Stoichiometry Chemistry 1010 Review Tutorial Stoichiometry and Lewis Structures April 9 th, 2013 Stoichiometry Stoichiometry involves MOLES Elements/compounds can only be compared side by side using moles 1 mol = 6.022 10 23 individual particles (Avogadro s #) applies to anything We cannot compare masses (or other properties) as they differ for each element or compound Types of Problems Mass Percent and Percent Composition Empirical and Molecular Formulas Combustion Analysis Dry Stoichiometry Percent Yield Limiting Reagent Titration (solution stoichiometry) 1
Mass Percent Mass Percent Example To find the mass percent of a particular element in a compound: Find the molar mass of the compound Find the TOTAL mass of the element in question in the compound Divide the mass of element by the mass of compound and multiply by 100 Find the mass percent of each element in C 6 H 12 O 6 Mass Element 100% = Percent Composition Mass Compound Empirical Formula from % Empirical Formula from % To find the empirical formula given percent composition: Assume 100g of sample (MUST state this!) With this assumption, percentages for each element equate to masses in grams Convert each mass to number of moles by diving by molar mass Divide each number of moles by the smallest number of moles to get a simple whole number ratio If you get a decimal, helpful to convert to a fraction Then you can see what number you have to multiply each fraction by 2
Empirical Formula Example Butanoate is 58.8% carbon, 9.87% hydrogen, and 31.3% oxygen by mass. What is its empirical formula? Don t forget you can t compare masses! Must look at moles! Molecular Formulas Normally, you will use the empirical formula provided, or one determined in a previous question The molecular mass of the compound will be provided Find the molar mass of the empirical formula: how many empirical formulas make up the molecular formula? Molecular Formulas Molecular Formula Example Molecular Mass Empirical Mass = Whole Number Ratio (Empirical) whole # = Molecular If you don t find a whole number, CHECK your work! A compound has the empirical formula CH 2 O. Its molar mass was determined to be 60.06 g/mol. What is its molecular formula? 3
Combustion Analysis You will be given the mass of an unknown compound (generally containing only C, H, and O) and told it is burned The masses of H 2 O and CO 2 (which are always products of combustion) produced will also be provided Since all C and H from compound form CO 2 and H 2 O, you can determine their amounts Combustion Analysis To find the empirical formula of unknown: Determine moles of CO 2 and H 2 O produced Moles = Mass/Molar Mass Find moles of C and H alone Molar ratios (eg. 1 mol C per mol CO 2 ) Turn moles of C and H into masses To find O - Law of Conservation of Mass Combustion Analysis Mass Unknown = Mass C + Mass H + Mass O Mass of O is easy to find! Must convert to MOLES to find empirical formula. Use moles of C, H, O to determine empirical formula as before (divide by smallest) Can be tangly, but remember you re just finding the number of moles of each element Combustion Analysis Example Vitamin C (ascorbic acid) is composed of carbon, hydrogen, and oxygen. Combustion of a 3.87 g sample yields 5.80 g of CO 2 and 1.58 g of H 2 O What is the empirical formula of Vitamin C? If the molar mass of Vitamin C is 176.1256 g/mol, what is its molecular formula? 4
Dry Stoichiometry For given a reaction, you will be provided with the amount of one substance and asked to find the amount of another Reaction MUST be balanced! Convert given mass of substance into moles Relate moles of what you want to moles of what you have (mole ratio) Convert moles of desired substance back to the required unit (e.g. mass in grams) Dry Stoichiometry Example Given the unbalanced reaction: NaHCO 3(s) Na 2 CO 2(s) + CO 2(s) + H 2 O (l) What mass of sodium carbonate is produced when 4.00 g of sodium bicarbonate is decomposed? Percent Yield Percent Yield Example Normally given as a secondary part of a longer question You will be given the mass of product actually obtained for a particular reaction To find the % yield, divide this value by the calculated (theoretical) mass and multiply by 100% Actual Mass 100% = Percent Yield Theoretical Mass If 2.12 g of sodium carbonate was produced in the previous reaction, what would the percent yield be? NaHCO 3(s) Na 2 CO 2(s) + CO 2(s) + H 2 O (l) 5
