If you have read and understood the previous post about Z, you will now find that the other four trends we mentioned are all quite intuitive. Let's consider them one at a time.
Atomic Radius: Remember that the size of the nucleus of an atom is extremely small relative to the size of the entire atom. So when we talk about an atom's size, what we are really considering is the size of its electron cloud.
Radius of Neutral Atoms: As we move from left to right across the periodic table, we find that Z is increasing. So what exactly does this mean? Well, as Z increases, the net positive pull of the nucleus increases. (Z is, after all, the number of protons that are not needed to cancel out with the core electrons.) As we increase the net positive charge of the nucleus, we will also increase its ability to pull on the electron cloud, even the valence electrons. (This should make sense to you from Coulomb's Law, also; if you increase the magnitude of Q, you also increase the magnitude of F between the two charges.) So as Z increases from left to right, those nuclei will pull their electrons in more and more tightly to the nucleus, and this will decrease the overall size of the atom. Thus, the smallest atoms are found on the right side of the table, and the largest are on the left.
As we go down a group, Z doesn't change very much. But what does change? Well, we are adding a new principle quantum number each time we move down a row, which means that we've added more core electrons. So even though we aren't changing the net pull of the nucleus, we ARE changing the distance between the nucleus and the valence electrons. In fact, we are greatly increasing it. (From Coulomb's Law, we know that increasing the distance between two charges decreases the force between them.) So as we move down the periodic table, we will continue to add more layers of electrons without significantly changing Z, making the atom grow larger and larger.
Radius of Cations: You may be asked on your test to compare the radius of a cation with that of a neutral atom. The most important thing that removing an electron does is that it decreases electron-electron repulsion within a shell. Recall from the previous post that we are assuming that valence electrons don't repel one another, but in reality, they actually do a little bit. So when you remove an electron to form a cation, you don't change your value of Z, but you do decrease repulsion among the remaining valence electrons, which allows the nucleus to pull them in even tighter. Thus, cations are always smaller than the neutral atom.
This effect is especially pronounced if you remove an electron from a Group I element, since you lose an entire shell when you do this. Thus, Li+ cation is MUCH smaller than neutral Li.
Radius of Anions: The effect of adding an electron to a neutral atom is the opposite of the effect of removing one. Again, you are not changing Z, but you ARE increasing repulsion among the valence electrons, and so the net effect is to increase the size of the atom. If you add an electron to a group VIII element, this effect is especially pronounced, since you will be adding another whole quantum level to the atom.
Ionization Energy (IE): IE is the energy required to remove an electron from a lone atom in the gaseous state. Note that the atom MUST be in the gas phase, and it MUST be alone (i.e., not bonded to anything else).
As we have seen, the value of Z increases from left to right as we go across the periodic table. This means that the atoms are able to hold on to their valence electrons more and more tightly. Thus, it gets harder to remove an electron as the value of Z increases, and IE increases from left to right.
Going down a group, Z doesn't change much. But, the size of the atom increases greatly, and because of this, so does the distance between the nucleus and the valence electrons. That decreases the magnitude of the force of attraction between the nucleus and the outermost electrons. Thus, IE decreases going down a group from top to bottom.
You may also get asked on your test to compare first and second IE values. In general, it always takes more energy to remove a second electron than it does to remove the first. This is because you are trying to remove an electron from an already positively charged species, and make it have a positive charge of even higher magnitude. So second IE is always greater than first IE, third IE is greater than second IE, and so on.
Electron Affinity (EA) and Electronegativity (EN): These two concepts are similar to one another, and they follow the same trends. EA measures the energy released when a lone gaseous atom accepts an electron. It is the opposite of IE, and as with IE, the atom must not be in a bond, and it must be in the gas phase. Note that EA values are given as the absolute values of the energy released. (That is, we usually report exothermic reactions, where energy is being released, as having a negative energy value. However, EA is always reported as the absolute value of that negative energy magnitude, and so EA values are always positive.) EN measures the ability of an atom in a bond with another atom to pull the bond's electron density toward itself.
As we travel from left to right across the periodic table, and Z increases, the ability of the atom to accept an electron will also increase. This is because the pull of the nucleus on the valence electrons increases from left to right. Thus, both EA and EN increase from left to right.
As we go down a group, Z doesn't change much, but the valence electrons get farther and farther from the nucleus. Because of this, the atom is less and less able to "handle" having another electron added to it. Thus, both EA and EN will decrease as we go from top to bottom down the group.
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Summary of Periodic Trends
Effective nuclear charge increases from left to right, and does not significantly change from top to bottom. The left to right trend occurs because the number of core electrons remains constant, while the number of net protons left over after cancelling out the core electrons increases. The top to bottom trend occurs because the net number of protons left over after cancelling out the core electrons remains constant.
Atomic radius decreases from left to right, and increases from top to bottom. The left to right trend occurs because Z is increasing, and the top to bottom trend occurs because each row adds another layer of core electrons without significantly changing Z.
Ionization energy increases from left to right, and decreases from top to bottom. The left to right trend occurs because Z is increasing, and the top to bottom trend occurs because each row adds another layer of core electrons without significantly changing Z, making the valence electrons farther from the nucleus.
Electron affinity increases from left to right, and decreases from top to bottom. The left to right trend occurs because Z is increasing, and the top to bottom trend occurs because each row adds another layer of core electrons without significantly changing Z, making the valence electrons farther from the nucleus.
Electronegativity increases from left to right, and decreases from top to bottom. The left to right trend occurs because Z is increasing, and the top to bottom trend occurs because each row adds another layer of core electrons without significantly changing Z, making the valence electrons farther from the nucleus.
This post is compliments of QofQuimica from SDN.
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