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Regents Chemistry Stoichiometry

Regents Chemistry Stoichiometry. 02. Atomic Mass Unit (amu). amu- unit used to for mass of an atom amu of Oxygen? 16amu Why not in grams? Mass of oxygen is really 2.7 × 10 -23 g Atoms are too small, number is too bulky. Find the mass of the following atoms. Mg = ______ Li = ______

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Regents Chemistry Stoichiometry

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  1. Regents Chemistry Stoichiometry

  2. 02 Atomic Mass Unit (amu) • amu- unit used to for mass of an atom • amu of Oxygen? 16amu • Why not in grams? • Mass of oxygen is really 2.7 × 10-23g • Atoms are too small, number is too bulky

  3. Find the mass of the following atoms • Mg = ______ • Li = ______ • Cl = ______ 4. Al = ______ 5. Ca = ______ 6. H = ______ 24 amu 23 amu 7 amu 20 amu 35 amu 1 amu

  4. Types of Elements Monoatomic Element • One atom of an element that is stable enough to stand on its own (VERY RARE) • not bonded to anything Diatomic Elements • Elements whose atoms always travel in pairs (Br,I,N,Cl,H,O,F) • Bonded to another atom of the same element.

  5. Formula Mass • mass of an atom, molecule, compound • O2 • This means that mass of O2 = 2 × _______ amu = ________ amu • Notice how we use rounded numbers Subscript = tells you the total number of atoms in the compound/molecule. 16 32

  6. Gram Formula Mass

  7. Formulas • Molecular – actual ratios of atoms in a molecule or compound. • ex: C4H10 or C6H12O6 • Empirical – simplest ratio of atoms in a compound • ex: C2H5 or CH2O • Structural – shows the arrangement of the atoms in the actual compound

  8. Example • The empirical formula is CH and the molecular mass is 26. What is the molecular formula? • C2H2 • C3H3 • C4H4 Step 1: find mass of empirical CH = 13amu Step 2: Divide molecular mass by empirical mass 26 (given)/13 = 2 Step 3: multiply empirical formula by answer in step 2 CH × 2 = C2H2

  9. Review - The molecular formula is C3H6. What is the empirical formula? ______________ - Which is an empirical formula? • C2H2 • H2O • H2O2 • C6H12O6 • What is the empirical formula of the compound whose molecular formula is P4O10? • PO • PO2 • P2O • P8O20

  10. Chemical Equation • Shows which bonds are broken and which bonds are built. • Numbers of atoms on the left side must equal number of atoms on the right side of the arrow • After the elements are correctly written, only the coefficient can be changed. • No coefficient means there is only one molecule H2 + O2 H20 2 atoms H 2 atoms O 2 hydrogen atoms 1 oxygen atom 2 2 4 = 4 = 2 =

  11. Fix this equation! • Formationof salt from sodium and chlorine gas. Na + Cl2 → NaCl Na Cl Na Cl Cl PRODUCES

  12. Modeling Conservation of Matter

  13. Review equations Mg + Cl2 MgCl2 Ca + HCl  CaCl2 + H2 Ca + H20  Ca(OH)2 + H2 Given the incomplete equation: 2N2O5(g)  Complete the balanced equation. • 2N2(g) + 3H2(g) • 2N2(g) + 2O2(g) • 4NO2(g) + O2(g) • 4NO(g) + 5O2(g)

  14. Reaction Types Four Basic Types • Synthesis • Decomposition • Single replacement (substitution) • Double replacement

  15. Synthesis • Formation of only ONE product from two reactants, but not always. Examples: N2(g) + 3H2(g) 2NH3(g) CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) because O2 combines with both metal and nonmetal to form two oxides.

  16. Decomposition • One reactant breaks apart to form several products. • AKA combustion if products are CO2 and H2O Examples: 2H2O2(aq) 2H2O(l) + O2(g) hydrogen peroxide decomposes over time to leave behind water.

  17. Single Replacement • A more active metal replaces a less active metal in a compound, or vice versa. Example 2Fe(s) + 6HCl(aq)  2FeCl3(aq) + 3H2(s) what happens when a metal becomes corroded by an acid where an element is reacting with a compound.

  18. Double replacement • Reaction between aqueous compounds. • Cations and anions switch position. • If an insoluble precipitate forms, the reaction is an end reaction, otherwise an aqueous mixture of ions Example AgNO3 (aq) + NaCl(aq)  NaNO3(aq) + AgCl(s)

  19. Review • Cl2(g) + 2NaBr(aq)  2NaCl(aq) + Br2(l) 2. FeCl3(aq) + 3NaOH(aq)  Fe(OH)3(s) + 3NaCl(aq) 3. 2Mg(s) + O2(g)  2MgO(s) 4. H2CO3(aq)  H2O(l) + CO3(g) SR DR S D/Combustion

  20. Gram Formula Mass (GFM) AKA • Molar Mass • Atomic mass • Molecular mass (only in covalent) • Mass of formula in grams

  21. Formula Mass • Sum of atomic masses in the molecule • What is the formula mass (or molecular mass) of K2CO3?

  22. gram formula mass (GFM) • GFM – describes the mass of one mole of a compound • To find GFM, add individual GAM for each element in the compound • H20 = 1.00 + 1.00 + 16.00 = 18.00 g • (NH4)2SO4 = N = 2 × 14 = 28 H = 8 × 1 = 8 S = 1 × 32 = 32 O = 4 × 16 = 64 132 g

