Oxoacids-Main Group Elements

Oxyacid, any oxygen-containing acid. Most covalent nonmetallic oxides react with water to form acidic oxides; that is, they react with water to form oxyacids that yield hydronium ions (H₃O⁺) in solution. There are some exceptions, such as carbon monoxide, CO, nitrous oxide, N₂O, and nitric oxide, NO.

The strength of an oxyacid is defined by the extent to which it dissociates in water (i.e., its ability to form H⁺ ions). In general, the relative strength of oxyacids can be predicted on the basis of the electronegativity and oxidation number of the central nonmetal atom. The acid strength increases as the electronegativity of the central atom increases. For example, because the electronegativity of chlorine (Cl) is greater than that of sulfur (S), which is in turn greater than that of phosphorus (P), it can be predicted that perchloric acid, HClO₄, is a stronger acid than sulfuric acid, H₂SO₄, which should be a stronger acid than phosphoric acid, H₃PO₄. For a given nonmetal central atom, the acid strength increases as the oxidation number of the central atom increases. For example, nitric acid, HNO₃, in which the nitrogen (N) atom has an oxidation number of +5, is a stronger acid than nitrous acid, HNO₂, where the nitrogen oxidation state is +3. In the same manner, sulfuric acid, H₂SO₄, with sulfur in its +6 oxidation state, is a stronger acid than sulfurous acid, H₂SO₃, where a +4 oxidation number of sulfur exists.

The salt of an oxyacid is a compound formed when the acid reacts with a base: acid + base → salt + water. This type of reaction is called neutralization, because the solution is made neutral.

Reactions with Hydrogen

Oxygen reacts with hydrogen to produce two compounds: water (H2O$H_{2}O$) and hydrogen peroxide (H2O2$H_2{O}_2$). Water is a versatile compound and participates in acid-base equilibrium and oxidation-reduction reactions. It can act as an acid, base, reducing agent, or oxidizing agent. Water's multifaceted abilities make it one of the most important compounds on earth. The reaction between hydrogen and oxygen to form water is given below:

2H2(g)+O2(g)→2H2O(l)(1.1)$$\begin{array}{}\text{(1.1)}& 2H_{2(g)}+O_{2(g)}\to 2H_2{O}_{(l)}\end{array}$$

Hydrogen peroxide's potent oxidizing abilities give it great industrial potential. The following equation shows the reaction of hydrogen and oxygen to form hydrogen peroxide:

H2+O2→H2O2(1.2)$$\begin{array}{}\text{(1.2)}& H_2+O_2\to H_2{O}_2\end{array}$$

The product of this reaction is called a peroxide because oxygen is in the O2−2$O_2^{2-}$ form (hydrogen has a +1 oxidation state). This concept is further explained regarding lithium below.

Reactions with Group 1 Elements

Reactions with Group 2 Elements

Reactions with Group 13 Elements

Reactions with Group 14 Elements

Group 14 is made up of both metals (toward the bottom of the group), metalloids, and nonmetals (at the top of the group). The oxides of the top of Group 4 elements are slightly acidic, and the acidity of the oxides decreases down the group.

• Non-metal: The non-metal carbon of Group 14 (and its compounds) burn to form CO2(g)$C{O}_{2(g)}$ and, in smaller amounts, CO(g)$C{O}_{(g)}$; both are acidic under different conditions. Carbon monoxide is only slightly soluble in water and does not react with it. Click here for more Information.
• Metalloid: The metalloid silicon reacts with oxygen to form only one stable compound, SiO​2$Si{O}_2$, which dissolves slightly in water and is weakly acidic(Figure 2).
• The three metals in this group have many different oxide compounds due to their extended octets. All of these oxides are amphoteric (exhibit both basic and acidic properties). For example:
  • Germanium: GeO$GeO$, GeO​2$Ge{O}_2$
  • Tin: SnO$SnO$, SnO​2$Sn{O}_2$
  • Led;PbO$PbO$, PbO​2$Pb{O}_{}2$, Pb​3O4$P{b}_{}3O_4$ 

Reactions with Group 15 Elements

The nitrogen family, Group 15, is capable of reacting with oxygen in many different ways. Nitrogen and phosphorus are nonmetallic, arsenic and antimony are metalloids, and bismuth is metallic.

Nitrogen 

Nitrogen reacts with oxygen to form many oxides ranging in oxidation states from +1 to +5: All these oxides are gases at room temperature except for N_​2O₅, which is solid. The nitrogen oxides are given below:

NO, N₂O, N₂O₃, NO₂, N₂O₅

All of these reactions are endothermic, requiring energy for oxygen to react directly with N₂(g). The oxides of nitrogen are acidic (because they are nonmetal oxides). N_​2O₃ and N_​2O₅ react with water to give acidic solutions of oxoacids. These reactions are shown below:

Nitrous acid:

N2O3(s)+H2O(l)→2HNO2(aq)(1.24)$$\begin{array}{}\text{(1.24)}& N_2{O}_3(s)+H_{2}O(l)\to 2HN{O}_2(aq)\end{array}$$

Nitric acid:

N2O5(s)+H2O(l)→2HNO3(aq)(1.25)$$\begin{array}{}\text{(1.25)}& N_2{O}_5(s)+H_{2}O(l)\to 2HN{O}_3(aq)\end{array}$$

Phosphorus

There are two forms of allotropes of phosphorus, white phosphorus and red phosphorus. Red phosphorus is less reactive than white phosphorus. Phosphorus reacts with oxygen, usually forming two oxides depending on the amount available oxygen: P₄O6 when reacted with a limited supply of oxygen, and P_​4O₁₀ when reacted with excess oxygen; the latter is shown below.

