Additional Empirical Formula Work

Additional Empirical formula work: Name

The empirical formula is defined as the simplest whole number ratio of the elements that make up a compound. This could also be the simplest ratio of molecules in a more complex molecule, such as a hydrate (EF = ratio of anhydrous salt to water) or a complex ion (moles of “central ion” to the “rest”.) Try these varied EF problems.

1. The cyanide ion is capable of forming complex ions with the copper (I) ion in two different ratios. Use the data given below to determine the formulas of these two complex ions, Cu(CN)x and Cu(CN)y. a) In the first lab experiment, it is found that a solution that contains 0.00786 g of the Cu+ ion is able to complex with the cyanide ion such that, when the complex ion Cu(CN)x is isolated, the complexed ion has a mass of 0.01430 g. What is the value of ‘x’ in the formula of this complex ion?

• Thought process: to find the formula of Cu(CN)x, we need to find the moles of copper ions, and the moles of cyanide ion.

Mass Cu(CN)x 0.01430 g (given in problem) Mass Cu+ 0.00786 g (given in problem) Mass CN–

• How can you calculate the grams of cyanide ion? Do so. • Now that you have grams of each (Cu+ and CN–), calculate moles of each.

• You now have the correct (but ugly!) ratio of moles of each. How to make it pretty? Do so.

• Write the formula of the complex ion using your findings. What is the correct charge on this ion?

b) A subsequent laboratory experiment finds that when 0.2630 g of Cu+ ion is placed in a more concentrated solution of cyanide, the complex ion formed has a mass of 0.6937 g. Using the technique from part a) above, find and write the correct formula (with correct charge) for the complex ion formed.

c) Calculate separately for each of the two complex ions above: Mass % Cu+ and mass % CN–. 2. By heating a hydrate, we can form the anhydrous salt. (Anhydrous = without water) I like to make a picture for these kinds of things:

In this pie chart, the entire pie represents the mass of the hydrate (the salt plus its associated water molecules). 1 40% • The two parts of the pie represent the mass of the water 1 2 molecules, and the mass of the anhydrous salt. You can find 2 60% out all three values by finding lab data (either by yourself in lab, or by interpreting a problem given to you.

Example Problem 2a: Alum is a hydrated salt that has a formula of KAl(SO4)2•XH2O. When 24.98 g of hydrated alum are heated, the salt is dehydrated, leaving behind 13.60 g of the dehydrated salt. Based on this information, fill out the following table (do work elsewhere) and answer the questions that follow. Mass of hydrated salt Mass of anhydrous salt Mass of associated water Moles of anhydrous salt Moles of associated water

What is the formula of the anhydrous salt?

What is the molar mass of the anhydrous salt?

What is the molar mass of water?

How many moles of anhydrous salt were in the sample?

How many moles of water were in the sample?

What is the correctly-written formula of the hydrated salt?

Example Problem 2b: Barium hydroxide is often found as a hydrated salt, Ba(OH)2•XH2O. When a 15.00 g sample of this hydrate were heated, 8.05 g of anhydrous salt remained. • What is the correct formula for this compound, based upon the given lab data? (It’s safe to assume that there is X moles per one mole of Ba(OH)2. If you find X isn’t exactly a whole number, this is because of either a lab error, or because sometimes hydrates can be slightly over- or under-hydrated. Please round your calculated X to the nearest whole number. • Calculate the percent composition of water twice: 1st, from your lab data. 2nd, from the formula of the hydrate with its X rounded to the nearest whole number. • Calculate the percent error of this lab data, using the “rounded formula” as the accepted value.