Chapter 4 Periodicity - eadfgfsferg PDF

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INQUIRY QUESTION Are there patterns in the properties of elements?

OUTCOMES

4 Periodicity Students: •

demonstrate, explain and predict the relationships in the observable trends in the physical and chemical properties of elements in periods and groups in the periodic table, including but not limited to: – state of matter at room temperature – electronic configurations and atomic radii – first ionisation energy and electronegativity – reactivity with water.

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Chemistry Stage 6 Syllabus © NSW Educational Standards Authority for and on behalf of the Crown in right of the State of New South Wales, 2017

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FIGURE 4.1 Dmitri Mendeleev (1834−1907) was professor of chemistry at the University of St Petersburg from 1867 to 1890. He developed his periodic table to organise chemistry for his students while writing a textbook for them.

Periodicity is the regular recurrence of events or properties. Sunrise is a periodic event as is high tide. They recur in a regular pattern – sunrise once and high tide twice in 24 hours. The regular occurrence of certain properties led Mendeleev (Figure 4.1) in 1869 to propose his periodic law which in modern terms states that the properties of the elements vary periodically with their atomic numbers. On the basis of this periodicity he gave the name ‘periodic table’ to the chart he had devised to systematise properties of the elements. In this chapter we shall look at the ways that certain properties vary periodically with atomic number. We shall summarise how these properties vary across periods and down groups of the table, and look for explanations in terms of the structure of the atoms of the elements.

4.1 Melting point and physical state One property of the elements that varies periodically with atomic number is melting point. Figure 4.2 shows its variation with atomic number. The curve passes through a series of minima and maxima. The minima correspond to the noble gases but the maxima are less well defined: the first two correspond to the elements carbon and silicon in group 14 while the others correspond to transition metals ingroup 6. FIGURE 4.2 Melting point of elements as a function of atomic number

C

Melting point (ºC)

4000

W Mo

3000 Cr 2000

Si

1000 0 Ne

He 0

10

20

30

40

Rn

Xe

Kr

Ar

50

60

70

80

90

Atomic number

We have already seen (in section 2.5) that most elements are solids at room temperature. The gaseous elements are on the right-hand side of the periodic table, mainly the noble gases and the halogens (groups17 and 18). In these two groups boiling point increases going down the group. In fact for the halogens the physical state at room temperature changes from gas ( for fluorine and chlorine) to liquid ( for bromine) to solid ( for iodine) down the group. There is a change from gas to solid going down groups 15 and 16 also. 97 8017 0408929

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State of matter at room temperature

However for the other groups there are no consistent trends in melting and boiling points down a group: the elements are all just solids. Remember from Investigation 2.3 that at room temperature a substance will be: ◗ ◗

◗

a solid if its melting point is greater than room temperature a liquid if its melting point is less than room temperature but its boiling point is greater than room temperature a gas if its boiling point is less than room temperature.

4.2 Atomic radius

FIGURE 4.3 Atomic radii of some elements as a function of atomic number

Atomic radius (picometres (10212 m))

When atomic radius is plotted against atomic number, as in Figure 4.3, the curve shows a distinctly periodic nature. The atomic radius passes through a set of sharp maxima corresponding to the alkali metals (Li, Na, K, Rb, Cs – group 1). The minima occur at the noble gases (Ne, Ar, Kr, Xe, Rn – group 18). This means that atomic radius decreases from left to right across any period of the table. Values in Figure4.4 further illustrate this.

Fr

Cs

Rb K Na 200 Li

Rn 100

Xe Kr

Ar He

0

10

Ne

20

30

40

50

60

70

80

90

Atomic number

FIGURE 4.4 Atomic radii (in picometres) of some elements. 1 picometre (pm) = 10−12 m.

76

H 37

He 50

Li 152

Be 112

B 88

C 77

N 70

O 66

F 68

Ne 70

Na 186

Mg 160

Al 143

Si 118

P 110

S 102

Cl 99

Ar 94

K 231

Ca 197

As 123

Se 116

Br 114

Kr 109

Rb 244

Sr 215

I 133

Xe 130

Cs 262

Ba 217

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Rn 140

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Figure 4.4 shows more clearly than does Figure 4.3 that atomic radius increases down any group of the periodic table. It decreases from left to right across any period of the periodic table and increases from top to bottom in any group.

Explaining trends in atomic radius To explain the decrease in atomic radius in going across a period the key facts are as follows. ◗

◗

In any period of the table all the outermost electrons are in the same main energy level and so have approximately the same energy. Going across a period the number of protons in the nucleus is increasing.

