PERIODIC TABLE TRENDS

If you were to look carefully at many of the properties of the elements, you would notice something besides the similarity of the properties within the groups. You would notice that many of these properties change in a fairly regular fashion that is dependent on the position of the element in the periodic table. As you compare elements from left to right across the periodic table, you will notice a trend or regular change in a number of properties. The same thing happens if you go up and down on the periodic table and compare the properties of the elements.

ATOMIC SIZE

The first of these properties is the atomic size. You know that each atom has a nucleus inside and electrons zooming around outside the nucleus. It should seem reasonable that the size of an atom depends on how far away its outermost (valence) electrons are from the nucleus. If they are very close to the nucleus, the atom will be very small. If they are far away, the atom will be quite a bit larger. So the atomic size is determined by how much space the electrons take up.

Measuring the size of atoms is, in some ways, like measuring the size of cotton balls or automobile tires. The value you get depends on the conditions under which they are measured. A "free" cotton ball has a different size than when it is in the package. The radius of the tire is different when measured to the top of the tire than when measured to the bottom of the tire resting on the ground. Different values for the sizes of atoms are obtained depending on both the method used and the conditions in which the atoms find itself - free or bonded to other atoms. The following table gives a variety of values collected from a variety of sources.Whichever set of values you choose to use, note the trends.

Atomic Sizes (in Angstroms, which is 10-10 meter) from Various Sources

 1.58 0.3 1.2 0.98 n.a. 4.10 1.52 2.80 1.12 2.34 0.88 1.82 0.77 1.50 0.70 1.5 1.30 0.66 1.40 1.14 0.64 1.35 1.02 n.a. 1.60 4.46 1.86 3.44 1.60 3.64 1.43 2.92 1.17 2.46 1.10 1.9 2.18 1.04 1.85 1.94 0.99 1.80 1.76 n.a. 1.92 5.54 2.31 4.46 1.97 4.18 1.60 4.00 1.46 3.84 1.31 3.70 1.25 3.58 1.29 3.44 1.26 3.34 1.25 3.24 1.24 3.14 1.28 3.06 1.33 3.62 1.22 3.04 1.22 2.66 1.21 2.0 2.44 1.17 2.00 2.24 1.14 1.95 2.06 n.a. 1.97 5.96 2.44 4.90 2.15 4.54 1.80 4.32 1.57 4.16 1.41 4.02 1.36 3.90 1.3 3.78 1.33 3.66 1.34 3.58 1.38 3.50 1.44 3.42 1.49 4.00 1.62 3.44 1.4 3.06 1.41 2.2 2.84 1.37 2.20 2.64 1.33 2.15 2.48 n.a. 2.17 6.68 2.62 5.56 2.17 5.48 1.88 4.32 1.57 4.18 1.43 4.04 1.37 3.94 1.37 3.84 1.34 3.74 1.35 3.66 1.38 3.58 1.44 3.52 1.52 4.16 1.71 3.62 1.75 3.26 1.46 3.06 1.4 2.86 1.4 2.68 n.a. 2.7 2.20 2.2
Atomic diameter computed using quantum mechanical calculations, Periodic Chart of the Atoms (1979), Sargent-Welch
Atomic radii and covalent radii, "Chemical Systems," Chemical Bond Approach Project (1964), McGraw-Hill
Van der Waals radii, Handbook of Chemistry and Physics, 65th Ed. (1984), CRC Press and "Chemical Systems"

Let's make some comparisons in a family and in a period. In a family--like from hydrogen to lithium to sodium on down--the atomic size increases. As you go down a group, the size increases. As you go across a period, as from lithium to neon, notice that the size decreases. You need to remember (or memorize) those trends.

Now let's talk about why that's the case and relate it back to the various factors presented earlier. Remember that the nuclear charge and the shielding electrons combine to make the effective nuclear charge. That is a very important factor when you are comparing elements in a period. As you go across a period, the nuclear charge increases and the number of energy levels stays the same. Consequently, the number of shielding electrons stays the same and the effective nuclear charge increases. As the effective nuclear charge increases, it pulls the electrons in closer and closer to the nucleus. So as you go across a period, the increase in the nuclear charge causes a decrease in the atomic size because the electrons in the valence energy level are pulled in closer and closer.

