1 CHAPTER

1.1 The periodic table Electron configurations Each electron in an atom, or monatomic ion, has potential energy arising from the attraction between its negative charge and the positive charge of the nucleus. Electrons in the atoms or monatomic ions of a particular element have energy values that are unique to that element. Each allowed energy level for an electron is represented by a main shell number (principal quantum number) using numbers 1, 2, 3, 4, 5, 6 and a subshell represented by the lowercase letters s, p, d or f. For a main shell (of number n), there are n subshells. The subshells corresponding to each of the first four main shells are listed in Table 1.1. Main shell

Subshells

1

1s

2

2s 2p

3

3s 3p 3d

4

4s 4p 4d 4f

Table 1.1 Electron shells and subshells.

The electron configuration of an atom or monatomic ion describes the number of electrons at each energy level. When writing electron configurations for atoms or monatomic ions the following principles apply:

• in the most stable state (the ground state) of any atom or ion the electrons ‘occupy’ subshells with the lowest available energy levels. They are ‘allocated’ to subshells in order of increasing energy as shown in the following energy sequence: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s lower energy

higher energy

• each subshell can accommodate a maximum number of electrons as shown in Table 1.2. Subshell

Maximum number of electrons

s

2

p

6

d

10

f

14

Table 1.2 Maximum number of electrons per subshell.

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Using subshell notation to write electron configurations for atoms For an atom the number of electrons is equal to the atomic number of the element. An atom’s electron configuration is written in energy sequence order for the subshells with the number of electrons ‘occupying’ each subshell shown as a superscript. This is illustrated in the following examples:

• sodium atom (11 electrons) • iron atom (26 electrons) • strontium atom (38 electrons)

1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 3p6 4s2 3d6 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2

Two electron configurations that do not conform Within the first 38 elements of the periodic table, the electron configurations of chromium and copper atoms do not conform to the principles for assigning electrons to subshells as described above. The electron configurations are as follows: 2

2

6

2

6

1

5

2

2

6

2

6

1

10

Cr (24 electrons)

1s 2s 2p 3s 3p 4s 3d

Cu (29 electrons)

1s 2s 2p 3s 3p 4s 3d

Using subshell notation to write electron configurations for monatomic ions For positive ions of the main group (groups I to VIII) elements 1. Determine the number of electrons for the ion. positive ions have less electrons than the atom of the element by the number equal to the numerical value of the charge on the ion. For example, because it has a 2+ charge, the Ca2+ ion has 18 electrons, 2 electrons less than the Ca atom. 2. Assign the 18 electrons to subshells as described for atoms above. The electron configuration for the Ca2+ ion is 1s2 2s2 2p6 3s2 3p6.

For negative ions of the main group (groups I to VIII) elements 1. Determine the number of electrons for the ion. negative ions have more electrons than the atom of the element by a number equal to the numerical value of the charge on the ion. For example, because it has a 1– charge, the Br– ion has 36 electrons, 1 electron more than the Br atom. 2. Assign the 36 electrons to subshells as described for atoms above. The electron configuration for the Br– ion is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6.

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For positive ions of the transition elements (for example Fe2+ and Fe3+) 1. Write the electron configuration for the atom of the element. 1s2 2s2 2p6 3s2 3p6 4s2 3d6

2. Decrease the number of electrons equal to the numerical value of the charge on the ion by deleting the 4s electrons first, with any subsequent deletions required being made from the 3d subshell. The Fe2+ ion

• has 2 electrons less than the Fe atom • the configuration for the Fe2+ ion is obtained by removing the 4s2 electrons from the configuration of the Fe atom

• the electron configuration for the Fe2+ ion is 1s2 2s2 2p6 3s2 3p6 3d6 The Fe3+ ion

• has 3 electrons less than the Fe atom • the configuration for the Fe3+ ion is obtained by removing the 4s2 electrons and one 3d electron from the configuration of the Fe atom

• the electron configuration for the Fe3+ ion is 1s2 2s2 2p6 3s2 3p6 3d5. Q1.1

Using subshell notation, write the electron configurations for the following atoms and ions (use a copy of the periodic table to find atomic number values): a. Argon atom:

...............................................................................................................................................................

b. Rubidium ion (Rb+):

...............................................................................................................................................................

c. Manganese atom:

...............................................................................................................................................................

d. Nickel ion (Ni2+):

...............................................................................................................................................................

