Question

You have a solution that contains both Al(NO3)3 and Ni(NO3)2. Al3+(aq) is colorless while Ni2+(aq) is a light green/blue. You start adding NH3 dropwise to this solution and you see first a white precipitate form. After a few more drops you start seeing a light green/blue precipitate form. What are the formulas of these precipitates and how did they form (what has adding NH3 got to do with their formation)? As you add more NH3, you start to see the light green/blue solid dissolve and the solution turns to a dark blue color. Nothing happens to the white precipitate. Provide an explanation for what might be occurring. Could you use what you just learned to separate a mixture and Al(NO3)3 and Ni(NO3)2? Describe what would you do to separate the two from each other?

Answer 1

Answer:

sorry I didn't studied this.

Answer 2

In this experiment you will explore characteristic reactions of a number of ions. The reactions you

observe will be used to identify unknown solutions next week. We will use this study to develop logical

schemes for analyzing unknowns, and also to learn and review some descriptive chemistry of the

elements.

INTRODUCTION:

You will be studying seven cations (Na1+, Mg2+, Ni2+, Cr3+, Zn2+, Ag1+, Pb2+) and four anions (NO3

1-

,Cl1-

,

I

1-

, SO4

2-

).

These ions may be separated from each other and identified using simple chemical and

physical properties. Characteristic colors of aqueous ions as well as solid compounds can be very useful

in identification. In general, we expect main group cations to be colorless in solution and to form white

solids. Transition metal cations often have characteristic colors both in the aqueous solution and in solid

compounds. Some ions that generally form colorless compounds can give brilliant colors if mixed with

the right partner. For example, both Pb(NO3)2 and KI dissolve to give colorless solutions, but when

mixed, a bright yellow precipitate of PbI2 is produced.

I. Reactions of cations with NaOH and NH3

Most metal ions react with aqueous OH1-

, hydroxide ion, to form gelatinous precipitates of the metal

hydroxides. For example, Al3+(aq) reacts to form Al(OH)3(s).

Al3+(aq) + 3 OH1-

(aq) → Al(OH)3(s) Eqn 1

Ag1+ is a rare exception, forming Ag2O(s) rather than AgOH. All other hydroxide precipitates you

encounter in this experiment are neutral hydroxides. Adding more OH1-

causes some insoluble

hydroxides to re-dissolve. For example, Al(OH)3(s) reacts to form the complex ion Al(OH)4

-

(aq). In lab,

if you observe a precipitate re-dissolving upon addition of more OH1-

, it is doing a reaction like Eqn 2.

Al(OH)3(s) + OH1-

(aq) → Al(OH)4

1-

(aq) Eqn 2

In this experiment metal hydroxide precipitates that re-dissolve in excess OH1-

form complex ions with

the general formula M(OH)4

n-

(aq).

Metals ions may react with NH3 to form either insoluble precipitates or soluble complex ions. If a

precipitate forms, it is the neutral hydroxide. For example, Al3+(aq) reacts with NH3(aq) to form

Al(OH)3(s). In lab, if you observe an ion reacting with NH3 to form a precipitate, it is doing a reaction

like that in Eqn 3.

Al3+(aq) + 3 NH3(aq) + 3 H2O(l) → Al(OH)3(s) + 3 NH4

1+(aq) Eqn 3

If no precipitate forms or is only briefly present, a complex ion with NH3 has formed, for example,

Zn(NH3)4

2+(aq), Ni(NH3)4

2+(aq), and Ag(NH3)2

1+(aq). See Eqn 4.

Zn2+(aq) + 4 NH3(aq) → Zn(NH3)4

2+(aq) Eqn 4

2