Now let's make comparison within a family such as hydrogen down to francium (Fr). It is true that the nuclear charge is increasing, but so is the number of shielding electrons. The number of shielding electrons increases by the same amount that the nuclear charge increases. So the effective nuclear charge felt by the valence electrons stays the same. There is no increase in the effective nuclear charge but there is an increase in the number of energy levels that are being used. Consequently, the electrons in the valence energy level are further and further away from the nucleus because they are in higher energy levels. Consequently, the important factor in a vertical comparison on the periodic table is the number of energy levels that are being used because the increase in the number of shielding electrons cancels out the increase in the nuclear charge.

To summarize, as you go across a period, the increase in the nuclear charge is the most important factor because the number of energy levels stays the same. As you go down a group, the increase in shielding electrons more or less cancels out the increase in nuclear charge, leaving the increase in the number of energy levels as the most important factor. This is true not only for atomic size but for other properties as well.

If you have a sharp eye and a good memory, you may have noticed that the trend shown here as you go from lithium through neon is slightly different than what was shown in the diagram of Lothar Meyer's atomic volumes. The reason for this is something that we will be getting into a little bit later in the course. It has to do with the way that atoms attract to one another. The amount of space taken up by a collection of atoms depends not only on the amount of space taken up by the individual atoms, but also on how much they compact when they combine with one another. In Meyer's diagram, there is first a decrease in volume as you go across the table and then an increase; whereas in this diagram, there is a decrease all the way across. Meyer was measuring two factors. One was the size of the individual atoms and the second was the compressibility of the atoms when they combine with more than one of themselves. In a sense, it would be like using atomic radii for the metals and an average of covalent and van der Waals radii for the nonmetals.

PRACTICE WITH COMPARING ATOMIC SIZES

Now try your hand at answering the following questions (also shown in exercise 6 in your workbook). Check your answers on the next page and then continue with the lesson.

For each of the following sets of atoms, decide which is larger, which is smaller, and why.

a.  Li, C, F

b.  Li, Na, K

c.  Ge, P, O

d.  C, N, Si

e.  Al, Cl, Br

Here are answers for the questions on the previous page.

a.  Li, C, F

All are in the same period and thus have the same number of energy levels. Therefore, the important factor is the nuclear charge. Li is the largest because it has the smallest nuclear charge and pulls the electrons toward the nucleus less than the others. F is the smallest because it has the largest nuclear charge and pulls the electrons toward the nucleus more than the others.

b.  Li, Na, K

All are in the same goup and thus have the same effective nuclear charge. Therefore, the important factor is the number of energy levels. Li is the smallest because it uses the smallest number of electron energy levels. K is the largest because it uses the largest number of electron energy levels.

c.  Ge, P, O

All are in different groups and periods, therefore both factors must be taken into account. Fortunately both factors reinforce one another. Ge is the largest because it uses the largest number of energy levels and has the smallest effective nuclear charge. O is the smallest because it uses the smallest number of energy levels and has the largest effective nuclear charge.

d.  C, N, Si

Not all are in the same group and period, so, again, both factors must be taken into account. C and N tie for using the smallest number of energy levels, but N has a higher effective nuclear charge. Therefore, N is the smallest. C and Si tie for having the lowest effective nuclear charge, but Si uses more energy levels. Therefore, Si is the largest.

e.  Al, Cl, Br

Not all are in the same group and period, so, again, both factors must be taken into account. Cl is the smallest because it has higher effective nuclear charge than Al and uses fewer energy levels than Br. Which is largest is less straightforward. Al has a lower effective nuclear charge (by four), but Br uses more energy levels (by one). Because the difference in effective nuclear charge is larger, it should be the more important factor in this case, making Al the largest.

Al and Br can also be compared to one another indirectly by comparing both to Cl. Both Al and Br are larger than Cl. Al is larger than Cl because it has lower effective nuclear charge (by four). Br is larger than Cl because it uses more energy levels (by one). Because Al is larger than Cl by four "steps" and Br is larger than Cl by only one "step", Al is likely the largest of the three.

QUICK QUIZ ON COMPARING ATOMIC SIZES

*** Na,Mg   O,S   S,F ***

IONIZATION ENERGY

Now on to another property. It's called ionization energy. It can be defined as being the energy required to remove the outermost electron from a gaseous atom. A "gaseous atom" means an atom that is all by itself, not hooked up to others in a solid or a liquid. When enough energy is added to an atom the outermost electron can use that energy to pull away from the nucleus completely (or be pulled, if you want to put it that way), leaving behind a positively charged ion. That is why it's called ionization, one of the things formed in the process is an ion. The ionization energy is the exact quantity of energy that it takes to remove the outermost electron from the atom.

