PERIODICITY - Wikispacesiknsphysicsib.wikispaces.com/file/view/periodicity.pdftopic 3 periodicity . ... chemistry: for use with the ib diploma programme standard level chapter 3 periodicity

T O P I C 3

PERIODICITY

PERIODIC TABLE

•  Mendeleev’s periodic table left gaps for elements that he believed should exist because the elements on either side of the gap matched the expected chemical properties of their groups.

PERIODIC TABLE

CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL

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• The long metal periods are labelled transition metals, lanthanides and actinides.

• There is diffi culty placing hydrogen in the table. It is placed at the top of group 1 because it has some similarities to other group 1 elements, but it is a non-metal.

• Unlike the other noble gases, helium has only 2 electrons in its valence shell, but as this is shell 1, this constitutes a full valence shell.

Group

Period 1

2

3

4

5

6

7

21 4 5 6 7 03

transition metals

57La

89Ac

58Ce

90Th

92U

transuranium elements

dividing line betweenmetals and non-metals halogens

noblegases

actinides89–103

lanthanides57–71

alkalimetals

alkaliearth

metals

1H

2He

Figure 3.1.4 The modern periodic table.

Groups 3 to 6 contain both metal and non-metal elements. Seven elements— boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te) and polonium (Po)—are called metalloids because they have characteristics of both metal and non-metal elements. Most of the elements are solid at room temperature and 1 atm pressure, except for the 11 gaseous elements—hydrogen, nitrogen, oxygen, fl uorine, chlorine, and all the group 0 elements—and the two liquid elements, bromine and mercury.

The arrangement of the elements in the periodic table is linked to their electron confi gurations. Elements with the same outer-shell electron confi guration belong to the same group of the periodic table. The alkali metals (group 1), for example, all have one electron in their outer shell. The electron arrangements of some of the alkali metals are:

Lithium 2,1Sodium 2,8,1Potassium 2,8,8,1Rubidium 2,8,18,8,1

The similarity in the number of outer-shell electrons of each of these alkali metals leads to similarities in their physical properties and their chemical reactions. They all react readily and need to be stored in oil, they have low melting points and are relatively soft compared to other metals, and all form ions with a charge of 1+.

Physical properties of the halogens

3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20. © IBO 2007

PERIODIC TABLE

•  Seven elements— boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te) and polonium (Po)—are called metalloids because they have characteristics of both metal and non-metal elements. •  Most of the elements are solid at room temperature

and 1 atm pressure, except for the 11 gaseous elements—hydrogen, nitrogen, oxygen, fluorine, chlorine, and all the group 0 elements—and the two liquid elements, bromine and mercury.

PERIODIC TABLE

•  Elements with the same outer-shell electron configuration belong to the same group of the periodic table. •  leads to similarities in their physical properties and

their chemical reactions

PHYSICAL PROPERTIES

•  Periodicity: The repetition of properties at regular intervals within the periodic table. •  Physical property: A characteristic that can be

determined without changing the chemical composition of the substance. [give examples]


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Terms and definitions

Alkali metals The name given to group 1 of the periodic table.

Alkaline earth metals The name given to group 2 of the periodic table.

Amphoteric Able to act as an acid or a base.

Atomic radius The distance from the centre of the nucleus to the outermost electron shell.

Chemical property A characteristic that is exhibited as one substance is chemically transformed into another.

Core charge The effective nuclear charge experienced by the outer-shell electrons. Core charge is the difference between the nuclear charge and the number of inner-shell electrons.

Electronegativity A measure of the attraction that an atom has for a shared pair of electrons when it is covalently bonded to another atom.

First ionization energy The amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Group A vertical column within the periodic table.

Halide ion A negative ion formed when a halogen atom gains one electron.

Halogens The name given to group 7 of the periodic table.

Highest oxide The compound formed with oxygen in which the element is in its highest possible oxidation state.

Noble gases The gaseous elements of group 0 of the periodic table, all of which have a full valence shell.

Oxidizing agent A substance which causes another substance to be oxidized by accepting electrons from it.

Period Horizontal row within the periodic table.

Periodicity The repetition of properties at regular intervals within the periodic table.

Physical property A characteristic that can be determined without changing the chemical composition of the substance.

Reducing agent A substance that causes another substance to be reduced by donating electrons to it.

Concepts

• The modern periodic table arranges elements horizontally in order of atomic number and vertically in groups with the same outer-shell electron confi guration and similar chemical reactivity.

Elements listed in order

Mendeleev’s table

Modern table

Elements with similarchemical properties

Elements with the samevalence-shell configuration

Elements in order ofincreasing atomic mass

Elements in order ofincreasing atomic number

• A group is a vertical column and a period is a horizontal row in the periodic table.


alkalimetals

alkaliearthmetals

57La

58Ce90Th

92U

89Ac

2He

1H

lanthanides57–71

actinides89–103

1Group 2 3 4 5 6 7 0

Period 1

2

3

4

5

6

7 dividing line betweenmetals and non-metals

halogensnoblegases

transition metals

• There are trends in the properties of the elements both in groups and across periods.

atomic radius decreases•

electronegativity increases••

first ionization energy increases••

ionic radius decreases•

• The patterns in the periodic table are a consequence of the valence-shell electron confi gurations of the elements. Elements in the same group behave similarly because they have the same valence-shell confi guration. Elements in the same period have electrons fi lling the same electron shell (energy level).

Chapter 3 Summary

PHYSICAL PROPERTIES

•  Example •  the melting point of the element sulfur can be found by

determining the temperature at which it turns from a solid to a liquid. The sulfur changes state only; its chemical composition does not alter.


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Terms and definitions

Alkali metals The name given to group 1 of the periodic table.

Alkaline earth metals The name given to group 2 of the periodic table.

Amphoteric Able to act as an acid or a base.

