Some Basic Concepts of Chemistry

To find the number of moles of methane ($CH_4$) required to produce 11g of $CO_2$ upon complete combustion, we first need the balanced chemical equation for the combustion of methane:

$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

From the balanced equation, we see that 1 mole of $CH_4$ produces 1 mole of $CO_2$. Therefore, the moles of $CH_4$ needed to produce a certain amount of $CO_2$ will be equal to the moles of $CO_2$ produced.

To find the moles of $CO_2$ produced from 11g of $CO_2$, we use the formula:

$\text{Moles} = \frac{\text{Mass}}{\text{Molar mass}}$

The molar mass of $CO_2$ is $44 \mathrm{~g/mol}$ (12 from carbon and 16$\times$2 from oxygen). Then, the moles of $CO_2$ produced from 11g of $CO_2$ can be calculated as:

$\text{Moles of } CO_2 = \frac{11 \mathrm{~g}}{44 \mathrm{g/mol}} = 0.25 \mathrm{~mol}$

Since the stoichiometry of the reaction between methane and oxygen is 1:1 for $CH_4$ and $CO_2$, the moles of $CH_4$ required to produce 11g of $CO_2$ is also 0.25 moles.

Therefore, the correct answer is Option B: 0.25.

To determine the molecular formula of the compound, we first calculate the empirical formula based on the given percentages of carbon, hydrogen, and oxygen. Then we use the molecular weight to find the molecular formula.

Step 1: Calculating the moles of each element per 100 g of the compound

• For carbon (%C = 42.1%), the moles of carbon are calculated as follows:

  $\text{Moles of C} = \frac{\% \text{ by weight of C}}{\text{Atomic weight of C}} = \frac{42.1 \text{ g}}{12.01 \text{ g/mol}}$

• For hydrogen (%H = 6.4%), the moles of hydrogen are:

  $\text{Moles of H} = \frac{\% \text{ by weight of H}}{\text{Atomic weight of H}} = \frac{6.4 \text{ g}}{1.008 \text{ g/mol}}$

• For oxygen, since it makes up the remainder, we find its percentage by subtracting the percentages of carbon and hydrogen from 100%:

  $\% \text{O} = 100 - 42.1 - 6.4 = 51.5\%$

  Then the moles of oxygen are:

  $\text{Moles of O} = \frac{\% \text{ by weight of O}}{\text{Atomic weight of O}} = \frac{51.5 \text{ g}}{16.00 \text{ g/mol}}$

Step 2: Dividing the calculated moles by the smallest number of moles to find the simplest whole number ratio:

This involves dividing each mole value by the smallest of the mole values obtained. However, since we didn't calculate the exact number of moles for each element in the above step, let's proceed conceptually based on given percentages:

To find the simplest whole number ratio, we would perform the following calculations:

$\text{Moles of C} \approx \frac{42.1}{12}$

$\text{Moles of H} \approx \frac{6.4}{1}$

$\text{Moles of O} \approx \frac{51.5}{16}$

Step 3: Determining the empirical formula:

Given the lack of exact numbers from the prior step, let's assume each mole value is divided by the smallest mole value among the C, H, and O calculations. This ratio generally defines the empirical formula, which constitutes the smallest whole number ratio of atoms in the compound.

Step 4: Using the molecular weight to find the molecular formula:

Once we have the empirical formula, we would calculate its mass and compare it to the molecular weight provided (342). The molecular formula is a multiple of the empirical formula that matches the given molecular weight. Since we haven't calculated the exact empirical formula mass, let's examine the options directly considering a molecular weight of 342:

Option A: $C_{14}H_{20}O_{10}$ - The molar mass would be $14 \times 12 + 20 \times 1 + 10 \times 16 = 168 + 20 + 160 = 348 \text{, which is close but not 342.}$

Option B: $C_{12}H_{22}O_{11}$ - The molar mass would be $12 \times 12 + 22 \times 1 + 11 \times 16 = 144 + 22 + 176 = 342 \text{, which matches the given molecular weight.}$

Thus, the correct molecular formula based on the molecular weight and composition given would be Option B: $\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}$.

