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A chemist adds 160.0 mL of a [tex]2.6 \, \text{mol/L}[/tex] nickel(II) chloride [tex]\left( \text{NiCl}_2 \right)[/tex] solution to a reaction.
Calculate the moles of [tex]\text{NiCl}_2[/tex] added to the flask. Round your answer to 2 significant digits.
2. Calculate the number of moles of nickel(II) chloride [tex]$(NiCl_2)$[/tex]:
The concentration of the solution is given as [tex]\( 2.6 \)[/tex] mol/L. We can use the formula:
[tex]\[
\text{Number of moles} = \text{Concentration} \times \text{Volume}
\][/tex]
Substituting the given values:
[tex]\[
\text{Number of moles of } NiCl_2 = 2.6 \text{ mol/L} \times 0.16 \text{ L} = 0.416 \text{ mol}
\][/tex]
3. Round the result to 2 significant digits:
The number of moles of nickel(II) chloride calculated is [tex]\( 0.416 \)[/tex]. When rounding this to 2 significant digits:
[tex]\[
0.416 \approx 0.42 \text{ mol}
\][/tex]
To summarize, the chemist added:
- Volume in liters: [tex]\( 0.16 \)[/tex] L - Number of moles of [tex]\( NiCl_2 \)[/tex]: [tex]\( 0.416 \)[/tex] mol - Rounded number of moles of [tex]\( NiCl_2 \)[/tex]: [tex]\( 0.42 \)[/tex] mol
So, the chemist added approximately [tex]\( 0.42 \)[/tex] moles of nickel(II) chloride to the reaction flask.
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