We just saw how the periodictable can help us quickly determine electron configurations and quantumnumbers. As you’ll see in this section, this is possible because of thespecial arrangement of elements in the periodic table. There willalmost definitely be at least one question about trends in the periodictable on the
SAT Chemistry test, so be sure to read this sectionclosely.
The Anatomy of the Periodic Table
As you are probably well aware,in the periodic table, elements are arranged in order of increasingatomic number. The 18 vertical columns of the table are called groups or families, while the seven horizontal rows are called periods and correspond to the seven principal quantum energy levels, n = 1 through n = 7.
On the right side of theperiodic table is a dividing line resembling a staircase. To the leftof the staircase lie the metals, and to the right of the staircase liethe nonmetals. Many of the elements that touch the staircase are calledmetalloids, and these exhibit both metallic and nonmetallic properties.Study the diagram below and memorize the names of the different typesof elements, because you will definitely see questions about thesegroupings on the test!
Metals aremalleable, ductile, and have luster; most of the elements on theperiodic table are metals. They oxidize (rust and tarnish) readily andform positive ions (cations). They are excellent conductors of both heat and electricity. The metals can be broken down into several groups.
Transition metals (alsocalled the transition elements) are known for their ability to refractlight as a result of their unpaired electrons. They also have severalpossible oxidation states. Ionic solutions of these metals are usuallycolored, so these metals are often used in pigments. The actinides andlanthanides are collectively called the rare earth elements and are filling the f orbitals. They are rarely found in nature. Uranium is the last naturally occurring element; the rest are man-made.
Nonmetals lie to theright of the staircase and do not conduct electricity well because theydo not have free electrons. All the elemental gases are included in thenonmetals. Notice that hydrogen is placed with the metals because ithas only one valence electron, but it is a nonmetal.
Here are some specific families you should know about, within the three main groups (metals, nonmetals, and metalloids):
Alkali metals (1A)—Themost reactive metal family, these must be stored under oil because theyreact violently with water! They dissolve and create an alkaline, orbasic, solution, hence their name.
Alkaline earth metals (2A)—These also are reactive metals, but they don’t explode in water; pastes of these are used in batteries.
Halogens (7A)—Known as the “salt formers,” they are used in modern lighting and always exist as diatomic molecules in their elemental form.
Noble gases (8A)—Knownfor their extremely slow reactivity, these were once thought to neverreact; neon, one of the noble gases, is used to make bright signs.
Now that you’re familiar withthe different groupings of the periodic table, it’s time to talk aboutthe ways we can use the periodic table to predict certaincharacteristics of elements.
Atomic Radius
Since in an atom there is noclear boundary beyond which the electron never strays, the way atomicradius is measured is by calculating the distance between the twonuclei of atoms when they are involved in a chemical bond. If the twobonded atoms are of the same element, you can divide the distance by 2to get the atom’s radius. That said, one of the two important thingsyou’ll need to know about atomic radii for the SAT II Chemistry exam isthat atomic radii decrease (
) moving across a period from left to right.But why? It seems as though the more protons you add, the more spacethe atom should take up, but this is not the case. The reason for thislies in the basic concept that opposite charges attract each other andlike charges repel each other. As you increase the number of protons inthe nucleus of the atom, you increase the effective nuclear charge of the atom (Zeff),and the nucleus pulls more strongly on the entire electron cloud. Thismakes the atomic radius decrease in size. The second thing you’ll needto know is that atomic radii increase moving down a group or family. This is easier to understand if you refer to the Bohr model. As you move down the table, the value of n increases as we add another shell. Remember that the principal quantum number, n,determines the size of the atom. As we move down a family, theattractive force of the nucleus dissipates as the electrons spend moretime farther from the nucleus.
One more thing about atomic size. As you know, when an atom loses an electron, a cation,or positive ion, is formed. When we compare the neutral atomic radiusto the cationic radius, we see that the cationic radius is smaller.Why? The protons in the nucleus hold the remaining electrons morestrongly. As you might expect, for negatively charged ions, or anions,the nuclear attractive force decreases (and there is enhancedelectron-electron repulsion), so the electrons are less tightly held bythe nucleus. The result is that the anion has a larger radius than theneutral atom.
The SAT II Chemistry test might ask you to compare the sizes of two atoms that are isoelectronic,meaning that they have the same number of electrons. In this case, youwould then consider the number of protons the two atoms possess.
Example
Which ion is larger, F– or O2-?
Explanation
Since these two atoms areisoelectronic and in the same period, the atom with more protons in itsnucleus will hold its electrons more tightly and be smaller. Fluoridewill be smaller since it has more protons (9, compared to oxide’s 8).
