Chapter 4

The reaction of iron(III) oxide with aluminum to give molten iron is known as the thermite reaction (page \(172)\). $$\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+2 \mathrm{Al}(\mathrm{s}) \rightarrow 2 \mathrm{Fe}(\ell)+\mathrm{Al}_{2} \mathrm{O}_{3}(\mathrm{s})$$ What amount of \(\mathrm{Al}\), in moles, is needed for complete reaction with 3.0 mol of \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) ? What mass of \(\mathrm{Fe},\) in grams, can be produced?

What mass of HCl, in grams, is required to react with \(0.750 \mathrm{g}\) of \(\mathrm{Al}(\mathrm{OH})_{3} ?\) What mass of water, in grams, is produced? $$\mathrm{Al}(\mathrm{OH})_{3}(\mathrm{s})+3 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{AlCl}_{3}(\mathrm{aq})+3 \mathrm{H}_{2} \mathrm{O}(\ell)$$

Like many metals, aluminum reacts with a halogen (here the orange-brown liquid \(\mathrm{Br}_{2}\) ) to give a metal halide, aluminum bromide. (The white solid on the lip of the beaker at the end of the reaction is \(\mathrm{Al}_{2} \mathrm{Br}_{6} .\) ) $$2 \mathrm{Al}(\mathrm{s})+3 \mathrm{Br}_{2}(\ell) \rightarrow \mathrm{Al}_{2} \mathrm{Br}_{6}(\mathrm{s})$$ What mass of \(\mathrm{Br}_{2}\), in grams, is required for complete reaction with \(2.56 \mathrm{g}\) of \(\mathrm{Al}\) ? What mass of white, solid \(\mathrm{Al}_{2} \mathrm{Br}_{6}\) is expected?

The balanced equation for the reduction of iron ore to the metal using CO is $$\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+3 \mathrm{CO}(\mathrm{g}) \rightarrow 2 \mathrm{Fe}(\mathrm{s})+3 \mathrm{CO}_{2}(\mathrm{g})$$ (a) What is the maximum mass of iron, in grams, that can be obtained from \(454 \mathrm{g}(1.00 \mathrm{lb})\) of iron(III) Oxide? (b) What mass of \(\mathrm{CO}\) is required to react with \(454 \mathrm{g}\) of \(\mathrm{Fe}_{2} \mathrm{O}_{3} ?\)

Methane, \(\mathrm{CH}_{4},\) burns in oxygen. (a) What are the products of the reaction? (b) Write the balanced equation for the reaction. (c) What mass of \(\mathrm{O}_{2}\), in grams, is required for complete combustion of \(25.5 \mathrm{g}\) of methane? (d) What is the total mass of products expected from the combustion of \(25.5 \mathrm{g}\) of methane?

The formation of water-insoluble silver chloride is useful in the analysis of chloride-containing substances. Consider the following unbalanced equation: $$\mathrm{BaCl}_{2}(\mathrm{aq})+\mathrm{AgNO}_{3}(\mathrm{aq}) \rightarrow \mathrm{AgCl}(\mathrm{s})+\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})$$ (a) Write the balanced equation. (b) What mass of \(\mathrm{AgNO}_{3}\), in grams, is required for complete reaction with 0.156 g of \(\mathrm{BaCl}_{2} ?\) What mass of AgCl is produced?

The metals industry was a major source of air pollution years ago. One common process involved "roasting" metal sulfides in the air: $$2 \mathrm{PbS}(\mathrm{s})+3 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{PbO}(\mathrm{s})+2 \mathrm{SO}_{2}(\mathrm{g})$$ If 2.50 mol of \(\mathrm{PbS}\) is heated in air, what amount of \(\mathrm{O}_{2}\) is required for complete reaction? What amounts of \(\mathrm{PbO}\) and \(\mathrm{SO}_{2}\) are expected?

Iron ore is converted to iron metal in a reaction with carbon. $$2 \mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+3 \mathrm{C}(\mathrm{s}) \rightarrow 4 \mathrm{Fe}(\mathrm{s})+3 \mathrm{CO}_{2}(\mathrm{g})$$ If 6.2 mol of \(\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})\) is used, what amount of \(\mathrm{C}(\mathrm{s})\) is needed, and what amounts of Fe and \(\mathrm{CO}_{2}\) are produced?

