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Notice how the chart is now broken into a group of two, another group of two, then a group of six, another group of two and then another group of six. If you were to look at a periodic table and slide the helium over to be next to the hydrogen and then proceed to read the chart from left to right you will notice that it is also set up to be a group of two, another group of two, then a group of six, another group of two and then another group of six.
What it comes down to is that quantum numbers are actually a mathematical representation of the periodic table.
The periodic table is broken into regions according to the quantum number table that we had earlier. An example is oxygen which is the fourth member of the 2p subshell so its address is written: 2p4.
Notice how the coefficients add up to eight and the atomic number of oxygen is eight, so you know you have it set up correctly. Orbgital notation is similar to ECN except that it breaks everything into its subshells and then represents every electron with its spin. In this diagram the x represents the element symbol and the numbers are where dots would be placed to represent electrons. Write the possible sublevels for each energy level in organized rows and columns like the image below. The electronic structure (electronic configuration) of an atom refers to the arrangement of electrons around its nucleus. There is something useful to remember; the maximum number of electrons that can be contained within a Principal Quantum Level, n, is given by 2n2 (2 x n2). Immediately, this may conflict with your understanding of the concept at GCSE; in the energy level n = 3, there is a maximum of 18 electrons. The Quantum Theory page refers to the Principal Quantum Number, n, arising from a solution of the Schrodinger equation, and the quantised nature of electrons in atoms.
The above plot shows the underlying increase in first ionisation enthalpy in going from left to right across a Period of the Periodic Table.
Although the Schrodinger equation is difficult to solve except in simple cases, its results can be described.

The hydrogen atom has only one electron, but the presence of several electrons creates complex mathematical problems because electrons repel each other.
In distributing the electrons in the ground states of atoms, each added electron will enter the orbitals in the order of increasing energy. The arrangement of electrons in atoms can be investigated further and related to their positions in the periodic table via the Aufbau Principle page. Notice the small differences between the two expressions above for the electron configuration of Scandium.
The n value is called the shell number, the value of l is related to the subshell where 0=s, 1=p, 2=d, and 3=f. However, to give the full correct electron configuration you need to include all of the completed shells below it. You go to the Noble gas that occurs before the element that you are looking for and place it in brackets.
A positive spin is an upward pointing arrow and a negative spin is a downward pointing arrow. Eight electrons enter into the n = 3 level, followed by two more being placed in the n = 4 level. Something to note is that the number of sub-levels within a Principal Quantum Level is the same as the Principal Quantum Number, n.
Evidence of sub-divisions with Principal Quantum Levels is suggested by the small 'unexpected' decrease in first I.E.
Both the exact position and exact energy of electrons in atoms cannot be known at the same time. Fortunately, the similarities between many-electron atoms and the hydrogen atom are sufficiently similar that we may use orbitals like those found in the H atom to describe their electron configurations.
Since all three orientations of the p orbitals is equivalent, which p orbital (px, py or pz) is then used is not important in 5B. Also notice that the 2p orbital has two lone electrons instead of a second pair of electrons.

At first examination, the spectral line in the hydrogen atomic emission spectrum, arising from the electron transition from n = 2 to n = 1, for example, appears to be a single line.
The n = 2 level has two sub-levels, offering an explanation for the two spectral lines observed above.
The electron can therefore move anywhere relative to the nucleus without changing its energy. You could also further divide the table by separating between the 0 and 1 in the second column.
This is becuase it must satisfy Hund's rule which states I have to have one electron in every orbital in a sublevel before a second one can be placed.
But at a very high resolution it is found that two lines exist at very close frequency (or energy). A solution of the Schrodinger equation does, however, describe the probability of finding the electron in a volume of space around the nucleus.
The diagonal rule, also shown above, is useful in predicting electron configurations of atoms. For the hydrogen atom in its ground state, the space in which the electron is likely to be found can be 'pictured' as a spherical cloud of negative charge, called its atomic orbital.
Though the diagonal rule does not flow smoothly from theoretical calculations, it does account for the order of elements in the periodic table. It says that: electrons stay unpaired in orbitals of the same energy until each such orbital has at least one electron in it.
But remember, the shapes of these atomic orbitals do not represent reality; they are pictorial representations of the mathematical solutions of the Schrodinger equation.

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