Hence, going across a period, the force of attraction between the nucleus and each electron in the outermost shell is increasing. This will draw each electron closer to the nucleus and so make the radius of the atom smaller. So atomic radius decreases from left to right across a period. To explain the increase in atomic radius going down a group, we note that the positive charge (number of protons) in the nucleus is increasing, but the energy of the outermost electrons is also increasing, because they are in higher electron energy levels. If an electron has more energy, it is better able to resist the stronger electrostatic attraction of the nucleus. These two effects, increasing nuclear charge and increasing electron energy, tend to cancel each other out. Another factor comes into play: the screening effect. This is the decrease in electrostatic force between a nucleus and an outermost electron brought about by completely filled electron shells between the nucleus and the outermost electron. This ‘sea’ of negative charge tends partly to cancel out the effect of the positive nucleus, so the nucleus is able to exert less attraction upon each outermost electron than if there were no intervening electron shells. This screening effect weakens the attraction between the nucleus and each outer electron and so the radius becomes larger. Hence atomic radius increases from top to bottom of a group in the table.

Electron configuration and atomic radii

4.3 Ionisation energy When an atom loses or gains an electron, it becomes an electrically charged species because the numbers of protons and electrons are no longer in balance. Such charged species, are called ions. When an atom loses an electron it becomes a positive ion. For example when a sodium atom loses an electron it becomes a sodium ion, which we write as Na+. This process of losing an electron is called ionisation. Energy is required to remove an electron from an atom; we call this energy the ionisation energy. The first ionisation energy,IA, of an element is the energy required to remove an electron from a neutral gaseous atom of the element. It is the energy change for the process: M(g) → M+(g) + e− where M is any element. The lower the ionisation energy, the easier it is to remove an electron. While ionisation energy can be measured in joules per atom (where a joule is the unit of energy), it is more commonly reported in kilojoules per mole of atoms (kJ mol−1). For the moment consider a mole of an element as a specific number of atoms, just like a dozen of things is always 12. A mole is 6 × 1023 atoms.

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The meaning and significance of the term ‘mole’ will be explained in chapter 7.

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Each element has several ionisation energies, called the first, second, third (and so on) ionisation energies. The second ionisation energy is the energy required to remove a second electron: M+(g) → M²+(g) + e− and so on. The second ionisation energy is always greater than the first. This is because more energy is needed to remove a negative electron from a positive species than from a neutral species. This is due to having to overcome greater electrostatic attraction.

INVESTIGATION

4.1

Trends in ionisation energies across a period and down a group Information and communication technology capability

AIM

To observe trends in first ionisation energy across period 3 and down group 1. MAT ERIAL S

• Graph paper or access to a spreadsheet program • Data for first ionisation energies for period 3 elements (Na to Ar) and for group 1 elements (Li to Cs) MET HOD

1 Plot a graph of first ionisation energies for period 3 elements, putting atomic number on the x-axis and first ionisation energy on the y-axis. 2 Similarly, for group 1 elements plot a graph of first ionisation energies against period number. ANALY SIS OF RE SULT S

1 Describe the trend in first ionisation energies across period 3. 2 Describe the trend in first ionisation energies down group 1. CONCLUSION

Offer an explanation for the trend in first ionisation energies: • across period 3. • down group 1.

Ionisation energy and the periodic table Ionisation energy shows quite distinct variation across and down the periodic table. In Figure 4.5 first ionisation energy is plotted against atomic number. Periodicity is obvious. The minimum values of IA are for the alkali metals: it is relatively easy to remove an electron from Li, Na, K, Rb and Cs. The noble gases have the maximum values of IA: it is difficult to remove an electron from these elements. This confirms that noble gas configurations are extremely stable. Elements with one or two extra electrons tend to lose them to form noble gas configurations. First ionisation energies provide further strong support for the periodic law that properties vary periodically with atomic number.

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FIGURE 4.5 First ionisation energy as a function of atomic number

He

First ionisation energy (kJ mol21)

2400

Ne

2000 Ar Kr

1600

Xe Rn 1200 800 400

Li 0

Na 10

K 20

Rb 30

40

Cs 50

60

70

80

90

Atomic number

Another way of presenting ionisation energy data is as shown in Figure 4.6. These data clearly demonstrate that in going across any period of the periodic table the first ionisation energy increases. In moving from left to right, there is less tendency to lose electrons. In going down any group of the table, ionisation energy decreases: in going down any group, electrons are lost more easily. Investigation 4.1 confirmed these trends.

H 1320

He 2380

Li 526

Be 905

B 810

C 1090

N 1410

O 1320

F 1690

Ne 2090

Na 504

Mg 740

Al 580

Si 790

P 1020

S 1000

Cl 1260

Ar 1526

K 425

Ca 600

As 953

Se 950

Br 1150

Kr 1360

Rb 410

Sr 560

I 1020

Xe 1180

Cs 380

Ba 510

FIGURE 4.6 First ionisation energies (in kJ mol−1) for some elements

Rn 1040

Ionisation energy and the drive to noble gas configurations Ionisation energies provide striking confirmation of the drive to noble gas configurations discussed in section 3.4. In Figure 4.7 successive ionisation energies are plotted against the number of the electrons being removed for some elements in groups 2 and 13. In Figure 4.7a, we note that it is relatively easy to remove one or two electrons from Be, Mg and Ca, but that to remove the third requires much more energy. In Figure 4.7b it is relatively easy to remove one, two and three electrons from the group 13 elements, but a far larger amount of energy is needed to remove the fourth electron.