Chromium metal reacts with oxygen to give chromium(III) oxide, \(\mathrm{Cr}_{2} \mathrm{O}_{3}\) (a) Write a balanced equation for the reaction. (b) What mass (in grams) of \(\mathrm{Cr}_{2} \mathrm{O}_{3}\) is produced if \(0.175 \mathrm{g}\) of chromium metal is converted completely to the oxide? (c) What mass of \(\mathrm{O}_{2}\) (in grams) is required for the reaction?

Ethane, \(\mathrm{C}_{2} \mathrm{H}_{6},\) burns in oxygen. (a) What are the products of the reaction? (b) Write the balanced equation for the reaction. (c) What mass of \(\mathrm{O}_{2}\), in grams, is required for complete combustion of 13.6 of ethane? (d) What is the total mass of products expected from the combustion of 13.6 g of ethane?

Sodium sulfide, \(\mathrm{Na}_{2} \mathrm{S},\) is used in the leather industry to remove hair from hides. The \(\mathrm{Na}_{2} \mathrm{S}\) is made by the reaction $$\mathrm{Na}_{2} \mathrm{SO}_{4}(\mathrm{s})+4 \mathrm{C}(\mathrm{s}) \rightarrow \mathrm{Na}_{2} \mathrm{S}(\mathrm{s})+4 \mathrm{CO}(\mathrm{g})$$ Suppose you mix \(15 \mathrm{g}\) of \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) and \(7.5 \mathrm{g}\) of \(\mathrm{C}\) Which is the limiting reactant? What mass of \(\mathrm{Na}_{2} \mathrm{S}\) is produced?

The compound \(\mathrm{SF}_{6}\) is made by burning sulfur in an atmosphere of fluorine. The balanced equation is $$\mathrm{S}_{8}(\mathrm{s})+24 \mathrm{F}_{2}(\mathrm{g}) \rightarrow 8 \mathrm{SF}_{6}(\mathrm{g})$$ Starting with a mixture of 1.6 mol of sulfur, \(S_{8,}\) and 35 mol of \(\mathrm{F}_{2}\), (a) Which is the limiting reagent? (b) What amount of \(\mathrm{SF}_{6}\) is produced?

Disulfur dichloride, \(S_{2} C l_{2},\) is used to vulcanize rubber. It can be made by treating molten sulfur with gaseous chlorine: $$\mathrm{S}_{8}(\ell)+4 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{s}_{2} \mathrm{Cl}_{2}(\ell)$$ Starting with a mixture of \(32.0 \mathrm{g}\) of sulfur and \(71.0 \mathrm{g}\) of \(\mathrm{Cl}_{2}\). (a) Which is the limiting reactant? (b) What is the theoretical yield of \(S_{2} C l_{2} ?\) (c) What mass of the excess reactant remains when the reaction is completed?

The reaction of methane and water is one way to prepare hydrogen for use as a fuel: $$\mathrm{CH}_{4}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \rightarrow \mathrm{CO}(\mathrm{g})+3 \mathrm{H}_{2}(\mathrm{g})$$ If you begin with 995 g of \(\mathrm{CH}_{4}\) and \(2510 \mathrm{g}\) of water, (a) Which reactant is the limiting reactant? (b) What is the maximum mass of \(\mathrm{H}_{2}\) that can be prepared? (c) What mass of the excess reactant remains when the reaction is completed?

Aluminum chloride, \(\mathrm{AlCl}_{3}\), is made by treating scrap aluminum with chlorine. $$2 \mathrm{Al}(\mathrm{s})+3 \mathrm{Cl}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{AlCl}_{3}(\mathrm{s})$$ If you begin with \(2.70 \mathrm{g}\) of \(\mathrm{Al}\) and \(4.05 \mathrm{g}\) of \(\mathrm{Cl}_{2}\). (a) Which reactant is limiting? (b) What mass of AlCl\(_{3}\) can be produced? (c) What mass of the excess reactant remains when the reaction is completed? (d) Set up an amounts table for this problem.

