1.2: Mole, Conversion Factors, Empirical & Molecular Formulas. Ms. Kiely Coral Gables Senior High IB Chemistry SL

Transcription

1 1.2: Mole, Conversion Factors, Empirical & Molecular Formulas Ms. Kiely Coral Gables Senior High IB Chemistry SL

2 TURN IN the Signed Syllabus and Topic 1 Exercises Bell-Ringer #2 What amount in grams is present in 2.1 mol of sodium hydroxide, NaOH?

3 1.2 Practice: NO CALCULATOR! How many moles of carbon dioxide are there in 66 g of carbon dioxide?

4 Answer How many moles of carbon dioxide are there in 66 g of carbon dioxide? 1.5 mol

5 Section 1.2: Empirical Versus Molecular Formulas Empirical Formula: formula showing the simplest ratio of numbers of atoms of each element in a compound Molecular Formula: formula showing all the atoms present in a molecule. The molecular formula of a compound is often a multiple of the empirical formula; however, in some cases the empirical formula and molecular formula of a compound are the same.

6 Empirical Versus Molecular Formulas Why are empirical formulas used? If we have an unknown substance that we know is a compound, we can experimentally determine its composition (what elements it is made of) by getting information on the different masses that exist in the sample. For example, if we notice there are many atoms that weigh approximately 12g and other atoms that approximately weight 1g then we can assume the sample is made up of carbon and hydrogen atoms, respectively, giving us an empirical formula of CH. CH, however, would be very unstable. The true molecular formula must be something more stable, such as CH₄, or C₂H₂, etc.

7 Solving for the Empirical Formula of a Compound 1) If you know how much of each element is present in grams: 1. Convert the mass of each element to moles. 2. Divide each result by the smallest mole amount you calculated in step Approximate to the nearest whole number. That whole number is now the subscript of each respective element in the compound. 2) If you only know the percent composition by mass of each element in the compound: 1. Assume you have 100g of the sample and convert each respective percentage to grams. 2. Convert the mass of each element to moles. 3. Divide each result by the smallest mole amount you calculated in step Approximate to the nearest whole number. That whole number is now the subscript of each respective element in the compound.

8 Example A sample of urea contains 1.210g N, 0.161g H, 0.480g C, and 0.640g O. What is the empirical formula of urea?

9 Example The mineral celestine consists mostly of a compound of strontium, sulfur, and oxygen. It is found by combustion analysis to have the composition 47.70% by mass Sr, 17.46% S, and the remainder is O. What is its empirical formula?

10

11 Molecular Formulas The molecular formula of a compound can be deduced from the empirical formula if the true molar mass of the compound is known: (M of empirical formula) = M of true compound = M of true compound M of empirical formula To obtain the molecular formula, obtain the value of x and multiply it against the subscripts present in the empirical formula.

12 Molecular Formulas Example Calomel has the empirical formula HgCl and a molar mass (M) of g mol ¹. What is its molecular formula?

13 Using Empirical Formula to Solve for Percent Composition Example: What is the percentage by mass of N, H, and O in the compound ammonium nitrate, NH₄NO₃? In this particular question, you are asking yourself What percentage of this compound is made up of Nitrogen? of hydrogen? of oxygen? Percent Composition: the relative amounts of the elements in a compound are expressed as the percent composition, known as the percent by mass of each element in the compound. Step 1: Calculate M of the compound. Step 2: For each element in the compound, total the mass of its atoms, divide by M of the whole compound, and multiply by 100. Step 3: Make sure that all percentages total to 100!

14 Using Empirical Formula to Solve for Percent Composition ANSWER: 34.99% N, 5.05% H, and 59.96% O