Electron configurations and the periodic table An element’s position on the modern periodic table is determined by its electron configuration. Horizontal rows of the periodic table are called periods and vertical columns are called groups. Periods are numbered from 1 to 7. The number of the period in which an element is placed is equal to the highest numbered main shell that is occupied by electrons. The number of electrons occupying the highest numbered main shell (outer shell) determines the vertical column (group) to which a main group element is assigned. For any main group element, the group number equals the number of electrons occupying its outer shell. This is illustrated by the examples in the following table:

2

2

6

2

3

Phosphorus has the electron configuration 1s 2s 2p 3s 3p . Its outer shell is the 3rd shell which is occupied by 5 electrons. Phosphorus is in period 3, group V. 2

2

6

2

6

2

Bromine has the electron configuration 1s 2s 2p 3s 3p 4s 3d

10

5

4p .

Its outer shell is the 4th shell which is occupied by 7 electrons in total. Bromine is in period 4, group VII.

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Fe (26 electrons)

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The following figure is an outline of the modern periodic table:

MAIN GROUP NUMBERS

I II

VIII III

IV

V

VI

VII

1

PERIOD NUMBER

2 3

TRANSITION ELEMENTS

4 5

*

6

**

7

*

LANTHANIDES

**

ACTINIDES

Figure 1.1 Outline of the modern periodic table.

• For all elements in group I, there is one electron in the highest energy s subshell. For example:

lithium (Li) rubidium (Rb)

1s22s1 1s22s22p63s23p64s23d104p65s1

• For all elements in group II, there are two highest energy electrons in an s subshell. For example:

magnesium (Mg) calcium (Ca)

1s22s22p63s2 1s22s22p63s23p64s2

Groups I and II form the s block of the periodic table.

• For all elements in group III, there is one electron in the highest energy p subshell. For example:

boron (B) gallium (Ga)

1s22s22p1 1s22s22p63s23p64s23d104p1

• For all elements in group IV, there are two electrons in the highest energy p subshell. For example:

silicon (Si) germanium (Ge)

1s22s22p63s23p2 1s22s22p63s23p64s23d104p2

• For elements in groups V to VIII there are in turn 3 to 6 electrons in the highest energy p subshell (helium being an exception at the top of group VIII). Groups III to VIII form the p block of the periodic table.

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• For the transition elements, the highest energy electrons are in a d subshell.

1s22s22p63s23p64s23d2 1s22s22p63s23p64s23d7

titanium (Ti) cobalt (Co)

The transition elements form the d block of the periodic table.

• For the lanthanides and actinides, the highest energy electrons are in an f subshell. The following is an example of an electron configuration of an element from the lanthanides: 1s22s22p63s23p64s23d104p65s24d105p66s24f 6

samarium, Sm, (a lanthanide)

The lanthanides and actinides form the f block of the periodic table.

The s, p, d and f blocks of the periodic table are summarised in Figure 1.2.

s-block

p-block

s1

p6 s2

p1

p2

p3

p4

p5

d-block d1 – 10

* ** f-block f 1 – 14

* **

Figure 1.2 s, p, d and f blocks of the periodic table. Q1.2

Q1.3

Using subshell notation, write the electron configurations for the following atoms for which the atomic number is given in brackets. From the electron configuration, determine the ‘block’ of the periodic table to which the element belongs. Electron configuration Periodic table block a. Arsenic, As (33):

........................................................

........................................................

b. Rubidium, Rb (37):

........................................................

........................................................

c. Krypton, Kr (36):

........................................................

........................................................

d. Cobalt, Co (27):

........................................................

........................................................

By inspection of a copy of the periodic table, determine the ‘main group’ and ‘block’ for each of the following elements. Main group number

Periodic table block

a. Francium:

........................................................

........................................................

b. Thorium:

n/a ........................................................

........................................................

c. Tungsten:

n/a ........................................................

........................................................

d. Thallium:

........................................................

........................................................

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The following are examples of electron configurations of elements from the first row (period 4) of the transition elements:

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Chemical properties of the elements and the periodic table The elements in each group of the s or p blocks of the periodic table display similar chemical properties to each other. They react with other elements and compounds forming products that conform to a common formula pattern.