In the thermite reaction, iron(III) oxide is reduced by aluminum to give molten iron. $$\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{s})+2 \mathrm{Al}(\mathrm{s}) \rightarrow 2 \mathrm{Fe}(\ell)+\mathrm{Al}_{2} \mathrm{O}_{3}(\mathrm{s})$$ If you begin with \(10.0 \mathrm{g}\) of \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) and \(20.0 \mathrm{g}\) of \(\mathrm{Al}\), (a) Which reactant is limiting? (b) What mass of Fe can be produced? (c) What mass of the excess reactant remains after the limiting reactant is consumed? (d) Set up an amounts table for this problem.

Aspirin, \(\mathrm{C}_{6} \mathrm{H}_{4}\left(\mathrm{OCOCH}_{3}\right) \mathrm{CO}_{2} \mathrm{H},\) is produced by the reaction of salicylic acid, \(\mathrm{C}_{6} \mathrm{H}_{4}(\mathrm{OH}) \mathrm{CO}_{2} \mathrm{H}\) and acetic anhydride, \(\left(\mathrm{CH}_{3} \mathrm{CO}\right)_{2} \mathrm{O}\) (page 182 ). $$\mathrm{C}_{6} \mathrm{H}_{4}(\mathrm{OH}) \mathrm{CO}_{2} \mathrm{H}(\mathrm{s})+\left(\mathrm{CH}_{3} \mathrm{CO}\right)_{2} \mathrm{O}(\ell) \rightarrow \mathrm{C}_{6} \mathrm{H}_{4}\left(\mathrm{OCOCH}_{3}\right) \mathrm{CO}_{2} \mathrm{H}(\mathrm{s})+\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}(\ell)$$ If you mix \(100 .\) g of each of the reactants, what is the maximum mass of aspirin that can be obtained?

In Example \(4.2,\) you found that a particular mixture of \(\mathrm{CO}\) and \(\mathrm{H}_{2}\) could produce \(407 \mathrm{g}\) \(\mathrm{CH}_{3} \mathrm{OH}\). $$\mathrm{CO}(\mathrm{g})+2 \mathrm{H}_{2}(\mathrm{g}) \rightarrow \mathrm{CH}_{3} \mathrm{OH}(\ell)$$ If only \(332 \mathrm{g}\) of \(\mathrm{CH}_{3} \mathrm{OH}\) is actually produced, what is the percent yield of the compound?

Ammonia gas can be prepared by the following reaction: $$\begin{aligned} \mathrm{CaO}(\mathrm{s})+2 \mathrm{NH}_{4} \mathrm{Cl}(\mathrm{s}) & \rightarrow \\ 2 \mathrm{NH}_{3}(\mathrm{g}) &+\mathrm{H}_{2} \mathrm{O}(\mathrm{g})+\mathrm{CaCl}_{2}(\mathrm{s}) \end{aligned}$$ If \(112 \mathrm{g}\) of \(\mathrm{CaO}\) and \(224 \mathrm{g}\) of \(\mathrm{NH}_{4} \mathrm{Cl}\) are mixed, the theoretical yield of \(\mathrm{NH}_{3}\) is \(68.0 \mathrm{g}\) (Study Question 12 ). If only \(16.3 \mathrm{g}\) of \(\mathrm{NH}_{3}\) is actually obtained, what is its percent yield?

The deep blue compound \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4}\) is made by the reaction of copper(II) sulfate and ammonia. $$\mathrm{CuSO}_{4}(\mathrm{aq})+4 \mathrm{NH}_{3}(\mathrm{aq}) \rightarrow \mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4}(\mathrm{aq})$$ (a) If you use \(10.0 \mathrm{g}\) of \(\mathrm{CuSO}_{4}\) and excess \(\mathrm{NH}_{3}\) what is the theoretical yield of \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4} ?\) (b) If you isolate \(12.6 \mathrm{g}\) of \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4},\) what is the percent yield of \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4} \mathrm{SO}_{4} ?\)