To find the molarity of the aqueous solution, we need to calculate the number of moles of NaCl and then divide it by the volume of the solution in liters.

First, determine the molar mass of NaCl:

$ \text{Molar mass of NaCl} = \text{Molar mass of Na} + \text{Molar mass of Cl} = 23 + 35.5 = 58.5 \, \text{g/mol} $

Next, calculate the moles of NaCl:

$ \text{Moles of NaCl} = \frac{5.85 \, \text{g}}{58.5 \, \text{g/mol}} = 0.1 \, \text{mol} $

Now, convert the volume of the solution from milliliters to liters:

$ 500 \, \text{mL} = 0.500 \, \text{L} $

Finally, calculate the molarity (M), which is the number of moles of solute per liter of solution:

$ M = \frac{\text{Moles of NaCl}}{\text{Volume of solution in liters}} = \frac{0.1 \, \text{mol}}{0.500 \, \text{L}} = 0.2 \, \text{M} $

Therefore, the molarity of the solution is 0.2 M. The correct option is D.

$\begin{aligned} & \mathrm{CaCO}_3(\mathrm{s}) \xrightarrow{\Delta} \mathrm{CaO}(\mathrm{s})+\mathrm{CO}_2(\mathrm{g}) \\ & \mathrm{MgCO}_3(\mathrm{s}) \xrightarrow{\Delta} \mathrm{MgO}(\mathrm{s})+\mathrm{CO}_2(\mathrm{g}) \end{aligned}$

Let the weight of $\mathrm{CaCO}_3$ be $\mathrm{x}$ gm

$\therefore$ weight of $\mathrm{MgCO}_3=(2.21-\mathrm{x}) \mathrm{~gm}$

Moles of $\mathrm{CaCO}_3$ decomposed $=$ moles of $\mathrm{CaO}$ formed

$\frac{\mathrm{x}}{100}=$ moles of $\mathrm{CaO}$ formed

$\therefore$ weight of $\mathrm{CaO}$ formed $=\frac{\mathrm{x}}{100} \times 56$

Moles of $\mathrm{MgCO}_3$ decomposed $=$ moles of $\mathrm{MgO}$ formed

$\frac{(2.21-x)}{84}=$ moles of $\mathrm{MgO}$ formed

$\therefore$ weight of $\mathrm{MgO}$ formed $=\frac{2.21-\mathrm{x}}{84} \times 40$

$\Rightarrow \frac{2.21-\mathrm{x}}{84} \times 40+\frac{\mathrm{x}}{100} \times 56=1.152$

$\therefore \mathrm{x}=1.1886 \mathrm{~g}=$ weight of $\mathrm{CaCO}_3$

& weight of $\mathrm{MgCO}_3=1.0214 \mathrm{~g}$

Mole fraction of $\mathrm{C=\frac{n_C}{n_A+n_B+n_C}}$

The quantity that changes with temperature from the given options is Molarity.

Molarity, which is denoted with a capital M, measures the number of moles of solute dissolved in one liter of solution. It is expressed as:

$ Molarity (M) = \frac{moles \ of \ solute}{liters \ of \ solution} $

Since the volume of liquids expands or contracts with temperature, the volume of the solution will change as the temperature changes, which in turn will affect the molarity. For example, if the temperature of a solution increases, its volume may expand. If the number of moles of solute remains constant, the increased volume will result in a decreased molarity because the solute is now dispersed in a greater volume of solution.

