SACE2 Chem Sample Chapter from Essentials Textbook

8/9/2019 SACE2 Chem Sample Chapter from Essentials Textbook

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C HA P T E R 1

1.1 The periodic tableElectron configurationsEach 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.

Elemental and

environmental chemistry1

CHAPTER

1 1s

2 2s

2p

3 3s

3p

3d

4 4s4p

4d

4f

Main shell Subshells

s 2

p 6

d 10

f 14

Subshell Maximum number of electrons

Table 1.1 Electron shells and subshells.

Table 1.2 Maximum number of electrons per subshell.

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 thelowest 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.


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2 SACE 2 Essentials Chemistry Workbook


CHAPTER 1

Using subshell notation to write electron configurations for atomsFor 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) 1s2 2s2 2p6 3s1

• iron atom (26 electrons) 1s2 2s2 2p6 3s2 3p6 4s2 3d6

• strontium atom (38 electrons) 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:

Cr (24 electrons) 1s22s

22p

63s

23p

64s

13d

5

Cu (29 electrons) 1s2 2s2 2p6 3s2 3p6 4s1 3d10

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 1s22s

22p

63s

23p

6.

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 1s22s

22p

63s

23p

64s

23d

104p

6.


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Electron configurations and the periodic tableAn 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 followingtable:

Chapter 1: Elemental and environmental chemistry 3


C HA P T E R 1

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+): ...............................................................................................................................................................

Phosphorus has the electron configuration 1s2

2s2

2p6

3s2

3p3.

Its outer shell is the 3rd shell which is occupied by 5 electrons. Phosphorus is in period 3, group V .

Bromine has the electron configuration 1s2

2s2

2p6

3s2

3p6

4s2

3d10

4p5.

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

For positive ions of the transition elements (for example Fe2+ and Fe3+ )

1. Write the electron configuration for the atom of the element.

Fe (26 electrons) 1s22s

22p

63s

23p

64s

23d

6

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 fromthe configuration of the Fe atom

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


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I VIII

II III IV V VI VII

The following figure is an 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) 1s22s

1

rubidium (Rb) 1s22s

22p

63s

23p

64s

23d

104p

65s

1

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

For example: magnesium (Mg) 1s22s

22p

63s

2

calcium (Ca) 1s22s

22p

63s

23p

64s

2

1

2

3

4

5

6

7

P E R I O D N U M B E R

TRANSITION ELEMENTS

MAIN GROUP NUMBERS

*

**

*

**

LANTHANIDES

ACTINIDES

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) 1s22s

22p

1

gallium (Ga) 1s22s

22p

63s

23p

64s

23d

104p

1

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

For example: silicon (Si) 1s22s

22p

63s

23p

2

germanium (Ge) 1s22s

22p

63s

23p

64s

23d

104p

2

• 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.



CHAPTER 1

Figure 1.1 Outline of the modern periodic table.


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

f1 –14



C HA P T E R 1

s-block p-block

s1 p6

p1 p2 p3 p4 p5s2

d-block

d1 –10

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

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

transition elements:

titanium (Ti) 1s22s

22p

63s

23p

64s

23d

2

cobalt (Co) 1s22s22p63s23p64s23d7

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:

samarium, Sm, (a lanthanide) 1s22s

22p

63s

23p

64s

23d

104p

65s

24d

105p

66s

24f

6

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.

f-block

*

**

*

**

Q1.2 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): ........................................................ ........................................................

Q1.3 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: ........................................................ ........................................................

c. Tungsten: ........................................................ ........................................................

d. Thallium: ........................................................ ........................................................

n/a

n/a


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CHAPTER 1

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 itscompounds.

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 formcovalent 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 ruleFor 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+

, 1s22s

22p

6, has a more stable configuration than the magnesium atom, Mg, 1s

22s

22p

63s

2.

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 haveconformed to the ‘octet rule’.

Expansion of the octetAtoms 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,

1s22s

22p

63s

23p

4, forms the sulfide ion, S

2– , ion, 1s

22s

22p

63s

23p

6.

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’.

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.

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.

Chemical properties of the elements and the periodic tableThe 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.


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C HA P T E R 1

Valence electronsElectrons 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.

• 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 1s22s

22p

63s

23p

64s

1. Its valence shell (the 4th shell) is ‘incomplete’,

containing only one electron. This is a less stable configuration than that for the K +

ion, 1s22s

22p

63s

23p

6.

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) 2NaCl(s)

2K (s) + S(s) K 2S(s)

Note that hydrogen is not included as a member of group I. Although its electron configuration

is 1s1, 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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The following are examples of reactions involving group II elements:

2Ca(s) + O2(g) 2CaO(s)

Mg(s) + S(s) MgS(s)



CHAPTER 1

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.

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 M

2+ion is more stable than that of the M atom.

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 s2 p

1outer shell electrons to form covalent bonds with other

non-metal atoms.

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 s2 p

1outer shell electrons in electron transfer reactions.

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.

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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C HA P T E R 1

Group IV elements


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 s2 p

2outer shell

electrons to form covalent bonds with other non-metal atoms.

