Synthesizing nickle ammonium chloride chemistry two
Today’s Lab: Synthesizing Ni(NH3)n Cl2• By Dr. Robert D. Craig, Ph.D.
Today’s Lab: Synthesizing Ni(NH3)n Cl2• Today You will be asked to start with• NiCl2.6H2O• And make Ni(NH3)nCl2 and• Determining the Mass Percent in NH3 in Ni(NH3)n Cl2 : (backtitration)The # of moles of NH3 = mole of HCl - # of mol of NaOH
Synthesizing Ni(NH3)n Cl21. Determining the Mass Percent in NH3 in Ni(NH3)n Cl22. Determining the Mass Percent Ni2+ Ion in Ni(NH3)nCl2 by spectrophotometry3. Determining the Empircal Formula and percent Yield fo Ni(NH3)n Cl2
How will You find this???• Mass of NH3 in Ni(NH3)nCl2 =• ( # of mol NH3) ( 17.03 gram ) =• (1 mole NH3 )• 1st titrate with HCl• 2 nd back titrate with NaOH
Mixed indicator????• Indicator:• From a pH of 3.8 to 5.4 – BromCresol Green• From a pH of 4.2 to 6.2 – Methyl Red• Must see this at equivalence point• Orange Yellow Green Blue
For Your assistance!!! And – remember to work in pairs• I will place the instructions up as a class. As you reach certain points• Look at power point for tips- You will need the addendum as well• First see figure 1 and figure 2 pg 42• Prepare a warm bath with a 600 mL beaker and tap water
Adjust a ring stand• Adjust a ring stand• Stabilize an Erlenmeyer in the beaker the Hot plate• Weight out 4.0 grams (see addendum) of NiCl2.6H2)• Record the mass of the solid and paper on Data sheet 1• Transfer the solid to a clean 125 mL Erlenmeyer flask.
What to do??• Add 10 mL of distilled water to the NiCl2- 6H2O in the flask• Place the flask in the 60oC water bath and clamp the flask in position, (see figure 2)
• Stir the mixture in the flask with a clean 125 mm glass stirring rod until the NiCl2.6H2O has dissolved.• Remove flash-attach into alternate ring stand• Allow to cool for 1-2 min.• Pg 42
Look here for help!! And BE CAREFULL!!• Slowly with strirring add 25 mL of concentrated NH3 solution to the NiCl2- 6H2O solution in the flask.• Cover the top of flask with a wet paper towel. After adding the concentrated nH3 solution, suspend the flask in the warm- water bath by clamping it to the ring stand.
Lab procedure• Make sure the water temperature is between 50 oC and 60oC• Leave the flask in the bath for 15 min. During this time periodically stir the mixture
the filtering apparatus• Note: For the filtering apparatus described below your lab instructor (Rob) will inform you whether you are to place a trap between your filter and the water aspirator and whether you are to place a Buchner funnel in a rubber Filter Vac, a bored stopper, or a rubber filtering adapter.
Look Here!!!• Do not turn off the aspirator before removing the tubing from the side arm of the filter flask• While the above reaction proceeds• Assemble your filtering apparatus as shown in figure 3
SEE FIGURE 3• (SEE FIGURE 3) PG 44• Clamp a 250 mL taped filtering flask to your other ring stand and place a Buchner funnel in the flask• Place a flat circle of filter paper in the funnel.
Can do this….• Attach a piece of pressure tubing to the flask• After isolating the solid disconnect and Reassemble
Be careful• WARNING• THE AMMONIA AND THE ETHYL ALCOHOL ARE EXTREMELY DANGEROUS-ANY SIGNS OF DIZZINESS, OR• HEAT EXHAUSTION-PLEASE LET ME KNOW• WHEN THESE CHEMICALS HIT THE SKIN-WASH IMMEDIATELY
Why are you doing this???• Two important tasks many chemist perform are the synthesis and analysis of compounds• Synthesis involves not only preparing the compoiund , but also maximizing the yield of pure product. After isoloating the product, the chemist must analyze it to ascertain its chemical composition
First . . . . .• Analyzing Ni(NH3)nCL2 Determining th Mass Percent NH3
A complicated titration• Ammonia is a base. Theoretically, you could determine th NH3 content of NI(NH3)nCL2 by titrating a know mass of the compound with an acid solution of known concentration, called a standardized solution.• However, in this case , the procedure gives inaccurate results for two (really 3) reasons.
