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III. Moles/Stoichiometry

III. Moles/Stoichiometry.

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III. Moles/Stoichiometry

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  1. III. Moles/Stoichiometry

  2. A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means (not physical means.) A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system. (3.1cc) Compounds are electrically neutral.

  3. Regents Question: 08/02 #6 Which species represents a chemical compound? (1) N2 (2) Na (3) NH4+ (4) NaHCO3 A compound is made up of 2 or more different elements. A compound is electrically neutral. þ

  4. A chemical formula is both qualitative and quantitative. It tells which elements are in the compound with symbols and how many of each with subscripts The formula for sulfuric acid is H2SO4 Element Number of atoms H 2 S 1 O 4 The subscript 1 is never written in a formula

  5. Compounds have more than one capital letter in the formula • Binary Compounds – Composed of just two elements (2 capital letters) • Water H2O 2 elements 3 atoms • Methane CH4 2 elements 5 atoms • Ammonia NH3 2 elements 4 atoms • Aluminum oxide Al2O3 2 elements 5 atoms • Other Compounds -more than 2 elements (more than 2 capital letters) • Potassium hydroxide KOH 3 elements 3 atoms • Calcium sulfate CaSO4 3 elements 6 atoms • Ammonium phosphate (NH4)3PO4 4 elements 20 atoms

  6. Regents Question: 06/03 #9 Which substance can be decomposed by a chemical change? (1) Co (2) CO (3) Cr (4) Cu þ

  7. Types of chemical formulas include empirical, molecular, and structural. (3.1ee)

  8. The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound. (3.3d) Hydrogen peroxide Sugar molecular formula: H2O2 C6H12O6 empirical formula: HO CH2O

  9. Regents Question: 06/03 #8 Which is an empirical formula? (1) P2O5 (2) P4O6 (3) C2H4 (4) C3H6 þ

  10. The molecular formula must be a whole number multiple of the empirical formula Molecular Empirical Multiple Formula Formula H2O2 HO 2 C3H6 CH2 3 C6H12O6 CH2O 6 Hg2O2 HgO 2

  11. The formula mass must be a whole number multiple of the empirical mass Molecular Empirical Multiple Formula Mass Formula Mass H2O2 34 HO 17 2 C3H6 42 CH2 14 3 C6H12O6 180 CH2O 30 6 Hg2O2 434 HgO 217 2

  12. To find the formula of a compound, use the criss-cross method • Write the symbol and oxidation state of the metal (lower electronegativiy if both are nonmetals) in the upper right hand corner of the symbol. Positive oxidation state is written first. • Repeat for nonmetal (-) C4+ O2- • Drop the sign C4 O2 • Criss Cross C2 O4 • Reduce if possible CO2 Although Carbon and Oxygen are both nonmetals, Carbon has a lower electronegativity and therefore uses a positive oxidation number

  13. When using the criss-cross method with polyatomic ions, use parenthesis ( ) • Write the symbol and oxidation state of the metal (+ ion) in the upper right hand corner of the symbol. • Repeat for nonmetal (- ion) (NH4) 1+ (PO4) 3- • Drop the sign (NH4) 1 (PO4) 3 • Criss Cross (NH4) 3 (PO4) 1 • Reduce if possible (numbers outside parenthesis only) • If there is no number following the parenthesis, drop the parenthesis. (NH4)3 PO4

  14. The oxidation states of the elements are available on the Periodic Table of the Elements

  15. Polyatomic ions are found on Reference Table E

  16. Naming Ionic Compounds • Compounds between a metal and a nonmetal • Name the metal • Name the nonmetal with an IDE ending • Ammonium chloride Aluminum oxide • Barium sulfide Zinc fluoride • If the negative ion is a polyatomic ion • Name the metal • Name the polyatomic ion • Sodium Nitrate Calcium phosphate

  17. The Stock System -Naming ionic compounds in which the metal has more than one oxidation state (polyvalent) • Check the Periodic Table to see if the metal has more than one oxidation state if it does… • Determine the oxidation state of the metal • Use the oxidation state of the nonmetal and the subscripts • The oxidation state of the metal is written as a roman numeral in parenthesis after the name of the metal • CuSO4 Copper(II) Sulfate

  18. Finding the oxidation state of a polyvalent metal • The net charge of a compound must be zero • The number of + charges must equal the total number of – charges • Multiply the oxidation state of the nonmetal by the subscript for the nonmetal • Divide that number by the subscript of the metal • PbO2 the oxidation state of oxygen is –2 we multiply the 2 by the subscript 2 and get 4. Divide the 4 by the subscript for the lead (1) and get 4. The oxidation state for the lead is 4 and the name is Lead(IV)oxide