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 s2 p

2outer shell electrons in electron

transfer reactions.

2+ ions are formed when the atoms of group IV elements lose only the p2

electrons of the s2 p

2outer

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, tometalloids for the rest of the group.

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.

Group V elements


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 s2 p

3outer 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 s2 p

3


3outer 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 s2 p

3to s

2 p

6. The resultant ions conform to the octet rule.


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Group VI elements


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 s2

p4

outer shell configuration to formcovalent 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 s2 p

4outer 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 s2 p

4electrons

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 s2 p

4to s

2 p




CHAPTER 1

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.

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.

Group VII elements


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


5outer 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 s2 p

5electrons 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 s2 p

5

to s2 p



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C HA P T E R 1

Limitations of the covalent bonding modelThere 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.

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.

Nitrogen +4 NO2

Nitrogen +2 NO

Chlorine +4 ClO2

Element Oxidation state Example

Table 1.3 Some anomalous oxidation states for nitrogen and chlorine.

Summary table of oxidation statesFor 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:

Hnon-metalox state

+1

Limetal

ox state+1

Bemetalloidox state

+2

Bnon-metalox states+3, –3

Cnon-metalox states+4, –4

Nnon-metalox states

+5, +3, –3

Onon-metalox state

–2

Fnon-metalox state

–1

Almetalloidox state

+3

Sinon-metalox states+4, –4

Pnon-metalox states

+5, +3, –3

Snon-metalox states

+6, +4, +2,–2

C1non-metalox states

+7, +5, +3,+1, –1

Gametalloidox state

+3

Gemetalloidox states+4, +2

Asmetalloidox states+5, +3

Senon-metalox states+4, +2

Brnon-metalox states

+7, +5, +3,+1, –1

Nametal

ox state+1

K metal

ox state+1

Rbmetal

ox state+1

Mgmetal

ox state+2

Cametal

ox state+2

TRANSITION

METALS

Srmetal

ox state+2

I VIII

II III IV V VI VII

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


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Electronegativities of the elementsThe 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

Li0.98

II

Be1.57

Na0.93

Mg1.31

K 0.82

Ca1.00

Rb0.82

Sr0.95

III IV V VI VII

B2.04

C2.55

N3.04

O3.44

F3.98

Al1.61

Si1.90

P2.19

S2.58

Cl3.16

Ga1.81

Ge2.01

As2.18

Se2.55

Br2.96

Low electronegativity(metals)

Intermediateelectronegativity(metalloids)

High electronegativity(non-metals)

The acidic/basic nature of oxidesAn oxide is a two-element compound with one of the elements being oxygen. Common examples are carbonmonoxide, 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.



CHAPTER 1

Note: H has electronegativity = 2.20

Acidic oxidesAcidic 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:

Oxyacids undergo complete or partial ionisation with water to produce hydronium ions.

For example:

H2CO3 + H2O H3O+ + HCO3 –

Figure 1.4 Electronegativity values of s and p block elements.

C HH

OO

O


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C HA P T E R 1

P4O10 P4O10 + 12OH–→ 4PO4

3–+ 6H2O P4O10 + 6H2O→ 4H3PO4

(phosphate) (orthophosphoric acid)

SO2 SO2 + 2OH–→ SO3

2–+ H2O SO2 + H2O→ H2SO3

(sulfite) (sulfurous acid)

SO3 SO3 + 2OH–→ SO4

2–+ H2O SO3 + H2O→ H2SO4

(sulfate) (sulfuric acid)

CO2 CO2 + 2OH–→ CO3

2–+ H2O CO2 + H2O→ H2CO3

(carbonate) (carbonic acid)

SiO2 SiO2 + 2OH–→ SiO32– + H2O NO REACTION(silicate)

Oxide *Reaction with hydroxide ions *Reaction with water

* 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.

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

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).

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

reactions occur).

Basic oxidesBasic 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 reactionsoccur).

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

Oxide *Reaction with hydrogen ions *Reaction with water

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

Table 1.5 Reactions of basic oxides.


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ZnO ZnO + 2H+→ Zn2+ + H2O ZnO + 2OH–→ ZnO2

2– + H2O(zincate)

Al2O3 Al2O3 + 6H+→ 2Al3+ + 3H2O Al2O3 + 2OH

–→ 2AlO2

– + H2O(aluminate)

Oxide Reaction with hydrogen ions Reaction with hydroxide ions

Table 1.6 Reactions of two amphoteric oxides.

Q1.5 Barium oxide, BaO, and lithium oxide, Li2O, are both basic oxides.

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

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

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

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

Amphoteric oxidesAmphoteric 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.



CHAPTER 1

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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C HA P T E R 1

B

Se

Molecular substancesCompounds 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.

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

III IV V VI VII

C N O F

Si P S Cl

As Br

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

Properties of molecular elements and compoundsElements 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.

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.

Q1.8 Predict whether the following elements or compounds are molecular or not:

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

Molecular (Yes/No)

SACE2 Chem Sample Chapter from Essentials Textbook

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