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koban [17]
3 years ago
7

If 35 ml of 6.0 m h2so4 was spilled, calculate the minimum mass of nahco3 that must be added to the spill to neutralize the acid

.
Chemistry
1 answer:
belka [17]3 years ago
6 0

The balanced equation between the H_2SO_4 and NaHCO_3 is:

H_2SO_4+2NaHCO_3\rightarrow Na_2SO_4+2CO_2+2H_2O

Formula of molarity is:

Molarity = \frac{Moles of solute}{Volume of solution in Liters}

Molarity = 6.0 M, Volume = 35 mL = 0.035 L

Substituting the values,

6 = \frac{Moles of solute}{0.035}

Moles of solute = 0.035 L\times 6 mol/L = 0.21 mole

So, number of moles of H_2SO_4 is 0.21 mole.

From the balanced equation it is clear that for 1 mole of H_2SO_4, 2 moles of NaHCO_3 are required.

Hence, 0.21 mole of H_2SO_4  = 2\times 0.21 mole = 0.42 mole of NaHCO_3

Molar mass of NaHCO_3 = 84.007 g/mol

So, the mass of NaHCO_3 = 84.007 g/mol \times 0.42 mol = 35.283 g




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Calculate the pH of 1.00 L of a buffer that contains 0.105 M HNO2 and 0.170 M NaNO2. What is the pH of the same buffer after the
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2.-pH =3.53

Explanation:

In the first part of this problem we can compute the pH of the buffer by making use of the Henderson-Hasselbach equation,

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where [A⁻] is the conjugate base anion concentration ( [NO₂⁻]), [HA] is the weak acid concentration,[HNO₂].

In the second part, our strategy has to take into account that some of the weak base NO₂⁻ will be consumed by reaction with the very strong acid HCl. Thus, first we will calculate the new concentrations, and then find the new pH similar to the first part.

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# mol HNO₂ produced = 0.012 mol

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pH = 3.40 + log {0.158/.0117}

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Notice there is little variation in the pH of the buffer. That is the usefulness of buffers.

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