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marysya [2.9K]
2 years ago
5

If a gas occupies 79.5 mL at -1.4°C, what temperature, in Kelvin, would it

Chemistry
1 answer:
Anna35 [415]2 years ago
5 0

Answer:

121 K

Explanation:

Step 1: Given data

  • Initial volume (V₁): 79.5 mL
  • Initial temperature (T₁): -1.4°C
  • Final volume (V₂): 35.3 mL

Step 2: Convert "-1.4°C" to Kelvin

We will use the following expression.

K = °C + 273.15 = -1.4°C + 273.15 = 271.8 K

Step 3: Calculate the final temperature of the gas (T₂)

Assuming ideal behavior and constant pressure, we can calculate the final temperature of the gas using Charles' law.

V₁/T₁ = V₂/T₂

T₂ = V₂ × T₁/V₁

T₂ = 35.3 mL × 271.8 K/79.5 mL = 121 K

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A chemist titrates 190.0 mL of a 0.8125 M ammonia (NH) solution with 0.3733 M HCl solution at 25 °C. Calculate the pH at equival
stealth61 [152]

Answer:

Approximately 4.92.

Explanation:

Initial volume of the solution: V = 190.0\; \rm mL = 0.1900\; \rm L.

Initial quantity of \rm NH_3:

\begin{aligned} n({\rm NH_3}) &= c({\rm NH_3}) \cdot V({\rm NH_3}) \\ &= 0.3733\; \rm mol \cdot L^{-1} \times 0.1900\; \rm L \\ &\approx 0.154375\; \rm mol\end{aligned}.

Ammonia \rm NH_3 reacts with hydrochloric \rm HCl acid at a one-to-one ratio:

\rm NH_3 + HCl \to NH_4 Cl.

Hence, approximately n({\rm HCl}) = 0.154375\; \rm mol of \rm HCl\! molecules would be required to exactly react with the \rm NH_3\! in the original solution and hence reach the equivalence point of this titration.

Calculate the volume of that 0.3733\; \rm mol \cdot L^{-1} \rm HCl solution required for reaching the equivalence point of this titration:

\begin{aligned}V({\rm HCl}) &= \frac{n({\rm HCl})}{c({\rm HCl})} \\ &\approx \frac{0.154375\; \rm mol}{0.3733\; \rm mol \cdot L^{-1}} \approx 0.413541\; \rm L\end{aligned}.

Hence, by the assumption stated in the question, the volume of the solution at the equivalence point would be approximately 0.413541\; \rm L + 0.1900\; \rm L \approx 0.6035\; \rm L.

If no hydrolysis took place, 0.154375\; \rm mol of \rm NH_4 Cl would be produced. Because \rm NH_4 Cl\! is a soluble salt, the solution would contain 0.154375\; \rm mol\! of \rm {NH_4}^{+} ions. The concentration of \rm {NH_4}^{+}\! would be approximately:

\begin{aligned}c({\rm {NH_4}^{+}}) &= \frac{n({\rm {NH_4}^{+}})}{V({\rm {NH_4}^{+}})}\\ &\approx \frac{0.154375\; \rm mol}{0.6035\; \rm L} \approx 0.255782\; \rm mol \cdot L^{-1}\end{aligned}.

However, because \rm NH_3 \cdot H_2O is a weak base, its conjugate \rm {NH_4}^{+} would be a weak base.

\begin{aligned}pK_{\rm a}({{\rm NH_4}}^{+}) &= pK_{\rm w} - pK_{\rm b}({\rm NH_3})\\ &\approx 13.99 - 4.75 = 9.25\end{aligned}.

Hence, the following reversible reaction would be take place in the solution at the equivalence point:

\rm {NH_4}^{+} \rightleftharpoons NH_3 + H^{+}.

Let x\; \rm mol \cdot L^{-1} be the increase in the concentration of \rm H^{+} in this solution because of this reversible reaction. (Notice that x \ge 0.) Construct the following \text{RICE} table:

\begin{array}{c|ccccc} \textbf{R}& \rm {\rm NH_4}^{+} & \rightleftharpoons & {\rm NH_3}& + & {\rm H}^{+}\\ \textbf{I} & 0.255782 \; \rm M \\ \textbf{C} & -x \;\rm M & & + x\;\rm M & & + x\; \rm M \\ \textbf{E} & (0.255782 - x)\; \rm M & & x\; \rm M & & x\; \rm M\end{array}.

Thus, at equilibrium:

  • Concentration of the weak acid: [{\rm {NH_4}^{+}}] \approx (0.255782 - x) \; \rm M.
  • Concentration of the conjugate of the weak acid: [{\rm NH_3}] = x\; \rm M.
  • Concentration of \rm H^{+}: [{\rm {H}^{+}}] \approx x\; \rm M.

\displaystyle \frac{[{\rm NH_3}] \cdot [{\rm H^{+}}]}{[{ \rm {NH_4}^{+}}]} = 10^{pK_\text{a}({\rm {NH_4}^{+}})}.

\displaystyle \frac{x^2}{0.255782 - x} \approx 10^{-9.25}

Solve for x. (Notice that the value of x\! is likely to be much smaller than 0.255782. Hence, the denominator on the left-hand side (0.255782 - x) \approx 0.255782.)

x \approx 1.19929 \times 10^{-5}.

Hence, the concentration of \rm H^{+} at the equivalence point of this titration would be approximately 1.19929 \times 10^{-5}\; \rm M.

Hence, the pH at the equivalence point of this titration would be:

\begin{aligned}pH &= -\log_{10}[{\rm {H}^{+}}] \\ &\approx -\log_{10} \left(1.19929 \times 10^{-5}\right) \approx 4.92\end{aligned}.

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