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# 6.3: Periodic Trends

Difficulty Level: At Grade Created by: CK-12

## Lesson Objectives

• Learn the periodic trends for atomic radius.
• Know the relationship of group number to valence electrons.
• Describe how ions are formed.
• Learn the periodic trends for ionization energy.
• Explain how multiple ionization energies are related to noble gas electron configurations.
• Describe electron affinity.
• Predict the effect that ion formation has on the size of an atom.
• Learn the periodic trends for electronegativity.

## Lesson Vocabulary

• anion
• cation
• electron affinity
• electronegativity
• ion
• ionization energy

So far, you have learned that the elements are arranged in the periodic table according to their atomic number and that elements in vertical groups share similar electron configurations and chemical properties. In this lesson, we will explore various measurable properties of elements and the variance of those properties among elements. Specifically, we will examine trends within periods and groups. A trend is a general increase or decrease in a particular measurable quantity. For example, as the calendar moves from August to December in the northern hemisphere, the trend is for the average daily temperature to decrease. That doesn’t mean that the temperature drops every single day, just that the overall direction is generally downward.

The size of an atom is defined by the edge of its orbital. However, orbital boundaries are fuzzy and in fact are variable under different conditions. In order to standardize the measurement of atomic radii, the distance between the nuclei of two identical atoms bonded together is measured. The atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together (Figure below).

The atomic radius (r) of an atom can be defined as one half the distance (d) between two nuclei in a diatomic molecule.

Atomic radii have been measured for elements, as shown in Figure below. The units for atomic radii are picometers, equal to 10-12 meters. As an example, the internuclear distance between the two hydrogen atoms in an H2 molecule is measured to be 74 pm. Therefore, the atomic radius of a hydrogen atom is 74/2 = 37 pm.

Atomic radii of the representative elements measured in picometers.

### Periodic Trend

As you can see from Figure above, the atomic radius of atoms generally decreases from left to right across a period. There are some small exceptions, such as the oxygen radius being slightly greater than the nitrogen radius. Within a period, protons are added to the nucleus as electrons are being added to the same principal energy level. These electrons are gradually pulled closer to the nucleus because of its increased positive charge. Since the force of attraction between nuclei and electrons increases, the size of the atoms decreases. The effect lessens as one moves further to the right in a period because of electron-electron repulsions that would otherwise cause the atom’s size to increase.

### Group Trend

As Figure above shows, the atomic radius of atoms generally increases from top to bottom within a group. As the atomic number increases down a group, there is again an increase in the positive nuclear charge. However, there is also an increase in the number of occupied principle energy levels. Higher principal energy levels consist of orbitals which are larger in size than the orbitals from lower energy levels. The effect of the greater number of principal energy levels outweighs the increase in nuclear charge and so atomic radius increases down a group.

Figure below shows a graph of atomic radius plotted versus atomic number. Each successive period is shown in a different color. You can see that as the atomic number increases within a period, the atomic radius decreases. As atomic number increases within a group, the atomic radius increases.

Graph of the atomic radii plotted against the atomic number.

## Valence Electrons

We have previously defined the valence electrons as those in the outermost principal energy level. Since the outermost principal energy level consists only of s and p sublevels, the maximum number of valence electrons is eight. Table below shows the relationship of valence electrons to group number for the representative elements.

Valence Electrons of Representative Elements
Group Number Outer Electron Configuration Number of Valence Electrons
1 $ns^1$ 1
2 $ns^2$ 2
13 $ns^2np^1$ 3
14 $ns^2np^2$ 4
15 $ns^2np^3$ 5
16 $ns^2np^4$ 6
17 $ns^2np^5$ 7
18 $ns^2np^6$ 8

You can see that the number of valence electrons is constant within a group because of the identical outer electron configuration. This is the reason why elements within a group share similar chemical properties.

