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Chapter 5: Problem 19

(a) What is work? (b) How do we determine the amount of work done, given the force associated with the work?

Short Answer

Expert verified
(a) Work is a concept in physics that relates the force applied on an object to its displacement. It is done when a force causes an object to move in the direction of the force and is a scalar quantity measured in Joules (J). (b) To determine the amount of work done, given the force associated with the work, use the formula: \(W = Fd\cos(\theta)\), where \(W\) represents the work done, \(F\) represents the force applied on the object, \(d\) represents the displacement of the object, and \(\theta\) represents the angle between the direction of the force and the direction of the displacement.

Step by step solution

01

Understanding Work

Work is a concept in physics that relates the force applied on an object to the displacement of the object. In simple terms, work is done when a force is applied on an object, causing it to move in the direction of the force. Work is a scalar quantity, as it has only magnitude and no direction. The unit of work is Joule (J), in the International System of Units (SI).
02

Formula for Work Done

To calculate the work done on an object, we use the formula: \[ W = Fd\cos(\theta) \] where \(W\) is the work done, \(F\) is the force applied on the object, \(d\) is the displacement of the object, and \(\theta\) is the angle between the direction of the force and the direction of the displacement.
03

Determining Work Done

Based on the given information, follow these steps to determine the amount of work done: 1. Identify the force (\(F\)) acting on the object. 2. Find the displacement (\(d\)) of the object in the direction of the force. 3. Determine the angle (\(\theta\)) between the direction of the force and the direction of the displacement. 4. Use the formula \(W = Fd\cos(\theta)\) to calculate the work done. By following these steps and applying the formula, you can determine the amount of work done given the force associated with the work.

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Most popular questions from this chapter

The complete combustion of acetic acid, \(\mathrm{CH}_{3} \mathrm{COOH}(l)\), to form \(\mathrm{H}_{2} \mathrm{O}(l)\) and \(\mathrm{CO}_{2}(g)\) at constant pressure releases \(871.7 \mathrm{~kJ}\) of heat per mole of \(\mathrm{CH}_{3} \mathrm{COOH}\). (a) Write a balanced thermochemical equation for this reaction. (b) Draw an enthalpy diagram for the reaction.

For each of the following compounds, write a balanced thermochemical equation depicting the formation of one mole of the compound from its elements in their standard states and use Appendix \(C\) to obtain the value of \(\Delta H_{f}^{\circ}:\) (a) \(\mathrm{NH}_{3}(g)\), (b) \(\mathrm{SO}_{2}(g)\) (c) \(\mathrm{RbClO}_{3}(s)\) (d) \(\mathrm{NH}_{4} \mathrm{NO}_{3}(s)\).

Consider the following reaction, which occurs at room temperature and pressure: $$ 2 \mathrm{Cl}(g) \longrightarrow \mathrm{Cl}_{2}(g) \quad \Delta H=-243.4 \mathrm{~kJ} $$ Which has the higher enthalpy under these conditions, \(2 \mathrm{Cl}(g)\) or \(\mathrm{Cl}_{2}(g) ?\)

Consider two solutions, the first being \(50.0 \mathrm{~mL}\) of \(1.00 \mathrm{M}\) \(\mathrm{CuSO}_{4}\) and the second \(50.0 \mathrm{~mL}\) of \(2.00 \mathrm{M} \mathrm{KOH}\). When the two solutions are mixed in a constant-pressure calorimeter, a precipitate forms and the temperature of the mixture rises from \(21.5^{\circ} \mathrm{C}\) to \(27.7^{\circ} \mathrm{C}\). (a) Before mixing, how many grams of Cu are present in the solution of \(\mathrm{CuSO}_{4} ?(\mathrm{~b})\) Predict the identity of the precipitate in the reaction. (c) Write complete and net ionic equations for the reaction that occurs when the two solutions are mixed. (d) From the calorimetric data, calculate \(\Delta H\) for the reaction that occurs on mixing. Assume that the calorimeter absorbs only a negligible quantity of heat, that the total volume of the solution is \(100.0 \mathrm{~mL}\), and that the specific heat and density of the solution after mixing are the same as that of pure water.

At one time, a common means of forming small quantities of oxygen gas in the laboratory was to heat \(\mathrm{KClO}_{3}\) : \(2 \mathrm{KClO}_{3}(s) \rightarrow 2 \mathrm{KCl}(s)+3 \mathrm{O}_{2}(g) \quad \Delta H=-89.4 \mathrm{~kJ}\) For this reaction, calculate \(\Delta H\) for the formation of (a) \(0.632 \mathrm{~mol}\) of \(\mathrm{O}_{2}\) and (b) \(8.57 \mathrm{~g}\) of \(\mathrm{KCl}\). (c) The decomposition of \(\mathrm{KClO}_{3}\) proceeds spontaneously when it is heated. Do you think that the reverse reaction, the formation of \(\mathrm{KClO}_{3}\) from \(\mathrm{KCl}\) and \(\mathrm{O}_{2}\), is likely to be feasible under ordinary conditions? Explain your answer.
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