Atomic radius The distance from the centre of the nucleus to the outermost electron shell.

Chemical property A characteristic that is exhibited as one substance is chemically transformed into another.

Core charge The effective nuclear charge experienced by the outer-shell electrons. Core charge is the difference between the nuclear charge and the number of inner-shell electrons.

Electronegativity A measure of the attraction that an atom has for a shared pair of electrons when it is covalently bonded to another atom.

First ionization energy The amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Group A vertical column within the periodic table.

Halide ion A negative ion formed when a halogen atom gains one electron.

Halogens The name given to group 7 of the periodic table.

Highest oxide The compound formed with oxygen in which the element is in its highest possible oxidation state.

Noble gases The gaseous elements of group 0 of the periodic table, all of which have a full valence shell.

Oxidizing agent A substance which causes another substance to be oxidized by accepting electrons from it.

Period Horizontal row within the periodic table.

Periodicity The repetition of properties at regular intervals within the periodic table.

Physical property A characteristic that can be determined without changing the chemical composition of the substance.

Reducing agent A substance that causes another substance to be reduced by donating electrons to it.

Concepts

• The modern periodic table arranges elements horizontally in order of atomic number and vertically in groups with the same outer-shell electron confi guration and similar chemical reactivity.

Elements listed in order

Mendeleev’s table

Modern table

Elements with similarchemical properties

Elements with the samevalence-shell configuration

Elements in order ofincreasing atomic mass

Elements in order ofincreasing atomic number

• A group is a vertical column and a period is a horizontal row in the periodic table.


alkalimetals

alkaliearthmetals

57La

58Ce90Th

92U

89Ac

2He

1H

lanthanides57–71

actinides89–103

1Group 2 3 4 5 6 7 0

Period 1

2

3

4

5

6

7 dividing line betweenmetals and non-metals

halogensnoblegases

transition metals

• There are trends in the properties of the elements both in groups and across periods.

atomic radius decreases•

electronegativity increases••

first ionization energy increases••

ionic radius decreases•

• The patterns in the periodic table are a consequence of the valence-shell electron confi gurations of the elements. Elements in the same group behave similarly because they have the same valence-shell confi guration. Elements in the same period have electrons fi lling the same electron shell (energy level).

Chapter 3 Summary

GROUPS 1 & 7

•  Alkali and halogens •  Electronegativity (define)?

GROUPS 1 & 7

•  Alkali and halogens •  Electronegativity: •  is a measure of the attraction an atom has for a shared pair

of electrons when it is covalently bonded to another atom. •  Metals à few electrons in their outer shell therefore they

lose it [low or high electronegativity???] •  Non-metals à will want to complete outer shell so they will

gain electron(s) [low or high electronegativity???]

GROUPS 1 & 7

•  Alkali and halogens •  Electronegativity: •  is a measure of the attraction an atom has for a shared pair

of electrons when it is covalently bonded to another atom. •  Metals à few electrons in their outer shell therefore they

lose it [low electronegativity] •  Non-metals à will want to complete outer shell so they will

gain electron(s) [high electronegativity]

GROUPS 1 & 7

82

TABLE 3.2.1 PROPERTIES OF THE ALKALI METALS AND THE HALOGENS

Element Atomic radius (10!12 m)

Ionic radius (10!12 m)

Electronegativity First ionization energy (kJ mol!1)

Melting point (°C)

Alkali metals

Lithium 152 68 1.0 519 181

Sodium 186 98 0.9 494 98

Potassium 231 133 0.8 418 64

Rubidium 244 148 0.8 402 39

Caesium 262 167 0.7 376 29

Halogens

Fluorine 58 133 4.0 1680 –219

Chlorine 99 181 3.0 1260 –101

Bromine 114 196 2.8 1140 –7

Iodine 133 219 2.5 1010 114

In fi gure 3.2.1, the arrows indicate that electronegativity increases from left to right across a period and from bottom to top in a group. If the electronegativities of two elements in the same period are compared, the element on the right will have the greater electronegativity. If the electronegativities of two elements in the same group are compared, the element that is higher in the group will have the greater electronegativity.

First ionization energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state. The outer-shell electron is the most easily removed, and so ionization energy is a measure of how tightly the outer-shell electrons are held in an atom. In general, electrons in metals are easily removed, so metals have low ionization energies. Non-metals hold electrons strongly, and so have high ionization energies. See fi gure 3.2.1.

Ionization of an atom can be represented by an equation. The following equation describes the fi rst ionization of sodium:

Na(g) " Na+(g) + e!

How can we explain the differences in electronegativity and fi rst ionization energy within a group? Members of group 1 are all metals and so have low fi rst ionization energies. Why do the fi rst ionization energies vary as we go down the group?

Two main factors determine how tightly an outer-shell electron is held. The force of electrostatic attraction between the positive protons in the nucleus and the outer-shell electron is directly related to the charges and inversely related to the distance between them. As the size of atoms increases, the attractive force on the outer-shell electron decreases. As the nuclear charge increases, the attractive force increases. There is, however, a complication: the outer-shell electrons are ‘shielded’ from the full nuclear charge by the inner-shell electrons. The concept of core charge is used to allow for this shielding. The effective nuclear charge felt by the outer-shell electrons, the core charge, may be found by subtracting the number of inner-shell electrons from the nuclear charge. For example, the core charge of some group 1 elements can be determined by:

Sodium: 11 protons and 10 inner-shell electrons = core charge of +1Potassium: 19 protons and 18 inner-shell electrons = core charge of +1Rubidium: 37 protons and 36 inner-shell electrons = core charge of +1

3.2.4Compare the relative electronegativity values of two or more elements based on their positions in the periodic table. © IBO 2007

Periodic trends: ionization energy

Compare across a period and down a group

GROUPS 1 & 7

82





Melting point (°C)

Alkali metals

Lithium 152 68 1.0 519 181

Sodium 186 98 0.9 494 98

Potassium 231 133 0.8 418 64

Rubidium 244 148 0.8 402 39

Caesium 262 167 0.7 376 29

Halogens

Fluorine 58 133 4.0 1680 –219

Chlorine 99 181 3.0 1260 –101

Bromine 114 196 2.8 1140 –7

Iodine 133 219 2.5 1010 114




Na(g) " Na+(g) + e!