Examples of ‘similar chemical properties’ The elements from group I all react with chlorine to form a chloride of formula type MCl. Group I elements also react with water to form a hydroxide of formula type MOH and hydrogen, H2. Each element in group V forms a compound with hydrogen of formula type XH3.

Electron configurations of the atoms of the s and p block elements can be used as a basis for explaining and predicting their chemical properties. The connection between the electron configuration of an element and its position on the periodic table can be used to make predictions about the properties of an element, including its metal/metalloid/non-metal nature, the charge(s) of its monatomic ion(s) and its likely oxidation state(s) in its compounds.

Metals:

Atoms of metals lose electrons in chemical reactions.

Non-metals:

Atoms of non-metals gain or share electrons in chemical reactions.

Metalloids:

Atoms of metalloids lose or share electrons in chemical reactions.

The similarity in chemical properties of the elements within each particular group is explained in terms of the similarity of their electron configurations. When elements react, their atoms either lose or gain electrons (to form positive or negative ions respectively) or they share electrons with those of other atoms (to form covalent bonds). The electron configurations of the resultant ions are more stable than the configurations of the atoms from which they have been formed. Similarly, when atoms share electrons they acquire more stable electron configurations.

The octet rule For period 1 and 2 elements an electron configuration in which the outer shell is ‘complete’ with its maximum number of electrons is more stable than a configuration with an incomplete outer shell. For example, the magnesium ion, Mg2+, 1s22s22p6, has a more stable configuration than the magnesium atom, Mg, 1s22s22p63s2. These complete, stable electron configurations are the same as for the noble gases of periods 1 and 2 and are often referred to as ‘noble gas configurations’. For the period 2 and 3 noble gases, the electron configurations show 8 electrons in the outer shell and when atoms attain 8 electrons in their outer shell they are said to have conformed to the ‘octet rule’.

Expansion of the octet Atoms of the period 3 elements from the p block often conform to the octet rule by accepting electrons to form negatively charged monatomic ions or by sharing electrons with other atoms. For example, the sulfur atom, S, 1s22s22p63s23p4, forms the sulfide ion, S2–, ion, 1s22s22p63s23p6. The atoms of the period 3 elements from groups V to VII can share all of their outer shell electrons and as a consequence acquire more than 8 electrons in their outer shells. The extra electrons above the octet are ‘accommodated’ in the previously unoccupied 3d subshell. This is referred to as the ‘expansion of the octet’.

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Valence electrons

• For s block elements, the valence shell electrons are the highest energy s subshell electrons. • For p block elements, the valence shell electrons are the ‘outer’ s and p subshell electrons.

Chemical reactions involving the s block elements Group I elements These elements readily form compounds by reacting with non-metal elements such as the halogens, oxygen and sulfur and with other oxidising agents such as water and the hydronium ion (present in dilute acid solutions). In these reactions, the atoms of the group I elements lose their s1 valence electrons to form an M+ ion. The electron lost is gained by the other reactant. The product compounds are ionic with formulae such as MC1, M2S, M2O and MOH. The configuration of the M+ ion is more stable than that of the M atom. The potassium atom has the configuration 1s22s22p63s23p64s1. Its valence shell (the 4th shell) is ‘incomplete’, containing only one electron. This is a less stable configuration than that for the K+ ion, 1s22s22p63s23p6. The outer shell for this ion is now the 3rd shell and it is complete (full). This ion conforms to the octet rule.

The charge on the monatomic ions of the group I elements is always 1+. Consequently the oxidation state of the group I elements in their compounds is always +1. The group I elements are classified as metals because their atoms lose electrons in chemical reactions.

The following are examples of reactions involving group I elements: 2Na(s) + Cl2(g) 2K(s) + S(s)

2NaCl(s) K2S(s)

Note that hydrogen is not included as a member of group I. Although its electron configuration 1 is 1s , its properties are quite different to those of other members of group I. Sometimes it is not shown at the top of group I but is given a separate ‘box’ of its own. In compounds with non-metals, hydrogen atoms share their one valence electron with valence electrons of the other non-metal atoms.

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Electrons lost or shared by atoms are those from the valence (outer) shells. Electrons gained are accepted into valence shells. Outer shell electrons are called valence electrons.

