Chapter 3

The equation for the oxidation of phosphorus in air is \(\mathrm{P}_{4}(\mathrm{s})+5 \mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{P}_{4} \mathrm{O}_{10}(\mathrm{s}) .\) Identify the reactants and products and the stoichiometric coefficients. To what do the designations s and g refer?

Write an equation from the following description: reactants are gaseous \(\mathrm{NH}_{3}\) and \(\mathrm{O}_{2},\) products are gaseous \(\mathrm{NO}_{2}\) and liquid \(\mathrm{H}_{2} \mathrm{O},\) and the stoichiometric coefficients are \(4,7,4,\) and \(6,\) respectively.

The equation for the reaction of phosphorus and chlorine is \(\mathrm{P}_{4}(\mathrm{s})+6 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{PCl}_{3}(\ell) .\) If you use 8000 molecules of \(P_{4}\) in this reaction how many molecules of \(\mathrm{Cl}_{2}\) are required to consume the \(\mathrm{P}_{4}\) completely?

The equation for the reaction of aluminum and bromine is \(2 \mathrm{Al}(\mathrm{s})+3 \mathrm{Br}_{2}(\ell) \rightarrow \mathrm{Al}_{2} \mathrm{Br}_{6}(\mathrm{s}) .\) If you use \(6.0 \times 10^{23}\) molecules of \(\mathrm{Br}_{2}\) in a reaction how many atoms of Al will be consumed?

Oxidation of \(1.00 \mathrm{g}\) of carbon monoxide, \(\mathrm{CO},\) produces 1.57 g of carbon dioxide, \(\mathrm{CO}_{2} .\) How many grams of oxygen were required in this reaction?

A 0.20 mol sample of magnesium burns in air to form 0.20 mol of solid \(\mathrm{MgO}\). What amount (moles) of oxygen \(\left(\mathrm{O}_{2}\right)\) is required for a complete reaction?

Write balanced chemical equations for the following reactions. (a) The reaction of aluminum and iron(III) oxide to form iron and aluminum oxide (known as the thermite reaction). (b) The reaction of carbon and water at high temperature to form a mixture of gaseous \(\mathrm{CO}\) and \(\mathrm{H}_{2}\) (known as water gas and once used as a fuel). (c) The reaction of liquid silicon tetrachloride and magnesium forming silicon and magnesium chloride. This is one step in the preparation of ultrapure silicon used in the semiconductor industry.

Write balanced chemical equations for the following reactions: (a) production of ammonia, \(\mathrm{NH}_{3}(\mathrm{g}),\) by combin\(\operatorname{ing} \mathrm{N}_{2}(\mathrm{g})\) and \(\mathrm{H}_{2}(\mathrm{g})\) (b) production of methanol, \(\mathrm{CH}_{3} \mathrm{OH}(\ell)\) by combining \(\mathrm{H}_{2}(\mathrm{g})\) and \(\mathrm{CO}(\mathrm{g})\) (c) production of sulfuric acid by combining sulfur, oxygen, and water

Balance the following equations: (a) \(\mathrm{Cr}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{3}(\mathrm{s})\) (b) \(\mathrm{Cu}_{2} \mathrm{S}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{SO}_{2}(\mathrm{g})\) (c) \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CH}_{3}(\ell)+\mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{CO}_{2}(\mathrm{g})\)

Balance the following equations: (a) \(\mathrm{Cr}(\mathrm{s})+\mathrm{Cl}_{2}(\mathrm{g}) \rightarrow \mathrm{CrCl}_{3}(\mathrm{s})\) (b) \(\mathrm{SiO}_{2}(\mathrm{s})+\mathrm{C}(\mathrm{s}) \rightarrow \mathrm{Si}(\mathrm{s})+\mathrm{CO}(\mathrm{g})\) (c) \(\mathrm{Fe}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \rightarrow \mathrm{Fe}_{3} \mathrm{O}_{4}(\mathrm{s})+\mathrm{H}_{2}(\mathrm{g})\)

