Electronic configuration subshells

There is no single best form of the periodic table since the choice depends on the purpose for which the table is used. Some forms emphasize chemical relations and valence, whereas others stress the electronic configuration of the elements or the dependence of the periods on the shells and subshells of the atomic structure. The most convenient form for our purpose is the so-called long form with separate panels for the lanthanide and actinide elements (see inside front cover). There has been a lively debate during the past decade as to the best numbering system to be used for the individual... 

FIGURE 2.2 When a main-group metal atom forms a cation, it loses its valence s-and p-electrons and acquires the electron configuration of the preceding noble-gas atom. The heavier atoms in Croups 1 S/lll and 14/IV retain their complete subshells of d-electrons. 

An effective way to determine the detailed electron configuration of any element is to use the periodic table to determine which subshell to fill next. Each s subshell holds a maximum of 2 electrons each p subshell holds a maximum of 6 electrons each d subshell holds a maximum of 10 electrons and each / subshell holds a maximum of 14 electrons (Table 17-5). These numbers match the numbers of elements in a given period in the various blocks. To get the electron configuration, start at hydrogen (atomic number = 1) and continue in order of atomic number, using the periodic table of Fig. 17-10. 

To determine the electron configuration in this manner, start with the noble gas of the previous period and use the subshell notation from only the period of the required element. Thus, for Fe, the notation for Ar (the previous noble gas) is included in the square brackets, and the 4s23db is obtained across the fourth period. It is suggested that you do not use this notation until you have mastered the full notation. Also, on examinations, use the full notation unless the question or the instructor indicates that the shortened notation is acceptable. 

To get the electron configuration of ions, a new rule is followed. We first write the electron configuration of the neutral atom. Then, for positive ions, we remove the electrons in the subshell with highest principal quantum number first. Note that these electrons might not have been added last, because of the n + / rule. Nevertheless, the electrons from the shell with highest principal quantum number are removed first. For negative ions, we add electrons to the shell of highest principal quantum number. (That shell has the electrons added last by the n +1 rule.)... 

A more detailed representation of the electron configuration of sodium is Na Is2 2s2 2p6 3s1 (subshell notation) instead of simply Na 2, 8, 1 (shell notation). 

Write down in full, using subshell notation, the electronic configuration of (a) the isolated fluorine atom (F has atomic number 9) and (b) the chloride ion CD (Cl has atomic number 17). (2)... 

Paramagnetism indicates unpaired electrons, which in turn are often associated with partially filled subshells. First we write the electron configurations of the elements, and then those of the ions. From those electron configurations, we determine whether the species is paramagnetic or diamagnetic. 

There are several ways of indicating the arrangement of the electrons in an atom. The most common way is the electron configuration. The electron configuration requires the use of the n and / quantum numbers along with the number of electrons. The principle quantum number, n, is represented by an integer (1,2,3. ..), and a letter represents the l quantum number (0 = s, 1 = p, 2 = d, and 3 = f). Any s-subshell can hold a maximum of two electrons, any p-subshell can hold up to six electrons, any d-subshell can hold a maximum of 10 electrons, and any f-subshell can hold up to 14 electrons. 

The even higher value of Be (greater than B) is due to the increased stability of the electron configuration of Be. Beryllium has a filled s-subshell. Filled subshells have an increased stability, and additional energy is required to pull an electron away. Give yourself 1 point for the filled subshell discussion. 

Electronic configurations are similar to electron arrangements. However, electronic configurations show the subshells that the electrons are in, whereas electron arrangements show only the number of electrons in each shell. 

For example, lithium has an electron arrangement 2,1, but its electronic configuration is Is 2s. The characters in red indicate the shell and subshell. The numbers in blue indicate the number of electrons in that subshell. So the two electrons in the first shell of lithium atoms are located in the Is subshell or Is orbital. The one electron in lithium s second shell is in the 2s subshell or 2s orbital. Now consider carbon. It has the electron arrangement 2, 4. The two electrons in the first shell go into the Is orbital. The next subshell to be filled is the 2s orbital, which holds a maximum of two electrons. The remaining two electrons go into the next available subshell, which is 2p. So carbon has an electronic configuration Is 2s 2pl... 

Likewise, the electron arrangement of sodium, 2,8,1, can be written as the electronic configuration Is 2s 2p 3s. This means that a sodium atom has two electrons in the Is subshell, two electrons in the 2s subshell, six electrons in the 2p subshell (two electrons in each of the 2p orbitals) and there is one electron in the 3s subshell. 