Watch these video lectures:

## Forming Ions

Many chemical compounds consist of particles called ions. An ion is an atom or group of bonded atoms that has a positive or negative charge. For our purposes, we will presently only consider monatomic ions, which are single atoms with an electrical charge. How do atoms obtain this charge? There are essentially two possibilities: (1) gaining or losing positively-charged protons, or (2) gaining or losing negatively-charged electrons. The nucleus of an atom is very stable and the number of protons is unchangeable in chemical reactions. As we have already seen, however, the electrons are capable of movement within an atom. Electron energy level transitions are very common and responsible for atomic emission spectra. It is the electrons which are either gained or lost to make ions (Figure below).

When an atom loses one or more electrons it becomes positively charged because it now has more protons than electrons. A positively charged ion is called a cation. The charge for a cation is written as a numerical superscript after the chemical symbol, followed by a plus sign. If the ion carries a single unit of charge the number “1” is assumed and not written. For example, a sodium atom that loses one electron becomes a sodium ion, written as Na+. A magnesium atom that loses two electrons becomes a magnesium ion, written as Mg2+. The magnesium ion carries a 2+ charge because it would now have two more protons than electrons.

When an atom gains one or more electrons it becomes negatively charged because it now has more electrons than protons. A negatively charged ion is called an anion. The anion charge is written in the same way as for cations except with a minus sign. A chlorine atom that gains one electron becomes a chloride ion, written as Cl-. Note that the names of monatomic anions have been given an “-ide” suffix. A sulfur atom that gains two electrons becomes a sulfide ion, written as S2-.

## Ionization Energy

To make an electron jump from a lower energy level to a higher energy level, there must be an input of energy. It stands to reason then, that removing the electron from the atom entirely requires even more energy. This is called an ionization process. Ionization energy is the energy required to remove an electron from an atom. An equation can be written to illustrate this process for a sodium atom.

$\text{Na} + \text{energy} \rightarrow \text{Na}^+ + e^-$

The equation shows that energy added to a sodium atom results in a sodium ion plus the removed electron (e-). The electron that is removed from an atom is always a valence electron, since this electron is in the outermost principal energy level and is furthest from the nucleus. Elements have differing ionization energies (Figure below), which are influenced by the size of the atom, the nuclear charge, and the electron energy levels. Ionization energies are measured in units of kilojoules per mole or kJ/mol.

Periodic table showing the first ionization energies of elements, measured in kJ/mol.

Graph of first ionization energy plotted against atomic number

### Period Trend

As can be seen from Figures above and above, the ionization energy of atoms generally increases from left to right across each row of the periodic table. The reason for this increase in ionization energy is the increase in nuclear charge. As more protons are added to the nucleus, the attraction for electrons increases. Because the electron that is to be removed is held more tightly to the nucleus, it is harder to remove it. This results in a larger ionization energy.

There are several exceptions to the general increase in ionization energy across a period. The elements of Group 13 (B, Al, etc.) have a lower ionization energy than the elements of Group 2 (Be, Mg, etc.). This is an illustration of a concept called electron shielding. Outer electrons are partially shielded from the attractive force of the protons in the nucleus by inner electrons (Figure below).

The shielding effect is shown by the interior electron cloud (in green) shielding the outer electron of interest from the full attractive force of the nucleus. A larger shielding effect results in a decrease in ionization energy.

To explain how shielding works, consider a lithium atom. It has three protons and three electrons – two in the first principal energy level and its valence electron in the second. The valence electron is partially shielded from the attractive force of the nucleus by the two inner electrons. Removing that valence electron becomes easier because of the shielding effect. There is also a shielding effect that occurs between sublevels within the same principal energy level. Specifically, an electron in the s sublevel is capable of shielding electrons in the p sublevel of the same principal energy level. This is because of the spherical shape of the s orbital. The reverse is not true – electrons in p orbitals do not shield electrons in s orbitals (Figure below).

The spherical 3s orbital exhibits a shielding effect on the dumbbell shaped 3p orbital, that is of slightly higher energy. This reduces the ionization energy of a 3p electron compared to a 3s electron.

The electron being removed from an Al atom is a 3p electron, which is shielded by the two 3s electrons as well as all the inner core electrons. The electron being removed from a Mg atom is a 3s electron, which is only shielded by the inner core electrons. Since there is a greater degree of electron shielding in the Al atom, it is slightly easier to remove the valence electron and its ionization energy is less than that of Mg. This is despite the fact that the nucleus of the Al atom contains one more proton than the nucleus of the Mg atom.