In your lab work on atomic spectra you observed that a gas would conduct electricity and emit light when it was subjected to a high voltage. When there is little or no voltage applied to the gas in the tubes, no light is emitted and the gas does not conduct electricity. One method for measuring the ionization energy of a gas is to slowly increase the voltage applied to it until it does conduct electricity and emit light. The voltage at which that occurs can be used to calculate the ionization energy.

If the ionization energy is high, that means it takes a lot of energy to remove the outermost electron. If the ionization energy is low, that means it takes only a small amount of energy to remove the outermost electron.

Let’s use your understanding of atomic structure to make some predictions. Think for a minute about how ionization energy would be affected by three of the factors we were talking about earlier: (1) nuclear charge, (2) number of energy levels, and (3) shielding.

As the nuclear charge increases, the attraction between the nucleus and the electrons increases and it requires more energy to remove the outermost electron and that means there is a higher ionization energy. As you go across the periodic table, nuclear charge is the most important consideration. So, going across the periodic table, there should be an increase in ionization energy because of the increasing nuclear charge.

Going down the table, the effect of increased nuclear charge is balanced by the effect of increased shielding, and the number of energy levels becomes the predominant factor. With more energy levels, the outermost electrons (the valence electrons) are further from the nucleus and are not so strongly attracted to the nucleus. Thus the ionization energy of the elements decreases as you go down the periodic table because it is easier to remove the electrons. Another way of looking at that is that if you are trying to take something from the first energy level, you have to take it past the second, the third, the fourth and so on, on the way out. But if something is already in the third or fourth energy level, it doesn't have to be taken as far to get away from the nucleus. It is already part way removed from the nucleus.

This table shows the measured values for the ionization energies of the first twenty elements. If you take a close look at what happens to the ionization energy as you go from left to right across the periodic table, you will find that there is not really a steady increase in ionization energy as I had indicated. You could describe the pattern you see there as being a few steps forward then one step back, repeating itself as you move across. Progress is made, but it is not steady.
 Ionization Energies (v) H 13.6 He 24.6 Li 5.4 Be 9.3 B 8.3 C 11.3 N 14.5 O 13.6 F 17.4 Ne 21.6 Na 5.1 Mg 7.6 Al 6.0 Si 8.2 P 10.5 S 10.4 Cl 13.0 Ar 15.8

 The periodic nature of ionization energy is emphasized in this diagram. With each new period the ionization energy starts with a low value. Within each period you will notice that the pattern is really kind of a zigzag pattern progressing up as you go across the periodic table. The zigs and zags on that graph correspond to the sublevels in the energy levels. So far in this lesson we have presumed that all the electrons in the second energy level are pretty much the same. Two factors make that not completely true. One factor is that because s and p orbitals have different shapes, the electrons in p orbitals have more energy and are further from the nucleus. The other factor is that when electrons are paired up in an orbital, they repel one another somewhat. Those two factors account for the zigzag nature of the increase in ionization energy. Nevertheless, as a general trend, from left to right across the periodic table, ionization energy does increase. Also as you go down the periodic table, the ionization energy does decrease for the reasons given.

Note that the trends in the periodic properties of atomic size and ionization energy are related. Going across the periodic table from left to right, the electrons are more tightly held by the nucleus, causing the atoms to be smaller and the ionization energy to be higher. As you go down the periodic table, the electrons are further from the nucleus, causing the atoms to be larger and the ionization energies to be lower.

PRACTICE WITH COMPARING IONIZATION ENERGIES

Please take some time now to do the following exercises (also shown in example 7 in your workbook). When you have done that, check your answers on the next page and then continue.