Compare across a period and down a group

GROUPS 1 & 7

•  First ionization energy •  is the energy required to remove one mole of electrons from

one mole of atoms in the gaseous state. •  ionization energy is a measure of how tightly the outer-shell

electrons are held in an atom. •  Metals à electrons easy or hard to remove?

GROUPS 1 & 7

•  First ionization energy •  is the energy required to remove one mole of electrons from

one mole of atoms in the gaseous state. •  ionization energy is a measure of how tightly the outer-shell

electrons are held in an atom. •  Metals à electrons easy to remove therefore less ionization

energy •  Non-metals à hold hard to their electrons therefore more

ionization energy

GROUPS 1 & 7

•  First ionization energy •  Equation example:

•  Factors that determine how tightly an outer-shell electron is held •  Attractive force between protons and electrons •  Distance between protons and electrons •  Hardest: more force and less distance

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Melting point (°C)

Alkali metals

Lithium 152 68 1.0 519 181

Sodium 186 98 0.9 494 98

Potassium 231 133 0.8 418 64

Rubidium 244 148 0.8 402 39

Caesium 262 167 0.7 376 29

Halogens

Fluorine 58 133 4.0 1680 –219

Chlorine 99 181 3.0 1260 –101

Bromine 114 196 2.8 1140 –7

Iodine 133 219 2.5 1010 114




Na(g) " Na+(g) + e!






GROUPS 1 & 7


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540

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519 900

494 736

418 590

402 548

632

636

376

661

669

531

648

653

760

653

694

770

716

699

762

762

724

841

757

745

887

736

803

866

745

732

891

908

866

1010

799

577

577

556

590

1090

786

762

707

716

1400

1060

966

833

703

1310(O– +844)

1000(S– +532)

941

870

812

1680

1260

1140

1010

920

2370

2080

1520

1350

1170

1040502

181

2.1

First ionizationenergy (kJ mol–1)

Electronegativity

1.0

0.9

0.8

0.8

0.7

0.7

1.5

1.2

1.0

1.0

0.9

0.9

1.3

1.2

1.1

1.1

1.5

1.4

1.3

1.6

1.6

1.5

1.6

1.8

1.7

1.5

1.9

1.9

1.8

2.2

2.2

1.8

2.2

2.2

1.8

2.2

2.2

1.9

1.9

2.4

1.6

1.7

1.9

2.0

1.5

1.6

1.7

1.8

2.5

1.8

1.8

1.8

1.8

3.0

2.1

2.0

1.9

1.9

3.5

2.5

2.4

2.1

2.0

4.0

3.0

2.8

2.5

2.2

510

La

Ac

H

Element

Li Be

Na Mg

K Ca

Rb Sr

Sc

Y

Hf

Ti

Zr

Ta

V

Nb

W

Cr

Mo

Re

Mn

Tc

Os

Fe

Ru

Ir

Co

Rh

Pt

Ni

Pd

Au

Cu

Ag

Hg

Zn

Cd

Tl

Ga

In

Pb

Ge

Sn

Bi

As

Sb

Po

Se

Te

At

Br

I

Rn

Kr

Al Si P S Cl Ar

B C N O F Ne

Xe

Cs Ba

Fr Ra

Increasing electronegativity and first ionization energy

Incr

easi

ng e

lect

rone

gativ

ity a

nd f

irst

ioni

zatio

n en

ergy

He

Figure 3.2.1 Electronegativity and first ionization energy values, showing trends within the periodic table. A similar table can be found in the IB Data Booklet © IBO 2007.

188

200

30

152 112

186 160

231 197

244 215

160

180

262

146

157

157

131

141

143

125

136

137

129

135

137

126

133

134

125

134

135

124

138

138

128

144

144

133

149

152

88

143

141

166

171

77

117

122

162

175

70

110

121

141

170

66

104

117

137

140

58

99

114

133

140217

270

154 (1–)

Atomicradius(10–12 m)

Ionicradius(10–12 m)

68 (1+)

98 (1+)

133 (1+)

148 (1+)

167 (1+)

30 (2+)

65 (2+)

94 (2+)

110 (2+)

34 (2+)

81 (3+)

93 (3+)

115 (3+)

90 (2+)68 (4+)

80 (4+)

81 (4+)

88 (2+)59 (5+)

70 (5+)

73 (5+)

63 (3+)

68 (4+)

68 (4+)

80 (2+)60 (4+)

76 (2+)64 (3+)

65 (4+)

67 (4+)

74 (2+)63 (3+)

86 (2+)

66 (4+)

72 (2+) 96 (1+)69 (2+)

126 (1+)

137 (1+)85 (3+)

74 (2+)

97 (2+)

127 (1+)110 (2+)

62 (3+)

81 (3+)

95 (3+)

53 (4+)272 (4–)

112 (2+)71 (4+)

120 (2+)84 (4+)

222 (3–)

245 (3–)

120 (3+)

202 (2–)

222 (2–)

196 (1–)

45 (3+) 42 (4+)271 (4–)

212 (3–) 190 (2–) 181 (1–)

16 (3+) 260 (4–) 171 (3–) 146 (2–) 133 (1–)

219 (1–)

220

La

Ac

H

Element

Li Be

Na Mg

K Ca

Rb Sr

Sc

Y

Hf

Ti

Zr

Ta

V

Nb

W

Cr

Mo

Re

Mn

Tc

Os

Fe

Ru

Ir Pt

Co

Rh

Ni

Pd

Au

Cu

Ag

Hg

Zn

Cd

Tl

Ga

In

Pb

Ge

Sn

Bi

As

Sb

Po

Se

Te

At

Br

I

Rn

Kr

Al Si P S Cl Ar

B C N O F Ne

Xe

Cs Ba

Fr Ra

Decreasing atomic and ionic radii

Incr

easi

ng a

tom

ic a

nd io

nic

radi

i

He

Figure 3.2.2 Atomic and ionic radii values showing trends within the periodic table. A similar table can be found in the IB Data Booklet © IBO 2007.