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Group II elements These elements also readily form compounds by reacting with non-metal elements such as the halogens, oxygen and sulfur and with oxidising agents such as water and the hydronium ion. In these reactions, the atoms of the group II elements lose their valence electrons to form an M2+ ion. The electrons lost are gained by the other reactant. The product compounds are ionic with formulae such as MCl2, MS, MO and M(OH)2. The configuration of the M2+ ion is more stable than that of the M atom.

The charge on the monatomic ions of the group II elements is always 2+. Consequently the oxidation state of the group II elements in their compounds is always +2. Except for beryllium, the group II elements are classified as metals because their atoms lose electrons in chemical reactions. Beryllium is classified as a metalloid.

The following are examples of reactions involving group II elements: 2Ca(s) + O2(g) Mg(s) + S(s)

2CaO(s) MgS(s)

Chemical reactions involving the p block elements Group III elements In their compounds – either they exhibit a covalence of 3 as in the case for boron in its compounds, for example BCl3. When boron forms compounds, its atoms share the s2p1 outer shell electrons to form covalent bonds with other non-metal atoms.

Covalence The covalence of an element is equal to the number of electrons that its atoms share when forming covalent bonds with other atoms. When boron shares its 3 outer shell electrons with, for example, electrons from three chlorine atoms, then boron is exhibiting a covalence of 3.

or

they exist as triple positive ions, such as Al3+, in compounds with non-metals. These ions are formed when atoms of the group III elements lose the s2p1 outer shell electrons in electron transfer reactions.

The charge on the monatomic ions of the group III elements is usually 3+. The oxidation state of the group III elements in their compounds is usually +3. In some compounds their oxidation state is –3. The elements range from a non-metal, boron, at the top of the group, to metalloids, aluminium and gallium, in the middle and metals at the bottom of the group.

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Group IV elements In their compounds –

or

they exist as 2+ or 4+ ions, such as Pb2+, Sn2+ or Pb4+, in compounds with non-metals. 4+ ions are formed when the atoms of group IV elements lose the s2p2 outer shell electrons in electron transfer reactions. 2+ ions are formed when the atoms of group IV elements lose only the p2 electrons of the s2p2 outer shell electrons in electron transfer reactions.

The charge on the monatomic ions of the group IV elements is usually 2+ or 4+. The oxidation state of the group IV elements in their compounds is either +4, +2 or –4. The elements range from the non-metals, carbon and silicon, at the top of the group, to metalloids for the rest of the group.

Group V elements In their compounds – either they exhibit a covalence of 3, as in the case of nitrogen in all of its compounds such as NH3, and of phosphorus and arsenic in some of their compounds, for example AsCl3. In these compounds, the nitrogen, phosphorus and arsenic atoms share only the p3 electrons from the s2p3 outer shell configuration to form covalent bonds with other non-metal atoms. In sharing in this way they are conforming to the octet rule. or

they exhibit a covalence of 5, as in the case of phosphorus and arsenic in some of their compounds, for example, AsCl5 and P4O10. In these compounds, the phosphorus and arsenic atoms share all of the s2p3 electrons from the s2p3 outer shell configuration to form covalent bonds with other non-metal atoms. In sharing in this way they are expanding the octet.

or

they exist as 3– ions, such as N3– or P3– in compounds with metals. 3– ions are formed when atoms of the group V elements gain three electrons into the p subshell thereby changing the outer shell configuration from s2p3 to s2p6. The resultant ions conform to the octet rule.

The charge on the monatomic ions of the group V elements is 3–. The oxidation state of the group V elements in their compounds is either +5, +3 or –3. The elements range from the non-metals, nitrogen and phosphorus, at the top of the group, to the metalloids arsenic and antimony in the middle and the metal bismuth at the bottom of the group.

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either they exhibit a covalence of 4, as in the case for carbon and silicon in their compounds, for example, CCl4 and SiH4. When carbon and silicon form compounds, their atoms share the s2p2 outer shell electrons to form covalent bonds with other non-metal atoms.