For each of the following sets of atoms, decide which has the highest and lowest ionization energies and why.

a.  Mg, Si, S

b.  Mg, Ca, Ba

c.  F, Cl, Br

d.  Ba, Cu, Ne

e.  Si, P, N

Here are answers to the exercises on the previous page.

a.  Mg, Si, S

All are in the same period and use the same number of energy levels. Mg has the lowest I.E. because it has the lowest effective nuclear charge. S has the highest I.E. because it has the highest effective nuclear charge.

b.  Mg, Ca, Ba

All are in the same group and have the same effective nuclear charge. Mg has the highest I.E. because it uses the smallest number of energy levels. Ba has the lowest I.E. because it uses the largest number of energy levels.

c.  F, Cl, Br

All are in the same group and have the same effective nuclear charge. F has the highest I.E. because it uses the smallest number of energy levels. Br has the lowest I.E. because it uses the largest number of energy levels.

d.  Ba, Cu, Ne

All are in different groups and periods, so both factors must be considered. Fortunately both factors reinforce one another. Ba has the lowest I.E. because it has the lowest effective nuclear charge and uses the highest number of energy levels. Ne has the highest I.E. because it has the highest effective nuclear charge and uses the lowest number of energy levels.

e.  Si, P, N

Si has the lowest I.E. because it has the lowest effective nuclear charge and is tied (with P) for using the most energy levels. N has the highest I.E. because it uses the fewest energy levels and is tied (with P) for having the highest effective nuclear charge.

QUICK QUIZ ON COMPARING IONIZATION ENERGIES

*** O,S  Ca,Cr  Al,N ***

TENDENCY TO LOSE ELECTRONS

Ionization energy measures how difficult it is for atoms to lose electrons but quite often we will want to talk about how easy it is for atoms to lose electrons. A low ionization energy means that it is easy for an atom to lose electrons. A high ionization energy means that it is hard for an atom to lose electrons.

Using this terminology, it gets harder for atoms to lose electrons as you go across the periodic table and it gets easier for atoms to lose electrons as you go down the periodic table.
 Ability to Lose Electrons H a r d E a s y

TENDENCY TO GAIN ELECTRONS

Next let's consider the opposite of losing electrons ( ionization of atoms), that is the gaining of electrons. Atoms can attract additional electrons if there is room for them in the valence energy level. When an extra electron moves into the valence shell, it can feel the attraction exerted by the effective nuclear charge. Because the effective nuclear charge is largest for the elements on the right side of the periodic table, those atoms provide the greatest attraction for electrons and have the greatest tendency to gain electrons.

Thus the tendency of atoms to gain electrons increases as we go from left to right across the periodic table. At least it increases until we get to the inert gases. There it drops off to zero because there is no room for additional electrons in the valence energy level. A new electron would have to start a new energy level, but there would not be an additional proton in the nucleus to provide any effective nuclear charge.
 Ability to Gain Electrons E a s y H a r d

As we look at elements going down the periodic table, the effective nuclear charge remains the same, so the increase in the number of energy levels is the important factor. The tendency of atoms to gain electrons decreases as we go down the periodic table. The reason for this is simply that with the larger atoms the added electron is not as close to the nucleus and therefore the attractive force exerted by the effective nuclear charge is not as powerful as it is in the smaller atoms.

PRACTICE COMPARING TENDENCIES TO GAIN ELECTRONS

For each of the following sets of atoms, decide which has the least and which has the greatest tendency to gain electrons and why.

a.  Li, C, N

b.  C, O, Ne

c.  Si, P, O

d.  K, Mg, P

e.  S, F, He

ANSWERS FOR COMPARING TENDENCIES TO GAIN ELECTRONS

Here are answers to the exercises on the previous page.

a.  Li, C, N

Li has the least tendency to gain electrons because it has the lowest effective nuclear charge (and all use the same number of energy levels). N has the greatest tendency to gain electrons because it has the highest effective nuclear charge (and all use the same number of energy levels).

b.  C, O, Ne

Ne has the lowest tendency to gain electrons because its outer energy level is full and there is no room for an additional electron. O has the greatest tendency to gain electrons because it has a higher effective nuclear charge than C (and both use the same number of energy levels).

c.  Si, P, O

O has the greatest tendency to gain electrons because it has the highest effective nuclear charge and uses the smallest number of energy levels. Si has the lowest tendency to gain electrons because it has the lowest effective nuclear charge and is tied (with P) for using the most energy levels.

d.  K, Mg, P

P has the greatest tendency to gain electrons because it has the highest effective nuclear charge and is tied (with Mg) for using the smallest number of energy levels. Neither Mg nor K have much attraction for electrons, but K has the lowest tendency to gain electrons because it has the lowest effective nuclear charge and uses the most energy levels.

e.  S, F, He

He has the lowest tendency to gain electrons because its outer energy level is full and there is no room for an additional electron. F has the greatest tendency to gain electrons because it has a higher effective nuclear charge and uses fewer energy levels than S.

