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Setler [38]
3 years ago
5

Consider the balanced chemical equation that follows. You are asked to determine how many moles of water you can form from 4 mol

es of hydrogen and excess oxygen. 2H2(g)+O2(g)→2H2O(l) Which of the following shows calculations for a correct way to solve this problem? View Available Hint(s) Consider the balanced chemical equation that follows. You are asked to determine how many moles of water you can form from 4 moles of hydrogen and excess oxygen. Which of the following shows calculations for a correct way to solve this problem? 4 mol H2×2 mol H2O2 mol H2=4 mol H2O 4 mol H2×2 mol H2O1 mol O2=8 mol H2O 2 mol H2×2 mol H22 mol H2O=2 mol H2O
Chemistry
1 answer:
kodGreya [7K]3 years ago
6 0

<u>Answer:</u> The correct answer is 4molH_2\times \frac{2molH_2O}{2molH_2}=4 molH_2O

<u>Explanation:</u>

We are given:

Moles of hydrogen gas = 4 moles

As, oxygen is given in excess. Thus, is considered as an excess reagent and hydrogen is considered as a limiting reagent because it limits the formation of products.

For the given chemical equation:

2H_2(g)+O_2(g)\rightarrow 2H_2O(l)

By Stoichiometry of the reaction:

2 moles of hydrogen produces 2 moles of water molecule.

So, 4 moles of hydrogen will produce = \frac{2molH_2O}{2molH_2}\times 4molH_2=4mol of water.

Hence, the correct answer is 4molH_2\times \frac{2molH_2O}{2molH_2}=4 molH_2O

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As the temperature of a liquid increases, its viscosity decreases.
5 0
2 years ago
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son4ous [18]

Answer: The molar enthalpy change is 73.04 kJ/mol

Explanation:

HCl+NaOH\rightarrow NaCl+H_2O

moles of HCl= molarity\times {\text {vol in L}}=0.415mol/L\times 0.1=0.0415mol

As NaOH is in excess 0.0415 moles of HCl reacts with 0.0415 moles of NaOH.

volume of water = 100.0 ml + 50.0 ml = 150.0 ml

density of water = 1.0 g/ml

mass of water = volume \times density=150.0ml\times 1.0g/ml=150.0g

q=m\times c\times \Delta T

q = heat released

m = mass  = 150.0 g

c = specific heat = 4.184J/g^0C

\Delta T = change in temperature = 4.83^0C

q=150.0\times 4.184\times 4.83

q=3031.3J

Thus 0.0415 mol of HCl produces heat = 3031.3 J

1 mol of HCL produces heat = \frac{3031.3}{0.0415}\times 1=73043.3J=73.04kJ

Thus molar enthalpy change is 73.04 kJ/mol

8 0
3 years ago
Consider 2H2 + O2 → 2H2O. To produce 1.2 g water, how many grams of H2 are required? Report to the correct number of significant
Elden [556K]

Answer:

0.133 mol (corrected to 3 sig.fig)

Explanation:

Take the atomic mass of H=1.0, and O=16.0,

no. of moles = mass / molar mass

so no. of moles of H2O produced = 1.2 / (1.0x2+16.0)

= 0.0666666 mol

From the equation, the mole ratio of H2:H2O = 2:2 = 1:1,

meaning every 1 mole of H2 reacted gives out 1 mole of water.

So, the no, of moles of H2 required should equal to the no, of moles of H2O produced, which is also  0.0666666 moles.

mass = no. of moles x molar mass

hence,

mass of H2 required = 0.066666666 x (1.0x2)

= 0.133 mol (corrected to 3 sig.fig)

3 0
3 years ago
A mole of oxygen and a mole of hydrogen (at STP) have all of the following in common EXCEPT
Drupady [299]

Answer:

Root mean squared velocity is different.

Explanation:

Hello!

In this case, since we have a mixture of oxygen and nitrogen at STP, which is defined as a condition whereas T = 298 K and P = 1 atm, we can infer that these gases have the same temperature, pressure, volume and moles but a different root mean squared velocity according to the following formula:

v_{rms}=\sqrt{\frac{3RT}{MM} }

Since they both have a different molar mass (MM), nitrogen (28.02 g/mol) and oxygen (32.02 g/mol), thus we infer that nitrogen would have a higher root mean squared velocity as its molar mass is less than that of oxygen.

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8 0
2 years ago
How do the equilibrium concentrations of the reactants compare to the equilibrium concentrations of the product?
Scorpion4ik [409]

Answer: It depends equilibrium constant K

Explanation:   You need to to have reaction formula.

If K >> 1 then concentrations of products are much bigger than

concentrations of reactants. If K < < 1, concentration of products is small.

6 0
3 years ago
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