Black smokers are found in the depths of the oceans. Thinking that the conditions in these smokers might be conducive to the formation of organic compounds, two chemists in Germany found the following reaction could occur in similar conditions. $$2 \mathrm{CH}_{3} \mathrm{SH}+\mathrm{CO} \rightarrow \mathrm{CH}_{3} \mathrm{COSCH}_{3}+\mathrm{H}_{2} \mathrm{S}$$ If you begin with \(10.0 \mathrm{g}\) of \(\mathrm{CH}_{3} \mathrm{SH}\) and excess \(\mathrm{CO}\). (a) What is the theoretical yield of \(\mathrm{CH}_{3} \mathrm{COSCH}_{3} ?\) (b) If \(8.65 \mathrm{g}\) of \(\mathrm{CH}_{3} \mathrm{COSCH}_{3}\) is isolated, what is its percent yield?

The reaction of methane and water is one way to prepare hydrogen for use as a fuel: $$\mathrm{CH}_{4}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \rightarrow \mathrm{CO}(\mathrm{g})+3 \mathrm{H}_{2}(\mathrm{g})$$ If this reaction has a \(37 \%\) yield under certain conditions, what mass of \(\mathrm{CH}_{4}\) is required to produce \(15 \mathrm{g}\) of \(\mathrm{H}_{2} ?\)

Methanol, \(\mathrm{CH}_{3} \mathrm{OH},\) can be prepared from carbon monoxide and hydrogen. $$\mathrm{CO}(\mathrm{g})+2 \mathrm{H}_{2}(\mathrm{g}) \rightarrow \mathrm{CH}_{3} \mathrm{OH}(\ell)$$ What mass of hydrogen is required to produce 1.0 L of \(\mathrm{CH}_{3} \mathrm{OH}(d=0.791 \mathrm{g} / \mathrm{mL})\) if this reaction has a \(74 \%\) yield under certain conditions?

A mixture of \(\mathrm{CuSO}_{4}\) and \(\mathrm{CuSO}_{4} \cdot 5 \mathrm{H}_{2} \mathrm{O}\) has a mass of \(1.245 \mathrm{g}\). After heating to drive off all the water, the mass is only \(0.832 \mathrm{g}\). What is the mass percent of \(\mathrm{CuSO}_{4} \cdot 5 \mathrm{H}_{2} \mathrm{O}\) in the mixture? (See page \(98 .\) )

A 2.634 -g sample containing impure \(\mathrm{CuCl}_{2} \cdot 2\) \(\mathrm{H}_{2} \mathrm{O}\) was heated. The sample mass after heating to drive off the water was 2.125 g. What was the mass percent of \(\mathrm{CuCl}_{2} \cdot 2 \mathrm{H}_{2} \mathrm{O}\) in the original sample?

A sample of limestone and other soil materials was heated, and the limestone decomposed to give calcium oxide and carbon dioxide. $$\mathrm{CaCO}_{3}(\mathrm{s}) \rightarrow \mathrm{CaO}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{g})$$ A \(1.506-\mathrm{g}\) sample of limestone-containing material gave \(0.558 \mathrm{g}\) of \(\mathrm{CO}_{2}\), in addition to \(\mathrm{CaO}\), after being heated at a high temperature. What was the mass percent of \(\mathrm{CaCO}_{3}\) in the original sample?

At higher temperatures, \(\mathrm{NaHCO}_{3}\) is converted quantitatively to \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) $$2 \mathrm{NaHCO}_{3}(\mathrm{s}) \rightarrow \mathrm{Na}_{2} \mathrm{CO}_{3}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g})$$ Heating a 1.7184 -g sample of impure \(\mathrm{NaHCO}_{3}\) gives \(0.196 \mathrm{g}\) of \(\mathrm{CO}_{2} .\) What was the mass percent of \(\mathrm{NaHCO}_{3}\) in the original 1.7184 -g sample?

