Let's go through some of the Periodic Properties that are influenced directly by the electron configuration:. The size of atoms increases going down in the periodic table. This should be intuitive since with each row of the table you are adding a shell n. What is not as intuitive is why the size decreases from left to right. But again the construction of the electron configuration gives us the answer. What are you doing as you go across the periodic table?
Answer, adding protons to the nucleus and adding electrons to the valence shell of the element. What is not changing as you cross a period? Answer, the inner shell electrons. So think of it this way, the inner shell electrons are a shield against the pull of the nucleus. As you cross a period and increase the number of protons in the nucleus you increase its pull but since you are only adding electrons to the new shell the shield is not increasing but remains the same all the way across.
This means the pull on the electrons being added to the valence shell is increasing steadily all the way across. What happens if you pull harder on the electrons? Well, they come closer to the nucleus and the size of the atom decreases.
Electronegativity may be the most important of the periodic properties you can learn and understand since so many other properties are depend on its value.
Electronegativity is an atoms ability to pull electrons towards itself. Electronegativity is generally expressed by the Pauling Scale and the values were determined experimentally. The table below shows the scale values for the elements. The electronegativity values increase from left to right and bottom to top in the periodic table excluding the Noble gases. The most electronegative element is Fluorine. From these electronegativity values we can derive the patterns of two other periodic properties: Ionization Energy and Electron Affinity.
Ionization energy is the amount of energy required to remove an electron from an atom. All ionization energies are positive values because all of these removals even those for elements that form positive ions require input of energy. Electronic configurations describe electrons as each moving independently in an orbital, in an average field created by all other orbitals. From electron configuration, an atoms' reactivity and potential for corrosion can be determined.
Knowledge of the electron configuration of different atoms is useful in understanding the structure of the periodic table of elements. The concept is also useful for describing the chemical bonds that hold atoms together. In bulk materials, this idea helps explain the peculiar properties of lasers and semiconductors. The electron configuration of an atom describes the orbitals occupied by electrons on the atom.
The basis of this prediction is a rule known as the Aufbau principle, which assumes that electrons are added to an atom, one at a time, starting with the lowest energy orbital, until all of the electrons have been placed in an appropriate orbital.
The electron configuration is used to describe the orbitals of an atom in its ground state, but it can also be used to represent an atom that has ionized into a cation or anion by compensating with the loss of or gain of electrons in their subsequent orbitals.
Many of the physical and chemical properties of elements can be correlated to their unique electron configurations. The most widespread application of electron configurations is in the rationalization of chemical properties, in both inorganic and organic chemistry.
In effect, electron configurations, along with some simplified form of molecular orbital theory, have become the modern equivalent of the valence concept, describing the number and type of chemical bonds that an atom can be expected to form.
A fundamental application of electron configurations is in the interpretation of atomic spectra. The electron configuration theory was proposed by Uhlig and is an extension of the adsorption theory of passivity.
Uhlig noted that a number of transition metals become passive at certain critical compositions when alloyed with a second metal. Before continuing, it's important to understand that each orbital can be occupied by two electrons of opposite spin which will be further discussed later. The following table shows the possible number of electrons that can occupy each orbital in a given subshell.
Using our example, iodine, again, we see on the periodic table that its atomic number is 53 meaning it contains 53 electrons in its neutral state. Its complete electron configuration is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5. If you count up all of these electrons, you will see that it adds up to 53 electrons.
Notice that each subshell can only contain the max amount of electrons as indicated in the table above. The word 'Aufbau' is German for 'building up'. The Aufbau Principle , also called the building-up principle, states that electron's occupy orbitals in order of increasing energy.
The order of occupation is as follows:. Another way to view this order of increasing energy is by using Madelung's Rule :. Figure 1. Madelung's Rule is a simple generalization which dictates in what order electrons should be filled in the orbitals, however there are exceptions such as copper and chromium.
This order of occupation roughly represents the increasing energy level of the orbitals. Hence, electrons occupy the orbitals in such a way that the energy is kept at a minimum.
That is, the 7s, 5f, 6d, 7p subshells will not be filled with electrons unless the lower energy orbitals, 1s to 6p, are already fully occupied.
Also, it is important to note that although the energy of the 3d orbital has been mathematically shown to be lower than that of the 4s orbital, electrons occupy the 4s orbital first before the 3d orbital. This observation can be ascribed to the fact that 3d electrons are more likely to be found closer to the nucleus; hence, they repel each other more strongly.
Nonetheless, remembering the order of orbital energies, and hence assigning electrons to orbitals, can become rather easy when related to the periodic table. To understand this principle, let's consider the bromine atom. Since bromine has 7 valence electrons, the 4s orbital will be completely filled with 2 electrons, and the remaining five electrons will occupy the 4p orbital. Hence the full or expanded electronic configuration for bromine in accord with the Aufbau Principle is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5.
If we add the exponents, we get a total of 35 electrons, confirming that our notation is correct. Hund's Rule states that when electrons occupy degenerate orbitals i. Furthermore, the most stable configuration results when the spins are parallel i. Nitrogen, for example, has 3 electrons occupying the 2p orbital. According to Hund's Rule, they must first occupy each of the three degenerate p orbitals, namely the 2p x orbital, 2p y orbital, and the 2p z orbital, and with parallel spins Figure 2.
The configuration below is incorrect because the third electron occupies does not occupy the empty 2p z orbital. Instead, it occupies the half-filled 2p x orbital. This, therefore, is a violation of Hund's Rule Figure 2. Figure 2. A visual representation of the Aufbau Principle and Hund's Rule. Note that the filling of electrons in each orbital p x , p y and p z is arbitrary as long as the electrons are singly filled before having two electrons occupy the same orbital.
Wolfgang Pauli postulated that each electron can be described with a unique set of four quantum numbers. Therefore, if two electrons occupy the same orbital, such as the 3s orbital, their spins must be paired. The way we designate electronic configurations for cations and anions is essentially similar to that for neutral atoms in their ground state.
The electronic configuration of cations is assigned by removing electrons first in the outermost p orbital, followed by the s orbital and finally the d orbitals if any more electrons need to be removed. In this case, all the 4p subshells are empty; hence, we start by removing from the s orbital, which is the 4s orbital. Hence, we can say that both are isoelectronic. The electronic configuration of anions is assigned by adding electrons according to Aufbau Principle.
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