There is another anomaly between Groups 15 and 16. Atoms of Group 16 (O, S, etc.) have lower ionization energies than atoms of Group 15 (N, P, etc.). Hund’s rule is behind the explanation. In a nitrogen atom, there are three electrons in the 2p sublevel and each are unpaired. In an oxygen atom, there are four electrons in the 2p sublevel and so one orbital contains a pair of electrons. It is that second electron in the orbital that is removed in the ionization of an oxygen atom. Since electrons repel each other, it is slightly easier to remove the electron from the paired set in the oxygen atom than it is to remove an unpaired electron from the nitrogen atom.

### Group Trend

The ionization energy of the representative elements generally decreases from top to bottom within a group. This trend is explained by the increase in size of the atoms within a group. The valence electron that is being removed in each atom is further from the nucleus for a larger atom. The attractive force between the valence electron and the nucleus weakens as the distance between them increases. This results in a lower ionization energy for the larger atoms within a group, even though the nuclear charge is increased. The shielding effect within a group increases because of a larger number of inner electrons. For all of the alkali metals, the single valence electron is shielded by all of the inner core electrons. This increase in shielding effect also makes it easier to remove the valence electron from the larger atom.

### Multiple Ionizations

So far, we have described the first ionization energy of atoms and its trends. However, in many cases multiple electrons can be removed from atoms. If an atom loses two electrons, it acquires a 2+ charge. If an atom loses three electrons, it acquires a 3+ charge, and so on. The ionization energies required are called the second ionization energy (IE2), third ionization energy (IE3), etc. The first six ionization energies are shown for the elements of the first three periods in Table below.

Ionization Energies (kJ/mol) of the First 18 Elements
Element IE1 IE2 IE3 IE4 IE5 IE6
H 1312
He 2373 5251
Li 520 7300 11,815
Be 899 1757 14,850 21,005
B 801 2430 3660 25,000 32,820
C 1086 2350 4620 6220 38,000 47,261
N 1400 2860 4580 7500 9400 53,000
O 1314 3390 5300 7470 11,000 13,000
F 1680 3370 6050 8400 11,000 15,200
Ne 2080 3950 6120 9370 12,200 15,000
Na 496 4560 6900 9540 13,400 16,600
Mg 738 1450 7730 10,500 13,600 18,000
Al 578 1820 2750 11,600 14,800 18,400
Si 786 1580 3230 4360 16,000 20,000
P 1012 1904 2910 4960 6240 21,000
S 1000 2250 3360 4660 6990 8500
Cl 1251 2297 3820 5160 6540 9300
Ar 1521 2666 3900 5770 7240 8800

Notice that the second ionization energy of an element is always higher than the first, the third is always higher than the second, and so on. This is because after one ionization, a positively charged ion is formed. Now there is a greater overall attractive force on the remaining electrons since the protons now outnumber the electrons. So to remove a second electron is more difficult.

The first ionization energies for the noble gases (He, Ne, Ar) are higher than those of any other element within that period. The noble gases have a full outer s and p sublevel, which gives them extra stability and means that they are unreactive. The stability of the noble gas electron configuration applies to other elements as well. Consider the element lithium, with its 1s22s1 electron configuration. Being an alkali metal, its first ionization energy is very low. After it loses its valence electron (the 2s electron), it becomes a lithium ion, Li+, and now has the electron configuration of 1s2. This is the electron configuration of the noble gas helium. The second ionization energy of lithium (shaded above) shows an extremely large jump compared to the first because the removal of a second electron requires breaking apart the noble gas electron configuration. The pattern continues across each period of the table. Beryllium shows a large jump after IE2, boron after IE3, and so on.

## Electron Affinity

In most cases, the formation of an anion by the addition of an electron to a neutral atom releases energy. This can be shown for chloride ion formation below:

$\text{Cl} + e^- \rightarrow \text{Cl}^- + \text{energy}$

The energy change that occurs when a neutral atom gains an electron is called its electron affinity. When energy is released in a chemical reaction or process, that energy is expressed as a negative number. Figure below shows electron affinities in kJ per mole for the representative elements.