In contrast, other measures of concentration such as:

Molality, which is measured as the number of moles of solute per kilogram of solvent (not affected by temperature since it's based on mass, not volume),

$ Molality (m) = \frac{moles \ of \ solute}{kilograms \ of \ solvent} $

Mole fraction, which is the ratio of the number of moles of one component to the total number of moles of all components in the solution (also not affected by temperature since the ratio of moles doesn't change with temperature),

$ Mole \ fraction \ (X) = \frac{moles \ of \ component}{total \ moles \ in \ mixture} $

and Mass percentage, which is the percentage by mass of a component in a mixture (again, not affected by temperature because it's based on mass),

$ Mass \ percentage \ $(%) $= \frac{mass \ of \ component}{total \ mass \ of \ mixture} \times 100\% $

do not change with temperature. Therefore, the correct answer to the question is:

Option B : Molarity

Statements given in (I) and (III) are correct, while statement given in (ii) is incorrect. The correct from is :

Manufacture of ammonia by Haber process is an exothermic reaction.

Determine Oxygen Mass:

The metal's mass is $ 10 \, \text{g} $. The resulting oxide weighs $ 11.6 \, \text{g} $. Thus, the mass of oxygen can be calculated as:

$ \text{Mass of oxygen} = 11.6 \, \text{g} - 10 \, \text{g} = 1.6 \, \text{g} $

Equivalent Mass of Oxygen:

It is given that the equivalent mass of oxygen is $ 8 \, \text{g} $.

Calculate Equivalent Mass of Metal (M):

Using the relationship for equivalent mass, which involves the mass of the substance and the mass of the reactive component (in this case, oxygen), we can set up the following equation:

$ \frac{\text{Equivalent mass of metal}}{\text{Equivalent mass of oxygen}} = \frac{\text{Mass of metal}}{\text{Mass of oxygen}} $

Plugging in the values:

$ \frac{x}{8} = \frac{10}{1.6} $

Solving for $ x $:

$ x = \frac{80}{1.6} = 50 $

Thus, the equivalent weight of the metal $ M $ is $ 50 $.

To find the remaining moles of $\mathrm{H}_2\mathrm{SO}_4$ in the flask, follow these steps:

Step 1: Calculate the Molar Mass

The molar mass of $\mathrm{H}_2\mathrm{SO}_4$ is determined by adding the atomic masses of its constituent elements:

Hydrogen ($H$): $2 \times 1 = 2$

Sulfur ($S$): $32$

Oxygen ($O$): $4 \times 16 = 64$

Thus, the total molar mass of $\mathrm{H}_2\mathrm{SO}_4$ is:

$ 2 + 32 + 64 = 98 \, \text{g/mol} $

Step 2: Convert Mass from mg to g

Given:

98 mg of $\mathrm{H}_2\mathrm{SO}_4$.

Since 1 g = 1000 mg, convert milligrams to grams:

$ 98 \, \text{mg} = \frac{98}{1000} \, \text{g} = 0.098 \, \text{g} $

Step 3: Determine Initial Molecules

The number of molecules in 98 mg of $\mathrm{H}_2\mathrm{SO}_4$ can be found using Avogadro’s number ($N_A = 6.02 \times 10^{23}$ molecules/mol):

98 grams of $\mathrm{H}_2\mathrm{SO}_4$ contains $6.02 \times 10^{23}$ molecules. Therefore, 0.098 grams will have:

$ \frac{6.02 \times 10^{23}}{1000} = 6.02 \times 10^{20} \text{ molecules} $

Step 4: Calculate Remaining Molecules

After removing $3.01 \times 10^{20}$ molecules, the remaining number of molecules is:

$ 6.02 \times 10^{20} - 3.01 \times 10^{20} = 3.01 \times 10^{20} \text{ molecules} $

Step 5: Convert Remaining Molecules to Moles

Using the proportion of molecules to moles:

$ 6.02 \times 10^{23} \text{ molecules} = 1 \text{ mol} $

Consequently, $3.01 \times 10^{20}$ molecules are:

$ \frac{3.01 \times 10^{20}}{6.02 \times 10^{23}} = 5 \times 10^{-4} \text{ moles} $

Thus, the remaining $\mathrm{H}_2\mathrm{SO}_4$ in the flask is $5 \times 10^{-4}$ moles.