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In their compounds – either they exhibit a covalence of 2, as in the case of oxygen in all of its compounds such as H2O, and of sulfur and selenium in some of their compounds, for example SF2. In these compounds the oxygen, sulfur and selenium atoms share only two of the p4 electrons from the s2p4 outer shell configuration to form covalent bonds with other non-metal atoms. In sharing in this way they are conforming to the octet rule. or

they exhibit a covalence of 4, as in the case of sulfur and selenium in some of their compounds, for example SO2 and SeF4. In these compounds, the sulfur and selenium atoms share all of the p4 electrons from the s2p4 outer shell configuration to form covalent bonds with other non-metal atoms. In sharing in this way they are expanding the octet. It must be noted that oxygen does not exhibit a covalence of 4.

or

they exhibit a covalence of 6, as in the case of sulfur and selenium in some of their compounds, for example SF6 and SeO3. In these compounds, the sulfur and selenium atoms share all of the s2p4 electrons from the outer shell configuration to form covalent bonds with other non-metal atoms. In sharing in this way they are expanding the octet. It must be noted that oxygen does not exhibit a covalence of 6.

or

they exist as 2- ions, such as O2– or S2– in compounds with metals. 2– ions are formed when the atoms of group VI elements gain two electrons into the p subshell thereby changing the outer shell configuration from s2p4 to s2p6. The resultant ions conform to the octet rule.

The charge on the monatomic ions of the group VI elements is 2–. The oxidation state of the group VI elements in their compounds is +6, +4, +2 or –2. An exception is oxygen with an oxidation number of –1 in H2O2. The elements range from the non-metals, oxygen and sulfur, at the top of the group, to the metalloids selenium and tellurium in the middle and the metal polonium at the bottom of the group.

Group VII elements In their compounds – either they exhibit a covalence of 1, as in the compounds such as HBr and CCl4. In these compounds, the group VII atoms share one of the p5 electrons from the s2p5 outer shell configuration to form a covalent bond with other non-metal atoms. In sharing in this way they are conforming to the octet rule. or

they exhibit a covalence of 3, 5 or 7 and in sharing in this way they are expanding the octet. It must be noted that fluorine does not exhibit a covalence of 3, 5 or 7.

In such compounds, the atoms of the group VII elements share electrons in the following ways: covalence of 3: three of the p5 electrons shared covalence of 5: all five of the p5 electrons shared covalence of 7: all seven of the s2p5 electrons shared or

they exist as 1– ions, in compounds with metals. 1– ions are formed when the atoms of group VII elements gain one electron into the p subshell, thereby changing the outer shell configuration from s2p5 to s2p6. The resultant ions conform to the octet rule.

The charge on the monatomic ions of the group VII elements is 1–. The oxidation state of the group VII elements in their compounds is +7, +5, +3, +1 or –1. The elements are all non-metals.

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Limitations of the covalent bonding model

Element

Oxidation state

Example

Nitrogen

+4

NO2

Nitrogen

+2

NO

Chlorine

+4

ClO2

Table 1.3 Some anomalous oxidation states for nitrogen and chlorine.

These oxidation states cannot be explained in terms of electron configurations and shared pairs of electrons. Other bonding models that are beyond the scope of this course must be used.

Summary table of oxidation states For the s and p block elements up to atomic number 38, Figure 1.3 summarises the likely oxidation states of the elements in their compounds and the metal/metalloid/non-metal nature of the elements:

I

VIII

H non-metal ox state +1 II

III

IV

V

VI

VII

Li metal ox state +1

Be metalloid ox state +2

B non-metal ox states +3, –3

C non-metal ox states +4, –4

N non-metal ox states +5, +3, –3

O non-metal ox state –2

F non-metal ox state –1

Na metal ox state +1

Mg metal ox state +2

Al metalloid ox state +3

Si non-metal ox states +4, –4

P non-metal ox states +5, +3, –3

S non-metal ox states +6, +4, +2, –2

C1 non-metal ox states +7, +5, +3, +1, –1

K metal ox state +1

Ca metal ox state +2

Ga metalloid ox state +3

Ge metalloid ox states +4, +2

As metalloid ox states +5, +3

Se non-metal ox states +4, +2

Br non-metal ox states +7, +5, +3, +1, –1

Rb metal ox state +1

Sr metal ox state +2

TRANSITION METALS

Figure 1.3 Oxidation states and nature of s and p block elements.

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There are some examples of the s and p block elements exhibiting oxidation states (and covalences) that are different from those given in the summaries above. Some of the more common of these are given in Table 1.3.