QUICK QUIZ ON COMPARING TENDENCIES TO GAIN ELECTRONS

Take some time now to answer the questions in exercise 8. Check them with the instructor and then we will continue.

*** C,O   Cl,Br   Na,N   F,Ne ***

ELECTRON AFFINITY

The electron affinity is a very specific measure of the tendency for atoms to gain electrons. Specifically, it is the amount of energy released by an electron when the electron joins or attaches to an isolated atom. This measurement is very sensitive to the difference between the types of orbitals and to the number of electrons in those orbitals. Consequently, the zigs and zags in a graph of electron affinity are even more pronounced than in the graph of ionization energy. Therefore we won't deal with the electron affinity trends - they get kind of lost in the fluctuations. You should remember what electron affinity is but don't worry about how it varies as you go across and up and down the periodic table.

IONIC SIZE

When atoms gain or lose electrons, the atom becomes an ion. When an atom gains an electron, it becomes a negatively charged ion that we call an anion. Anions are larger in size than their parent atoms because they have one or more additional electrons, but without an additional proton in the nucleus to help moderate the size.

When an atom loses an electron, it becomes a positively charged ion called a cation. Cations are smaller than their parent atoms because they have lost electrons (sometimes the entire outermost energy level) and the electrons that remain behind simply don't take up as much room.

 Sizes of Atoms and Ions Cations Anions Na 1.86 Na+ 0.95 Mg 1.60 Mg2+ 0.65 O 0.74 O2- 1.40 F 0.71 F- 1.36 K 2.27 K+ 1.33 Ca Ca2+ S 1.03 S2- 1.84 Cl 0.99 Cl- 1.81 Values from Hein, Best, Pattison, Arena, "College Chemistry," 5th Ed., 1993, Brooks/Cole

There is a chart on the wall of the lab that shows not only the sizes of atoms, but also the sizes of ions. You might want to take a look at that. It is on the south wall of the lab. Also note that when comparing cations and anions, the anions are larger.

PRACTICE COMPARING IONIC SIZES

Try your hand at the making the folowing comparisons (also shown in exercise 9 in your workbook), based on your understanding of ionic size comparisons and without reference to the wall chart, except to check your answers. Answers also follow on the next page.

For each of the following sets of atoms and ions, decide which is the smallest and which is the largest.

a.  Na, Na+

b.  Cl, Cl-

c.  Na+, Cl-

d.  H+, H, H-

e.  Fe2+, Fe3+

f.  F-, Ne, Na+

Try your hand at the making the folowing comparisons (also shown in exercise 9 in your workbook), based on your understanding of ionic size comparisons and without reference to the wall chart, except to check your answers. Answers also follow on the next page.

For each of the following sets of atoms and ions, decide which is the smallest and which is the largest.

a.  Na is largest, Na+ is smallest.

b.  Cl is smallest, Cl- is largest.

c.  Na+ is smallest, Cl- is largest.

d.  H+ is smallest, H, H- is largest.

e.  Fe2+ is largest, Fe3+ is smallest.

f.  F- is largest, Ne, Na+ is smallest.

QUICK QUIZ ON COMPARING IONIC SIZES

*** sm Mg,Mg2+  sm S,S2-  lg K+,Cl- ***

REACTIVITY

Now let's consider something that is more directly observable. Let's look at the reactivity of some of the elements and relate that reactivity to some of the properties we have just talked about.

Reactivity means just what it says. A highly reactive element reacts very easily, maybe even violently, with lots of other elements or compounds.

In your lab work for this lesson, you will make observations comparing the reactivities of some of the elements by observing the reactions of a few. If your schedule permits, it would be best to do this portion of the lesson after doing the lab work. The "Post-Lab Discussions" of parts I and II of the lab work can be done on the computers in the lab immediately after completing your observations in each part.

PRE-LAB COMMENTS ON METAL REACTIVITY TRENDS

You will compare the reactivities of the metallic elements Na, K, Mg, and Ca to each other when they are added to water. You will need to get the instructor to help you because some of these elements react quite violently with water and require special precautions. So take the time now to observe the reaction of these four elements with water by doing the first part of exercise 10 in your workbook.