The decomposition reaction is,

$ \underset{2 \text { moles }}{2 \mathrm{MMOO}_4} \longrightarrow \underset{1 \text { moles }}{\mathrm{K}_2 \mathrm{MnO}_4}+\underset{1 \text { moles }}{\mathrm{MnO}_2}+\underset{1 \text { moles }}{\mathrm{O}_2 \uparrow} $

For 0.1 moles of $\mathrm{KMnO}_4$

$ =0.1 / 2=0.05 \mathrm{moles} $

(same for $\mathrm{K}_2 \mathrm{MnO}_4$ and $\mathrm{MnO}_2$ )

Molar mass of products,

$ \begin{gathered} \mathrm{K}_2 \mathrm{MnO}_4=197 \mathrm{~g} / \mathrm{mol} \\ \mathrm{MnO}_2=87 \mathrm{~g} / \mathrm{mol} \end{gathered} $

Weight of $\mathrm{K}_2 \mathrm{MnO}_4=0.05 \times 197$

$ =9.85 \mathrm{~g} $

Weight of $\mathrm{MnO}_2=0.05 \times 87$

$ =4.35 \mathrm{~g} $

Weight of residue $=$ Weight of $\mathrm{K}_2 \mathrm{MnO}_4$ and $\mathrm{MnO}_2$

$ 9.85+4.35=14.2 \mathrm{~g} $

To determine the normality of the resulting solution with respect to $\mathrm{NO}_3^{-}$, we begin by examining the individual solutions and their normalities:

For $\mathrm{Ca(NO_3)_2}$:

Molarity ($M$): $\frac{M}{10}$

$n$-factor: 2 (because $\mathrm{Ca(NO_3)_2}$ provides two $\mathrm{NO}_3^{-}$ ions per formula unit)

Normality ($N$): Calculated as $N = M \times n\text{-factor} = \frac{1}{10} \times 2 = 0.2 \, \text{N}$

For $\mathrm{KNO_3}$:

Molarity ($M$): $\frac{M}{10}$

$n$-factor: 1 (because $\mathrm{KNO_3}$ provides one $\mathrm{NO}_3^{-}$ ion per formula unit)

Normality ($N$): Calculated as $N = M \times n\text{-factor} = \frac{1}{10} \times 1 = 0.1 \, \text{N}$

To find the normality of the mixed solution with respect to $\mathrm{NO}_3^{-}$, use the formula for mixing solutions:

$ N_{\text{result}} = \frac{V_1 N_1 + V_2 N_2}{V_1 + V_2} $

$V_1 = 100 \, \text{mL}$, $N_1 = 0.2 \, \text{N}$ for $\mathrm{Ca(NO_3)_2}$

$V_2 = 200 \, \text{mL}$, $N_2 = 0.1 \, \text{N}$ for $\mathrm{KNO_3}$

Plugging in the values:

$ N_{\text{result}} = \frac{100 \times 0.2 + 200 \times 0.1}{100 + 200} = \frac{20 + 20}{300} = \frac{40}{300} = 0.13 \, \text{N} $

Thus, the normality of the resulting solution with respect to $\mathrm{NO}_3^{-}$ is 0.13 N.

To determine the value of 'x', we need to analyze two reactions involving metal hydrogen carbonate ($ \text{MHCO}_3 $).

Reaction I: Decomposition of $ \text{MHCO}_3 $

$ 2 \, \text{MHCO}_3 \xrightarrow{\Delta} \text{CO}_2 + \text{M}_2\text{CO}_3 + \text{H}_2\text{O} $

In this reaction, 2 moles of $ \text{MHCO}_3 $ decompose to give 1 mole of $ \text{CO}_2 $.

For $ x \, \text{g} $ of $ \text{MHCO}_3 $, the number of moles is $ \frac{x}{84} $.

This results in $ \frac{x}{2 \times 84} $ moles of $ \text{CO}_2 $.

