A 360.0 −mL buffer solution is 0.150 M in HF and 0.150 M in NaF.
a) What mass of NaOH can this buffer neutralize before the pH rises above 4.00? = 1.6
b)If the same volume of the buffer were 0.370 M in HF and 0.370 M in NaF, what mass of NaOH could be handled before the pH rises above 4.00?
Sol.
(a) Let Conc. of NaOH added be x
When , NaOH added , then , conc. of F- increases and conc. of HF decreases .
So ,
Final Conc. of F- = Inital Conc. of F- + Conc. of NaOH added
= 0.150 + x
and , Final Conc. of HF = Initial Conc. of HF - Conc. of NaOH added = 0.150 - x
Also , pKa of HF = 3.17
So , Using Henderson - Hasselbalch equation ,
pH = pKa + log ( [F-] / [HF] )
4.00 = 3.17 + log ( (0.150 + x ) / ( 0.150 - x ) )
or , ( 0.150 + x ) / ( 0.150 - x ) = 100.83 = 6.76
0.150 + x = 6.76 ( 0.150 - x ) = 1.014 - 6.76 x
7.76 x = 0.864
x = 0.111
Therefore , Conc. of NaOH added = 0.111 M
As Volume of solution = 360.0 mL = 0.360 L
So , Moles of NaOH added = 0.111 × 0.360 = 0.03996 mol
As Molar Mass of NaOH = 40 g/mol
So , Mass of NaOH added = 0.03996 × 40 = 1.5984 g
(b)
Now , Final Conc. of F- = Inital Conc. of F- + Conc. of NaOH added = 0.370 + x
and , Final Conc. of HF = Initial Conc. of HF - Conc. of NaOH added = 0.370 - x
Also , pKa of HF = 3.17
So , Using Henderson - Hasselbalch equation ,
pH = pKa + log ( [F-] / [HF] )
4.00 = 3.17 + log ( (0.370 + x ) / ( 0.370 - x ) )
or , ( 0.370 + x ) / ( 0.370 - x ) = 100.83 = 6.76
0.370 + x = 6.76 ( 0.370 - x ) = 2.5012 - 6.76 x
7.76 x = 2.1312
x = 0.275
Therefore , Conc. of NaOH added = 0.275 M
As Volume of solution = 360.0 mL = 0.360 L
So , Moles of NaOH added = 0.275 × 0.360 = 0.099 mol
As Molar Mass of NaOH = 40 g/mol
So , Mass of NaOH added = 0.099 × 40 = 3.96 g
A 360.0 −mL buffer solution is 0.150 M in HF and 0.150 M in NaF. a)...
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A 130.0 −mL buffer solution is 0.100 M in NH3 and 0.135 M in NH4Br. Part A What mass of HCl can this buffer neutralize before the pH falls below 9.00? Express the mass in grams to three significant figures. Part B If the same volume of the buffer were 0.270 M in NH3 and 0.390 M in NH4Br, what mass of HCl could be handled before the pH falls below 9.00?
A 130.0 −mL buffer solution is 0.100 M in NH3 and 0.135 M in NH4Br. What mass of HCl can this buffer neutralize before the pH falls below 9.00? Express the mass in grams to three significant figures. If the same volume of the buffer were 0.265 M in NH3 and 0.400 M in NH4Br, what mass of HCl could be handled before the pH falls below 9.00? Express the mass in grams to three significant figures.
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A 120.0 −mL buffer solution is 0.110 M in NH3 and 0.125 M in NH4Br. You may want to reference (Pages 740 - 751) Section 17.2 while completing this problem. Part A What mass of HCl can this buffer neutralize before the pH falls below 9.00? Express the mass in grams to three significant figures. If the same volume of the buffer were 0.270 M in NH3 and 0.395 M in NH4Br, what mass of HCl could be handled before...
This is a two-part question if someone could please help! 110.0 −mL buffer solution is 0.110 M in NH3and 0.130 M in NH4Br. Part A: What mass of HCl can this buffer neutralize before the pH falls below 9.00? Part B: If the same volume of the buffer were 0.255 M in NH3 and 0.395 M in NH4Br, what mass of HCl could be handled before the pH falls below 9.00?