GROUPS 1 & 7

•  Atomic radii and ionic radii (what will the trend be across a period and down a group)

GROUPS 1 & 7

•  Atomic radii and ionic radii •  Group: increases as electron shells increase •  Period: decrease as the charges increase and hence the

attractive force increases

GROUPS 1 & 7


CH

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83

540

669

1310

519 900

494 736

418 590

402 548

632

636

376

661

669

531

648

653

760

653

694

770

716

699

762

762

724

841

757

745

887

736

803

866

745

732

891

908

866

1010

799

577

577

556

590

1090

786

762

707

716

1400

1060

966

833

703

1310(O– +844)

1000(S– +532)

941

870

812

1680

1260

1140

1010

920

2370

2080

1520

1350

1170

1040502

181

2.1

First ionizationenergy (kJ mol–1)

Electronegativity

1.0

0.9

0.8

0.8

0.7

0.7

1.5

1.2

1.0

1.0

0.9

0.9

1.3

1.2

1.1

1.1

1.5

1.4

1.3

1.6

1.6

1.5

1.6

1.8

1.7

1.5

1.9

1.9

1.8

2.2

2.2

1.8

2.2

2.2

1.8

2.2

2.2

1.9

1.9

2.4

1.6

1.7

1.9

2.0

1.5

1.6

1.7

1.8

2.5

1.8

1.8

1.8

1.8

3.0

2.1

2.0

1.9

1.9

3.5

2.5

2.4

2.1

2.0

4.0

3.0

2.8

2.5

2.2

510

La

Ac

H

Element

Li Be

Na Mg

K Ca

Rb Sr

Sc

Y

Hf

Ti

Zr

Ta

V

Nb

W

Cr

Mo

Re

Mn

Tc

Os

Fe

Ru

Ir

Co

Rh

Pt

Ni

Pd

Au

Cu

Ag

Hg

Zn

Cd

Tl

Ga

In

Pb

Ge

Sn

Bi

As

Sb

Po

Se

Te

At

Br

I

Rn

Kr

Al Si P S Cl Ar

B C N O F Ne

Xe

Cs Ba

Fr Ra

Increasing electronegativity and first ionization energy

Incr

easi

ng e

lect

rone

gativ

ity a

nd f

irst

ioni

zatio

n en

ergy

He

Figure 3.2.1 Electronegativity and first ionization energy values, showing trends within the periodic table. A similar table can be found in the IB Data Booklet © IBO 2007.

188

200

30

152 112

186 160

231 197

244 215

160

180

262

146

157

157

131

141

143

125

136

137

129

135

137

126

133

134

125

134

135

124

138

138

128

144

144

133

149

152

88

143

141

166

171

77

117

122

162

175

70

110

121

141

170

66

104

117

137

140

58

99

114

133

140217

270

154 (1–)

Atomicradius(10–12 m)

Ionicradius(10–12 m)

68 (1+)

98 (1+)

133 (1+)

148 (1+)

167 (1+)

30 (2+)

65 (2+)

94 (2+)

110 (2+)

34 (2+)

81 (3+)

93 (3+)

115 (3+)

90 (2+)68 (4+)

80 (4+)

81 (4+)

88 (2+)59 (5+)

70 (5+)

73 (5+)

63 (3+)

68 (4+)

68 (4+)

80 (2+)60 (4+)

76 (2+)64 (3+)

65 (4+)

67 (4+)

74 (2+)63 (3+)

86 (2+)

66 (4+)

72 (2+) 96 (1+)69 (2+)

126 (1+)

137 (1+)85 (3+)

74 (2+)

97 (2+)

127 (1+)110 (2+)

62 (3+)

81 (3+)

95 (3+)

53 (4+)272 (4–)

112 (2+)71 (4+)

120 (2+)84 (4+)

222 (3–)

245 (3–)

120 (3+)

202 (2–)

222 (2–)

196 (1–)

45 (3+) 42 (4+)271 (4–)

212 (3–) 190 (2–) 181 (1–)

16 (3+) 260 (4–) 171 (3–) 146 (2–) 133 (1–)

219 (1–)

220

La

Ac

H

Element

Li Be

Na Mg

K Ca

Rb Sr

Sc

Y

Hf

Ti

Zr

Ta

V

Nb

W

Cr

Mo

Re

Mn

Tc

Os

Fe

Ru

Ir Pt

Co

Rh

Ni

Pd

Au

Cu

Ag

Hg

Zn

Cd

Tl

Ga

In

Pb

Ge

Sn

Bi

As

Sb

Po

Se

Te

At

Br

I

Rn

Kr

Al Si P S Cl Ar

B C N O F Ne

Xe

Cs Ba

Fr Ra

Decreasing atomic and ionic radii

Incr

easi

ng a

tom

ic a

nd io

nic

radi

i

He

Figure 3.2.2 Atomic and ionic radii values showing trends within the periodic table. A similar table can be found in the IB Data Booklet © IBO 2007.

GROUPS 1 & 7

•  Melting and boiling points: •  differ in the trends they exhibit. •  Depends on the type of bonding exhibited by each group. •  The stronger the bonding within a substance, the higher the

melting point.