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Electronegativities of the elements The relative ability of an atom to attract electrons to itself is called its electronegativity. The higher the electronegativity, the stronger the attraction for electrons. Metal atoms have lower electronegativity values than metalloids, which in turn have lower electronegativities than non-metals. Using the periodic table, two clear trends for electronegativities of the s and p block elements can be described as follows: Group I

Group VII

electronegativities increase across each period

Period 1

Period 4

electronegativities decrease down each group

The periodic table can then be divided into regions of high, intermediate and low electronegativities as shown on part of the periodic table in Figure 1.4. The numerical values beneath the symbols of the elements are Pauling electronegativity values. (Linus Pauling, an American chemist, developed a scale of electronegativity values last century.)

I

II

III

IV

V

VI

VII

Li 0.98

Be 1.57

B 2.04

C 2.55

N 3.04

O 3.44

F 3.98

Na 0.93

Mg 1.31

Al 1.61

Si 1.90

P 2.19

S 2.58

Cl 3.16

K 0.82

Ca 1.00

Ga 1.81

Ge 2.01

As 2.18

Se 2.55

Br 2.96

Rb 0.82

Sr 0.95

Low electronegativity (metals) Intermediate electronegativity (metalloids) High electronegativity (non-metals)

Note: H has electronegativity = 2.20

Figure 1.4 Electronegativity values of s and p block elements.

The acidic/basic nature of oxides An oxide is a two-element compound with one of the elements being oxygen. Common examples are carbon monoxide, CO, iron (III) oxide, Fe2O3, and nitrogen dioxide, NO2. Oxides can be classified as acidic, amphoteric or basic on the basis of their reactivity, or lack of reactivity, with acids and bases.

Acidic oxides Acidic oxides react with hydroxide ions to produce oxyanions (negatively charged ions of the element and oxygen) and water molecules. Examples of oxyanions are carbonate, CO32–, sulfate, SO42– and aluminate, AlO2–. If soluble in water, acidic oxides react with water to form oxyacids (acids consisting of the element combined with hydrogen and oxygen). Examples of oxyacids are carbonic acid, H2CO3, sulfuric acid, H2SO4 and orthophosphoric acid, H3PO4. These oxyacids consist of molecules with covalent hydroxyl groups (O — H) as part of their structure. For example, the structure of H2CO3 is: O H

C O

H O

Oxyacids undergo complete or partial ionisation with water to produce hydronium ions. For example: H2CO3 + H2O

H3O+ + HCO3–

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Acidic oxides are the oxides of non-metals. They are the oxides of the elements with high electronegativity and are covalent molecular (such as CO2 and SO3) or continuous covalent compounds (such as SiO2).

Oxide

*Reaction with hydroxide ions

P4O10

P4O10 + 12OH– → 4PO43– + 6H2O

*Reaction with water P4O10 + 6H2O → 4H3PO4

(phosphate)

SO2

SO2 + 2OH– → SO32– + H2O

(orthophosphoric acid)

SO2 + H2O → H2SO3

(sulfite)

SO3

SO3 + 2OH– → SO42– + H2O

(sulfurous acid)

SO3 + H2O → H2SO4

(sulfate)

CO2

CO2 + 2OH → CO3 + H2O –

2–

(sulfuric acid)

CO2 + H2O → H2CO3

(carbonate)

SiO2

(carbonic acid)

SiO2 + 2OH– → SiO32– + H2O

NO REACTION

(silicate) * Note: The oxidation numbers of the elements are unchanged in these reactions.

Table 1.4 Reactions of acidic oxides. Q1.4

One of the oxides of chlorine is Cl2O. It is an acidic oxide with a corresponding oxyanion, CIO– (hypochlorite), and a corresponding oxyacid, HCl0 (hypochlorous acid). a. Write an equation for the reaction of Cl2O with hydroxide ions. ...................................................................................................................................................................................................

b. Write an equation for the reaction of Cl2O with water. ...................................................................................................................................................................................................

Basic oxides Basic oxides react with acids (or hydrogen ions) to produce positively charged metal ions and water molecules. When reacting in this way, the solid oxides appear to dissolve in the acid. If soluble in water, basic oxides react with water to form metal ions and hydroxide ions in solution. Basic oxides are the oxides of metals. They are the oxides of elements with low electronegativity and are ionic compounds consisting of metal ions and oxide, O2–, ions. Table 1.5 summarises the reactions of some basic oxides with hydrogen ions and with water (where reactions occur). Oxide

*Reaction with hydrogen ions

*Reaction with water

Na2O

Na2O + 2H+ → 2Na+ + H2O

Na2O + H2O → 2Na+ + 20H–

MgO

MgO + 2H+ → Mg2+ + H2O

MgO + H2O → Mg2+ + 20H–

CuO

CuO + 2H+ → Cu2+ + H2O

NO REACTION

Fe2O3

Fe2O3 + 6H+ → 2Fe3+ + 3H2O

NO REACTION

* Note: The oxidation numbers of the elements are unchanged in these reactions.

