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University of Maine

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Problem 135

The rocket fuel hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)$ is made by the three-step

Raschig process, which has the following overall equation:

$\mathrm{NaOCl}(a q)+2 \mathrm{NH}_{3}(a q) \longrightarrow \mathrm{N}_{2} \mathrm{H}_{4}(a q)+\mathrm{NaCl}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$

What is the percent atom economy of this process?

Answer

29,5 percent

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## Discussion

## Video Transcript

percent. Adam Economy is described as the massive Adams in the desired product, divided by the massive Adams in all reactive. So in this reaction, the desired product is hydrazine and to age for and are to reactant are these compounds here? So we simply use the periodic table to find the massive each Adam multiplying by any coefficients present to find the percent Adam economy. So we have in our product to nitrogen ins plus four hydrogen ins divided by one sodium plus one oxygen, plus the chlorine plus two nitrogen plus six hydrogen, which gives US 29.4% adam economy.

## Recommended Questions

The industrial production of hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)$ by the Raschig process is the topic of the Focus On feature for Chapter 4 on www.masteringchemistry.com. The following chemical equation represents the overall process, which actually involves three consecutive reactions.

$2 \mathrm{NH}_{3}(\mathrm{aq})+\mathrm{Cl}_{2}(\mathrm{g})+2 \mathrm{NaOH}(\mathrm{aq}) \longrightarrow$

$$\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq})+2 \mathrm{NaCl}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})$$

(a) Use the definition of percent atom economy (AE) from exercise 130 to calculate, to the nearest percent, the percent AE for the Raschig process.

(b) Propose a reaction for the synthesis of $\mathrm{N}_{2} \mathrm{H}_{4}$ that has percent AE of $100 \%$.

Hydrazine, $\mathrm{N}_{2} \mathrm{H}_{4}$ , emits a large quantity of energy when it reacts with oxygen, which has led to hydrazine's use as a fuel for rockets:

$$\mathrm{N}_{2} \mathrm{H}_{4}(l)+\mathrm{O}_{2}(g) \rightarrow \mathrm{N}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g)$$

How many moles of each of the gaseous products are produced when 20.0 g of pure hydrazine is ignited in the presence of 20.0 g of pure oxygen? How many grams of each product are produced?

Rocket Fuel The exothermic reaction between liquid hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{2}\right)$ and liquid hydrogen peroxide $\left(\mathrm{H}_{2} \mathrm{O}_{2}\right)$

is used to fuel rockets. The products of this reaction are nitrogen gas and water.

a. Write the balanced chemical equation.

b. How much hydrazine, in grams, is needed to produce 10.0 mol of nitrogen gas?

Hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)$ is used as a fuel in liquid-fueled rockets. When hydrazine reacts with oxygen gas, nitrogen gas and water vapor are produced. Write a balanced equation and use bond energies from Table 8.5 to estimate $\Delta H$ for this reaction.

Hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)$ is a fuel used by some spacecraft. It is normally oxidized by $\mathrm{N}_{2} \mathrm{O}_{4}$ according to the equation:

$$

\mathrm{N}_{2} \mathrm{H}_{4}(l)+\mathrm{N}_{2} \mathrm{O}_{4}(g) \longrightarrow 2 \mathrm{N}_{2} \mathrm{O}(g)+2 \mathrm{H}_{2} \mathrm{O}(g)

$$

Calculate $\Delta H_{\mathrm{rxn}}^{\circ}$ for this reaction using standard enthalpies of

formation.

Hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)$ is used as a rocket fuel because it reacts very exothermically with oxygen to form nitrogen gas and water vapor. The heat released and the increase in number of moles of gas provide thrust. Calculate the enthalpy of reaction.

The hydrazine fuel cell is based on the reaction

$$\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq})+\mathrm{O}_{2}(\mathrm{g}) \longrightarrow \mathrm{N}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(1)$$

The theoretical $E_{\text {cell }}^{\circ}$ of this fuel cell is 1.559 V. Use this information and data from Appendix D to calculate a value of $\Delta_{f} G^{\circ}$ for $\left[\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq})\right]$

Hydrazine, $\mathrm{N}_{2} \mathrm{H}_{4},$ decomposes according to the following reaction: $$3 \mathrm{N}_{2} \mathrm{H}_{4}(l) \longrightarrow 4 \mathrm{NH}_{3}(g)+\mathrm{N}_{2}(g)$$ (a) Given that the standard enthalpy of formation of hydrazine is $50.42 \mathrm{kJ} / \mathrm{mol},$ calculate $\Delta H^{\circ}$ for its decomposition. (b) Both hydrazine and ammonia burn in oxygen to produce $\mathrm{H}_{2} \mathrm{O}(l)$ and $\mathrm{N}_{2}(g) .$ Write balanced equations for each of these processes and calculate $\Delta H^{\circ}$ for each of them. On a mass basis (per kg), would hydrazine or ammonia be the better fuel?

Hydrazine, $\mathrm{N}_{2} \mathrm{H}_{4},$ a base like ammonia, can react with an acid such as sulfuric acid.

$2 \mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \longrightarrow 2 \mathrm{N}_{2} \mathrm{H}_{5}^{+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})$

What mass of hydrazine reacts with $250 .$ mL of $0.146 \mathrm{M}$ $\mathrm{H}_{2} \mathrm{SO}_{4} ?$

Hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)$ and dinitrogen tetroxide $\left(\mathrm{N}_{2} \mathrm{O}_{4}\right)$ form a self-igniting mixture that has been used as a rocket propellant. The reaction products are $\mathrm{N}_{2}$ and $\mathrm{H}_{2} \mathrm{O}$ . (a) Write a balanced chemical equation for this reaction. (b) What is being oxidized, and what is being reduced? (c) Which substance serves as the reducing agent and which as the oxidizing agent?