Tuesday
September 23, 2014

Homework Help: Chemistry

Posted by a Canadian on Sunday, March 3, 2013 at 6:05pm.

Can someone check my answers please?
I had specific trouble with 7, 13, and 27 (I'm pretty sure my other answers are correct; some reassurance would be nice though). And I have a small question about 8.

1. The ΔHf of an element in its standard state is defined to be

0 kJ/mol

4. Which substance has a standard enthalpy of formation, ΔHf, equal to zero?
a) gold, Au(s)
b) water H2O (l)
c) carbon monoxide, CO(s)
d) zinc, Zn(g)
e) water, H2O

I chose a) because Au is an element and (s) is its standard state.

5. Which of he following statements are true?
I. The reaction vessel cools when an endothermic reaction occurs.
II. An endothermic reaction has a negative value of ΔH.
III. Heat is liberated when an exothermic reaction occurs.

a) I and II
b) I, II, and III
c) I and III only
d) II and III only
e) none of them

I chose c).

6. Which of the following processes are exothermic?
I. boiling water
II. freezing water
III. condensing steam
IV. melting ice

I chose II and III because the reaction needs to release heat, i.e. become cooler, in order for water to freeze and for steam to condense.(?)

7. Which factor does not affect the rate of a chemical reaction in aqueous solution?
a) the enthalpy change of the reaction
b) the activation energy of the reaction
c) the collision frequency of the reacting particles
d) the relative orientation of the colliding particles
e) the temperature of the solution

What's the significance of it being an aqueous solution? How does that change things?
I THINK the answer is a)? Because even if the enthalpy change is super high or super low, the activation energy (which I believe DOES have an effect) won't be affected, and the activation energy is what matters, right...?

8. Which statement about an activated complex is true?
a) It is a stable substance.
b) It has a lower chemical potential energy, or enthalpy, than reactants or products.
c) It occurs only in endothermic reactions.
d) It occurs at the transition state of the reaction.
e) It always breaks down to form product molecules.

I chose d) because I'm sure of that, but why is e) false?

9. A catalyst changes the
I. mechanism of a reaction
II. enthalpy change of a reaction
III. activation energy of a reaction

I said I and III.

10. The overall rate of any chemical reaction is most closely related to
a) the number of steps in the reaction mechanism
b) the overall reaction
c) the fastest step in the reaction mechanism
d) the slowest step in the reaction mechanism
e) the average rate of all the steps in the reaction mechanism

I chose d).

13. In a chemical reaction, bonds are formed and broken.
a) How would you characterize the enthalpy change of bond breaking?

I'm not totally sure how to answer this question because isn't "enthalpy change" the difference in the enthalpies of the reactants and the products? So can the breaking of bonds have an "enthalpy change"...? Anyway, I just wrote:
Bond breaking results in energy absorption so its enthalpy change is positive?

b) How wold you characterize the enthalpy change of bond formation?
Again, I'm not sure how to answer, but:
Bond formation results in energy being released, so its enthalpy change is negative.

c) State the relationship between the enthalpy change of the overall reaction (exothermic and endothermic) and bond breakage and formation.

The enthalpy change of the overall reaction depends on whether more energy is released from bonds forming or if more energy is absorbed from bonds breaking. If more is absorbed, then it's endothermic; if more is released, then it's exothermic.

19. C4H10(g) + 6.5 O2(g) -> 4CO2(g) + 5H2O(l)
a) Write a separate balanced chemical equation for the formation of C4H10, CO2, and H2O, directly from the elements in their standard states.
4C(s) + 5H2(g) -> C4H10(g)
C(s) +O2(g -> CO2(g)
H2(g) + 1/2 O2(g) -> H2O(l)
b) Algebraically combine these equations to get the balanced chemical equation for the complete combustion of C4H10.
I reversed 4C(s) + 5H2(g) -> C4H10(g),
multiplied C(s) +O2(g -> CO2(g) by 4,
and multiplied H2(g) + 1/2 O2(g) -> H2O(l) by 5. (then added the equations)

27. A student dissolves 1.96 g of NaOH in 100.0 mL water in a coffee-cup calorimeter. The initial temperature of the water is 23.4 C. After the NaOH dissolves, the temperature of the water rises to 28.7 C.
a) Determine the enthalpy of dissolution of sodium hydroxide, in kJ/mol NaOH. Assume the heat capacity of the calorimeter is negligible.

Q = mc(ΔT)
= (100g)(4.184 J/gC)(28.7 C - 23.4 C)
=2217.52 J
=2.21752 kJ

n=m/MM
= 1.96 NaOH/(23+16+1.01)
= 0.049 mol

ΔH = -Q/n
= -2.21752 kJ/0.049 mol
= -45.3 kJ/mol

For this one, the book actually gives the answer: 47 kJ/mol.
Firstly, I got a negative (book's typo, or my error?). Secondly, I'm about 2 off, so I think I might've missed something in my calculations? I know 2 is a small difference, but it could mean I missed something important... (That's why i tried it again carrying all the decimals... Still got 45 though.)

b) Suppose that the heat capacity of the calorimeter was not negligible. Explain how the value of ΔH that you calculated in part (a) would compare with the actual ΔH.

I really have no idea.

Would the value of ΔH, calculated when the ΔH of the calorimeter was assumed to be negligible, be higher or lower than the actual ΔH value?
And would it be because the amount of heat absorbed by the calorimeter was not taken into account, or because the amount of heat LOST through the calorimeter was not taken into account?
If you somehow included the heat capacity of the calorimeter into the calculations, how would the result change?

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