Question

What is the I– concentration just as AgCl begins to precipitate when 1.0 M AgNO3 is slowly added to a solution containing 0.0500 M Cl– and 0.0500 M I– ? (Ksp values: AgCl = 1.8×10−10 ; AgI = 1×10−16 ) (the halide ions could come from sodium salts...)

Answer #1

A solution is 0.10 M Pb(NO3)2 and 0.10 M AgNO3. If solid NaCl is
added to the solution, what is [Ag+] when PbCl2 begins to
precipitate? (Ksp PbCl2 = 1.7 x 10-5; AgCl = 1.8 x 10-10) A
solution is 0.10 M Pb(NO3)2 and 0.10 M AgNO3. If solid NaCl is
added to the solution, what is [Ag+] when PbCl2 begins to
precipitate? (Ksp PbCl2 = 1.7 x 10-5; AgCl = 1.8 x 10-10)

A solution is 0.0010 M in both Ag+ and Au+. Some solid NaCl is
added slowly until the second solid compound just begins to
precipitate. What is the concentration of Au+ ions at this point?
Ksp for AgCl = 1.8 x 10-10 and for AuCl is
2.0 x 10-13

1. a. 0.0500 M of AgNO3 is used to titrate a 25.00-mL containing
0.1000 M sodium chloride (NaCl) and 0.05000 M potassium iodide
(KI), what is the pAg of the solution after 15.00 mL of AgNO3 is
added to the solution? Ksp, AgCl (s) = 1.82 x 10-10; Ksp, AgI(s) =
8.3*10-17.
b. Same titration as in (a), what is the pAg of the solution
after 25.00 mL of AgNO3 is added to the above solution?
c. Same titration as...

a) A 0.1500 M of AgNO3 solution was employed to titrate a 25.00
mL of 0.1250 M of NaI and 0.2500 M NaCl. Given that Ksp, AgI(s) =
8.3*10-17, and Ksp, AgCl(s) = 1.8*10-10, please calculate the
concentration of Ag+ ion after 6.00 mL of AgNO3 was added.
Answer: 1.2 x 10-15 M
b) please calculate the pAg after 100.00 mL of AgNO3 was
added
Answer: 1.35
I just want to see how they got that.

Calculate the concentration (in M) of Ag+ when AgIO3 just begins
to precipitate from a solution that is 0.0275 M in IO3−. (Ksp =
3.17 ✕ 10−8)

A solution is 0.10 M Pb(NO3)2 and 0.10 M AgNO3. If solid NaCl is
added to the solution, what is [Ag+] when PbCl2 begins to
precipitate? (Ksp PbCl2 = 1.7 x 10-5; AgCl = 1.8 x 10-10)

. (8) A 10.00-mL portion of a 0.50 M AgNO3 (aq) solution is
added to 100.0 mL of a solution that is 0.010 M in Cl- a) Will AgCl
(s) (Ksp = 1.8X10-10) precipitate from this solution? If so, how
many moles will precipitate and what will be the concentrations of
the ions after precipitation?

a) A 0.1500 M of AgNO3 solution was employed to
titrate a 25.00 mL of 0.1250 M of NaI and 0.2500 M NaCl. Given that
Ksp, AgI(s) = 8.3*10-17, and Ksp, AgCl(s) =
1.8*10-10, please calculate the concentration of
Ag+ ion after 6.00 mL of AgNO3 was added.
b) please calculate the pAg after 100.00 mL of AgNO3
was added

Sodium sulfate is slowly added to a solution containing 0.0500 M
Ca2 (aq) and 0.0220 M Ag (aq). What will be the concentration of
Ca2 (aq) when Ag2SO4(s) begins to precipitate? Solubility-product
constants. (can be found here:
https://sites.google.com/site/chempendix/Ksp )
[Ca^2+] = __________ M

What is the equilibrium constant for the dissolution of lead(II)
chromate in Na2S2O3? For PbCrO4, Ksp = 2.0 x 10–16; for
Pb(S2O3)34–, Kf = 2.2 x 106. (Use E for the power of
10)
Solid AgNO3 is slowly added to a solution that contains 0.24 M
of Cl− and 0.10 M of Br− (assume volume does not change). What is
[Br−] (in M) when AgCl(s) starts to precipitate? Ksp of AgCl is 1.8
x 10−10. Ksp for AgBr is 5.0...

ADVERTISEMENT

Get Answers For Free

Most questions answered within 1 hours.

ADVERTISEMENT

asked 5 minutes ago

asked 18 minutes ago

asked 24 minutes ago

asked 31 minutes ago

asked 33 minutes ago

asked 38 minutes ago

asked 38 minutes ago

asked 47 minutes ago

asked 53 minutes ago

asked 57 minutes ago

asked 1 hour ago

asked 1 hour ago