of the following quantum number sets describes a possible element?
A. n = 2; l = 2; ml = 1; ms = +½
B. n = 2; l = 1; ml = -1; ms = +½
C. n = 2; l = 0; ml = -1; ms = -½
D. n = 2 ; l = 0; ml = 1; ms = -½
3. What is the maximum number of electrons allowed in a single atomic energy level in terms of the principal quantum number n ?
A. 2n
B. 2n + 2
C. 2n2
D. 2n2 + 2
4. Which of the following equations describes the maximum number of electrons that can fill a subshell?
A. 2l + 2
B. 4l + 2
C. 2l2
D. 2l2 + 2
5. Which of the following substances is most likely to be diamagnetic?
A. Hydrogen
B. Iron
C. Cobalt
D. Sulfur
6. An electron returns from an excited state to its ground state, emitting a photon at = 500 nm. What would be the magnitude of the energy change if this process were repeated such that a mole of these photons were emitted?
A. 3.98 × 10-19 J
B. 3.98 × 10-21 J
C. 2.39 × 105 J
D. 2.39 × 103 J
7. Suppose an electron falls from n = 4 to its ground state, n = 1. Which of the following effects is most likely?
A. A photon is absorbed.
B. A photon is emitted.
C. The electron gains velocity.
D. The electron loses velocity.
8. Which of the following compounds is not a possible isotope of carbon?
A. 6C
B. 12C
C. 13C
D. 14C
9. According to the Heisenberg uncertainty principle, which of the following properties of a particle can an observer measure simultaneously?
I. Position
II. Momentum
III. Velocity
A. I and II
B. I and III
C. II and III
D. I, II, and III
10. Which of the following electronic transitions would result in the greatest gain in energy for a single hydrogen electron, assuming that its ground state is n = 1?