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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
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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.
transuranium elements
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.
transuranium elements
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
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
• 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
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
GROUPS 1 & 7
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL
CH
APT
ER 3
PE
RIO
DIC
ITY
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
• 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
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL
CH
APT
ER 3
PE
RIO
DIC
ITY
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
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL
CH
APT
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PE
RIO
DIC
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85
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
2.0
1.5
1.0
0.5
0
2000
1500
1000
500
0
element
Paul
ing
elec
tron
egat
ivity
valu
e
first
ioni
zatio
n en
ergy
(kJ
mol
–1)
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.
250
200
150
100
50
0
element
Na Mg Al Si P S Cl
ato
mic
rad
ius
(10
–12
m)
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.
250
200
150
100
50
0
element
Na Mg Al Si P Si Cl
ioni
c ra
dius
(10
–12
m)
TRENDS ACROSS PERIOD 3
86
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
TRENDS ACROSS PERIOD 3
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL
CH
APT
ER 3
PE
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DIC
ITY
85
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
2.0
1.5
1.0
0.5
0
2000
1500
1000
500
0
element
Paul
ing
elec
tron
egat
ivity
valu
e
first
ioni
zatio
n en
ergy
(kJ
mol
–1)
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.
250
200
150
100
50
0
element
Na Mg Al Si P S Cl
ato
mic
rad
ius
(10
–12
m)
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.
250
200
150
100
50
0
element
Na Mg Al Si P Si Cl
ioni
c ra
dius
(10
–12
m)
TRENDS ACROSS PERIOD 3
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL
CH
APT
ER 3
PE
RIO
DIC
ITY
85
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
2.0
1.5
1.0
0.5
0
2000
1500
1000
500
0
element
Paul
ing
elec
tron
egat
ivity
valu
e
first
ioni
zatio
n en
ergy
(kJ
mol
–1)
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.
250
200
150
100
50
0
element
Na Mg Al Si P S Cl
ato
mic
rad
ius
(10
–12
m)
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.
250
200
150
100
50
0
element
Na Mg Al Si P Si Cl
ioni
c ra
dius
(10
–12
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.
TRENDS ACROSS PERIOD 3
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL
CH
APT
ER 3
PE
RIO
DIC
ITY
85
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
2.0
1.5
1.0
0.5
0
2000
1500
1000
500
0
elementPa
ulin
g el
ectr
oneg
ativ
ityva
lue
first
ioni
zatio
n en
ergy
(kJ
mol
–1)
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.
250
200
150
100
50
0
element
Na Mg Al Si P S Cl
ato
mic
rad
ius
(10
–12
m)
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.
250
200
150
100
50
0
element
Na Mg Al Si P Si Cl
ioni
c ra
dius
(10
–12
m)
TRENDS ACROSS PERIOD 3
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL
CH
APT
ER 3
PE
RIO
DIC
ITY
85
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
2.0
1.5
1.0
0.5
0
2000
1500
1000
500
0
elementPa
ulin
g el
ectr
oneg
ativ
ityva
lue
first
ioni
zatio
n en
ergy
(kJ
mol
–1)
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.
250
200
150
100
50
0
element
Na Mg Al Si P S Cl
ato
mic
rad
ius
(10
–12
m)
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.
250
200
150
100
50
0
element
Na Mg Al Si P Si Cl
ioni
c ra
dius
(10
–12
m)
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)
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME STANDARD LEVEL
CH
APT
ER 3
PE
RIO
DIC
ITY
87
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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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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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
CHEMICAL PROPERTIES TRENDS – DOWN A GROUP
• 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
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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.
CHEMICAL PROPERTIES TRENDS – DOWN A GROUP
• 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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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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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
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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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
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.