GROUPS 1 & 7

•  Melting and boiling points: •  Group 1: •  All metals therefore metallic bonding •  Gets weaker as you go down as same number of delocalized

electrons but more distance therefore weaker electrostatic force à hence lower melting/boiling points

•  Group 7: •  Diatomic – therefore non-polar (why? – what forces will it

have??)

GROUPS 1 & 7

•  Melting and boiling points: •  Group 1: •  All metals therefore metallic bonding •  Gets weaker as you go down as same number of

delocalized electrons but more distance therefore weaker electrostatic force à hence lower melting/boiling points

•  Group 7: •  Diatomic – therefore non-polar •  Van der waals forces •  More electrons down a group so more force and therefore

higher melting/boiling point

GROUPS 1 & 7

84

The core charge remains constant within a group. This means that within a group the only factor affecting the electrostatic attraction between outer-shell electrons and the nucleus is the distance of the outer shell from the nucleus. As the atomic number increases within a group (going down the group), the attractive force between the nucleus and the outer-shell electrons decreases.

Atomic radii and ionic radii increase going down groups 1 and 7 because there in an increase in the number of electron shells surrounding the nucleus as you go down these groups (indeed any group). The electron shells account for most of the volume of the atom, so an extra electron shell increases the atomic or ionic radius.

Both fi rst ionization energy and electronegativity decrease down groups 1 and 7 because of the decreasing electrostatic attraction between the outer-shell electrons and the nucleus. This is due to the increasing distance of the outer shell from the nucleus. The smaller the attraction of the outer shell to the nucleus, the easier it is to remove an electron from the outer shell (fi rst ionization energy) and the harder it is to attract an electron to the outer shell (electronegativity).

The melting points of group 1 and group 7 elements differ in the trends they exhibit. This can be attributed to the type of bonding exhibited by each group. You will recall from chapter 2 that the strength of the bonding within a substance governs its melting and boiling points.

The stronger the bonding within a substance, the higher the melting point.

1194

1320

14

454 1551

371 922

337 1112

312 1042

1814

1780

302

1933

2125

2503

1973

2741

3269

2130

2890

3680

1517

2445

3453

1808

2583

3327

1768

2239

2683

1726

1825

2045

1357

1235

1338

693

594

234

2573

936

303

429

577

4100

1683

1211

505

601

63

317

1090

904

545

55

392

490

723

527

54

172

266

387

575983

300

20

Melting point (K)

Boiling point (K)

1600

1156

1047

961

952

950 1413 3470

3243

1363

1757

1657

2023

3104

3611

3730

3560

4650

5470

3650

5015

5698

2755

4885 5150

5930 5900

2235 3023

4173

5300

3143

4000 3413

4403

3005

4100

2840

2485

3080

1180

1038

630

2676

2353

1730

3103

2543

2013

886

2023

1833 1235 610 211

958

1263

332

2740 2628 553 718 239

3931 5100 77 90 85

458

25

84

117

161

202

121

87

27

1

4

166

973

La

Ac

H

Element

Li Be

Na Mg

K Ca

Rb Sr

Sc

Y

Hf

Ti

Zr

Ta

V

Nb

W

Cr

Mo

Re

Mn

Tc

Os

Fe

Ru

Ir Pt

Co

Rh

Ni

Pd

Au

Cu

Ag

Hg

Zn

Cd

Tl

Ga

In

Pb

Ge

Sn

Bi

As

Sb

Po

Se

Te

At

Br

I

Rn

Kr

Al Si P S Cl Ar

B C N O F Ne

Xe

Cs Ba

Fr Ra

Incr

easi

ng m

eltin

g po

int Increasing m

elting point

He

Figure 3.2.3 Melting and boiling points of elements, showing trends within the periodic table. A similar table can be found in the IB Data Booklet © IBO 2007.

3.2.2Describe and explain the trends in atomic radii, ionic radii, first ionization energy, electronegativities and melting points for the alkali metals (Li ! Cs) and the halogens (F ! I). © IBO 2007

Periodic trends: atomic radii

GROUPS 1 & 7


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The members of group 1 are all metals. As the atomic number increases down the group, the size of the positive metallic ion increases; however, the number of delocalized electrons does not change, nor does the charge on the ion. This results in a weaker electrostatic attraction between the ions and the delocalized electrons and a weaker metallic bond. Consequently, the melting point decreases down the group.

The group 7 elements, the halogens, are non-metals. They exist as diatomic, non-polar molecules between which the only intermolecular bonds are van der Waals’ forces. As the atomic number of the halogen molecules increases, the strength of the van der Waals’ forces increases signifi cantly and the melting point also increases. See table 2.5.1 page 59.

Properties of elements within periods are more variable than properties within groups. Consider the trends shown in fi gures 3.2.5 to 3.2.7. Marked variations occur because the electron confi gurations and core charges differ for each element in the period.

Recall that the core charge of an atom may be found by subtracting the number of inner-shell electrons from the nuclear charge. For example, across period 3 the core charge changes as shown:

Sodium: 11 protons and 10 inner-shell electrons = core charge of +1Aluminium: 13 protons and 10 inner-shell electrons = core charge of +3Chlorine: 17 protons and 10 inner-shell electrons = core charge of +7

3.5

3.0

2.5

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1.5

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Na Mg Al Si P S Cl Na Mg Al Si P S Cl Arelement

a b

Figure 3.2.5 (a) Electronegativities of period 3 elements. (b) First ionization energies of period 3 elements.

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Figure 3.2.6 Atomic radius decreases across period 3 elements.

atomic sizeincreases

•

electronegativitydecreases

•

first ionizationenergy decreases

•

melting pointdecreases(group 1)

•

melting pointincreases(group 7)

•

Figure 3.2.4 Trends in properties within groups of the periodic table.

