How To Solve Electronic Configuration Of Elements
Glendora Spink
The electron configuration of an element describes how electrons are distributed in its atomic orbitals. Electron configurations of atoms follow a standard notation in which all electron-containing atomic subshells (with the number of electrons they hold written in superscript) are placed in a sequence. For example, the electron configuration of sodium is 1s22s22p63s1.
However, the standard notation often yields lengthy electron configurations (especially for elements having a relatively large atomic number). In such cases, an abbreviated or condensed notation may be used instead of the standard notation. In the abbreviated notation, the sequence of completely filled subshells that correspond to the electronic configuration of a noble gas is replaced with the symbol of that noble gas in square brackets. Therefore, the abbreviated electron configuration of sodium is [Ne]3s1 (the electron configuration of neon is 1s22s22p6, which can be abbreviated to [He]2s22p6).
This notation for the distribution of electrons in the atomic orbitals of atoms came into practice shortly after the Bohr model of the atom was presented by Ernest Rutherford and Niels Bohr in the year 1913.
It is important to note that there exist many exceptions to the Aufbau principle such as chromium and copper. These exceptions can sometimes be explained by the stability provided by half-filled or completely filled subshells.
The atomic number of hydrogen is 1. Therefore, a hydrogen atom contains 1 electron, which will be placed in the s subshell of the first shell/orbit. The electron configuration of hydrogen is 1s1, as illustrated below.
The electronic configuration of an element is a symbolic notation of the manner in which the electrons of its atoms are distributed over different atomic orbitals. While writing electron configurations, a standardized notation is followed in which the energy level and the type of orbital are written first, followed by the number of electrons present in the orbital written in superscript. For example, the electronic configuration of carbon (atomic number: 6) is 1s22s22p2.
Electron configurations provide insight into the chemical behaviour of elements by helping determine the valence electrons of an atom. It also helps classify elements into different blocks (such as the s-block elements, the p-block elements, the d-block elements, and the f-block elements). This makes it easier to collectively study the properties of the elements.
The electronic configuration of copper is [Ar]3d104s1. This configuration disobeys the aufbau principle due to the relatively small energy gap between the 3d and the 4s orbitals. The completely filled d-orbital offers more stability than the partially filled configuration.
Electron configurations are the summary of where the electrons are around a nucleus. As we learned earlier, each neutral atom has a number of electrons equal to its number of protons. What we will do now is place those electrons into an arrangement around the nucleus that indicates their energy and the shape of the orbital in which they are located. Here is a summary of the types of orbitals and how many electrons each can contain:
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. The effect of the nucleus pulling on the electrons being added across a period is called the effective nuclear charge and is calculated as ZEff = #protons - Core # Electrons.
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.
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. The more electronegative the element, the higher the ionization eneregy.
The Electron Affinity of an element is the amount of energy gained or released with the addition of an electron. The electronegativity and Electron Affinity increases in the same pattern in the periodic table. Left to right and bottom to top.
The Bohr model was a one-dimensional model that used one quantum number to describe thedistribution of electrons in the atom. The only information that was important was the sizeof the orbit, which was described by the n quantum number. Schrdinger's modelallowed the electron to occupy three-dimensional space. It therefore required threecoordinates, or three quantum numbers, to describe the orbitals in which electronscan be found.
There is only one orbital in the n = 1 shell because there is only one way inwhich a sphere can be oriented in space. The only allowed combination of quantum numbersfor which n = 1 is the following.
Before we can use these orbitals we need to know the number of electrons that canoccupy an orbital and how they can be distinguished from one another. Experimentalevidence suggests that an orbital can hold no more than two electrons.
To distinguish between the two electrons in an orbital, we need a fourth quantumnumber. This is called the spin quantum number (s) because electrons behaveas if they were spinning in either a clockwise or counterclockwise fashion. One of theelectrons in an orbital is arbitrarily assigned an s quantum number of +1/2, theother is assigned an s quantum number of -1/2. Thus, it takes three quantum numbersto define an orbital but four quantum numbers to identify one of the electrons that canoccupy the orbital.
Because of the force of attraction between objects of opposite charge, the mostimportant factor influencing the energy of an orbital is its size and therefore the valueof the principal quantum number, n. For an atom that contains only one electron,there is no difference between the energies of the different subshells within a shell. The3s, 3p, and 3d orbitals, for example, have the same energy in ahydrogen atom. The Bohr model, which specified the energies of orbits in terms of nothingmore than the distance between the electron and the nucleus, therefore works for thisatom.
The hydrogen atom is unusual, however. As soon as an atom contains more than oneelectron, the different subshells no longer have the same energy. Within a given shell,the s orbitals always have the lowest energy. The energy of the subshells graduallybecomes larger as the value of the angular quantum number becomes larger.
A very simple device can be constructed to estimate the relativeenergies of atomic orbitals. The allowed combinations of the n and l quantumnumbers are organized in a table, as shown in the figure below and arrows are drawn at 45degree angles pointing toward the bottom left corner of the table.
The order of increasing energy of the orbitals is then read off by following thesearrows, starting at the top of the first line and then proceeding on to the second, third,fourth lines, and so on. This diagram predicts the following order of increasing energyfor atomic orbitals.
The electron configuration of an atom describes the orbitals occupied byelectrons on the atom. The basis of this prediction is a rule known as the aufbauprinciple, which assumes that electrons are added to an atom, one at a time, startingwith the lowest energy orbital, until all of the electrons have been placed in anappropriate orbital.
The next element has two electrons and the second electron fills the 1s orbitalbecause there are only two possible values for the spin quantum number used to distinguishbetween the electrons in an orbital.
To answer this, we need to understand the concept of degenerate orbitals. Bydefinition, orbitals are degenerate when they have the same energy. The energy ofan orbital depends on both its size and its shape because the electron spends more of itstime further from the nucleus of the atom as the orbital becomes larger or the shapebecomes more complex. In an isolated atom, however, the energy of an orbital doesn'tdepend on the direction in which it points in space. Orbitals that differ only in theirorientation in space, such as the 2px, 2py, and 2pzorbitals, are therefore degenerate.
Electrons fill degenerate orbitals according to rules first stated by Friedrich Hund. Hund'srules can be summarized as follows.
- One electron is added to each of the degenerate orbitals in a subshell before two electrons are added to any orbital in the subshell.
- Electrons are added to a subshell with the same value of the spin quantum number until each orbital in the subshell has at least one electron.
There is something unusually stable about atoms, such as He and Ne, that have electronconfigurations with filled shells of orbitals. By convention, we therefore writeabbreviated electron configurations in terms of the number of electrons beyond theprevious element with a filled-shell electron configuration. Electron configurations ofthe next two elements in the periodic table, for example, could be written as follows.
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