  19. Regents Question: 06/03 #19 Which formula correctly represents the composition of iron (III) oxide? (1) FeO3 (2) Fe2O3 (3) Fe3O (4) Fe3O2 þ

  20. Naming molecular compounds – compounds between nonmetals • Use a prefix to indicate how many of each element mono 1 (never used for first element) di 2 tri 3 tetra 4 • CO2 Carbon dioxide • CO Carbon monoxide • N2O4 Dinitrogen tetraoxide • BCl3 Boron trichloride

  21. In all chemical reactions there is a conservation of mass, energy, and charge. (3.3a) • Matter cannot be created nor destroyed, only changed from one form to another • Energy cannot be created nor destroyed, only changed from one form to another

  22. A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. (3.3c) • A mole of molecules is made up of 6.02 x 1023 molecules • A chemical equation is balanced to ensure the conservation of matter. The types and number of atoms on each side of the equation (before and after the reaction) must be equal. • The total mass before the reaction must equal the total mass after the reaction has taken place.

  23. Chemical equations must be balanced so that mass can be conserved. Word equation: hydrogen + oxygen water Chemical equation: 2H2 + O2 2H2O

  24. 2H2 + O2 2H2O Regents Question: 08/02 #9 If an equation is balanced properly, both sides of the equation must have the same number of (1) atoms (2) coefficients (3) molecules (4) moles of molecules þ

  25. How many on the left? How many on the right? What do I do? • Balance the reaction Fe + O2 Fe2O3 • There are 2 oxygen atoms on the left and 3 on the right. To get them equal I need to multiply the left by 3 and the right by 2. These multiples are called coefficients and are placed in front of the formula. The yield sign and the plus separate the formulas. • Fe + 3O2 2Fe2O3 • Now the number of iron atoms has to be balanced. There is one iron on the left and four on the right. Use a coefficient of 4 in front of the Fe on the left. • 4Fe + 3O2 2Fe2O3 • Now it is correctly balanced

  26. Regents Question: 08/03 #47 • Given the unbalanced equation: • _Al + _CuSO4 _Al2(SO4)3 + _Cu • When the equation is balanced using the smallest whole-number coefficients, what is the coefficient of Al? • 1 • (2) 2 • (3) 3 • (4) 4 þ

  27. Regents Question: 01/03 #48 • Given the unbalanced equation: • _Fe2O3 + _CO _Fe + _CO2 • When the equation is correctly balanced using the smallest whole-number coefficients, what is the coefficient of CO? • 1 • 2 • 3 • 4 Try a 3 in front of CO2 þ

  28. Regents Question: 06/03 #59 Given the reaction between two different elements in the gaseous state Box A below represents a mixture of the two reactants before the reaction occurs. The product of this reaction is a gas. Draw the system after the reaction has gone to completion, based on the Law of Conservation of Matter. One reactant is in excess. Box B – System after reaction Box A – System before reaction

  29. Use the coefficients to predict the amount of reactant consumed or product formed(Mole-Mole problems) • The ratio of the coefficients is a ratio of the moles taking part in a reaction. • Reactants (on the left of the arrow) are consumed • Products (on the right of the arrow) are formed • Given the number of moles of any substance in a reaction, you can use the coefficients to find the number of moles of any other substance.

  30. Place the moles given over the coefficient for that substance and set up a proportion with the coefficients. unknown given • How many moles of oxygen are produced when 4 moles of KClO3 react according to the equation: 2 KClO3 2 KCl + 3 O2 4 X 2 KClO3 2 KCl + 3 O2 4 = X 2 3 cross multiply to get 2X = 12 X=6 moles

  31. Regents Question: 01/03 #42 • Given the equation: • 2C2H2(g) + 5O2(g) 4 CO2(g) + 2H2O(g) • How many moles of oxygen are required to react completely with 1.0 mole of C2H2? • 2.5 • 2.0 • 5.0 • 10 þ

  32. Regents Question: 06/02 #37 • Given the reaction: • 6CO2 +6H2O C6H12O6 + 6O2 • what is the total number of moles of water needed to make 2.5 moles of C6H12O6 ? • 2.5 • 6.0 • 12 • 15 þ

  33. Regents Question: 06/03 #20 Given the reaction: PbCl2 (aq) + Na2CrO4 (aq) PbCrO4 (s) + 2NaCl(aq) What is the total number of moles of NaCl formed when 2 moles of Na2CrO4 react completely? (1) 1 mole (2) 2 moles (3) 3 moles (4) 4 moles þ

  34. The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram formula mass) of a substance equals one mole of that substance. (3.3e) Formula mass of H2SO4 H 2 x 1 = 2 S 1 x 32 = 32 O 4 x 16 = 64 98 amu

  35. The formula mass represents the mass of one molecule of a substance while the gram formula mass represents the mass of a mole (6.02 x 1023 molecules) of that substance. • The calculation for formula mass and gram formula mass (GFM) are the same, the difference is in the units. • Formula Mass of H2SO4 = 98 amu • Gram Formula Mass of H2SO4 = 98 g If they tell you to use the mass rounded to the nearest tenth then use tenths for each mass as well as your final answer.