Trends across period 3

3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization engergy and electronegativities for elements across period 3. © IBO 2007

Interactive periodic table

Figure 3.2.7 Ionic radius varies across period 3 elements.

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TRENDS ACROSS PERIOD 3

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The core charge increases across the period. This means that the outer-shell electrons of chlorine therefore experience a greater attraction to the nucleus than does the outer-shell electron of sodium.

Atomic radius decreases from left to right across period 3 due to the increasing attraction experienced by the outer-shell electrons. These outer-shell electrons are all in the third electron shell of the atoms; however, as the core charge increases, the electrostatic attraction between the outer-shell electrons and the nucleus increases. This has the effect of pulling the electrons in closer to the nucleus and making the atom smaller.

The trend in ionic radii is not as clear as for atomic radii (see fi gure 3.2.6). For the metals (sodium to aluminium) in period 3, the ionic radius decreases across the period. Silicon can be represented as a positive (Si4+) or negative (Si4!) ion. For the non-metals, the ionic radius decreases from the phosphorus (P3!) to the chloride ion (Cl!) (fi gure 3.2.7). A positive ion has a smaller ionic radius than the original atom, due to the loss of the valence electrons, and a negative ion has a larger ionic radius than the original atom, since the addition of extra negative charges introduces more electron–electron repulsion. Negative ions have a larger radius than positive ions, as they have one more shell of

electrons than the positive ions.

The increase in fi rst ionization energy and electronegativity from left to right across period 3 can also be explained by the increasing core charge. As the core charge increases, it becomes increasingly diffi cult to remove an electron from the outer shell of the atom (fi rst ionization energy). Similarly, the increasing electrostatic attraction of outer-shell electrons for the nucleus results in a greater power of attraction for electrons in the outer shell (electronegativity).

1 State and explain how:a the fi rst ionization energy of strontium compares with that

of magnesiumb the electronegativity of selenium compares with that of oxygen.

2 a Explain what is meant by the term core charge.b How is core charge used to explain the trend in atomic radius of the

period 3 elements?

3 Explain why the electronegativity of fl uorine is higher than that of magnesium.

4 a Compare the atomic radii of magnesium and chlorine.b Explain the difference you have described.

5 a Compare the fi rst ionization energy of phosphorus and chlorine. b Explain the difference you have described.

6 a State the trend in the ionic radii from Na+ to Al3+.b State the trend in the ionic radii from Si4! to Cl!.c Explain the trend in part b.

11+

innerelectronsshield thevalenceelectronfrom thenucleus

electronattractedby aneffectivecharge of +1

this electronexperiencesa strongerattractionthan theelectron in sodium

17+

sodiumcore charge = 11 – 10 = +1

chlorinecore charge = 17 – 10 = +7

Figure 3.2.8 Core charge can be used to explain trends within the periodic table.

• atomic radius decreases• electronegativity increases• first ionization energy increases• ionic radius decreases

Figure 3.2.9 Trends in properties across a period.

Section 3.2 Exercises

Summary of periodic trends

WORKSHEET 3.3 Periodic table trends



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a b


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•


•


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•


•






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1.0

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a b


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•


•


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•






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Na Mg Al Si P Si Cl

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(10

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m)The increase in first ionization energy and

electronegativity from left to right across period 3 can also be explained by the increasing core charge.



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a b


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a b


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A positive ion has a smaller ionic radius than the original atom, due to the loss of the valence electrons, and a negative ion has a larger ionic radius than the original atom, since the addition of extra negative charges introduces more electron–electron repulsion. Negative ions have a larger radius than positive ions, as they have one more shell of electrons than the positive ions.

CHEMICAL PROPERTIES TRENDS – DOWN A GROUP

•  The reactivity of the alkali metals with water increases down group 1.

•  All three alkali metals produce an

alkaline solution when they react with water. (Phenolphthalein is used to test for the alkaline solution)


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7 For each of the following pairs, determine whether the atomic/ionic radius of the fi rst particle listed is larger than (>), the same size (=), or smaller than (<) the radius of the second particle.

Just as the physical properties exhibit a periodicity within groups and periods, so do the chemical properties of the elements and some of their compounds. The chemical properties relate to the electron arrangement of the elements.

The alkali metals are well known for their reactive nature. Their tendency to react sometimes violently with water means that they must be stored under oil. Potassium, in particular, is seldom found in secondary school laboratories because of its diffi culty in storage and its violent reaction with water.

The reactivity of the alkali metals with water increases down group 1. While lithium metal fl oats on the surface of the water and reacts slowly, producing some hydrogen gas, sodium reacts more violently, whizzing around on the surface of the water on a layer of hydrogen gas in what is sometimes described as ‘hovercraft’ motion.

2Li(s) + H2O(l) ! Li2O(aq) + H2(g)

2Na(s) + H2O(l) ! Na2O(aq) + H2(g)

The sodium can be forced to ignite by slowing its progress on the surface of the water by placing it on paper towel (fi gure 3.3.1). The white smoke in fi gure 3.3.1 is sodium oxide being carried off as a smoke. Potassium burns spontaneously in water, producing a violet fl ame.

2K(s) + H2O(l) ! K2O(aq) + H2(g)

All three alkali metals produce an alkaline solution when they react with water.

Li2O(s) + H2O(l) ! 2LiOH(aq)

Na2O(s) + H2O(l) ! 2NaOH(aq)

K2O(s) + H2O(l) ! 2KOH(aq)

Evidence for this alkaline solution can be seen in fi gure 3.3.1. Phenolphthalein indicator has been added to the water before the reaction with sodium and has turned pink due to the presence of sodium hydroxide.

The increase in reactivity can be explained by the decrease in electrostatic attraction between the outer-shell electron and the positive nucleus of the alkali metal. The further the outer shell is from the nucleus, the more easily it is lost in the reaction with water, and the more spectacular the result.

