Passage:Electron Configuration

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Classically, electrons were thought to orbit the atomic nucleus, much like the planets around the Sun. Explaining the behavior of the electrons that "orbit" an atom was one of the driving forces behind the development of quantum mechanics as many phenomena can not be explained by the classical view. In quantum mechanics, atomic orbitals are described as wave functions over space, indexed by the quantum numbers n, l, and m.


Fundamental to the theory of quantum mechanics is the understanding that the electron is subject to both particle-like and wave-like properties. Formally, the quantum state of a particular electron is defined by its wave function, a complex-valued function of space and time. According to the Copenhagen interpretation of quantum mechanics, the position of a particular electron is not well defined until an act of measurement causes it to be detected. The probability that the act of measurement will detect the electron at a particular point in space is proportional to the square of the absolute value of the wavefunction at that point.


Electron configuration is intimately related to the structure of the periodic table. The chemical properties of an atom are largely determined by the arrangement of the electrons in its outermost "valence" shell therefore elements in the same table group are chemically similar because they contain the same number of "valence" electrons.


Quantized energy levels result from the relation between a particle's energy and its wavelength. The energy of the electron state is mainly determined by the electrostatic interaction of the (negative) electron with the (positive) nucleus. If the potential energy is set to zero when the electron is infinity far away from the nucleus, the usual convention, then a bound electron will have a negative potential. The energy level of an electron around a nucleus is given by :


E_n = - h c R_{\infty} \frac{Z^2}{n^2} \


where R_{\infty} \ is the Rydberg constant (typically between 1eV and 103 eV), Z is the charge of the atom's nucleus, n \ is the principal quantum number, e is the charge of the electron, h is Planck's constant (6.63x10-34 Js, and c is the speed of light.


Electrons are able to move from one energy level to another by emission or absorption of a quantum of energy, in the form of a photon. Because no more than two electrons may exist in a given atomic orbital an electron may only leap to another orbital if there is a vacancy there.



1. What states that no two electrons can share the same four quantum numbers?

Hund's rule
→ Hund's rule states that when there are multiple orbitals of the same energy level, electrons are added to each orbital with the same spin before they are paired up in an orbital with opposite spin.
The Pauli exclusion principle
The Aufbau principle
→ The Aufbau principle determines the filling order of electron orbitals based on the increasing energy of them.
The Rydberg equation
→ This is purely meant to confuse the reader as Rydberg is stated in the passage.

2. Selenium (Se) is likely to have which of the following properties
          i) a brittle character
          ii) a high melting point
          iii) a high ionization energy compared to other elements of the same period

i and iii.
i and ii.
ii, and iii.
iii.
→ High melting points are a character of metals, not nonmetals like selenium.

3. The electronic configuration for palladium, Pd, is?

[Kr] 5s1 4d7
[Kr] 5s2 4d8
[Kr] 5s0 4d10
[Kr] 5s2 4d10
→ Palladium has 10 valence electrons to be placed in its 5s and 4d orbitals. Because halffilled and full orbitals have a unique lowered energy, Palladium prefers to fill the 4d orbital (10 electrons) instead of filling the 5s and then partially filling the 4d.

4. How many electrons are needed to fill the 3p subshell?

2
6
8
18
→ The p subshell has three different orbitals each capable of being filled by two electrons. Thus a total of six can be placed in the 3p subshell (the 3 in 3p is meaningless in answering this problem). One should always know: s holds two electrons, p holds six, d holds 10.

5. Which of the following elements is likely to have an unusual configuration of electrons?

Fe
As
K
Cu
→ Cu exists in a location in the periodic table where it is more likely to fill the d orbital and leave the s empty. All elements in this group (column) will obey this odd pattern.

6. The energy released when an atom holds on to an additional electron is known as?

Electronegativity
→ Electronegativity describes how well an atom holds an electron in a bond.
Ionization energy
→ Ionization energy is used to describe with an electron is removed.
Electron affinity
oxidation state

7. Which of the following arrangements correctly orders the elements in increasing atomic radii?

K+, Cl, Mg, K
K+, K, Cl, Mg
Mg, Cl, K+, K
Cl, Mg, K+, K
→ In general, atoms decrease in size across a period (left to right) but increase down groups (columns). K+is equivalent to Argon in electron filling, however it has more protons, so would be just slightly smaller due to an increases pull. K on the other hand, has to put its extra electron (compared to the + cation) in new, larger shell, therefore it is going to be much larger. The other atoms, Mg, and Cl, are on the same period as Argon, therefore they will decrease in size from left to right. Thus Mg, followed by Cl, followed by the equivalent of Argon, K+ is the order for that period.

8. Atoms with different atomic numbers but the same atomic mass are said to be different _____.

ions
→ Ions are the same element but with different charge.
elements
isotopes
→ Isotopes have the same atomic number but different atomic masses and thus have a different number of neutrons.
compounds
→ Compounds are made up of multiple atoms in bonds.

9. Assuming the energy of a photon is 7.10 x 10^-18 J, what must the wavelength of this photon be?

7 nm
14.0
28.0 nm
56.0 nm

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