Electron affinities (in kJ/mol) of the first five periods of the representative elements. Electron affinities are negative numbers because energy is released.

The elements of the halogen group (Group 17) gain electrons most readily, as can be seen from their large negative electron affinities. This means that more energy is released in the formation of a halide ion than for the anions of any other elements. Considering electron configuration, it is easy to see why. The outer configuration of all halogens is ns2np5. The addition of one more electron gives the halide ions the same electron configuration as a noble gas, which we have seen is particularly stable.

Period and group trends for electron affinities are not nearly as regular as for ionization energy. In general, electron affinities increase (become more negative) from left to right across a period and decrease (become less negative) from top to bottom down a group. However, there are many exceptions, owing in part to inherent difficulties in accurately measuring electron affinities.

Figure below shows the comparison of ion sizes to atom sizes for Groups 1, 2, 13, 16 and 17. The atoms are shown in gray. Groups 1, 2, and 13 are metals and form cations, shown in red. Groups 16 and 17 are nonmetals and form anions, shown in blue.

Atomic and ionic radii of the elements in Periods 1-5 and Groups 1,2, 13, 16, and 17: Atoms are shown in gray. The most common ion is shown in either red (for cations) or blue (for anions).

The removal of electrons always results in a cation that is considerably smaller than the parent atom. When the valence electron(s) are removed, the resulting ion has one fewer occupied principal energy level, so the electron cloud that remains is smaller. Another reason is that the remaining electrons are drawn closer to the nucleus because the protons now outnumber the electrons.

The addition of electrons always results in an anion that is larger than the parent atom. When the electrons outnumber the protons, the overall attractive force that the protons have for the electrons is decreased. The electron cloud also spreads out because more electrons results in greater electron-electron repulsions.

### Trends

Period and group trends for ion radii are similar to the trends for atomic radii and for the same reasons. Left to right across the second period, the cations decrease in size because of greater nuclear charge. Starting in Group 15, a nitrogen atom becomes more stable by gaining three electrons to become a nitride ion, N3-, and acquire a noble gas electron configuration. The nitride ion is larger than the previous cations, but the anions then decrease in size in Groups 16 and 17. Both types of ions increase in size from top to bottom within a group due to an increase in the number of occupied principal energy levels.

## Electronegativity

Valence electrons of both atoms are always involved when those two atoms come together to form a chemical bond. Chemical bonds are the basis for how elements combine with one another to form compounds. When these chemical bonds form, atoms of some elements have a greater ability to attract the valence electrons involved in the bond than other elements. Electronegativity is a measure of the ability of an atom to attract the electrons when the atom is part of a compound. Electronegativity differs from electron affinity because electron affinity is the actual energy released when an atom gains an electron. Electronegativity is not measured in energy units, but is rather a relative scale. All elements are compared to one another, with the most electronegative element, fluorine, being assigned an electronegativity value of 3.98. Fluorine attracts electrons better than any other element. Figure below shows electronegativity values for all elements.

The electronegativity scale was developed by Nobel Prize winning American chemist Linus Pauling. The largest electronegativity (3.98) is assigned to fluorine and all other electronegativities measurements are on a relative scale.

Since metals have few valence electrons, they tend to increase their stability by losing electrons to become cations. Consequently, the electronegativities of metals are generally low. Nonmetals have more valence electrons and increase theirs stability by gaining electrons to become anions. The electronegativities of nonmetals are generally high.

### Trends

Electronegativities generally increase from left to right across a period. This is due to an increase in nuclear charge. Alkali metals have the lowest electronegativities, while halogens have the highest. Because most noble gases do not form compounds, they do not have electronegativities. Note that there is little variation among the transition metals. Electronegativities generally decrease from top to bottom within a group due to the larger atomic size.