Limiting Reagent Similar to a regular stoichiometry problem, except you will be given amounts of two or more reactants and asked to determine the amount of product that would be produced May be followed by a percent yield question You must work with ALL reactants to determine how much product each would produce Limiting Reagent To begin, convert all reactant masses into moles (n = m/mm) Determine the amount of product that would result if each reactant was used up completely (multiply by mole ratio comes from BALANCED equation) # mol product moles reactant # mol reactant = moles Product Mole ratio (from coefficients) Limiting Reagent Limiting reagent is the reactant that gives the smallest amount of product This smallest amount of product is the actual amount of product produced (in moles) Use this value to determine the mass (or other desired unit) of product produced Amounts of excess reagents can be determined by subtracting the moles used (mole ratio again!) from initial moles of each excess reactant Limiting Reagent Example Ammonia gas can be prepared as follows: 2NH 4 Cl (s) + CaO (s) 2NH 3(g) + H 2 O (g) + CaCl 2(s) If 112 g of CaO reacts with 224 g of NH 4 Cl, what mass of NH 3 would be produced? What mass of excess reagent will remain after complete reaction? If 38.7 g of NH 3 is obtained, find the % yield 6
Titrations Titration Example A known amount of a solution with known concentration is added to a known amount of a second solution, having unknown concentration Identical to other stoichiometry problems, except we now determine moles using concentrations and volumes instead of masses and molar masses A 25.00 ml aliquot of Ba(OH) 2 solution is titrated with standardized 0.3500 mol/l HBr (aq). At the equivalence point, 15.39 ml of HBr solution was added. What is the concentration of the Ba(OH) 2 solution? Lewis Structures Lewis Structures Useful for showing valence electrons in atoms, ions, and compounds Use dots, arranged in pairs, to depict valence electrons 7
Molecular Lewis Structures Molecular compounds contain two or more non-metals, held together by covalent bonds Electron sharing Lewis structures are constructed a little differently, but in a step-wise manner that is the SAME for all covalent compounds and ions! Drawing Lewis Structures 1. Count total valence electrons for all atoms Groups 1 and 2 have 1 and 2 valence electrons respectively For other main group elements (Grps 13-18), use Group # - 10: Group 14 has 4 valence e - Any ionic charge MUST be accounted for! For negatively charged ions, add a number of electrons corresponding to the charge For positively charged ions, subtract a number of electrons corresponding to the charge Drawing Lewis Structures 2. Determine the central atom typically the least electronegative atom Exceptions: Hydrogen is NEVER the central atom Carbon is ALWAYS the central atom 3. Attach remaining atoms to central atom via single bonds (bonding pairs of electrons) Drawing Lewis Structures 4. Subtract bonding pair electrons from total valence electrons (2 e - per single bond) 5. Assign remaining electrons to terminal atoms, fulfilling octet (8e - total each) 6. Assign any leftover electrons to the central atom (in pairs) Don t forget to indicate overall charge when drawing ions! 8
Drawing Lewis Structures Lewis Structure Examples 7. Calculate formal charge for each atom: FC = (valence e - ) (unbonded e - ) (½ bonding e - ) [FC = (valence e - ) (unbonded e - ) (bonds)] 8. Reduce formal charges by making double or triple bonds wherever possible If you have a choice, take e - from the less EN atom 9. Check for resonance structures (more than one possible Lewis Structure) Draw Lewis Structures for the following: Dichloromethane (CH 2 Cl 2 ) Nitrite ion (NO 2- ) Chlorine difluoride cation (ClF 2+ ) Preparation for Final Review old tests, assignments, and tutorial notes. Do the online practice sheets! If you re having trouble with a particular topic get some FREE help with it Chemistry 1010 Help Centre: C-2022 Mon-Thurs, 9am-4:30pm; Fri, 9am-3:30pm Chemistry Resource Room: C-2010 Open Sat, April 13 and Sun, April 14: 10am-6pm A Few Pointers There are many parts of a question, other than the right answer, that are worth marks Show any/all formulas you use Keep track of sig figs! It s good to keep an extra digit or two until the end, but be sure to indicate this! Always give final answers with proper sig figs Show all workings!!! (This includes molar mass) Read and re-read the question make sure you know what exactly is expected/required Show all units, and double check them 9
Good Luck on Exams!!! Have a great summer! 10