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FIGURE 4.7 Successive ionisation energies for elements of a group 2 and b group 13

b

a

16 000

Ionisation energy (kJ mol21)

Ionisation energy (kJ mol21)

16 000 Be 12 000

Mg

8000

Ca 4000

1st

2nd

3rd

12 000

B Al

8000 Ga 4000

1st

4th

2nd

Electron removed

3rd

4th

Electron removed

Table 4.1 shows a similar situation for group 1: it is easy to remove one electron but quite difficult to remove a second electron. These observations support the ‘lose’ part of the claim on page 56 that elements tend to lose or gain electrons to obtain the electron configuration of the nearby noble gas. Once one, two or three electrons have been removed from elements in groups 1, 2 and 13 respectively, the ions have noble gas configurations. It is extremely hard to remove further electrons from these atoms because they would have to come from lower energy shells, which are more strongly held electrostatically to the nucleus. TABLE 4.1 Successive ionisation energies (kJ mol−1) for group 1 elements ELEMENT

FIRST

SECOND

THIRD

FOURTH

Li

526

7300

11 800

Na

504

4570

6920

9550

K

425

3080

4400

5880

Rb

410

2660

3900

5100

Cs

380

2430

3400

4900

Explaining trends in ionisation energies A first step in explaining the trends in ionisation energies across and down the periodic table is to note the very close parallel between atomic radius and first ionisation energy: as atomic radius increases, first ionisation energy decreases. Trends in atomic radii were explained in terms of the strength of the attractive force between the nucleus and each electron in the outermost energy level. The stronger the electrostatic attraction the smaller the atomic radius. The stronger the electrostatic attraction, the greater the ionisation energy: more energy is required to remove an electron. On page 79 it was explained that in going across a period the electrostatic attraction between the nucleus and each outermost electron increases. This means that the first ionisation energy increases from left to right across a period. It was also explained that this electrostatic attraction decreases going down a group, so the ionisation energy also decreases from top to bottom of a group.

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KEY CONCEPTS

●

Periodicity is the regular recurrence of events or properties.

●

The periodic law is that the properties of the elements vary periodically with their atomic numbers.

●

Atomic radius decreases from left to right across any period of the periodic table and increases from top to bottom in any group.

●

The screening effect is the decrease in electrostatic force between a nucleus and an outermost electron brought about by completely filled electron shells between the nucleus and the outermost electron.

●

When an atom loses or gains an electron it becomes an electrically charged species because the numbers of protons and electrons are no longer in balance. Such charged species are called ions.

●

The first ionisation energy, IA , of an element is the energy required to remove an electron from a neutral gaseous atom of the element. It is the energy change for the process: M(g) → M+(g) + e−

●

First ionisation energies increase from left to right across any period of the periodic table and decrease from top to bottom in any group.

●

As atomic radius increases, first ionisation energy decreases.

CHECK YOUR UNDERSTANDING

1 Describe how melting point varies across a period of the periodic table. 2 Identify where gases (at room temperature) are in the periodic table.

4.1

3 Outline the way that values of successive ionisation energies support the theory that, in forming compounds, elements try to obtain noble gas configurations.

4.2

4.3

4 Propose an explanation for the: a increase in first ionisation energy going from left to right across a period. b decrease in first ionisation energy going from top to bottom of a group. 5 a Use a data book or search the internet for melting and boiling points of the halogens and for boiling points of the noble gases. b Draw graphs of these melting and boiling points against period number. c

Identify any trends you see in these properties.

6 a Use a data book or search the internet for melting points of the elements of groups 2 and 14. b For each of these groups prepare tables of period number and melting point. c

Identify trends (if any) that you see in these data.

Answer questions 7 and 8 without consulting any tables or diagrams. 7 Four elements of atomic numbers 17, 18, 19 and 20 have first ionisation energies of 1500, 600, 400 and 1300 kJ mol−1 (listed in random order). Which first ionisation energy belongs to which element? Explain how you reached your decision. 8 Three elements of group 1 have first ionisation energies of 380, 530 and 430 kJ mol−1. List the elements in order of increasing atomic number. 9 Four elements have the following successive ionisation energies (in kJ mol−1). A: 500 4600 6900 9500 B: 600 1150 4900 6500 C: 2100 3960 6100 9400 D: 380 2400 3400 4400

To which group of the periodic table does each belong? Justify your decisions. 10 The claim was made on page 80 that first ionisation energy increases as atomic radius decreases. Demonstrate that this is true by plotting (on one graph) first ionisation energy versus atomic radius for the main-group elements of periods 3 and 4, and Rb, Cs, Sr and Ba. Use data from Figure 4.6 and Table 4.1.

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4.4 Electronegativit...