Reaction II: Reaction of $ \text{CO}_2 $ with $ \text{MOH} $

$ \text{CO}_2 + \text{MOH} \longrightarrow \text{MHCO}_3 $

In this reaction, 1 mole of $ \text{CO}_2 $ reacts with 1 mole of $ \text{MOH} $ to form 1 mole of $ \text{MHCO}_3 $.

From Reaction I, we have $ \frac{x}{2 \times 84} $ moles of $ \text{CO}_2 $.

According to the problem, this amount of $ \text{CO}_2 $ completely reacts with 0.2 moles of $ \text{MOH} $.

Set the moles of $ \text{CO}_2 $, $ \frac{x}{2 \times 84} $, equal to 0.2:

$ \frac{x}{2 \times 84} = 0.2 $

Solving for $ x $:

$ x = 0.2 \times 2 \times 84 = 33.6 $

Thus, the value of $ x $ is $ 33.6 $.

To determine the concentration of a 6% urea solution with a total weight of 1000 g, follow these steps:

Calculate the amount of urea:

The solution contains 6% urea by mass. Thus, the mass of urea is:

$ \text{Urea} = 1000 \, \text{g} \times \frac{6}{100} = 60 \, \text{g} $

The molar mass of urea (CH₄N₂O) is calculated using the atomic masses, which are:

Carbon (C): 12 u

Nitrogen (N): 14 u

Oxygen (O): 16 u

Hydrogen (H): 1 u

Therefore, the molar mass of urea is:

$ \text{Molar mass of urea} = 12 + (4 \times 1) + (2 \times 14) + 16 = 60 \, \text{g/mol} $

Convert the mass of urea to moles:

$ \text{Moles of urea} = \frac{60 \, \text{g}}{60 \, \text{g/mol}} = 1 \, \text{mol} $

Calculate the volume of water:

The remaining mass of the solution is water:

$ \text{Water} = 1000 \, \text{g} - 60 \, \text{g} = 940 \, \text{g} $

Using the density of water (1 g/mL), convert this mass to volume:

$ \text{Volume of water} = \frac{940 \, \text{g}}{1 \, \text{g/mL}} = 940 \, \text{mL} = 0.94 \, \text{L} $

Calculate the concentration of the urea solution:

The concentration in molarity ($ \text{mol/L} $) is given by:

$ \text{Concentration} = \frac{\text{Moles of urea}}{\text{Volume of solution in liters}} = \frac{1 \, \text{mol}}{0.94 \, \text{L}} = 1.0638 \, \text{M} $

Rounding to three significant figures gives:

$ \text{Concentration} \approx 1.064 \, \text{mol/L} $

Therefore, the concentration of the urea solution is approximately $1.064 \, \text{mol/L}$.

Given reaction is

$ \underset{\substack{\text { 1 mole }}}{\mathrm{Xe}(g)}+\underset{5 \text { moles }}{2 \mathrm{~F}_2(g)} \xrightarrow[7 \text { bar }]{873 \mathrm{~K}} \mathrm{XeF}_4(s) $

The ratio of $\mathrm{Xe}: \mathrm{Fe}_2$ is $1: 5$.

To find the concentration of nitric acid in mol/L, we'll use the provided data about its density and weight percentage.

Given:

Density of the solution = $1.5 \, \text{g/mL}$

Weight percentage of nitric acid = 68%

Calculate the mass of the solution in 1L:

Since the density is given as $1.5 \, \text{g/mL}$, the mass of 1 liter (1000 mL) of the solution would be:

$ \text{Mass of solution} = 1.5 \, \text{g/mL} \times 1000 \, \text{mL} = 1500 \, \text{g} $

Determine the mass of nitric acid in the solution:

The weight percentage indicates that 68% of the solution's mass is nitric acid. Therefore, the mass of nitric acid is:

$ \text{Mass of } \text{HNO}_3 = \frac{68}{100} \times 1500 \, \text{g} = 1020 \, \text{g} $

