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Periodic Trends • Introduction • 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. And of course, that is what you will do next. 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.
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.
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.
Atomic RadiusAtomic radius => distance from the center of an atom's nucleus to its outer most electron • Atomic Radii • 1) As you move down a group, atomic radius increases. • WHY? - The number of energy levels increases as you move down a group as the number of electrons increases. Each subsequent energy level is further from the nucleus than the last. Therefore, the atomic radius increases as the group and energy levels increase. • 2) As you move across a period, atomic radius decreases. • WHY? - As you go across a period, electrons are added to the same energy level. At the same time, protons are being added to the nucleus. The concentration of more protons in the nucleus creates a "higher effective nuclear charge." In other words, there is a stronger force of attraction pulling the electrons closer to the nucleus resulting in a smaller atomic radius.
Atomic RadiusAtomic radius => distance from the center of an atom's nucleus to its outer most electron Common Student Misconceptions 1. “The radius of an atom increases in a period with the addition of electrons.” From the data given on atomic radii, we see that with the addition of electrons, the atomic radii decrease within a period. This is due to the increase in the positive charge of the nucleus and its attraction for the electrons in the atom. Although the number of electrons in the atom increases as one goes across a period and although the electrons repel each other potentially causing the atoms to grow in size, this effect is more than opposed by the increase in the number of protons attracting these electrons. Since these electrons are added to the same shell, the number of protons attracting them causes the size of the atoms to decrease. 2. “All properties show periodic characteristics.” There are deviations in all periods. 3. “Elements are arranged according to atomic weight.” While this is true in general, the periodic properties are characteristic of atomic numbers.
Practice with Comparing Atomic Size • Now try your hand at answering the following questions. Check your answers below 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
Answers for Comparing Atomic Sizes • Here are answers for the questions above. • 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 group 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.
Atomic RadiusAtomic radius => distance from the center of an atom's nucleus to its outer most electron
Atomic Radius • ATOMIC RADIUS • Measures of atomic radius • Unlike a ball, an atom doesn't have a fixed radius. The radius of an atom • can only be found by measuring the distance between the nuclei of two • touching atoms, and then halving that distance. • As you can see from the diagrams, the same atom could be found to have • a different radius depending on what was around it. • The left hand diagram shows bonded atoms. The atoms are pulled closely • together and so the measured radius is less than if they are just touching. • This is what you would get if you had metal atoms in a metallic structure, • or atoms covalently bonded to each other. The type of atomic radius • being measured here is called the metallic • radius or the covalent • radius depending on the bonding.
Ionic radius • 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 acation. 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.
Ionic Radius • Ionic Radii • 1) Anions (negative ions) are larger than their respective atoms. • WHY? • Electron-electron repulsion forces them to spread further apart. Electrons outnumber protons; the protons cannot pull the extra electrons as tightly toward the nucleus. • 2) Cations (positive ions) are smaller than their respective atoms. • WHY? • There is less electron-electron repulsion, so they can come closer together. Protons outnumber electrons; the protons can pull the fewer electrons toward the nucleus more tightly. If the electron that is lost is the only valence electron so that the electron configuration of the cation is like that of a noble gas, then an entire energy level is lost. In this case, the radius of the cation is much smaller than its respective atom.
Practice Comparing Ionic Sizes • Try your hand at the making the following 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 this 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+
Answers for Comparing Ionic Sizes • Here are the answers to the questions above. 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
Electronegativity • Electronegativity—the strength with which an atom pulls on the electrons it shares in a bond (the attractions that hold molecules together) is its electronegativity. The trend is the same as the electron affinity trends.
Electronegativity • ELECTRONEGATIVITY • This page explains what electronegativity is, and how and why it varies • around the Periodic Table. It looks at the way that electronegativity • differences affect bond type and explains what is meant by polar bonds • and polar molecules. • If you are interested in electronegativity in an organic chemistry context, • you will find a link at the bottom of this page. • What is electronegativity • Definition • Electronegativity is a measure of the tendency of an atom to attract a • bonding pair of electrons. • The Pauling scale is the most commonly used. Fluorine (the most • electronegative element) is assigned a value of 4.0, and values range • down to caesium and francium which are the least electronegative at 0.7.
Trends in electronegativity across a period • As you go across a period the electronegativity increases. The chart shows • electronegativities from sodium to chlorine - you have to ignore argon. It • doesn't have an electronegativity, because it doesn't form bonds. Trends in electronegativity down a group • As you go down a group, electronegativity decreases. (If it increases up to • fluorine, it must decrease as you go down.) The chart shows the patterns • of electronegativity in Groups 1 and 7.
Explaining the patterns in electronegativity • The attraction that a bonding pair of electrons feels for a particular nucleus depends on: • ● the number of protons in the nucleus; • ● the distance from the nucleus; • ● the amount of screening by inner electrons.
Why does electronegativity increase across a period? • Consider sodium at the beginning of period 3 and chlorine at the end • (ignoring the noble gas, argon). Think of sodium chloride as if it were • covalently bonded. • Both sodium and chlorine have their bonding electrons in the 3-level. The • electron pair is screened from both nuclei by the 1s, 2s and 2p electrons, • but the chlorine nucleus has 6 more protons in it. It is no wonder the • electron pair gets dragged so far towards the chlorine that ions are • formed. • Electronegativity increases across a period because the number of charges • on the nucleus increases. That attracts the bonding pair of electrons more • strongly.
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.
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 below 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
Answers to Comparing Ionization Energies • Here are answers to the exercises above. • 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.
Ionization energy and the Loss of electrons • 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.
Ionization energy and gaining electrons • 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.
IE affect on Gaining Electrons • 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. Check your answers below and then continue with the lesson. • 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 above. • 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.
