Elemental and environmental chemistry

1 CHAPTER 1 Elemental and environmental chemistry 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.

2 SACE 2 Essentials Chemistry Workbook 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) 1s 2 2s 2 2p 6 3s 1 iron atom (26 electrons) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 strontium atom (38 electrons) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 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) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Cu (29 electrons) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 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 Ca 2+ 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 Ca 2+ ion is 1s 2 2s 2 2p 6 3s 2 3p 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 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6.

Chapter 1: Elemental and environmental chemistry 3 For positive ions of the transition elements (for example Fe 2+ and Fe 3+ ) 1. Write the electron configuration for the atom of the element. Fe (26 electrons) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 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 Fe 2+ ion has 2 electrons less than the Fe atom the configuration for the Fe 2+ ion is obtained by removing the 4s 2 electrons from the configuration of the Fe atom the electron configuration for the Fe 2+ ion is 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 The Fe 3+ ion has 3 electrons less than the Fe atom the configuration for the Fe 3+ ion is obtained by removing the 4s 2 electrons and one 3d electron from the configuration of the Fe atom the electron configuration for the Fe 3+ ion is 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5. 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 (Ni 2+ ):... 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: Phosphorus has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 3. 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 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5. Its outer shell is the 4th shell which is occupied by 7 electrons in total. Bromine is in period 4, group VII.

4 SACE 2 Essentials Chemistry Workbook The following figure is an outline of the modern periodic table: 1 I MAIN GROUP NUMBERS II III IV V VI VII VIII 2 PERIOD NUMBER 3 4 5 TRANSITION ELEMENTS 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) 1s 2 2s 1 rubidium (Rb) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 For all elements in group II, there are two highest energy electrons in an s subshell. For example: magnesium (Mg) 1s 2 2s 2 2p 6 3s 2 calcium (Ca) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 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) 1s 2 2s 2 2p 1 gallium (Ga) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 For all elements in group IV, there are two electrons in the highest energy p subshell. For example: silicon (Si) 1s 2 2s 2 2p 6 3s 2 3p 2 germanium (Ge) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 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: Elemental and environmental chemistry 5 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) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 cobalt (Co) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 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) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 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. s-block p-block s 1 p 6 s 2 p 1 p 2 p 3 p 4 p 5 d-block d 1 10 * ** * ** f-block f 1 14 Figure 1.2 s, p, d and f blocks of the periodic table. 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:... n/a... c. Tungsten:... n/a... d. Thallium:......

6 SACE 2 Essentials Chemistry Workbook 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, H 2. Each element in group V forms a compound with hydrogen of formula type XH 3. 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/ nature, the charge(s) of its monatomic ion(s) and its likely oxidation state(s) in its compounds. Metals: Non-metals: Metalloids: Atoms of metals lose electrons in chemical reactions. Atoms of s gain or share electrons in chemical reactions. 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, Mg 2+, 1s 2 2s 2 2p 6, has a more stable configuration than the magnesium atom, Mg, 1s 2 2s 2 2p 6 3s 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 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, 1s 2 2s 2 2p 6 3s 2 3p 4, forms the sulfide ion, S 2, ion, 1s 2 2s 2 2p 6 3s 2 3p 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.

Chapter 1: Elemental and environmental chemistry 7 Valence electrons 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. 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 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 s 1 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, M 2 S, M 2 O and MOH. The configuration of the M + ion is more stable than that of the M atom. The potassium atom has the configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 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, 1s 2 2s 2 2p 6 3s 2 3p 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) + Cl 2(g) 2NaCl (s) 2K (s) + S (s) K 2 S (s) Note that hydrogen is not included as a member of group I. Although its electron configuration is 1s 1, 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 s, hydrogen atoms share their one valence electron with valence electrons of the other atoms.

8 SACE 2 Essentials Chemistry Workbook Group II elements These elements also readily form compounds by reacting with 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 M 2+ ion. The electrons lost are gained by the other reactant. The product compounds are ionic with formulae such as MCl 2, MS, MO and M(OH) 2. The configuration of the M 2+ 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) + O 2(g) 2CaO (s) Mg (s) + S (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 BCl 3. When boron forms compounds, its atoms share the s 2 p 1 outer shell electrons to form covalent bonds with other 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 Al 3+, in compounds with s. These ions are formed when atoms of the group III elements lose the s 2 p 1 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, boron, at the top of the group, to metalloids, aluminium and gallium, in the middle and metals at the bottom of the group.