  36. Regents Question: 06/02 #34 • A compound has a gram formula mass of 56 grams per mole. What is the molecular formula for this compound? • CH2 • C2H4 • C3H6 • C4H8 þ

  37. Regents Question: 06/02 #41 • The gram formula mass of NH4Cl is • 22.4 g/mole • 28.0 g/mole • 53.5 g/mole • 95.5 g/mole N 1 x 14.0 = H 4 x 1.0 = Cl 1 x 35.5 = _______ g/mole þ

  38. Grams/GFM = MolesTo convert grams into moles, divide the grams by the gram formula massTo convert moles into grams, multiply by the moles by the gram formula massThe unit for gram formula mass is grams/mole • How many moles in 49 grams of H2SO4? 49 g = X moles X= 0.50 moles 98 g/mol • How many grams are contained in 2.00 moles of H2SO4? X = 2.00 moles X = 196 grams 98 g/mol

  39. The percent composition by mass of each element in a compound can be calculated mathematically. (3.3f) Calculate the formula mass and then divide the component of the element you are looking for by the total mass of the formula then multiply by 100%. Formula mass of H2SO4 % of Oxygen in H2SO4 H 2 x 1 = 2 64/98 x 100% = 65% S 1 x 32 = 32 O 4 x 16 = 64 98

  40. Regents Question: 01/03 #8 • What is the percent by mass of oxygen in H2SO4? [ formula mass = 98] • 16% • 33% • 65% • 98% 64/98 x 100 þ

  41. Regents Question: 06/02 #7 • The percent by mass of hydrogen in NH3 is equal to • 17 x 100 • 1 • (2) 17 x 100 • 3 • 1 x 100 • 17 • (4) 3 x 100 • 17 þ

  42. Regents Question: 06/03 #10 • The percent by mass of calcium in the compound calcium sulfate (CaSO4 ) is approximately • 15% • (2) 29% • (3) 34% • (4) 47% þ

  43. A hydrate is a compound which has water trapped in its crystal structure. We can determine the percentage of water in a hydrate. • CuSO45H2O – for every copper(II) sulfate there are 5 trapped water molecules • Find the formula mass including the water • Cu 1 x 64 = 64 • S 1 x 32 = 32 • O 4 x 16 = 64 • H 10 x 1 = 10 Water contributes 90 to • O 5 x 16 = 80 the 250 250 90/250 x 100% = 36%

  44. Regents Question: 06/02 #36 • What is the total number of oxygen atoms in the formula MgSO4•7H2O? [The • represents seven units of H2O attached to one unit of MgSO4] • 11 • 7 • 5 • 4 þ

  45. Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement. (3.2b) • Synthesis – one product from several reactants • Decomposition – several products from one reactant • Single replacement – an element and a compound form a different element and a different compound • Double replacement – two compounds for two different compounds by switching ions

  46. Examples of different types of reactions • Synthesis: A + B C • Na(s) + Cl2(g) 2NaCl(s) • Decomposition: C A + B • 2KClO3(s) 2KCl(s) + 3O2(g) • Single replacement: AB + C AC + B • Zn(s) + CuSO4(aq) ZnSO4(aq) + Cu(s) • Double replacement: AB + CD AD + BC • AgNO3(aq) + KCl(aq) KNO3(aq) + AgCl(s) Check Table F for insoluble compounds that form precipitates when solutions are mixed. Precipitate – not soluble in water

  47. Check Table F to see if there is a precipitate during a double replacement reaction. AgCl is a precipitate because it is insoluble in water.

  48. A single replacement reaction will take place when: • A more active metal replaces a less active metal from its compound • A more active nonmetal replaces a less active nonmetal from its compound • Criss Cross to find the new formula • Check Table J for the activity of metals and nonmetals. Zn +PbCl2 ZnCl2 + Pb Pb + ZnCl2 No Reaction Zinc is more active than lead

  49. Hydrogen is on Table J to show which metals will react with acids (HCl, H2SO4, HNO3) to produce hydrogen. Zn + 2HCl ZnCl2 + H2 A more active metal will replace a less active metal from its compound during a single replacement reaction. • Zn + CuSO4 Cu + ZnSO4 Zinc replaces copper because zinc is more active than copper. • Cu + ZnSO4 No Reaction Copper cannot replace zinc

  50. Regents Question: 01/03 #46 • According to Reference Table J, which of these metals will react most readily with 1.0 M HCl to produce H2(g)? • Ca • K • Mg • Zn þ

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