Balance the following equations, and name each reactant and product: (a) \(\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+\mathrm{Mg}(\mathrm{s}) \rightarrow \mathrm{MgO}(\mathrm{s})+\mathrm{Fe}(\mathrm{s})\) (b) \(\mathrm{AlCl}_{3}(\mathrm{s})+\mathrm{NaOH}(\mathrm{aq}) \rightarrow\) \(\mathrm{Al}(\mathrm{OH})_{3}(\mathrm{s})+\mathrm{NaCl}(\mathrm{aq})\) (c) \(\mathrm{NaNO}_{3}(\mathrm{s})+\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \rightarrow\) \(\mathrm{Na}_{2} \mathrm{SO}_{4}(\mathrm{s})+\mathrm{HNO}_{3}(\mathrm{aq})\) (d) \(\mathrm{NiCO}_{3}(\mathrm{s})+\mathrm{HNO}_{3}(\mathrm{aq}) \rightarrow\) \(\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{CO}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell)\)

Balance the following equations, and name each reactant and product: (a) \(\mathrm{SF}_{4}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{SO}_{2}(\mathrm{g})+\mathrm{HF}(\ell)\) (b) \(\mathrm{NH}_{3}(\mathrm{aq})+\mathrm{O}_{2}(\mathrm{aq}) \rightarrow \mathrm{NO}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell)\) (c) \(\mathrm{BF}_{3}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{HF}(\mathrm{aq})+\mathrm{H}_{3} \mathrm{BO}_{3}(\mathrm{aq})\)

Identify each of the following statements as either true or false. (a) At equilibrium the rates of the forward and reverse reactions are equal. (b) When a reaction reaches equilibrium the forward and reverse reactions cease to occur. (c) Chemical reactions always proceed toward equilibrium.

Equal amounts of two acids-HCl and \(\mathrm{HCO}_{2} \mathrm{H}\) (formic acid)-are placed in aqueous solution. When equilibrium has been achieved, the HCl solution has a much greater electrical conductivity than the HCO_A solution. Which reaction is more product-favored at equilibrium? $$\mathrm{HCl}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})$$ $$\mathrm{HCO}_{2} \mathrm{H}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}(\mathrm{aq})+\mathrm{HCO}_{2}^{-}(\mathrm{aq})$$

Two aqueous solutions were prepared, one containing 0.10 mol of boric acid \(\left(\mathrm{H}_{3} \mathrm{BO}_{3}\right)\) in \(200 \mathrm{mL}\) and the second containing 0.10 mol phosphoric acid \(\left(\mathrm{H}_{3} \mathrm{PO}_{4}\right)\) in \(200 \mathrm{mL}\). Both were weak conductors of electricity, but the \(\mathrm{H}_{3} \mathrm{PO}_{4}\) solution was a noticeably stronger conductor. Write equations to describe the equilibrium in each solution, and explain the observed difference in conductivity.

What is an electrolyte? How can you differentiate experimentally between a weak electrolyte and a strong electrolyte? Give an example of each.

Name and give the formulas of two acids that are strong electrolytes and one acid that is a weak electrolyte. Name and give formulas of two bases that are strong electrolytes and one base that is a weak electrolyte.

Which compound or compounds in each of the following groups is (are) soluble in water? (a) \(\mathrm{CuO}, \mathrm{CuCl}_{2}, \mathrm{FeCO}_{3}\) (b) \(\mathrm{AgI}, \mathrm{Ag}_{3} \mathrm{PO}_{4}, \mathrm{AgNO}_{3}\) (c) \(\mathrm{K}_{2} \mathrm{CO}_{3}, \mathrm{KI}, \mathrm{KMnO}_{4}\)

The following compounds are water-soluble. What ions are produced by each compound in aqueous solution? (a) KOH (b) \(\mathrm{K}_{2} \mathrm{SO}_{4}\) (c) \(\mathrm{LiNO}_{3}\) (d) \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}\)

The following compounds are water-soluble. What ions are produced by each compound in aqueous solution? (a) KI (b) \(\mathrm{Mg}\left(\mathrm{CH}_{3} \mathrm{CO}_{2}\right)_{2}\) (c) \(\mathrm{K}_{2} \mathrm{HPO}_{4}\) (d) \(\mathrm{NaCN}\)

Decide whether each of the following is watersoluble. If soluble, tell what ions are produced when the compound dissolves in water. (a) \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) (b) \(\mathrm{CuSO}_{4}\) (c) NiS (d) \(\mathrm{BaBr}_{2}\)