Calculate the molar mass of nitric acid ($\text{HNO}_3$):

Using the atomic masses provided:

$ \text{Molar mass of } \text{HNO}_3 = 1 \, (\text{H}) + 14 \, (\text{N}) + 16 \times 3 \, (\text{O}) = 63 \, \text{g/mol} $

Find the concentration in mol/L:

To find the number of moles of $\text{HNO}_3$ in 1L:

$ \text{Number of moles} = \frac{\text{Mass of } \text{HNO}_3}{\text{Molar mass } \text{HNO}_3} = \frac{1020 \, \text{g}}{63 \, \text{g/mol}} = 16.20 \, \text{mol/L} $

Therefore, the concentration of nitric acid is approximately $16.20 \, \text{mol/L}$.

Carbon $=92.3 / 12=7.7$

Hydrogen $=7.7 / 1=7.7$

Ratio $\mathrm{C}: \mathrm{H}=1$ : 1

So, emperical formula $=\mathrm{CH}$

In question, 39 g of hydrocarbon is given, it means 3 moies of CH is needed.

Also, $2 \mathrm{Na}+2 \mathrm{H}_2 \mathrm{O} \longrightarrow \mathbf{2 N a O H}+\mathbf{H}_2$

$2 \mathrm{H}_2 \mathrm{O}$ gives $1 \mathrm{~mol} \mathrm{H}_2$

so, $\frac{3}{2} \mathrm{~mol}$ of $\mathrm{H}_2 \mathrm{O}$ gives $\frac{1}{2} \times 1.5 \mathrm{~mol}$ of $\mathrm{H}_2$

$ =0.75 \mathrm{~mol} \mathrm{H}_2 $

So, all $\mathrm{H}_2 \mathrm{O}$ is consumed to liberate 0.75 mole of $\mathrm{H}_2$.

Thus, weight of $\mathrm{O}_2$ consumed

$ =\frac{15}{4} \times 32=120 \mathrm{~g} $

To determine the concentration of a $ \mathrm{CaCO}_3 $ solution in molarity (mol/L), given its concentration is 1000 ppm, we proceed as follows:

Mass of $\mathrm{CaCO}_3$ in 1 Liter:

Since 1000 ppm means 1000 mg of $\mathrm{CaCO}_3$ per liter, it converts to grams as:

$ \frac{1000 \times 1000 \, \text{mg}}{1000000} = 1 \, \text{g} $

Calculation of Moles:

The molar mass of $\mathrm{CaCO}_3$ is calculated using the atomic masses:

$\mathrm{Ca} = 40 \, \text{u}, \mathrm{C} = 12 \, \text{u}, \mathrm{O} = 16 \, \text{u}$.

Therefore, the molar mass of $\mathrm{CaCO}_3 = 40 + 12 + (16 \times 3) = 100 \, \text{g/mol}$.

Now, calculate the number of moles ($n$):

$ n = \frac{\text{Mass of } \mathrm{CaCO}_3}{\text{Molar mass}} = \frac{1}{100} = 0.01 \, \text{mol} $

Determine Molarity:

The molarity (concentration in mol/L) is given by:

$ \text{Molarity} = \frac{\text{Number of moles}}{\text{Volume of solution in L}} = \frac{0.01}{1} = 0.01 \, \text{mol/L} = 10^{-2} \, \text{mol/L} $

Thus, the concentration of the $\mathrm{CaCO}_3$ solution is $10^{-2} \, \text{mol/L}$.

The weight percent of chlorine in the compound is given as 55.0%. This implies that the weight percent of the metal is 45.0%.