Nickel(II) sulfide, NiS, occurs naturally as the relatively rare mineral millerite. One of its occurrences is in meteorites. To analyze a mineral sample for the quantity of NiS, the sample is dissolved in nitric acid to form a solution of \(\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}\) $$\begin{aligned} \mathrm{NiS}(\mathrm{s})+& 4 \mathrm{HNO}_{3}(\mathrm{aq}) \rightarrow \\ & \mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+2 \mathrm{NO}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{S}(\mathrm{s}) \end{aligned}$$ The aqueous solution of \(\mathrm{Ni}\left(\mathrm{NO}_{3}\right)_{2}\) is then reacted with the organic compound dimethylglyoxime \(\left(\mathrm{C}_{4} \mathrm{H}_{8} \mathrm{N}_{2} \mathrm{O}_{2}\right)\) to give the red solid \(\mathrm{Ni}\left(\mathrm{C}_{4} \mathrm{H}_{7} \mathrm{N}_{2} \mathrm{O}_{2}\right)_{2}\). Suppose a \(0.468-\mathrm{g}\) sample containing millerite produces \(0.206 \mathrm{g}\) of red, solid \(\mathrm{Ni}\left(\mathrm{C}_{4} \mathrm{H}_{7} \mathrm{N}_{2} \mathrm{O}_{2}\right)_{2}\) What is the mass percent of NiS in the sample?

The aluminum in a 0.764-g sample of an unknown material was precipitated as aluminum hydroxide, \(\mathrm{Al}(\mathrm{OH})_{3},\) which was then converted to \(\mathrm{Al}_{2} \mathrm{O}_{3}\) by heating strongly. If \(0.127 \mathrm{g}\) of \(\mathrm{Al}_{2} \mathrm{O}_{3}\) is obtained from the 0.764-g sample, what is the mass percent of aluminum in the sample?

Styrene, the building block of polystyrene, consists of only \(\mathrm{C}\) and \(\mathrm{H}\). If 0.438 g of styrene is burned in oxygen and produces 1.481 g of \(\mathrm{CO}_{2}\) and \(0.303 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O},\) what is the empirical formula of styrene?

Mesitylene is a liquid hydrocarbon. Burning \(0.115 \mathrm{g}\) of the compound in oxygen gives \(0.379 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.1035 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O}\). What is the empirical formula of mesitylene?

Naphthalene is a hydrocarbon that once was used in mothballs. If 0.3093 g of the compound is burned in oxygen, \(1.0620 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.1739 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O}\) are isolated. (a) What is the empirical formula of naphthalene? (b) If a separate experiment gave \(128.2 \mathrm{g} / \mathrm{mol}\) as the molar mass of the compound, what is its molecular formula?

Azulene is a beautiful blue hydrocarbon. If \(0.106 \mathrm{g}\) of the compound is burned in oxygen, \(0.364 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.0596 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O}\) are isolated. (a) What is the empirical formula of azulene? (b) If a separate experiment gave \(128.2 \mathrm{g} / \mathrm{mol}\) as the molar mass of the compound, what is its molecular formula?

An unknown compound has the formula \(\mathrm{C}_{x} \mathrm{H}_{y} \mathrm{O}_{z}\) You burn 0.0956 g of the compound and isolate \(0.1356 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.0833 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O} .\) What is the empirical formula of the compound? If the molar mass is \(62.1 \mathrm{g} / \mathrm{mol},\) what is the molecular formula?

An unknown compound has the formula \(\mathrm{C}_{x} \mathrm{H}_{y} \mathrm{O}_{z}\) You burn \(0.1523 \mathrm{g}\) of the compound and isolate \(0.3718 \mathrm{g}\) of \(\mathrm{CO}_{2}\) and \(0.1522 \mathrm{g}\) of \(\mathrm{H}_{2} \mathrm{O} .\) What is the empirical formula of the compound? If the molar mass is \(72.1 \mathrm{g} / \mathrm{mol},\) what is the molecular formula?

Nickel forms a compound with carbon monoxide, \(\mathrm{Ni}_{x}(\mathrm{CO})_{y},\) To determine its formula, you carefully heat a 0.0973-g sample in air to convert the nickel to \(0.0426 \mathrm{g}\) of \(\mathrm{NiO}\) and the CO to \(0.100 \mathrm{g}\) of \(\mathrm{CO}_{2}\) What is the empirical formula of \(\mathrm{Ni}_{\mathrm{x}}(\mathrm{CO})_{y} ?\)

To find the formula of a compound composed of iron and carbon monoxide, \(\mathrm{Fe}_{x}(\mathrm{CO})_{y^{\prime}}\) the compound is burned in pure oxygen to give \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) and \(\mathrm{CO}_{2} .\) If you burn \(1.959 \mathrm{g}\) of \(\mathrm{Fe}_{x}(\mathrm{CO})_{y}\) and obtain \(0.799 \mathrm{g}\) of \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) and \(2.200 \mathrm{g}\) of \(\mathrm{CO}_{2},\) what is the empirical formula of \(\mathrm{Fe}_{x}(\mathrm{CO})_{y} ?\)

If 6.73 g of \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) is dissolved in enough water to make \(250 .\) mL of solution, what is the molar concentration of the sodium carbonate? What are the molar concentrations of the \(\mathrm{Na}^{+}\) and \(\mathrm{CO}_{3}^{2-}\) ions?

Some potassium dichromate \(\left(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\right), 2.335 \mathrm{g}\) is dissolved in enough water to make exactly \(500 .\) mL of solution. What is the molar concentration of the potassium dichromate? What are the molar concentrations of the \(\mathrm{K}^{+}\) and \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}\) ions?

What is the mass of solute, in grams, in \(250 . \mathrm{mL}\) of a 0.0125 M solution of \(\mathrm{KMnO}_{4} ?\)

What is the mass of solute, in grams, in \(125 \mathrm{mL}\) of a \(1.023 \times 10^{-3} \mathrm{M}\) solution of \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) ? What is the molar concentration of the \(\mathrm{Na}^{+}\) and \(\mathrm{PO}_{4}^{3-}\) ion?

What volume of \(0.123 \mathrm{M} \mathrm{NaOH},\) in milliliters, contains \(25.0 \mathrm{g}\) of \(\mathrm{NaOH} ?\)

Identify the ions that exist in each aqueous solution, and specify the concentration of each ion. (a) \(0.25 \mathrm{M}\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}\) (b) \(0.123 \mathrm{M} \mathrm{Na}_{2} \mathrm{CO}_{3}\) (c) \(0.056 \mathrm{M} \mathrm{HNO}_{3}\)

Identify the ions that exist in each aqueous solution, and specify the concentration of each ion. (a) \(0.12 \mathrm{M} \mathrm{BaCl}_{2}\) (b) \(0.0125 \mathrm{M} \mathrm{CuSO}_{4}\) (c) \(0.500 \mathrm{M} \mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\)

An experiment in your laboratory requires \(500 .\) mL of a \(0.0200 \mathrm{M}\) solution of \(\mathrm{Na}_{2} \mathrm{CO}_{3} .\) You are given solid \(\mathrm{Na}_{2} \mathrm{CO}_{3},\) distilled water, and a \(500 .\) -mL volumetric flask. Describe how to prepare the required solution.

What mass of oxalic acid, \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4},\) is required to prepare \(250 .\) mL of a solution that has a concentration of \(0.15 \mathrm{M} \mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4} ?\)

If you dilute \(25.0 \mathrm{mL}\) of \(1.50 \mathrm{M}\) hydrochloric acid to \(500 . \mathrm{mL},\) what is the molar concentration of the dilute acid?

If \(4.00 \mathrm{mL}\) of \(0.0250 \mathrm{M} \mathrm{CuSO}_{4}\) is diluted to \(10.0 \mathrm{mL}\) with pure water, what is the molar concentration of copper(II) sulfate in the diluted solution?

Which of the following methods would you use to prepare 1.00 L of \(0.125 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4} ?\) (a) Dilute \(20.8 \mathrm{mL}\) of \(6.00 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) to a volume of 1.00 L. (b) Add \(950 .\) mL of water to \(50.0 \mathrm{mL}\) of \(3.00 \mathrm{M}\) \(\mathrm{H}_{2} \mathrm{SO}_{4}\).

Which of the following methods would you use to prepare \(300 .\) mL of \(0.500 \mathrm{M} \mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7} ?\) (a) Add \(30.0 \mathrm{mL}\) of \(1.50 \mathrm{M} \mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) to \(270 . \mathrm{mL}\) of water. (b) Dilute \(250 . \mathrm{mL}\) of \(0.600 \mathrm{M} \mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) to a volume of \(300 .\) mL.