  23. Gram formula Mass (molar mass)= mass in grams • Mass of 6.02 x 1023 particles (1 mole of particles). • If you weigh 6.02 x 1023 particles (1 mole of particles) of K2CO3, the weight on the scale will be 138 grams

  24. Molecular Mass H20 = 1.01 + 1.01 + 16.00 = 18.02 g/mol (NH4)2SO4 = N = 2 × 14.00 = 28 H = 8 × 1.01 = 8.08 S = 1 × 32.00 = 32 O = 4 × 16.00 = 64 132.08 g/mol

  25. Determine the molecular Formula of the following HNO3 H N O 1×1 1×14 3×16 1 + 14 + 48 63amu (NH4)2CO3 N H C O 2×14 8×1 1×12 3×16 28 + 8 + 12 + 48 96amu

  26. % Composition HNO3 63amu H N O 1/63 14/63 48/63 ×100 ×100 ×100 1.59% 22.22% 76.19% Needs to total 100% (NH4)2CO3 96amu N H C O 28/96 8/96 12/96 48/96 ×100 ×100 ×100 ×100 29.17% 8.33% 12.50% 50.00% Needs to total 100%

  27. % Given instead • Cmpd is 86% C and 14% H. What is the empirical formula? C H 86/12 14/1 7.17 C 14 H 7.17/7.17 14/7.17 1 2 CH2 % of water in Na2CO3 • 10H2O (formula mass = 286)? H2O = 2 + 16 = 18amu 10 × 18am = 180 180/286 × 100 = 62.94%

  28. Rule 1 mole (of molecules) Equals 1 gram molecular mass Equals 6.02 x 1023 molecules Equals 22.4 liters (for gases at STP)

  29. Example • What is the formula mass of NO2? 46 gram • What is the mass of 2 moles of NO2? 2 X 46 gram = 92 grams • What is the mass of 12 x 1023 molecules of NO2? 12 x 1023 / 6.02 x 1023 = 2 2 x 46 = 92 grams • What is the mass of 44.8 liters of NO2? 44.8 / 22.4 = 2 2 x 46 = 92 grams

  30. Density Density = mass / volume Usually expressed for gases in grams/liter Gram formula mass = density at STP (g/L) x 22.4 liters A piece of aluminum has a mass of grams and a volume of ml. What is its density? 1.35 g/l x 22.4 l = 30.24 grams

  31. Another example Which gas has a density of 1.70g/l at STP? • F2 • He • N2 • SO2 Total the mass of each choice to find the answer.

  32. Mole-Mole Problems • Answers how many moles of one element or compound react with a given number of moles of another element or compound. • How many moles of Ca are needed to react completely with 6 moles of H2O in the following reaction: Step 1: Balance equation Step 2: Cross out molecules not needed. Step 3: Write mole number on top of given formula and an x on the unknown Step 4: Write mole number on bottom of formula from balanced equation Step 5: set up proportions x 6 Ca + H2O  Ca(OH)2 + H2 2 1 mole 2 mole

  33. Review/Practice • Given the reaction: CH4 + O2 CO2 + H2O • How many moles of oxygen are needed for the complete combustion of 3.0 moles of CH4? • 6.0 moles • 2.0 moles • 3.0 moles • 4.0 moles • What amount of oxygen is needed to completely react with 1 mole of CH4? • 2 moles • 2 atoms • 2 grams • 2 molecules

  34. Avogadro’s number • measured to be approximately 6.022 x 1023 (to 4 s.f) • Chemists use the mole in the same way that grocers use the dozen for groups of 12 and stationers use the ream for groups of 500.  •  we can use the mole without being overly concerned about exactly how many objects it represents

  35. Mole • smallest measurable mass of matter contains trillions of atoms, so chemists use a unit of amount called the mole (abbreviated mol). • one mole is the number of atoms in 12 g of carbon-12 • one atom of tin-120 has a mass of 120 u, it follows that one mole of tin-120 atoms will have ten times the mass of one mole of carbon-12 atoms, i.e. 120 g.  • the mass in grams of one mole of atoms of any element will be numerically equivalent to its atomic mass in g/mol.

  36. DIMO

  37. Dimensional Analysis

  38. Review • Which gas sample contains a total of 3.0 x 1023 molecules? • 71g of Cl2 • 2.0 g of H2 • 14g of N2 • 38g of F2 A sample of an unknown gas at STP has a density of 0.630 g/l. What is the gram molecular mass of this gas? • 2.81g • 14.1g • 22.4g • 63.0g • Which quantity represents 0.500 mole at STP? • 22.4L of Ar • 11.2L of N2 • 32.0L of H2 • 44.8L of He

  39. Using Avogadro’s number • how many atoms are in a sample of silicon that has a mass of 5.23 g.  • Moles can be number of atoms or particles in a molecule

  40. Mole • mass of 1 atom = mass of a mole of atoms / 6.022 x 1023 • mass of 1 C atom = 12.01 g / 6.022 x 1023 C atoms • mass of 1 C atom = 1.994 x 10-23 g

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