P4O10+6H2O→4H3PO(1.26)$$\begin{array}{}\text{(1.26)}& P_4{O}_{10}+6H_{2}O\to 4H_{3}PO\end{array}$$

On rare occasions, P_​4O₇, P_​4O₈, and P_​4O₉ are also formed, but in small amounts.
Both P_​4O₄ and P_​4O₁₀ react with water to generate oxoacids. Reactions are shown below.
Phosphorous acid:

P4O6(l)+6H2O(l)→4H3PO3(aq)(1.27)$$\begin{array}{}\text{(1.27)}& P_4{O}_6(l)+6H_{2}O(l)\to 4H_{3}P{O}_3(aq)\end{array}$$

Phosphoric acid:

P4O10(s)+6H2O(l)→4H3PO4(aq)(1.28)$$\begin{array}{}\text{(1.28)}& P_4{O}_{10}(s)+6H_{2}O(l)\to 4H_{3}P{O}_4(aq)\end{array}$$

Other Group 15 Elements

Arsenic, antimony and bismuth react with oxygen when burned. The common oxidation states for arsenic, antimony, and bismuth are +3 and +5. There are two main types of oxides for each element: 

• Arsenic: As_​2O₃, As_​2O₅
• Antimony: Sb_​2O₃, Sb_​2O₅
• Bismuth: Bi_​2O₃, Bi_​2O₅

There are other oxides, such as Sb_​4O₁₀, that are not formed directly through reaction with oxygen. Arsenic(III) oxide and antimony(III) oxide are amphoteric, whereas bismuth(III) oxide acts only as a base (this is because it is the most metallic element in the group).

Reactions with Group 16 Elements

The elements in Group 16 include oxygen, sulfur, selenium, tellurium, and polonium. Oxygen reacts with the elements in its own group to form various oxides, mostly in the form of AO₂ and AO₃. 

Oxygen

Although oxygen is located in Group 16, it is unique in its extreme electronegativity; this allows it to readily gain electrons and create hydrogen bonds. Because it is the smallest element in its group, it is capable of forming double bonds. It has no d-orbitals, and cannot expand its valence shell.

Oxygen is capable of reacting with itself, forming allotropes. One of oxygen's allotropes, ozone (O₃), is formed when oxygen gas, O₂, is subjected to ultraviolet light.

Sulfur

Sulfur dioxide, SO₂, and sulfur trioxide, SO₃, are the only common sulfur oxides. 

S(s)+O2(g)−→ΔSO2(g)(1.29)$$\begin{array}{}\text{(1.29)}& S(s)+O_2(g)\stackrel{\mathrm{\Delta }}{\to }S{O}_2(g)\end{array}$$

Sulfur's reaction with oxygen produces the oxides mentioned above as well as oxoacids. All are powerful oxidizing agents. SO₂ is mainly used to make SO₃, which reacts with water to produce sulfuric acid (recall that nonmetals form acidic oxides). These sequential reactions are shown below:

2SO2(g)+O2(g)⇌2SO3(g)(1.30)$$\begin{array}{}\text{(1.30)}& 2S{O}_2(g)+O_2(g)\rightleftharpoons 2S{O}_3(g)\end{array}$$

2SO3(g)+H2O(l)→H2SO4(aq)(1.31)$$\begin{array}{}\text{(1.31)}& 2S{O}_3(g)+H_{2}O(l)\to H_{2}S{O}_4(aq)\end{array}$$

Selenium and tellurium 

Selenium and tellurium adopt compounds of the forms AO₂, AO₃, and AO.

Reactions with Group 17 Elements

The elements in Group 17 include fluorine, chlorine, bromine, and iodine. These elements are called halogens, from Greek roots translating to "salt formers." The halogens react with oxygen, but many of the resulting compounds are unstable, lasting for only moments at a time. They range in structure from X_​2O to X_​2O₇, where X represents a halogen. Their extended octets allow them to bond with many oxygen atoms at a time.

Fluorine:

The most electronegative element adopts the -1 oxidation state. Fluorine and oxygen form OF₂, which is known as oxygen fluoride.

2F2+2NaOH→OF2+2NaF+H2O(1.32)$$\begin{array}{}\text{(1.32)}& 2F_2+2NaOH\to O{F}_2+2NaF+H_{2}O\end{array}$$

Other Halogens

The other halogens form oxoacids instead of oxides. For example:

Oxidation state of halogen	Chlorine	Bromine	Iodine
+1	HOCl	HOBr	HOI
+3	HClO2	–—	–—
+5	HClO3	HBrO3	HIO HIO3
+7	HClO4	HBrO4	HIO4; H5IO6

Reactions with Group 18 Elements

The Group 18 noble gases include helium, neon, krypton, xenon, and radon. Noble gases are chemically inert with the exception of xenon, which reacts with oxygen to form XeO_​3 and XeO_​4 at low temperatures and high pressures. The ionization energy of xenon is low enough for the electronegative oxygen atom to capture electrons. XeO_​3 is highly unstable, and is known to spontaneously detonate in a clean, dry environment.