As you know, the size of an atom is determined by the location of its outermost electrons , that is, the electrons located on the highest energy level. Simply put, the further away the outermost electrons are from the nucleus, the bigger the atomic size. As you know, electrons are located in orbitals that correspond to various energy levels.
The higher the energy level, the further away from the nucleus the electrons it contains will be. Potassium , "K" , is located in period 4, group 1 of the periodic table , which means that it has one electron on its highest energy level. The electron configuration for a neutral potassium atom looks like this.
The size of a potassium atom is determined by how far that solitary electron that resides on the highest energy level, i. Now the outermost electrons are located on the third energy level, closer to the nucleus. Atomic radii decrease from left to right across a row and increase from top to bottom down a column.
Trends in atomic size result from differences in the effective nuclear charges Z eff experienced by electrons in the outermost orbitals of the elements.
As we described in Chapter 2 , for all elements except H, the effective nuclear charge is always less than the actual nuclear charge because of shielding effects.
The greater the effective nuclear charge, the more strongly the outermost electrons are attracted to the nucleus and the smaller the atomic radius. The atoms in the second row of the periodic table Li through Ne illustrate the effect of electron shielding.
Although electrons are being added to the 2 s and 2 p orbitals, electrons in the same principal shell are not very effective at shielding one another from the nuclear charge. In contrast, the two 2 s electrons in beryllium do not shield each other very well, although the filled 1 s 2 shell effectively neutralizes two of the four positive charges in the nucleus. Consequently, beryllium is significantly smaller than lithium. Similarly, as we proceed across the row, the increasing nuclear charge is not effectively neutralized by the electrons being added to the 2 s and 2 p orbitals.
The result is a steady increase in the effective nuclear charge and a steady decrease in atomic size. The atomic radius of the elements increases as we go from right to left across a period and as we go down the periods in a group. The increase in atomic size going down a column is also due to electron shielding, but the situation is more complex because the principal quantum number n is not constant.
As we saw in Chapter 2, the size of the orbitals increases as n increases, provided the nuclear charge remains the same. In group 1, for example, the size of the atoms increases substantially going down the column. It may at first seem reasonable to attribute this effect to the successive addition of electrons to ns orbitals with increasing values of n. However, it is important to remember that the radius of an orbital depends dramatically on the nuclear charge.
As a consequence the radii of the lower electron orbitals in Cesium are much smaller than those in lithium and the electrons in those orbitals experience a much larger force of attraction to the nucleus. That force depends on the effective nuclear charge experienced by the the inner electrons. In fact, the effective nuclear charge felt by the outermost electrons in cesium is much less than expected 6 rather than This means that cesium, with a 6 s 1 valence electron configuration, is much larger than lithium, with a 2 s 1 valence electron configuration.
The effective nuclear charge changes relatively little for electrons in the outermost, or valence shell, from lithium to cesium because electrons in filled inner shells are highly effective at shielding electrons in outer shells from the nuclear charge. The same dynamic is responsible for the steady increase in size observed as we go down the other columns of the periodic table.
Irregularities can usually be explained by variations in effective nuclear charge. Electrons in the same principal shell are not very effective at shielding one another from the nuclear charge, whereas electrons in filled inner shells are highly effective at shielding electrons in outer shells from the nuclear charge. On the basis of their positions in the periodic table, arrange these elements in order of increasing atomic radius: aluminum, carbon, and silicon. Given: three elements.
Asked for: arrange in order of increasing atomic radius. A Identify the location of the elements in the periodic table. Determine the relative sizes of elements located in the same column from their principal quantum number n.
Then determine the order of elements in the same row from their effective nuclear charges. If the elements are not in the same column or row, use pairwise comparisons. B List the elements in order of increasing atomic radius. A These elements are not all in the same column or row, so we must use pairwise comparisons.
On the basis of their positions in the periodic table, arrange these elements in order of increasing size: oxygen, phosphorus, potassium, and sulfur. An ion is formed when either one or more electrons are removed from a neutral atom cations to form a positive ion or when additional electrons attach themselves to neutral atoms anions to form a negative one.
The designations cation or anion come from the early experiments with electricity which found that positively charged particles were attracted to the negative pole of a battery, the cathode, while negatively charged ones were attracted to the positive pole, the anode.
Ionic compounds consist of regular repeating arrays of alternating positively charged cations and negatively charges anions.
As illustrated in Figure 3. A variety of methods have been developed to divide the experimentally measured distance proportionally between the smaller cation and larger anion. These methods produce sets of ionic radii that are internally consistent from one ionic compound to another, although each method gives slightly different values.
Thus despite minor differences due to methodology, certain trends can be observed. A comparison of ionic radii with atomic radii Figure 3. When one or more electrons is removed from a neutral atom, two things happen: 1 repulsions between electrons in the same principal shell decrease because fewer electrons are present, and 2 the effective nuclear charge felt by the remaining electrons increases because there are fewer electrons to shield one another from the nucleus.
Gray circles indicate the sizes of the ions shown; colored circles indicate the sizes of the neutral atoms, previously shown in Figure 3. Source: Ionic radius data from R. Cations are always smaller than the neutral atom, and anions are always larger. Because most elements form either a cation or an anion but not both, there are few opportunities to compare the sizes of a cation and an anion derived from the same neutral atom.
Anions: Non-metals tend to gain electrons to make stable anions. So in a likewise but opposite manner - we ADD electrons to the valence shell thus increasing electron repulsions which means the resulting anion is bigger than the atom from which they came. The more electrons you add, the bigger the anion gets.
This is illustrated in the diagram below starting on the left with a neutral atom. Here's a figure from Wikipedia showing the neutral atomic radii vs the ionic radii sizes for some cations and anions. The reaction with energy shown is. The energy needed to do this must overcome the attraction of the outermost electron to the nucleus.
All atoms have a wide variety of energies needed to do this, but they DO follow a trend that is easily seen on the periodic table. Much like all the trends, the two extremes of this property are at the bottom left smallest IE and the top right largest IE.
Going down a column, IE's decrease. Going across rows, IE's increase. Electron affinity is the amount of energy released when one electron is added to a neutral atom A in order to form a —1 anion.
You can think of EA as the "desire of an electron" by an atom. If the atom "wants" the electron a lot, then the EA is big. Less desire is smaller energy and there is even no desire and the numbers go to zero and even negative. The trends on the periodic table are not as pronounced as with other trends they're a bit janky - but in general, the upper right corner has the largest EAs while the lower left corner has the lowest values.
I'm including this for the purpose of pointing out this is a real measurement and the recognition of EA is more important for our studies than the actual values.
Move on to electronegativity now. Electronegativity is a relative scale from zero to four that measures the "desire" or "pull" on electron pairs.
0コメント