POST-LAB DISCUSSION OF METAL REACTIVITY TRENDS

Let's take a look at your observations for the reactions of these four metals and take a look at what kinds of trends you were able to observe and see what kinds of projections and explanations we can come up with.

 Let's take a look at how those elements compare with one another on the periodic table. Sodium and magnesium are both in period 3. In period 4 we have potassium and calcium. In group Ia are sodium and potassium and in group IIa are magnesium and calcium. We have 2 elements in each of 2 groups and also 2 periods, so we can make some comparisons.

Within period 3, what happened? Which was more reactive? Was it sodium or magnesium? It was sodium. The more reactive metal was on the left and the least reactive was on the right. In period 4, what about the comparison of potassium with calcium? Well, again the most reactive was on the left and the reactivity got less as you went across. The vertical comparison in both group Ia and group IIa is that the reactivity increased as you went down the group. So the reactivity of metals decreases as you go from left to right and it increases as you go down on the periodic table.

RELATING METALLIC REACTIVITY TRENDS TO ATOMIC STRUCTURE

Let’s consider how metals react. Metals react by losing electrons. They have a low ionization energy so it's fairly easy for them to lose electrons.

 As you go across a period, the nuclear charge will increase; the number of energy levels will stay the same, so there is a stronger and stronger attraction for the electrons. The electrons are being held more tightly as you go across a period. It becomes more and more difficult to lose electrons and consequently the reactivity of the metals decreases as you go from left to right across the periodic table. As you go down the periodic table, the nuclear charge increases but so does the number of shielding electrons. Consequently the dominant factor is that we have more and more energy levels and the electrons are further and further away from the nucleus. Thus it is easier for those electrons to come off.

Those are the reasons for the pattern of reactivity that is seen for the metals.

Does this trend work for elements beyond the ones we have just looked at? You can check this out by taking a look at the reactions shown on a video tape called "Close-Up on Chemistry" by Julie and James Ealy produced by the American Chemical Society. It is available for you to look at on a monitor over in the lab.

PRE-LAB COMMENTS ON NONMETAL REACTIVITY TRENDS

Next, you should do part 2 of this experiment which deals with the reactivity of nonmetals. You will compare bromine with iodine by observing how each reacts with the chemical 1-octene. There aren't as many of the nonmetals and it is also more difficult to come up with good simple reactions for you to work with to make these comparisons. Consequently, we will have to deal with fewer observations and more inferences. So take some time now to do part 2 of this experiment.

POST-LAB DISCUSSION OF NONMETAL REACTIVITY TRENDS

You should have noticed in this case that bromine reacted more readily than iodine. Notice that this is the opposite of what we found with the metals. With the metals, the element that was further down on the periodic table was more reactive. But with nonmetals over on the right side of the periodic table, the element which is further up is most reactive. This points out that there is something fundamentally different about the nonmetals compared to the metals. There is a fundamental difference between the way that metals and nonmetals react.

RELATING NONMETALLIC REACTIVITY TRENDS TO ATOMIC STRUCTURE

Nonmetals usually react by gaining electrons, rather than by losing electrons like the metals do. Let’s review how atomic structure affects the ability to gain electrons. From your observations in the lab you know that as you go down a nonmetallic group in the periodic table, the elements become less reactive. You also know that as you go down a group on the periodic table, the number of energy levels is the most predominant factor. If an electron comes into an atom that has a large number of energy levels, it will be further away from the nucleus and not be attracted as strongly as it would be in a smaller atom with fewer energy levels. For example, iodine is attracting an electron into its fifth energy level. Bromine is attracting an electron into its fourth energy level. Bromine does a better job of attracting electrons, and thus is more reactive, because it allows the new electron to get closer to the nucleus where the force of attraction is stronger. Following this line of reasoning and extending it to other atoms, we would expect chlorine to be even more reactive and fluorine to be even more reactive still.

The reactivity of the nonmetals ties in well with the concept of electron affinity and the tendency to gain electrons. With nonmetals the greater the tendency to gain electrons, the more reactive it is. This argument should hold true whether we are talking about nonmetals within a family or within a period. As you go across a period, there is a greater nuclear charge and thus the electrons should be attracted more readily by elements that are further to the right and the tendency to gain electrons will increase. Thus the reactivity of the nonmetals should increase as you go from left to right across the periodic table, up to but not including the inert gases.