Chapter 1: Elemental and environmental chemistry 9 Group IV elements In their compounds either they exhibit a covalence of 4, as in the case for carbon and silicon in their compounds, for example, CCl 4 and SiH 4. When carbon and silicon form compounds, their atoms share the s 2 p 2 outer shell electrons to form covalent bonds with other atoms. or they exist as 2+ or 4+ ions, such as Pb 2+, Sn 2+ or Pb 4+, in compounds with s. 4+ ions are formed when the atoms of group IV elements lose the s 2 p 2 outer shell electrons in electron transfer reactions. 2+ ions are formed when the atoms of group IV elements lose only the p 2 electrons of the s 2 p 2 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 s, 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 NH 3, and of phosphorus and arsenic in some of their compounds, for example AsCl 3. In these compounds, the nitrogen, phosphorus and arsenic atoms share only the p 3 electrons from the s 2 p 3 outer shell configuration to form covalent bonds with other 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, AsCl 5 and P 4 O 10. In these compounds, the phosphorus and arsenic atoms share all of the s 2 p 3 electrons from the s 2 p 3 outer shell configuration to form covalent bonds with other atoms. In sharing in this way they are expanding the octet. or they exist as 3 ions, such as N 3 or P 3 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 s 2 p 3 to s 2 p 6. 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 s, 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.

10 SACE 2 Essentials Chemistry Workbook Group VI elements In their compounds either they exhibit a covalence of 2, as in the case of oxygen in all of its compounds such as H 2 O, and of sulfur and selenium in some of their compounds, for example SF 2. In these compounds the oxygen, sulfur and selenium atoms share only two of the p 4 electrons from the s 2 p 4 outer shell configuration to form covalent bonds with other 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 SO 2 and SeF 4. In these compounds, the sulfur and selenium atoms share all of the p 4 electrons from the s 2 p 4 outer shell configuration to form covalent bonds with other 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 SF 6 and SeO 3. In these compounds, the sulfur and selenium atoms share all of the s 2 p 4 electrons from the outer shell configuration to form covalent bonds with other 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 O 2 or S 2 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 s 2 p 4 to s 2 p 6. 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 H 2 O 2. The elements range from the s, 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 CCl 4. In these compounds, the group VII atoms share one of the p 5 electrons from the s 2 p 5 outer shell configuration to form a covalent bond with other 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 p 5 electrons shared covalence of 5: all five of the p 5 electrons shared covalence of 7: all seven of the s 2 p 5 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 s 2 p 5 to s 2 p 6. 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 s.

Chapter 1: Elemental and environmental chemistry 11 Limitations of the covalent bonding model 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. Element Oxidation state Example Nitrogen +4 NO 2 Nitrogen +2 NO Chlorine +4 ClO 2 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/ nature of the elements: I H +1 Li metal +1 II III IV V VI VII Be metalloid +2 B s +3, 3 C s +4, 4 N s +5, +3, 3 O 2 F 1 VIII Na metal +1 Mg metal +2 Al metalloid +3 Si s +4, 4 P s +5, +3, 3 S s +6, +4, +2, 2 C1 s +7, +5, +3, +1, 1 K metal +1 Ca metal +2 TRANSITION METALS Ga metalloid +3 Ge metalloid s +4, +2 As metalloid s +5, +3 Se s +4, +2 Br s +7, +5, +3, +1, 1 Rb metal +1 Sr metal +2 Figure 1.3 Oxidation states and nature of s and p block elements.

