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raketka [301]
1 year ago
6

Determine the limiting reactant and calculate the number of grams of sulfur dioxide, so2, that can be formed when 27.3 g of meth

ane thiol, ch3sh, reacts with 38.6 g of oxygen, o2.
Chemistry
1 answer:
lyudmila [28]1 year ago
6 0

When Methane thiol, CH₃SH, reacts with O₂, 25.8 grams of Sulfur dioxide, SO₂ is formed and limiting reactant is Oxygen, O₂.

The balanced chemical reaction is:

<em>CH₃SH(g) + 3 O₂(g) → 2 H₂O(g) + CO₂(g) + SO₂(g)</em>

<em />

             Molecular weights:       Actual/given mass:

CH₃SH = 48.11 g/mol                    27.3 g      

O₂         = 32.01 g/mol                    38.6 g  

H₂O      = 18.01 g/mol            

CO₂      = 44.01 g/mol                  

SO₂       = 64.06 g/mol

Number of moles = <u>    Given mass     </u>

                                  Molecular mass

So,

moles of methane thiol = 27.3 ÷ 48.11= 0.56 moles

moles of oxygen = 32.01÷38.6 = 0.82 moles

From the reaction,

1 mole of CH₃SH reacts with 3 moles of O₂  to give 1 mole of SO₂

Thus, 0.56 moles of CH₃SH reacts to form 1 × 0.56 = 0.56 moles of SO₂

Mass of SO₂ produced is 0.56 × 64.06 g = 35.86 g

moles of SO₂ = 35.86 ÷ 64.06 = 0.55 moles

So,

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Density is d= m/vol. If a material has a mass of 65.5 g and a volume of 32.5 ml, it has a density of
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Using the given formula, the density of the material is 2.015 g/mL

<h3>Calculating Density </h3>

From the question, we are to determine the density of the material

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Density = Mass / Volume

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Mass = 65.5 g

and volume = 32.5 mL

Putting the parameters into the equation,

Density = 65.5/32.5

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Ethyl butyrate, CH3CH2CH2CO2CH2CH3, is an artificial fruit flavor commonly used in the food industry for such flavors as orange
SIZIF [17.4K]

Answer:

A. 10.0 grams of ethyl butyrate would be synthesized.

B. 57.5% was the percent yield.

C. 7.80 grams of ethyl butyrate would be produced from 7.60 g of butanoic acid.

Explanation:

CH_3CH_2CH_2CO_2H(l)+CH_2CH_3OH(l)+H^+\rightarrow CH_3CH_2CH_2CO_2CH_2CH_3(l)+H_2O(l)

A

Moles of butanoic acid = \frac{7.60 g}{88 g/mol}=0.08636 mol

According to reaction ,1 mole of butanoic acid gives 1 mol of ethyl butyrate,then 0.08636 mol of butanoic acid will give :

\frac{1}{1}\times 0.08636 mol=0.08636 mol of ethyl butyrate

Mass of 0.08636 moles of ethyl butyrate =

0.08636 mol × 116 g/mol = 10.0 g

Theoretical yield = 10.0 g

Experimental yield = ?

Percentage yield of the reaction = 100%

Yield\%=\frac{\text{Experimental yield}}{\text{Theoretical yield}}\times 100

100\%=\frac{\text{Experimental yield}}{10.0 g}\times 100

Experimental yield = 10.0 g

10.0 grams of ethyl butyrate would be synthesized.

B

Theoretical yield of ethyl butyrate  = 10.0 g

Experimental yield ethyl butyrate = 5.75 g

Percentage yield of the reaction = ?

Yield\%=\frac{\text{Experimental yield}}{\text{Theoretical yield}}\times 100

=\frac{5.75 g}{10.0 g}\times 100=57.5\%

57.5% was the percent yield.

C

Moles of butanoic acid = \frac{7.60 g}{88 g/mol}=0.08636 mol

According to reaction ,1 mole of butanoic acid gives 1 mol of ethyl butyrate,then 0.08636 mol of butanoic acid will give :

\frac{1}{1}\times 0.08636 mol=0.08636 mol of ethyl butyrate

Mass of 0.08636 moles of ethyl butyrate =

0.08636 mol × 116 g/mol = 10.0 g

Theoretical yield = 10.0 g

Experimental yield = ?

Percentage yield of the reaction = 78.0%

Yield\%=\frac{\text{Experimental yield}}{\text{Theoretical yield}}\times 100

78.0\%=\frac{\text{Experimental yield}}{10.0 g}\times 100

Experimental yield = 7.80 g

7.80 grams of ethyl butyrate would be produced from 7.60 g of butanoic acid.

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