Reasons here• First, aqueous Ni(NH3)Cl2 solutions slowly evolve NH3,• So some NH3 may evaporate before the sample is completely titrated. Second, the procedure takes longer than normal because the bonds between NH3 and Ni2+ ion must be broken before NH3 can react with the titrant. This bond breaking is slow and the titration reaction itself is fast• So some of the indicator may react with the acid before all the NH3 is released , causing a premature end point.
What is a backtitration?• Goto handout – Rob – upload power point now!!
Volitile sample• To minimize NH3 evoluiton and ensure complete disruption of all NH3-Ni2- bonds , you will use an indirection titration method- a back titration.• You will add a known volume of standard hydrochloric acid solution (HCl) that contains more moles of HCl than ther are moles of NH3**** in the Ni(NH3)nCl2
Not accurate• The HCl wil react with the Ni(NH3)nCl2 sample as shown in equation 5• Ni(NH3)nCl2 (aq, bluish purple) + NHCl (aq) n NH4+ 9aq) + Ni2=(aq, green) = (n + 2) Cl- 9aq)
See manual• The excess acid and a favorable equilibrium constant drive the reaction in Equation 5 to completion• Then you will titrate the excess HCl with standard sodium hydroxide soltution (NaOH) in a process called back tritration• The back titration reaction is shown in Equation 6
neutralization• The back titration reaction is shown in Equation 6• HCl(aq) + NaOH(aq) NaCl(aq) + H2O (l0
Your job• You wil detect the equivalence point of the back titration, the point at which the number of moles of NaOH added is stoichimetrically equivalent to the number of mole of HCl present, by observing an indicator color change.
Equivalence pointThe point at which the indicator changes color is the end point Of the titration . In order to select the proper indicator for a particular titration, one that will produce an end point that is close to the equivalence point, your must consider chemical behavior of the species present in solution at the equivalence point.
Need better indicator• In this case , if the titration mixture contained only products shown in equation 6, The equivalence point would be at pH 7• Hence, phenolphthalein, which changes color at pH 8 , would be a satisfoactory indicator
complications• However, the reaction misture also conatins ammoniumion (NH4+) which hydrolyzes as shown in equation 7, causing the titration mixture to be acidic at the equivalence point.• NH4+(aq) + H2O(aq) NH3l(aq) + H3O+ (Eq &
Mixed indicator• Therefore , phenolphthalein is not a good indicator for this titration, Because the end point would occur after more than an equivalent amount of NaOH had been added. Instead, you will use a mixed indicator solution, composed of bromocresol green and methyl red.
Mixed indicator• Bromcresol green changes form jelow to blue over a pH range of 3.8 to 5.4 and methyl red changes from red to yellow over a pH range of 4.2 to 6.2
Mixed indicator• The mixed indicator changes from rose to green at pH 5.1 which is the pH at the equivalence point of the titration .• Due to the presence of green Ni2+ ion in the misture the color change you will see is orange-yellow to green-blue
Final calculation• From the volume and concentration of HCL solution added you will calculate the number of moles of HCl, added using equation 8
In a nutshell• Record the results of following calculations on Data Sheet 1• Calculate the mass of NiCl2.6H20 used in the synthesis• Calculate the mass of synthesized Ni(NH3)nCl2• Determining the Mass Percent in NH3 in Ni(NH3)n Cl2• Calculate the mass of Ni(NH3)nCl2 sample titrated• Calculate the number of moles of HCl added to Ni(NH3)nCl2 solutioin• Calculate the number of moles of NaOH required for the titration• Calculate the number of moles of NH3 present in the Ni(NH3)n Cl2 sample• Calculate the mass percent NH3 in Ni(NH3)n Cl2• Calculate the mean mass percent NH3 in Ni(NH3)n Cl2