Now, we need to find the molar mass of the compound. At STP (standard temperature and pressure), 1 mole of any gas occupies 22.4 L. We know that 100 mL (or 0.1 L) of the metal chloride vapor weighs 0.57 g. Therefore, the molar mass of the compound (MClx) is:

$M_{MCl_x} = \frac{0.57 \, \text{g}}{0.1 \, \text{L}} \times 22.4 \, \text{L/mol} = 127.68 \, \text{g/mol}$

Given the weight percent of the metal and chlorine, we can find the molar mass of the metal (MM) using the equation:

$MM = M_{MCl_x} \times \frac{45.0}{55.0} = 104.04 \, \text{g/mol}$

Subtracting the atomic mass of chlorine from the molar mass of the metal gives the molar mass of the metal:

$MM - x \cdot 35.5 = 104.04 \, \text{g/mol}$

Where 'x' is the number of chlorine atoms in the compound.

Solving this equation for 'x' gives:

$x = \frac{104.04 - MM}{35.5}$

We find that 'x' is approximately 2. Therefore, the molecular formula of the compound is MCl₂.

So, the correct answer is MCl₂.

Firstly, let's calculate the molarity of the oxalic acid solution given in the Assertion (A).

The molarity (M) is defined as the number of moles of solute per liter of solution. The number of moles of solute is given by the mass of the solute divided by its molar mass.

So, the number of moles of hydrated oxalic acid is $\frac{3.1500 \, \text{g}}{126 \, \text{g/mol}} = 0.025 \, \text{mol}$

And the volume of the solution is 250.0 mL, or 0.250 L.

Therefore, the molarity of the solution is $\frac{0.025 \, \text{mol}}{0.250 \, \text{L}} = 0.1 \, \text{M}$

This matches the molarity given in the Assertion (A), so the Assertion is true.

The Reason (R) gives the molar mass of hydrated oxalic acid as 126 g/mol, which was used in the calculation above. Therefore, the Reason is also true.

Moreover, the Reason (R) is the correct explanation of Assertion (A), as it provides the necessary information (the molar mass of hydrated oxalic acid) to calculate the molarity of the solution.

So, the correct option is C: Both A and R are true and R is the correct explanation of A.

A. 16 g of $\mathrm{CH_4~(g)}$

The molar mass of $\mathrm{CH_4}$ (Methane) is 16 g/mol. Therefore, 16 g of $\mathrm{CH_4}$ is equivalent to 1 mole of $\mathrm{CH_4}$. Furthermore, each molecule of $\mathrm{CH_4}$ has 10 electrons (6 from Carbon and 4 from Hydrogen). As a result, 1 mole of $\mathrm{CH_4}$ (or $6.022 \times 10^{23}$ molecules of $\mathrm{CH_4}$) would have $10 \times 6.022 \times 10^{23} = 60.2 \times 10^{23}$ electrons. This matches with (II).

B. 1 g of $\mathrm{H_2~(g)}$

The molar mass of $\mathrm{H_2}$ (Hydrogen) is 2 g/mol. Therefore, 1 g of $\mathrm{H_2}$ is equivalent to 0.5 moles of $\mathrm{H_2}$. The volume that a given quantity of gas occupies is proportional to the number of moles of gas. At standard temperature and pressure (STP), 1 mole of any ideal gas occupies 22.4 liters. Therefore, 0.5 moles of gas would occupy $0.5 \times 22.4 = 11.2$ liters. The closest match is (IV) with 11.4 liters (the small discrepancy may be due to rounding or slightly different conditions than STP).

C. 1 mole of $\mathrm{N_2~(g)}$

The molar mass of $\mathrm{N_2}$ (Nitrogen) is 28 g/mol. Therefore, 1 mole of $\mathrm{N_2}$ weighs 28 g. This matches with (I).

D. 0.5 mol of $\mathrm{SO_2~(g)}$

The molar mass of $\mathrm{SO_2}$ (Sulfur Dioxide) is 64 g/mol. Therefore, 0.5 moles of $\mathrm{SO_2}$ weigh $0.5 \times 64 = 32$ g. This matches with (III).