Decide whether each of the following is watersoluble. If soluble, tell what ions are produced when the compound dissolves in water. (a) \(\mathrm{NiCl}_{2}\) (b) \(\operatorname{Cr}\left(\mathrm{NO}_{3}\right)_{3}\) (c) \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\) (d) \(\mathrm{BaSO}_{4}\)

Balance the equation for the following precipitation reaction, and then write the net ionic equation. Indicate the state of each species \((\mathrm{s}, \ell, \mathrm{aq}\) or g). $$\mathrm{CdCl}_{2}+\mathrm{NaOH} \rightarrow \mathrm{Cd}(\mathrm{OH})_{2}+\mathrm{NaCl}$$

Balance the equation for the following precipitation reaction, and then write the net ionic equation. Indicate the state of each species \((\mathrm{s}, \ell, \mathrm{aq}\) or g). $$\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}+\mathrm{Na}_{2} \mathrm{CO}_{3} \rightarrow \mathrm{NiCO}_{3}+\mathrm{NaNO}_{3}$$

Predict the products of each precipitation reaction. Balance the equation, and then write the net ionic equation. (a) \(\mathrm{NiCl}_{2}(\mathrm{aq})+\left(\mathrm{NH}_{4}\right)_{2} \mathrm{S}(\mathrm{aq}) \rightarrow\) (b) \(\mathrm{Mn}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{Na}_{3} \mathrm{PO}_{4}(\mathrm{aq}) \rightarrow\)

Predict the products of each precipitation reaction. Balance the equation, and then write the net ionic equation. (a) \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{KBr}(\mathrm{aq}) \rightarrow\) (b) \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{KF}(\mathrm{aq}) \rightarrow\) (c) \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(\mathrm{aq}) \rightarrow\)

Write a balanced equation for the ionization of nitric acid in water.

Write a balanced equation for the ionization of perchloric acid in water.

Phosphoric acid can supply one, two, or three \(\mathrm{H}_{3} \mathrm{O}^{+}\) ions in aqueous solution. Write balanced equations (like those for sulfuric acid on page 142 ) to show this successive loss of hydrogen ions.

Write a balanced equation for reaction of the basic oxide, magnesium oxide, with water.

Write a balanced equation for the reaction of sulfur trioxide gas with water.

Complete and balance the equations for the following acid-base reactions. Name the reactants and products. (a) \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}(\mathrm{aq})+\mathrm{Mg}(\mathrm{OH})_{2}(\mathrm{s}) \rightarrow\) (b) \(\mathrm{HClO}_{4}(\mathrm{aq})+\mathrm{NH}_{3}(\mathrm{aq}) \rightarrow\)

Write a balanced equation for the reaction of barium hydroxide with nitric acid.

Write a balanced equation for the reaction of aluminum hydroxide with sulfuric acid.

Write an equation that describes the equilibrium that exists when nitric acid dissolves in water. Identify each of the four species in solution as either Bronsted acids or Bronsted bases. Does the equilibrium favor the products or the reactants?

Write an equation that describes the equilibrium that exists when the weak acid benzoic acid \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CO}_{2} \mathrm{H}\right)\) dissolves in water. Identify each of the four species in solution as either Bronsted acids or Bronsted bases. Does the equilibrium favor the products or the reactants? (In acting as an acid, the \(\left.-\mathrm{CO}_{2} \mathrm{H} \text { group supplies } \mathrm{H}^{+} \text {to form } \mathrm{H}_{3} \mathrm{O}^{+} .\right)\)

Write two chemical equations, one that shows \(\mathrm{H}_{2} \mathrm{O}\) reacting (with HBr) as a Bronsted base and a second that shows \(\mathrm{H}_{2} \mathrm{O}\) reacting (with \(\mathrm{NH}_{3}\) ) as a Bronsted acid.

Write two chemical equations, one in which \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) is a Bronsted acid (in reaction with the carbonate ion, \(\mathrm{CO}_{3}^{2-}\) ), and a second in which \(\mathrm{HPO}_{4}^{2-}\) is a Bronsted base (in reaction with acetic acid, \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\) ).

