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Mrrafil [7]
2 years ago
13

How many grams of water will be produced when 1.6 moles of ethanol (CH3CH2OH) are burned completely? Enter a number only (no uni

ts) and express your answer using the correct number of significant figures.
Chemistry
1 answer:
AfilCa [17]2 years ago
4 0

Answer:

86

Explanation:

The reaction that takes place is:

  • C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O

First we <u>convert moles of ethanol to moles of water</u>:

  • 1.6 mol ethanol * \frac{3molH_2O}{1molEthanol} = 4.8 mol H₂O

Then we <u>convert moles of water to grams of water</u>, using its molar mass:

  • 4.8 mol H₂O * 18 g/mol = 86.4 g

So 84.6 grams of water will be produced.

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Question 3) A 1.00 L buffer solution is 0.250 M in HF and 0.250 M in LiF. Calculate the pH of the solution after the addition of
Masja [62]

The pH of the solution after adding 0.150 moles of solid LiF is 3.84

<u>Explanation:</u>

We have the chemical equation,

HF (aq)+NaOH(aq)->NaF(aq)+H2O

To find how many moles have been used in this

c= n/V=> n= c.V

nHF=0.250 M⋅1.5 L=0.375 moles HF

Simillarly

nF=0.250 M⋅1.5 L=0.375 moles F

nHF=0.375 moles - 0.250 moles=0.125 moles

nF=0.375 moles+0.250 moles=0.625 moles

[HF]=0.125 moles/1.5 L=0.0834 M

[F−]=0.625 moles/1.5 L=0.4167 M

To determine the problem using the Henderson - Hasselbalch equation

pH=pKa+log ([conjugate base/[weak acid])

Find the value of Ka

pKa=−log(Ka)

pH=−log(Ka) +log([F−]/[HF]

pH= -log(3.5 x 10 ^4)+log(0.4167 M/0.0834 M)

pH=-log(3.5 x 10 ^4)+log(4.996)

pH= -4.54+0.698

pH=-(-3.84)

pH=3.84

The pH of the solution after adding 0.150 moles of solid LiF is 3.84

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A reactant decomposes with a half-life of 29.5 s when its initial concentration is 0.229 M. When the initial concentration is 0.
Dmitrij [34]

Answer :

The order of reaction is, 0 (zero order reaction).

The value of rate constant is, 0.00388Ms^{-1}

Explanation :

Half life : It is defined as the time in which the concentration of a reactant is reduced to half of its original value.

The general expression of half-life for nth order is:

t_{1/2}\propto \frac{1}{[A_o]^{n-1}}

or,

\frac{t_{1/2}_1}{t_{1/2}_2}=\frac{[A_2]^{n-1}}{[A_1]^{n-1}}

or,

n=\left(\frac{\log\frac{(t_{1/2})_1}{(t_{1/2})_2}}{\log\frac{(A)_2}{(A)_1}}\right )+1           .............(1)

where,

t_{1/2} = half-life of the reaction

n = order of reaction

[A] = concentration

As we are given:

Initial concentration of A = 0.229 M

Final concentration of A = 0.639 M

Initial half-life of the reaction = 29.5 s

Final half-life of the reaction = 82.3 s

Now put all the given values in the above formula 1, we get:

n=\left (\frac{\log \frac{29.5}{82.3}}{\log\frac{0.639}{0.229}}\right )+1

n=0.000196\approx 0

Thus, the order of reaction is, 0 (zero order reaction).

Now we have to determine the rate constant.

To calculate the rate constant for zero order the expression will be:

t_{1/2}=\frac{[A_o]}{2k}

When,

t_{1/2} = 29.5 s

[A_o] = 0.229 M

29.5s=\frac{0.229M}{2k}

k=0.00388Ms^{-1}

Thus, the value of rate constant is, 0.00388Ms^{-1}

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