1. Unit 3: Stoichiometry Part 1

2. Atomic Masses Atomic mass – (atomic weight) – The atomic mass of an element indicates how heavy, on average, an atom of an element is when compared to an atom of another element (weighted avg. of all isotopes) Atomic mass units – (amu) – the units for atomic masses on the periodic table

3. The Carbon-12 Scale Mass of one 12 C atom = 12 amu (exactly) Note that 12 C and C-12 mean the same thing

4. Atomic Masses and Isotopic Abundances: Mass spectrometer – a device used to experimentally determine the atomic mass of an atom Isotopic abundances – the percentage of each isotope that exists in nature (also, determined using the mass spec.)

5. Figure 3.1 – Mass Spectrometer A mass spectrometer is used to determine atomic masses

6. Figure 3.2 – Mass Spectrum of Cl The area under the peak in the mass spectrogram gives the isotopic abundance

7. Atomic Mass Calculations:

8. Example 3

9. Masses of Individual Atoms; Avogadro’s Number: Avogadro’s Number – The number of atoms that is equal to the atomic mass of any element N A = 6.02 X 10 23 For Example: 6.02x10 23 H atoms in 1.008 grams of H (atomic mass of H = 1.008)

10. Figure 3.3 – One Mole of Several Substances

11. Example 2

12. The Mole Mole – equal to Avogadro’s Number, equal to 6.02x10 23 particles of a substance

13. Specialized units: The correct name for a particle of a substance based on the type of matter Atom – the representative particle for an element example – Fe, S, etc. ion – the representative particle for a charged particle example – Na +1, Cl -1, NH 4 +1,etc. Molecule – the representative particle for a molecular compound (made up of non-metals) example – CO 2, CH 4, etc. Formula unit – the representative particle for an ionic compound (metal and non-metal or polyatomic ion) example – KCl, MgSO 4, etc.

14. Molar mass: Molar mass – (MM) – the mass of 1 mole of a substance; equal to the atomic mass on the periodic table Round to the tenths place from the Periodic Table to simplify calculations

15. Calculating molar mass: 1. find the mass of the element on the periodic table 2. multiply by the number of atoms of that element in the formula (distribute parenthesis) 3. sum the relative masses of the individual elements

16. Molar Masses of Some Substances

17. Practice: Example: Determine the molar masses of the following substances. AluminumCaSO 4 (NH 4 ) 3 P

18. The Significance of the Mole In the laboratory, substances are weighed on balances, in units of grams The mole allows us to relate the number of grams of a substance to the number of atoms or molecules of a substance

19. Mole-Gram Conversions Molar mass can be used like any other conversion factor: Molar mass (g) = 1 mole

20. Example 3

21. Mass Relations in Chemical Formulas Percent composition formula percent composition - the number of grams of each element in 100 g of the compound Part x 100 = % composition whole

22. 2 types of % Comp. problems: 1.Given data as masses of elements, etc. a.Use the masses of each element and the total mass of the compound

23. 2 types of % Comp. problems: 2. Given the chemical formula of the compound a. use the relative molar masses of each element and the molar mass of the compound

24. Example 4

25. Subscripts: 1.Represent the atom ratio in a compound 2.Represent the mole ratio in a compound

26. Diatomic Elements: Elements that due to there chemical reactivity exist only as molecules of 2 atoms in nature. 7 diatomic elements: Br I N Cl H O F H O N Cl Br I F

27. Simplest Formula from Chemical Analysis (Empirical Formula): Simplest formula – (empirical formula) – the simplest whole number ratio of the atoms in a compound

28. Calculating the empirical formula: 1. can be determined from masses of the individual elements or the % composition of the elements in a compound 2. if %’s are given consider the sample to be of 100 grams and so the %’s become the masses in grams: 25.6% = 25.6 g 3. convert the mass of each element to moles 4. divide each number of moles by the smallest number of moles of all of the answers to #3 5. *If the answers to #4 are whole numbers, these are the subscripts in the empirical formula. * If any of the answers to #4 is not a whole number, convert all answers to a common fraction. Multiply each fraction by the denominator resulting in a whole number and these are the subscripts in the empirical formula.

29. Common Fractions: 0.250.330.500.660.75 ¼1/3½2/3¾

30. Example 3.5 – Simplest Formula from Masses of Elements

31. Example 3.6 – Simplest Formula from Mass Percents The compound that gives vinegar its sour taste is acetic acid, which contains the elements carbon, hydrogen, and oxygen. When 5.00g of acetic acid is analyzed it is found to contain 2.00g of carbon, 0.336g of hydrogen, and 2.66g of oxygen. What is the empirical formula of acetic acid?

32. Molecular Formula: Molecular formula – the true ratio of the atoms in a compound as it exists naturally; may be the same as the empirical formula; will relate to the empirical formula in whole number ratios

33. Molecular Formula: Calculating the molecular formula: 1. find the molar mass of the empirical formula 2. divide the molecular molar mass (given in the question) by the empirical molar mass MMM EMM 3. multiply each subscript in the empirical formula by the answer to #2, these are the subscripts for the molecular formula

34. Example 3.7