Table 1.5 Reactions of basic oxides.

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Table 1.4 below summarises the reactions of some acidic oxides with hydroxide ions and with water (where reactions occur).

14

SACE 2 Essentials Chemistry Workbook Q1.5

Barium oxide, BaO, and lithium oxide, Li2O, are both basic oxides.

CHAPTER 1

a. Write an equation for the reaction of BaO with hydrogen ions. ...................................................................................................................................................................................................

b. Write an equation for the reaction of Li2O with water. ...................................................................................................................................................................................................

Amphoteric oxides Amphoteric oxides display basic character by reacting with acids (or hydrogen ions) to produce positively charged monatomic ions and water molecules. When reacting in this way, the solid oxides appear to ‘dissolve’ in the acid. Amphoteric oxides also display acidic character by reacting with hydroxide ions to produce oxyanions and water molecules. When reacting in this way, the solid oxides appear to dissolve in the hydroxide solution. Amphoteric oxides do not react with water. Table 1.6 summarises the reactions of two amphoteric oxides with hydrogen ions and with hydroxide ions. Oxide

Reaction with hydrogen ions

ZnO

ZnO + 2H+ → Zn2+ + H2O

Al2O3

Al2O3 + 6H → 2Al + 3H2O

Reaction with hydroxide ions ZnO + 2OH– → ZnO22– + H2O (zincate)

+

3+

Al2O3 + 2OH– → 2AlO2– + H2O (aluminate)

Table 1.6 Reactions of two amphoteric oxides.

Q1.6

Lead oxide, PbO, is an amphoteric oxide. a. Write an equation for the reaction of PbO with hydrogen ions. ...................................................................................................................................................................................................

b. Write an equation for the reaction of PbO with hydroxide ions, given that it forms the plumbate ion, PbO22–. ...................................................................................................................................................................................................

Q1.7

When rubidium oxide is mixed with water the resulting solution has a pH greater than 7. Explain, with the aid of an equation, why the solution has a pH greater than 7. .......................................................................................................................................................................................................... .......................................................................................................................................................................................................... .......................................................................................................................................................................................................... ..........................................................................................................................................................................................................

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Chapter 1: Elemental and environmental chemistry

15

Molecular substances

Small molecules are formed when atoms of non-metal elements covalently bond to each other. These non-metal elements are located in the top right-hand section of the periodic table: III

IV

V

VI

VII

B

C

N

O

F

Si

P

S

Cl

As

Se

Br

Hydrogen atoms can also form small molecules with atoms of the elements shown above.

There is a small number of non-metal elements and non-metal – non-metal compounds that are not molecular. For example, silicon dioxide, SiO2, diamond, C, and silicon carbide, SiC, are not molecular. These substances consist of continuous lattices of atoms bonded to each other by covalent bonds. They are commonly referred to as continuous covalent substances.

Properties of molecular elements and compounds Elements and compounds which consist of small molecules, have low melting and boiling points and are usually gases or liquids at room temperature. Those that are solids generally have melting points below 200˚C. Molecular elements and compounds are poor conductors of electricity in the solid, liquid and gaseous state. Q1.8

Predict whether the following elements or compounds are molecular or not: Molecular (Yes/No) Element 1

Boiling point 58˚C, melting point –7˚C

Element 2

Boiling point 2,680˚C, melting point 1,410˚C

Compound 1

Compound of calcium and sulfur

Compound 2

Formula Cl20

Compound 3

Liquid at room temperature. This liquid is a non-conductor of electricity

Compound 4

Silver chloride

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CHAPTER 1

Compounds and elements consisting of molecules are described as ‘molecular substances’. Molecules consisting of 10 or less atoms per molecule may be considered as ‘small molecules’. Some common examples are CO2, SO2, H2S, Cl2, NH3, O3, CFCl3 and H2SO4.