12 SACE 2 Essentials Chemistry Workbook 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 s. 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 Low electronegativity (metals) Li 0.98 Na 0.93 K 0.82 Be 1.57 Mg 1.31 Ca 1.00 B 2.04 Al 1.61 Ga 1.81 C 2.55 Si 1.90 Ge 2.01 N 3.04 P 2.19 As 2.18 O 3.44 S 2.58 Se 2.55 F 3.98 Cl 3.16 Br 2.96 Intermediate electronegativity (metalloids) High electronegativity (s) Rb 0.82 Sr 0.95 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, Fe 2 O 3, and nitrogen dioxide, NO 2. 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, CO 2 3, sulfate, SO 2 4 and aluminate, AlO 2. 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, H 2 CO 3, sulfuric acid, H 2 SO 4 and orthophosphoric acid, H 3 PO 4. These oxyacids consist of molecules with covalent hydroxyl groups (O H) as part of their structure. For example, the structure of H 2 CO 3 is: O H C H O O Oxyacids undergo complete or partial ionisation with water to produce hydronium ions. For example: H 2 CO 3 + H 2 O H 3 O + + HCO 3

Chapter 1: Elemental and environmental chemistry 13 Acidic oxides are the oxides of s. They are the oxides of the elements with high electronegativity and are covalent molecular (such as CO 2 and SO 3 ) or continuous covalent compounds (such as SiO 2 ). Table 1.4 below summarises the reactions of some acidic oxides with hydroxide ions and with water (where reactions occur). Oxide *Reaction with hydroxide ions *Reaction with water P 4 O 10 P 4 O 10 + 12OH 4PO 3 4 + 6H 2 O P 4 O 10 + 6H 2 O 4H 3 PO 4 (phosphate) (orthophosphoric acid) SO 2 SO 2 + 2OH SO 2 3 + H 2 O SO 2 + H 2 O H 2 SO 3 (sulfite) (sulfurous acid) SO 3 SO 3 + 2OH SO 2 4 + H 2 O SO 3 + H 2 O H 2 SO 4 (sulfate) (sulfuric acid) CO 2 CO 2 + 2OH CO 2 3 + H 2 O CO 2 + H 2 O H 2 CO 3 (carbonate) (carbonic acid) SiO 2 SiO 2 + 2OH SiO 2 3 + H 2 O (silicate) * Note: The oxidation numbers of the elements are unchanged in these reactions. Table 1.4 Reactions of acidic oxides. NO REACTION Q1.4 One of the oxides of chlorine is Cl 2 O. 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 Cl 2 O with hydroxide ions.... b. Write an equation for the reaction of Cl 2 O 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, O 2, 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 Na 2 O Na 2 O + 2H + 2Na + + H 2 O Na 2 O + H 2 O 2Na + + 20H MgO MgO + 2H + Mg 2+ + H 2 O MgO + H 2 O Mg 2+ + 20H CuO CuO + 2H + Cu 2+ + H 2 O NO REACTION Fe 2 O 3 Fe 2 O 3 + 6H + 2Fe 3+ + 3H 2 O NO REACTION * Note: The oxidation numbers of the elements are unchanged in these reactions. Table 1.5 Reactions of basic oxides.

14 SACE 2 Essentials Chemistry Workbook Q1.5 Barium oxide, BaO, and lithium oxide, Li 2 O, are both basic oxides. a. Write an equation for the reaction of BaO with hydrogen ions.... b. Write an equation for the reaction of Li 2 O 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 Reaction with hydroxide ions ZnO ZnO + 2H + Zn 2+ + H 2 O ZnO + 2OH ZnO 2 2 + H 2 O (zincate) Al 2 O 3 Al 2 O 3 + 6H + 2Al 3+ + 3H 2 O Al 2 O 3 + 2OH 2AlO 2 + H 2 O (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, PbO 2 2.... 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.............

Chapter 1: Elemental and environmental chemistry 15 Molecular substances 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 CO 2, SO 2, H 2 S, Cl 2, NH 3, O 3, CFCl 3 and H 2 SO 4. Small molecules are formed when atoms of elements covalently bond to each other. These 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 elements and compounds that are not molecular. For example, silicon dioxide, SiO 2, 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 Element 2 Compound 1 Boiling point 58 C, melting point 7 C Boiling point 2,680 C, melting point 1,410 C Compound of calcium and sulfur Compound 2 Formula Cl 2 0 Compound 3 Compound 4 Liquid at room temperature. This liquid is a non-conductor of electricity Silver chloride