Thus, the correct matches are:

A - II

B - IV

C - I

D - III

At STP, one mole of any gas occupies 22.4 liters. Therefore, 2.8375 liters of oxygen gas at STP is equal to 0.125 moles. The number of molecules of oxygen gas in 0.125 moles is calculated by multiplying the number of moles by Avogadro's number, which is $6.022 × 10^{23} $ molecules/mol. This gives us a total of $7.527 × 10^{22}$ molecules of oxygen gas in 2.8375 liters at STP.

Here is the calculation:

Significant figures are the digits in a number that carry meaningful information about its precision. In other words, they are the numbers that are known with certainty.

Here is how we calculate significant figures for each number:

• A. 0.00253: There are 3 significant figures here (253). Leading zeroes do not count as significant figures.




• B. 1.0003: There are 5 significant figures here. All non-zero digits are significant, and zeroes between non-zero digits or at the end of the number and after the decimal point are significant.




• C. 15.0: There are 3 significant figures here. The zero after the decimal point counts as it indicates the precision of the measurement.




• D. 163: There are 3 significant figures here. All non-zero digits are significant.

So, comparing these, we see that options A (0.00253), C (15.0), and D (163) each have 3 significant figures.

The reaction between $\mathrm{HBr}$ and $\mathrm{Ba(OH)_2}$ is a neutralization reaction where one $\mathrm{Ba(OH)_2}$ reacts with two $\mathrm{HBr}$ to form $\mathrm{BaBr_2}$ and $2\mathrm{H_2O}$:

$\mathrm{Ba(OH)_2} + 2\mathrm{HBr} \rightarrow \mathrm{BaBr_2} + 2\mathrm{H_2O}$

We can use the stoichiometry of the reaction to find the volume of $\mathrm{HBr}$ required.

The number of moles of $\mathrm{Ba(OH)_2}$ is given by its molarity times the volume in liters:

$n(\mathrm{Ba(OH)_2}) = 0.01 \, \mathrm{M} \times \frac{10.0 \, \mathrm{mL}}{1000 \, \mathrm{mL/L}} = 0.0001 \, \mathrm{moles}$

From the balanced equation, we can see that the reaction requires $2$ moles of $\mathrm{HBr}$ for each mole of $\mathrm{Ba(OH)_2}$:

$n(\mathrm{HBr}) = 2 \times n(\mathrm{Ba(OH)_2}) = 2 \times 0.0001 \, \mathrm{moles} = 0.0002 \, \mathrm{moles}$

Then, we find the volume of the $\mathrm{HBr}$ solution by dividing the number of moles by the molarity:

$V(\mathrm{HBr}) = \frac{n(\mathrm{HBr})}{0.02 \, \mathrm{M}} = \frac{0.0002 \, \mathrm{moles}}{0.02 \, \mathrm{mol/L}} = 0.01 \, \mathrm{L} = 10 \, \mathrm{mL}$

Therefore, the volume of the $0.02 \, \mathrm{M}$ $\mathrm{HBr}$ solution required to neutralize the $\mathrm{Ba(OH)_2}$ solution is $10.0 \, \mathrm{mL}$.

On mixing both $\mathrm{\left[ {Co{{(N{H_5})}_5}S{O_4}} \right]Br}$ and $\mathrm{\left[ {Co{{(N{H_3})}_5}Br} \right]S{O_4}}$ becomes 0.01 molar.

$\therefore$ Moles of y and z formed will also be 0.01 both.

${C_x}{H_y} + \left( {x + {y \over 4}} \right){O_2} \to xC{O_2} + {y \over 2}{H_2}O$

${y \over 2} = 3$

$y = 6$

$x + {y \over 4} = {{19} \over 2}$

$x = {{19} \over 2} - {3 \over 2} = 8$

So, formula is ${C_8}{H_6}$.

(1) Molality of solution is defined as number of moles of solute present per kg of solvent.

$m = {{no.\,of\,moles\,(solute)} \over {weight(in\,kg) of solvent}}$

(2) Molarity of solution is defined as number of moles of solute present per litre of solution.

$M = {{no.\,of\,moles\,(solute)} \over {v(in\,litre\,of\,solution)}}$