Consider the electronic configuration of carbon again Is 2s 2pl Remember, there are three different p orbitals in the 2p subshell the p orbital lies on the x-axis the p orbital lies on the y-axis and the p orbital lies on the z-axis. The different p orbitals are degenerate. To obey Hund s rule, these degenerate orbitals must be filled singly before spin pairing occurs. To obey the Pauli exclusion principle, when an orbital is full with two electrons, these electrons must have opposite spins. This is not shown using spectroscopic notation, but is seen when orbital box notation is used. 

For every element, the electronic configuration must agree with the electron arrangement as given in the SQA Data Booklet. Looking at the electron arrangements in the SQA Data Booklet, you can see that there should be two electrons in the 4s orbital before the 3d subshell starts to fill. You should be able to write the electronic configurations for all the elements up to krypton, atomic number 36. 

The diagram above shows the electronic configuration for carbon in orbital box notation. The two electrons in the p subshell are in different orbitals, but have parallel spins, and the electrons sharing the same orbitals in the Is and 2s subshells have opposite spins. The diagram also suggests that one of the 2p orbitals is empty. In reality, there is no such thing as an empty orbital. If an orbital is empty, then it does not exist. However, it is acceptable to show empty orbitals in this type of notation. 

This can be explained in terms of the relative stability of different electronic configurations and thus provides evidence for these electronic configurations. To help you understand this, you have to appreciate that there is a special stability associated with a filled subshell or a half-filled subshell - for example, the p subshell when it contains three or six electrons. Likewise, the d subshell is most stable when it contains five or ten electrons. The more stable the electronic configuration, then the more difficult it is to remove an electron and therefore the ionisation energy is higher. 

When we consider the electronic configurations of the elements from scandium to zinc, we are usually filling the 3d subshell according to the aufbau principle. Once again, the electronic configuration has to fit in with the electron arrangement given in the SQA Data Booklet. 

However, when any transition metal atom forms an ion, the electrons that are lost first are those in the outer subshell, the 4s electrons. Therefore the electronic configuration of the Co ion is [Ar] 3dT... 

The d block transition metals are metals with an incomplete d subshell in at least one of their ions. Try to explain why Sc and Zn are often considered not to be transition metals. Consider the electronic configurations of the Fe + and Fe ions in both spectroscopic and orbital box notations. Use these notations to explain why Fe(lll) compounds are more stable than Fe(ll) compounds. 

The electronic configuration of N is is 2s 2p and for 0 it is is 2s 2p . if you write these in orbital box notation, you will see that the electron to be removed from N is from a half-full 2p subshell. As half-full subshells are also fairly stable, then more energy is required to remove an outer electron from N than from 0. Therefore the first ionisation energy of N is slightly greater than that of 0. 

Scandium always forms Sc + ions and, because the 4s electrons are lost first, the electronic configuration of the Sc + ion is Is 2s 2p 3s 3p so there are no electrons in the d subshell. 

For brevity, many chemists record the electron configuration of an atom by giving only its outermost subshell, like As for potassium or 4/ for calcium. These electrons are most distant from the positive nucleus and, therefore, are most easily transferred between atoms in chemical reactions. These are the valence electrons. 

The same type of subshell is used to describe the electron configurations of elements in the underlying rows. (See Figure 4-4.)... 

The anomalous electronic configuration of chromium and copper is interpreted as the displacement of 1 electron from an r orbital into a d orbital these 2 elements have only 1 electron in the As subshell because the second electron was promoted into a id subshell. This example warns you that there are exceptions to the general pattern of electronic configurations of... 

You may recall from the discussion of electron transfer (see Table 4-4) that a stable configuration precisely filled an r-type subshell and a />-type subshell. Only five elements have atoms with their valence y>-subshells filled these are the inert gases in the far right column of the periodic table. Their lack of chemical reactivity is explained by their stable electron configurations. 

Symbol Nd atomic number 60 atomic weight 144.24 a rare earth lanthanide element a hght rare earth metal of cerium group an inner transition metal characterized by partially filled 4/ subshell electron configuration [Xe]4/35di6s2 most common valence state -i-3 other oxidation state +2 standard electrode potential, Nd + -i- 3e -2.323 V atomic radius 1.821 A (for CN 12) ionic radius, Nd + 0.995A atomic volume 20.60 cc/mol ionization potential 6.31 eV seven stable isotopes Nd-142 (27.13%), Nd-143 (12.20%), Nd-144 (23.87%), Nd-145 (8.29%), Nd-146 (17.18%), Nd-148 (5.72%), Nd-150 (5.60%) twenty-three radioisotopes are known in the mass range 127-141, 147, 149, 151-156. 

Electron configuration of an atom indicates its extranuclear structure that is, arrangement of electrons in shells and subshells. Chemical properties of elements (their valence states and reactivity) can be predicted from electron configuration. 

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