## Metallic and Nonmetallic Character

Metallic character refers to the level of reactivity of a metal. Metals tend to lose electrons in chemical reactions, as indicated by their low ionization energies. Within a compound, metal atoms have relatively low attraction for electrons, as indicated by their low ionization energies. By following the trend summary in Figure below, you can see that the most reactive metals would reside in the lower left portion of the periodic table. The most reactive metal is cesium, which is not found in nature as a free element. It reacts explosively with water and will ignite spontaneously in air. Francium is below cesium in the alkali metal group, but is so rare that most of its properties have never been observed.

Nonmetals tend to gain electrons in chemical reactions and have a high attraction for electrons within a compound. The most reactive nonmetals reside in the upper right portion of the periodic table. Since the noble gases are a special group because of their lack of reactivity, the element fluorine is the most reactive nonmetal. It is not found in nature as a free element. Fluorine gas reacts explosively with many other elements and compounds and is considered to be one of the most dangerous known substances.

Summary of periodic trends within periods and groups.

Look at the reactivity of metals in the form of Sumo Wrestlers at http://freezeray.com/flashFiles/ReactivitySumo.htm!

## Lesson Summary

• Atomic radius generally decreases from left to right across a period and increases from top to bottom within a group.
• The number of valence electrons varies from one to eight among the groups of representative elements.
• Metals tend to lose electrons easily to form positively-charged cations, while nonmetals tend to gain electrons to form negatively-charged anions.
• Ionization energy is the energy required to remove an electron. It generally increases from left to right across a period and decreases from top to bottom within a group.
• Electron affinity is the energy released when an atom gains an electron and is highest for the halogen group.
• Cations are smaller than their parent atom, while anions are larger than their parent atom.
• Electronegativity is a measure of the attraction for electrons. It generally increases from left to right across a period and decreases from top to bottom within a group.
• Metallic reactivity is greatest for elements in the lower left area of the periodic table, while nonmetallic reactivity is greatest for elements in the upper right area.

## Lesson Review Questions

### Reviewing Concepts

1. How does atomic radius change from left to right across a period? Explain.
2. How does atomic radius change from top to bottom within a group? Explain.
1. How does ionization energy change from left to right across a period? Explain.
2. How does ionization energy change from top to bottom within a group? Explain.
3. Why is the second ionization energy of an element always larger than the first?
4. Why are most electron affinities negative numbers?
1. Which group of elements has the highest electronegativities?
2. Which group of elements has the lowest electronegativities?

### Problems

1. Which element in each pair below has the larger atomic radius?
1. K, Na
2. S, Cl
3. F, Br
4. Ba, Cs
2. Which equation below shows the second ionization of an alkaline earth metal?
1. $\text{Sr} \rightarrow \text{Sr}^+ + e^-$
2. $\text{Sr}^+ \rightarrow \text{Sr}^{2+} + e^-$
3. $\text{Cs} \rightarrow \text{Cs}^+ + e^-$
4. $\text{Cs}^+ \rightarrow \text{Cs}^{2+} + e^-$
3. Which element in each pair has the higher first ionization energy?
1. Mg, P
2. Se, O
3. Li, Rb
4. Ne, N
4. Why does the third ionization energy (IE3) of calcium show an unusually large jump in value compared to the second ionization energy?
5. Which atom/ion of each pair is larger?
1. Na, Na+
2. Br, Br
3. Be2+, Ca2+
4. F-, O2-
6. Which element in each pair has a higher electronegativity value?
1. F, Cl
2. P, S
3. Sr, Be
4. Al, Na
7. Which equation below represents the electron affinity of a halogen atom?
1. $\text{Cl} \rightarrow \text{Cl}^+ + e^-$
2. $\text{S} + e^- \rightarrow \text{S}^-$
3. $\text{Br} + e^- \rightarrow \text{Br}^-$
4. $\text{Na} \rightarrow \text{Na}^+ + e^-$
8. Arrange the elements in order of increasing metallic character: K, Al, Cs, Na.
9. Arrange the elements in order of increasing nonmetallic character: O, P, F, N.
10. Explain why the Na2+ ion is very unlikely to form.

Aug 02, 2012

Sep 21, 2015