The increasing reactivity of the alkali metals can also be seen in their reaction with halogens. As lithium, sodium and potassium all lose electrons easily (are good reducing agents) and the halogens gain electrons easily (are good oxidizing agents, see chapter 10), these reactions are predictably violent.

First particle >, =, < Second particlea Chlorine atom (Cl) Chloride ion (Cl")

b Aluminium ion (Al3+) Aluminium atom (Al)

c Calcium atom (Ca) Sulfur atom (S)

d Sodium ion (Na+) Fluoride ion (F")

e Magnesium ion (Mg2+) Calcium ion (Ca2+)

f Sulfide ion (S2") Potassium ion (K+)

3.3 CHEMICAL PROPERTIES OF ELEMENTS AND THEIR OXIDES

Trends in chemical properties within a group

3.3.1Discuss the similarities and differences in the chemical properties of elements in the same group. © IBO 2007

Figure 3.3.1 Sodium will burn with a yellow flame if its motion on the surface of water is stilled.

DEMO 3.2Reactions of group 1 and group 2 elements with water


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2Li(s) + H2O(l) ! Li2O(aq) + H2(g)

2Na(s) + H2O(l) ! Na2O(aq) + H2(g)


2K(s) + H2O(l) ! K2O(aq) + H2(g)




















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2Li(s) + H2O(l) ! Li2O(aq) + H2(g)

2Na(s) + H2O(l) ! Na2O(aq) + H2(g)


2K(s) + H2O(l) ! K2O(aq) + H2(g)




















•  The increasing reactivity of the alkali metals can also be seen in their reaction with halogens. •  reactions are predictably violent. •  The product in each case is an ionic

compound

88

The product in each case is an ionic compound. The reactions all follow the same pattern, since the alkali metals all form ions with a 1+ charge and the halogens all form ions with a 1! charge.

2Na(s) + Cl2(g) " 2NaCl(s)

2K(s) + I2(g) " 2KI(s)

2Li(s) + Br2(g) " 2LiBr(s)

The smaller the halogen atom, the greater is its ability to gain electrons. This can be explained by the closeness of the outer shell to the nucleus. When halogens are mixed with halide salts such as potassium iodide, KI, the ability to react depends on the relative electron attracting strength of the halogen and how easily the halide

ion will lose its electron. The larger the halide ion, the less strongly the outer-shell electrons are attracted to the nucleus (due to distance from the nucleus) and so the easier it is to remove an electron from the ion.

A list of the halogens in order of their electron attracting ability and the halide ions in order of their ability to lose an electron can be seen in fi gure 3.3.3.

Fluorine is the most reactive halogen, but its reactions are too violent to perform in a school laboratory. The next most reactive halogen chlorine, Cl2, will reduce bromide and iodide ions to bromine and iodine; bromine, Br2 will only reduce iodide ions and iodine cannot reduce any of the halide ions.

Cl2(g) + 2I!(aq) " 2Cl!(aq) + I2(s)

Cl2(g) + 2Br!(aq) " 2Cl!(aq) + Br2(s)

Br2(g) + 2I!(aq) ! 2Br!(aq) + I2(s)

THEORY OF KNOWLEDGEIn a Chemistry textbook, chlorine is described as a yellow-green gas at room temperature with a pungent, irritating odour. It is approximately two and a half times denser than air. When chlorine gas is inhaled, depending on the level of exposure, it can cause irritation to the eyes, skin and throat, a cough, chest tightness, wheezing and severe chemical burns.

During World War I, chlorine gas was deployed as a chemical weapon. Dulce et decorum est, written by Wilfred Owen in 1918 is one of the best known poems written in English from World War I.

Gas! Gas! Quick, boys! — An ecstasy of fumbling,Fitting the clumsy helmets just in time;But someone still was yelling out and stumbling,And fl ound’ring like a man in fi re or lime … Dim, through the misty panes and thick green light,As under a green sea, I saw him drowning.In all my dreams, before my helpless sight,He plunges at me, guttering, choking, drowning.

• Read the stanza of the poem and comment on the claim that art conveys no knowledge or literal truths that can be verifi ed.

• Consider the language used by the chemist and the poet to describe the effects of chlorine gas. Which is more precise, specifi c and direct? Which is more suggestive and leaves itself open to the readers’ interpretation? In what other ways might the use of language in English differ from that in science?

Figure 3.3.2 Potassium burns spontaneously in water with a violet flame.

Reactions with oxygenSodium and potassium in water

Figure 3.3.3 Chlorine has the greatest ability of these three halogens to gain an electron and the iodide ion has the greatest ability to lose an electron.

increasingabilityto gain

electrons

increasingability

to lose anelectron

Cl2 + 2e– 2Cl–

Br2 + 2e– 2Br–

I2 + 2e– 2I–

Figure 3.3.4 The reaction between chlorine and potassium iodide displaces red-brown iodine from the solution.


•  The smaller the halogen atom, the greater is its ability to gain electrons.

88

The product in each case is an ionic compound. The reactions all follow the same pattern, since the alkali metals all form ions with a 1+ charge and the halogens all form ions with a 1! charge.

2Na(s) + Cl2(g) " 2NaCl(s)

2K(s) + I2(g) " 2KI(s)

2Li(s) + Br2(g) " 2LiBr(s)

The smaller the halogen atom, the greater is its ability to gain electrons. This can be explained by the closeness of the outer shell to the nucleus. When halogens are mixed with halide salts such as potassium iodide, KI, the ability to react depends on the relative electron attracting strength of the halogen and how easily the halide

ion will lose its electron. The larger the halide ion, the less strongly the outer-shell electrons are attracted to the nucleus (due to distance from the nucleus) and so the easier it is to remove an electron from the ion.

A list of the halogens in order of their electron attracting ability and the halide ions in order of their ability to lose an electron can be seen in fi gure 3.3.3.

