• For your test, may want to learn to write down the following chart from memory: During the test, when you are given the atomic number and asked for an electron configuration, you can use the chart to write down the filled levels and "subtract off" the electrons for that level from the atomic number, until you have used all of the electrons. Here is an example: Let's try the atomic number 14: Look at the chart and see that 1s2 is first so you write down: 1s2 then subtract 2 from 14 14 - 2 = 12 Look at the chart and see that 2s2 is next so write down 2s2 after the 1s2: 1s2 2s2 Then subtract 2 from 12 10 - 2 = 10 Look at the chart and see that 2p6 is next so write down 2p6 after the 2s2: 1s2 2s2 2p6 Then subtract 6 from 10 10 - 6 = 4 Look at the chart and see that 3s2 is next to write down 3s2 after 2p6 is next: 1s2 2s2 2p6 3s2 Then subtract 2 from 4 4 - 2 = 2 Look at the chart and see that 3p6 is next but you only have 2 electrons left so you write 3p2 after 3s2: 1s2 2s2 2p6 3s2 3p2 This is the electron configuration for atomic number 14. Because you have used all of the electrons, you stop.
  • 1) "Electron configuration was first conceived of under the Bohr model of the atom, and it is still common to speak of shells and subshells despite the advances in understanding of the quantum-mechanical nature of electrons. An electron shell is the set of allowed states electrons may occupy which share the same principal quantum number, n (the number before the letter in the orbital label). An electron shell can accommodate [2*n^2] electrons, i.e. the first shell can accommodate 2 electrons, the second shell 8 electrons, the third shell 18 electrons, etc. The factor of two arises because the allowed states are doubled due to electron spin—each atomic orbital admits up to two otherwise identical electrons with opposite spin, one with a spin +1/2 (usually noted by an up-arrow) and one with a spin -1/2 (with a down-arrow). A subshell is the set of states defined by a common azimuthal quantum number, l, within a shell. The values l = 0, 1, 2, 3 correspond to the s, p, d, and f labels, respectively. The number of electrons which can be placed in a subshell is given by [2*(2^l + 1)]. This gives two electrons in an s subshell, six electrons in a p subshell, ten electrons in a d subshell and fourteen electrons in an f subshell. The numbers of electrons that can occupy each shell and each subshell arise from the equations of quantum mechanics, in particular the Pauli exclusion principle, which states that no two electrons in the same atom can have the same values of the four quantum numbers." Source and further information: 2) "Each shell is composed of one or more subshells, which are themselves composed of atomic orbitals. For example, the first (K) shell has one subshell, called "1s"; the second (L) shell has two subshells, called "2s" and "2p"; the third shell has "3s", "3p", and "3d"; and so on. The various possible subshells are shown in the following table: Subshell label â„“ Max electrons Shells containing it Historical name s 0 2 Every shell sharp p 1 6 2nd shell and higher principal d 2 10 3rd shell and higher diffuse f 3 14 4th shell and higher fundamental g 4 18 5th shell and higher The first column is the "subshell label", a lowercase-letter label for the type of subshell. For example, the "4s subshell" is a subshell of the fourth (N) shell, with the type ("s") described in the first row. The second column is the azimuthal quantum number of the subshell. The precise definition involves quantum mechanics, but it is a number that characterizes the subshell. The third column is the maximum number of electrons that can be put into a subshell of that type. For example, the top row says that each s-type subshell ("1s", "2s", etc.) can have at most two electrons in it. The fourth column says which shells have a subshell of that type. For example, looking at the top two rows, every shell has an s subshell, while only the second shell and higher have a p subshell (i.e., there is no "1p" subshell). The final column gives the historical origin of the labels s, p, d, and f. They come from early studies of atomic spectral lines. The other labels, namely g, h and i, are an alphabetic continuation following the last historically originated label of f." Source and further information:
  • Electronic Structure of Atoms Each electron in an atom is described by four different quantum numbers. Three of these quantum numbers (n, l, and m) represent the three dimensions to space in which an electron could be found. A wave function for an electron gives the probability of finding the electron at various points in space. A wave function for an electron in an atom is called an atomic orbital. The fourth quantum number (ms) refers to a certain magnetic quality called spin. n-The Principal Quantam Number The n quantam number relates to the size of the atomic orbital. n can have any positive integer value from 1 to 7. The smaller the n, the lower the energy, the higher the value of n, the higher the energy. In the case of any single-electron atom, or hydrogen atom, n is the only quantum number which determines the energy. The size of an orbital depends on n. The larger the orbital, the larger the value of n. Orbitals of the same quantum state belong the the same shell. To use an analogy for n, why not relate it to the size of a computer, where larger values would represent larger houses. Computer l-The angular momentum quantum number l can have any integer value from 0 to 3. This quantum number distinguishes orbitals of a given n value which have different states. Or, the secondary quantum number gives the shape of the orbital so the analogy can be made to the shape of the computer with larger values associated with computers with more components. Computer M-magnetic quantum number The third quantum number has to do with the orientation of an orbital in a magnetic field. Because of this, we can relate its values to different directions the computer might be facing. Computer The final quantum number is the spin quantum number, it describes the spin orientation of an electron. The electron configuration of an atom is the particular distribution of electrons among available shells. It is described by a notation that lists the subshell symbols, one after another. Each symbol has a subscript on the right giving the number of electrons in that subshell. For example, a configuration of the lithium atom (atomic number 3) with two electrons in the 1s subshell and one electron in the 2s subshell is written 1s22s1. sublevel orbital maximum # of electrons s 1 2 p 3 6 d 5 10 f 7 14 The notation for electron configuration gives the number of electrons in each subshell. The number of electrons in an atom of an element is given by the atomic number of that element. On the left we have a diagram to show how the orbitals of a subshell are occupied by electrons. On the right there is a diagram for the filling order of electrons in a subshell. fillingorder" Here are some examples that show how to use the filling order diagram to complete the electron configuration for a certain substance. Element # of Electrons in Element Electron Configuration He 2 1s2 Li 3 1s22s1 Be 4 1s22s2 O 8 1s22s22p4 Cl 17 1s22s22p63s23p5 K 19 1s22s22p63s23p64s1 explanation Often times you will be asked to find the electron configuration for something that looks like this: 53I The 53 denotes the number of electrons in an atom of iodine. You would now proceed to do the electron configuration by looking at the filling order chart. With increasing atomic number, the electron configuration of the atoms display a periodic variation. Because of this the elements show periodic variations of both physical and chemical behavior. The periodic law is a law stating that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. We are going to be looking at three physical properties of an atom: atomic radius, ionization energy, and electron affinity. Atomic Radius The size of the electron cloud increases as the principal quantum number increases. Therefore, as you look down the periodic table, the size of atoms in each group is going to increase. When you look across the periodic table, you see that all the atoms in each group have the same principal quantum number. However, for each element, the positive charge on the nucleus increases by one proton. This means that the outer electron cloud is pulled in a little tighter. One periodic property of atoms is that they tend to decrease in size from left to right across a period of the table. So finally we have a good definition for how the atomic radii increases: the atomic radii increases top to bottom and right to left in the periodic table. Ionization Energy The energy needed to remove the most loosely held electron from an atom is known as ionization energy. Ionization energies are periodic. The ionization energy tends to increase as atomic number increases in any horizontal row or period. In any column or group, there is a gradual decrease in ionization energy as the atomic number increases. Metals typically have a low ionization energy. Nonmetals typically have a high ionization energy. Electron Affinity The attraction of an atom for an electron is called electron affinity. Metals have low electron affinities while nonmetals have high electron affinities. The general trend as you go down a column is a decreasing tendancy to gain electrons. As you go across a row there is also a trend for a greater attraction for electrons.

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