Balance the following equations, and then write the net ionic equation. (a) \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{CO}_{3}(\mathrm{aq})+\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq}) \rightarrow \mathrm{CuCO}_{3}(\mathrm{s})+\mathrm{NH}_{4} \mathrm{NO}_{3}(\mathrm{aq})\) (b) \(\mathrm{Pb}(\mathrm{OH})_{2}(\mathrm{s})+\mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{PbCl}_{2}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\ell)\) (c) \(\mathrm{BaCO}_{3}(\mathrm{s})+\mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{BaCl}_{2}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{CO}_{2}(\mathrm{g})\) (d) \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}(\mathrm{aq})+\mathrm{Ni}(\mathrm{OH})_{2}(\mathrm{s}) \rightarrow \mathrm{Ni}\left(\mathrm{CH}_{3} \mathrm{CO}_{2}\right)_{2}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell)\)

Balance the following equations, and then write the net ionic equation: (a) \(\mathrm{Zn}(\mathrm{s})+\mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{H}_{2}(\mathrm{g})+\mathrm{ZnCl}_{2}(\mathrm{aq})\) (b) \(\mathrm{Mg}(\mathrm{OH})_{2}(\mathrm{s})+\mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{MgCl}_{2}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell)\) (c) \(\mathrm{HNO}_{3}(\mathrm{aq})+\mathrm{CaCO}_{3}(\mathrm{s}) \rightarrow \mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{CO}_{2}(\mathrm{g})\) (d) \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{S}(\mathrm{aq})+\mathrm{FeCl}_{2}(\mathrm{aq}) \rightarrow \mathrm{NH}_{4} \mathrm{Cl}(\mathrm{aq})+\mathrm{FeS}(\mathrm{s})\)

Balance the following equations, and then write the net ionic equation. Show states for all reactants and products \((s, \ell, g, a q).\) (a) the reaction of silver nitrate and potassium iodide to give silver iodide and potassium nitrate (b) the reaction of barium hydroxide and nitric acid to give barium nitrate and water (c) the reaction of sodium phosphate and nickel(II) nitrate to give nickel(II) phosphate and sodium nitrate

Balance each of the following equations, and then write the net ionic equation. Show states for all reactants and products (s, \(\ell, \mathrm{g},\) aq). (a) the reaction of sodium hydroxide and iron(II) chloride to give iron(II) hydroxide and sodium chloride (b) the reaction of barium chloride with sodium carbonate to give barium carbonate and sodium chloride (c) the reaction of ammonia with phosphoric acid

Write balanced net ionic equations for the following reactions: (a) the reaction of nitrous acid (a weak acid) and sodium hydroxide in aqueous solution (b) the reaction of calcium hydroxide and hydrochloric acid

Write balanced net ionic equations for the following reactions: (a) the reaction of aqueous solutions of silver nitrate and sodium iodide (b) the reaction of aqueous solutions of barium chloride and potassium carbonate

Siderite is a mineral consisting largely of iron(II) carbonate. Write an overall, balanced equation for its reaction with nitric acid, and name the products.

The mineral rhodochrosite is manganese(II) carbonate. Write an overall, balanced equation for the reaction of the mineral with hydrochloric acid, and name the products.

Write an overall, balanced equation for the reaction of \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{S}\) with \(\mathrm{HBr},\) and name the reactants and products.

Write an overall, balanced equation for the reaction of \(\mathrm{Na}_{2} \mathrm{SO}_{3}\) with \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H},\) and name the reactants and products.

Determine the oxidation number of each element in the following ions or compounds. (a) \(\mathrm{BrO}_{3}^{-}\) (b) \(\mathrm{C}_{2} \mathrm{O}_{4}^{2-}\) (c) \(\mathrm{F}^{-}\) (d) \(\mathrm{CaH}_{2}\) (e) \(\mathrm{H}_{4} \mathrm{SiO}_{4}\) (f) \(\mathrm{HSO}_{4}^{-}\)

Determine the oxidation number of each element in the following ions or compounds. (a) \(\mathrm{PF}_{6}\) (b) \(\mathrm{H}_{2} \mathrm{AsO}_{4}^{-}\) (c) \(\mathrm{UO}^{2+}\) (d) \(\mathrm{N}_{2} \mathrm{O}_{5}\) (e) \(\mathrm{POCl}_{3}\) (f) \(\mathrm{XeO}_{4}^{2-}\)