Ionization Energy (IE)
The ionization energy of anatom is the energy required to remove an electron from the atom in thegas phase. Although removing the first electron from an atom requiresenergy, the removal of each subsequent electron requires even moreenergy. This means that the second IE is usually greater than thefirst, the third IE is greater than the second, and so on. The reasonit becomes more difficult to remove additional electrons is thatthey’re closer to the nucleus and thus held more strongly by thepositive charge of the protons.
Ionization energies differsignificantly, depending on the shell from which the electron is taken.For instance, it takes less energy to remove a p electron than an s electron, even less energy to extract a d electron, and the least energy to extract an f electron. As you can probably guess, this is because s electrons are held closer to the nucleus, while felectrons are far from the nucleus and less tightly held. You’ll needto remember two important facts about ionization energy for the test.The first is that ionization energy increases as we move across a period.
The reason for this, as is the casewith periodic trends in atomic radii, is that as the nucleus becomesmore positive, the effective nuclear charge increases its pull on theelectrons and it becomes more difficult to remove an electron.
The second thing you’ll need to remember is that ionization energy decreases as you move down a group or family.The increased distance between electrons and the nucleus and increasedshielding by a full principal energy level means that it requires lessenergy to remove an electron. Shielding occurs when the innerelectrons in an atom shield the outer electrons from the full charge ofthe nucleus. Keep in mind that this phenomenon is only important as youmove down the periodic table! Here are the values for the firstionization energies for some elements:
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There are some importantexceptions to the above two ionization energy trends in the periodictable, so make sure you study these closely:
- When electron pairing first occurs within an orbital,electron-electron repulsions increase, so that removing an electrontakes less energy (it’s easier); thus the IE drops at this time. Forexample, less energy is required to remove an electron from oxygen’svalence in spite of an increasing Zeff because oxygen’s p4 electronis the first to pair within the orbital. The repulsion created lowersthe amount of energy required to remove either electron.
- There is also a drop in ionization energy from s2 to p1—also in spite of an increasing Zeff. This drop is due to the fact that you are removing a p electron rather than an s electron. The p electrons are less tightly held because they do not penetrate the electron cloud toward the nucleus as well as an s electron does.
Example
Which of the following elements has the highest ionization energy: K, Ca, Ga, As, or Se?
Explanation
The answer is arsenic, or As.Since IE increases as we move across a period, you may have chosen Se.However, there is a drop in IE in spite of increasing Zeff due to the increased electron-electron repulsion in the family that contains oxygen, since they are np4.
Electron Affinity
An atom’s electron affinityis the amount of energy released when an electron is added to the atomin its gaseous state—when an electron is added to an atom, the atomforms a negative ion. Most often, energy is released as anelectron is added to an atom, and the greater the attraction betweenthe atom and the electron added, the more negative the atom’s electronaffinity.
For the SAT II Chemistry test, remember that electron affinity becomes more negative as we move across a period.This means that it’s easier to add an electron to elements, the fartherto the right you travel on the periodic table. Why? Again, this isbecause the higher Zeff increases the nuclearattraction for the incoming electron. Important exceptions to this ruleare the noble gases: He, Ne, Ar, Kr, and Xe. They have electronaffinities that are positive (meaning very low), because if they wereto accept another electron, that electron would have to go into a new,higher-energy subshell, and this is energetically unfavorable.
Electron affinities do not change very much as you go down a group.This is because the lower electron-nucleus attraction that’s seen as wego down a group is pretty evenly counterbalanced by a simultaneouslowering in electron-electron repulsion. Remember that there is noclear trend for electron affinity as you go down a group on theperiodic table—this fact could come up in a synthesis of knowledgequestion!
Electronegativity
Electronegativity is a measureof the attraction an atom has for electrons when it is involved in achemical bond. Elements that have high ionization energy and highelectron affinity will also have high electronegativity since theirnuclei strongly attract electrons. Electronegativity increases fromleft to right as we move across a period and decreases as we move downany group or family.
By now, these trends shouldmake sense. You know that ionization energies tend to decrease withincreasing atomic number in a group, although there isn’t a significantchange in electron affinity, so it makes sense that atoms’ attractionfor electrons in a bond would also increase as their Zeff increased. We will discuss the concept of electronegativity further in the next section, when we discuss chemical bonding.
Here’s a summary of the trends we discussed in this section. Make sure to memorize them!
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本帖最后由 端木·宇 于 2008-6-17 22:56 编辑 ]