Fluorine is the most reactive halogen, but its reactions are too violent to perform in a school laboratory. The next most reactive halogen chlorine, Cl2, will reduce bromide and iodide ions to bromine and iodine; bromine, Br2 will only reduce iodide ions and iodine cannot reduce any of the halide ions.

Cl2(g) + 2I!(aq) " 2Cl!(aq) + I2(s)

Cl2(g) + 2Br!(aq) " 2Cl!(aq) + Br2(s)

Br2(g) + 2I!(aq) ! 2Br!(aq) + I2(s)

THEORY OF KNOWLEDGEIn a Chemistry textbook, chlorine is described as a yellow-green gas at room temperature with a pungent, irritating odour. It is approximately two and a half times denser than air. When chlorine gas is inhaled, depending on the level of exposure, it can cause irritation to the eyes, skin and throat, a cough, chest tightness, wheezing and severe chemical burns.

During World War I, chlorine gas was deployed as a chemical weapon. Dulce et decorum est, written by Wilfred Owen in 1918 is one of the best known poems written in English from World War I.

Gas! Gas! Quick, boys! — An ecstasy of fumbling,Fitting the clumsy helmets just in time;But someone still was yelling out and stumbling,And fl ound’ring like a man in fi re or lime … Dim, through the misty panes and thick green light,As under a green sea, I saw him drowning.In all my dreams, before my helpless sight,He plunges at me, guttering, choking, drowning.

• Read the stanza of the poem and comment on the claim that art conveys no knowledge or literal truths that can be verifi ed.

• Consider the language used by the chemist and the poet to describe the effects of chlorine gas. Which is more precise, specifi c and direct? Which is more suggestive and leaves itself open to the readers’ interpretation? In what other ways might the use of language in English differ from that in science?

Figure 3.3.2 Potassium burns spontaneously in water with a violet flame.

Reactions with oxygenSodium and potassium in water

Figure 3.3.3 Chlorine has the greatest ability of these three halogens to gain an electron and the iodide ion has the greatest ability to lose an electron.

increasingabilityto gain

electrons

increasingability

to lose anelectron

Cl2 + 2e– 2Cl–

Br2 + 2e– 2Br–

I2 + 2e– 2I–

Figure 3.3.4 The reaction between chlorine and potassium iodide displaces red-brown iodine from the solution.

OXIDES OF PERIOD 3

•  Each of the period 3 elements reacts with oxygen. •  metals and oxygen à large electronegativity

difference à produces an ionic compound •  non-metals form a covalent compound with

oxygen (less electronegativity difference). •  The ionic nature of the oxides decreases from left to

right across period 3.

OXIDES OF PERIOD 3


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Each of the period 3 elements reacts with oxygen. This reaction was one of the periodic properties that gave Mendeleev confi dence in the periodic law. The formulas of some of the highest oxides of the period 3 elements and the reaction of the oxide with water are shown in table 3.3.1.

As we saw in chapter 2, the large difference in electronegativity between the metals and oxygen produces an ionic compound, whereas the non-metals form a covalent compound with oxygen. The ionic nature of the oxides decreases from left to right across period 3.

TABLE 3.3.1 THE OXIDES OF SOME PERIOD 3 ELEMENTS AND THEIR REACTION WITH WATER

Element Highest oxide Reaction Nature of aqueous solutionSodium Na2O Na2O (s) + H2O(l) ! 2NaOH(aq) Alkaline

Magnesium MgO MgO(s) + H2O(l) ! Mg(OH)2(aq) Alkaline

Phosphorus P4O10 P4O10(s) + 6H2O(l) ! 4H3PO4(aq) Acidic

Sulfur SO3 SO3(g) + H2O(l) ! H2SO4(l) Acidic

The periodicity of properties is very clear in these reactions. Sodium burns in air to produce sodium oxide, Na2O, which reacts easily with water to produce a strongly alkaline solution of sodium hydroxide, NaOH. To the right of sodium in period 3 is magnesium. Magnesium burns with a bright white fl ame in air to produce magnesium oxide, MgO (see fi gure 3.3.5), which then reacts with water to make a solution of quite alkaline magnesium hydroxide, Mg(OH)2.

The decrease in metallic nature from left to right across period 3 can be seen from the very different behaviour of aluminium oxide, Al2O3, compared with the group 1 and group 2 metal oxides. Aluminium oxide can behave as an acid or base—it is amphoteric (see chapter 9). Aluminium oxide does not dissolve in water, but will react with acids and bases.

Acting as a base: Al2O3(s) + 6HCl(aq) ! 2AlCl3(aq) + 3H2O(l)

Acting as an acid: Al2O3(s) + 2NaOH(aq) + 3H2O(l) ! 2NaAl(OH)4(aq)

sodium aluminate

The oxides of period 3 elements 3.3.2Discuss the changes in nature from ionic to covalent and from basic to acidic of the oxides across period 3. © IBO 2007

Periodic trends: acid–base properties of oxides

DEMO 3.3Acidic and basic properties of oxides

Figure 3.3.5 Magnesium burns in air with a bright white flame to form magnesium oxide.

Figure 3.3.6 Sulfur burns in air with a blue flame to form sulfur dioxide.

WORKSHEET 3.4 Chemical trends


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sodium aluminate








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OXIDES OF PERIOD 3

•  Aluminium oxide can behave as an acid or base—it is amphoteric •  Aluminium oxide does not dissolve in water, but will

react with acids and bases.


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OXIDES OF PERIOD 3

•  We can generalize this trend by stating that: •  non-metals form acidic oxides •  metals form basic oxides •  aluminium forms an amphoteric oxide.

PERIODICITY - Wikispacesiknsphysicsib.wikispaces.com/file/view/periodicity.pdftopic 3 periodicity . ... chemistry: for use with the ib diploma programme standard level chapter 3 periodicity

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