If 3.13 mol of an ideal gas has a pressure of 2.33 atm and a volume of 72.31 L, what is the temperature of the sample in degrees Celsius?

Answer:

E: the point where the acid and base have been added in proper stoichiometric amounts

Explanation:

Equivalence point in titration is simply the point where the amounts of acid and base used just sufficiently reacts chemically to cause neutralization whereas the endpoint is the point where the indicator of the titration changes colour.

The Equivalence point occurs before the endpoint.

Thus, option E is correct.

Answer:

1.5 × 10³ g

Explanation:

Step 1: Given and required data

Step 2: Calculate the temperature change

ΔT = 28.5°C - 22.0 °C = 6.5 °C

Step 3: Calculate the mass (m) of water

We will use the following expression.

Q = c × m × ΔT

m = Q / c × ΔT

m = 41,840 J / (4.184 J/g.°C) × 6.5 °C = 1.5 × 10³ g

Answer:

Following are the solution to the given question:

Explanation:

Each method through KHP is somewhat more precise since we have weighed that requisite quantity, we exactly know the KHP intensity appropriately. Its initial 6 M HCl concentration was never considered mandatory. They have probably prepared 6 M HCl solution although long ago and could have changed its concentration over even a period.

Answer:

I dont know

Explanation:

good luck

"0.0340" mol of CH₄ are in sealed flask.

Methane would also be a greenhouse gas, therefore its existence tends to affect humanity's surface temp as well as weather patterns framework; it is released into the atmosphere from such a wide assortment of life forms as well as biogenic.

According to the question,

Volume, V = 800 mL or, 0.800 L

Temperature, T = 22°C or, 295

Pressure, P = [tex]\frac{780}{760}[/tex] = 1.03 atm

As we know the relation,

The gram of moles will be will be:

→ n = [tex]\frac{PV}{RT}[/tex]

By substituting the values, we get

     = [tex]\frac{1.03\times 0.800}{0.08206\times 295}[/tex]

     = [tex]\frac{0.824}{242.077}[/tex]

     = 0.0340

Thus the response above is correct.

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Answer:

H2O2 reduces itself to H2O and also oxidizes to O2 simultaneously thereby acting both as an oxidizing and reducing agent .

Explanation:

When

H2O2 acts as an oxidizing agent

H2O2 + 2e- 2H+--->   2H2O

Reducing agent

H2O2 --> O2 + 2e + 2H+

H2O2 reduces itself to H2O and also oxidizes to O2 simultaneously thereby acting both as an oxidizing and reducing agent .

Answer:

Refraction

Explanation:

The 3.13 g of K would be needed to completely react with the remaining [tex]Cl_2[/tex].

To determine which reactant is limiting, we need to calculate the amount of product that can be formed from each reactant and compare them. The reactant that produces less product is the limiting reactant, since the reaction cannot proceed further once it is consumed.

The balanced chemical equation for the reaction between solid potassium and chlorine gas is:

2 K(s) + [tex]Cl_2[/tex](g) -> 2 KCl(s)

From the equation, we can see that 2 moles of K react with 1 mole of [tex]Cl_2[/tex] to form 2 moles of KCl.

First, we need to convert the masses of K and [tex]Cl_2[/tex] into moles:

moles of K = 15.20 g / 39.10 g/mol = 0.388 mol

moles of [tex]Cl_2[/tex] = 2.830 g / 70.90 g/mol = 0.040 mol

Now, we can use the mole ratio from the balanced equation to calculate the theoretical yield of KCl from each reactant:

Theoretical yield of KCl from K: 0.388 mol K x (2 mol KCl / 2 mol K) = 0.388 mol KCl

Theoretical yield of KCl from [tex]Cl_2[/tex]: 0.040 mol [tex]Cl_2[/tex] x (2 mol KCl / 1 mol [tex]Cl_2[/tex]) = 0.080 mol KCl

We can see that the theoretical yield of KCl from K is 0.388 mol, while the theoretical yield of KCl from [tex]Cl_2[/tex] is 0.080 mol. Therefore, the limiting reactant is [tex]Cl_2[/tex], since it produces less product.

To determine how much more of the limiting reactant would be needed to completely consume the excess reactant, we can use the stoichiometry of the balanced equation.

We know that 1 mole of [tex]Cl_2[/tex] reacts with 2 moles of K to produce 2 moles of KCl. Therefore, the amount of additional K needed to react with the remaining [tex]Cl_2[/tex] can be calculated as follows:

moles of K needed = 0.040 mol [tex]Cl_2[/tex] x (2 mol K / 1 mol [tex]Cl_2[/tex])

                                = 0.080 mol K

This means that 0.080 moles of K would be needed to completely consume the remaining [tex]Cl_2[/tex]. We can convert this to a mass by multiplying by the molar mass of K:

mass of K needed = 0.080 mol K x 39.10 g/mol

                              = 3.13 g K

Therefore, The 3.13 g of K would be needed to completely react with the remaining.

Example verification:

Suppose we had an additional 0.50 g of [tex]Cl_2[/tex] in the reaction. Would all of the K be consumed, or would there still be excess K?

Moles of additional [tex]Cl_2[/tex] = mass of [tex]Cl_2[/tex] / molar mass of [tex]Cl_2[/tex]

Moles of additional [tex]Cl_2[/tex] = 0.50 g / 70.90 g/mol

Moles of additional [tex]Cl_2[/tex] = 0.0070 mol

The theoretical yield of KCl that can be formed from the additional [tex]Cl_2[/tex] is:

0.0070 mol [tex]Cl_2[/tex] x (2 mol KCl / 1 mol [tex]Cl_2[/tex]) x (74.55 g KCl / 1 mol KCl) = 1.04 g KCl

Therefore, the total amount of KCl that can be formed from all of the [tex]Cl_2[/tex] is:

5.95 g + 1.04 g = 6.99 g

The amount of K that would be needed to completely consume all of the [tex]Cl_2[/tex].

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Answer:

316.227766

Explanation: