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CHE 111 - Module 4

0. CHE 111 - Module 4. CHAPTER 4 & 5 LECTURE NOTES. 0. Stoichiometry & Balancing Equations. Remember we stated in the previous chapter that stoichiometry is the study of the quantitative relationships between the amounts of reactants and products in chemical reactions.

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CHE 111 - Module 4

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  1. 0 CHE 111 - Module 4 CHAPTER 4 & 5 LECTURE NOTES

  2. 0 Stoichiometry & Balancing Equations • Remember we stated in the previous chapter that stoichiometry is the study of the quantitative relationships between the amounts of reactants and products in chemical reactions. • We use BALANCED equations to understand stoichiometric relationships of the elements and compounds within a chemical reaction.

  3. 0 The Balanced Equation 2Al(s) + 3Br2(l) Al2Br6(s) 2mol of Al : 3mol of Br2 : 1mol of Al2Br6 2 atoms of Al = 2 atoms of Al 6 atoms of Br = 6 atoms of Br The number of the same atom of each element must be equal on each side of the equation.

  4. 0 A Closer Look at the Equation 2Al(s) + 3Br2(l) Al2Br6(s) • The chemicals on the left are the reactants and the right are the products. • The coefficient in front of the chemical denotes the stoichiometric relationship. • The numerical subscriptrepresents the number of atoms present in the molecule. • The letter subscripted denotes the phase of matter.

  5. For example the following is balanced. CH4 + 2O2 CO2 + 2H2O Try to balance the following: Fe2S3 + O2 Fe2O3 + S Al+ H2SO4 Al2(SO4)3 + H2 Ca + Al2Br6 CaBr2 + Al 0 Balancing Equations

  6. Balanced Equations Check your answer from the previous slide: • 2Fe2S3 + 3O2 2Fe2O3 + 6S • 2Al+ 3H2SO4 Al2(SO4)3 + 3H2 • 3Ca + Al2Br6 3CaBr2 + 2Al

  7. 0 Types of Reactions • Combination Reactions • Decomposition Reactions • Displacement (Single-Replacement) Reactions • Metathesis (Double-Replacement) Reactions • Combustion Reactions

  8. 0 Combination Reactions • A combination reaction is a reaction where two substances chemically combine to form another substance. A + B AB 2Na(s) + Cl2(g) 2NaCl(s) P4(s) + 6Cl2(g) 4PCl3(s)

  9. 0 Decomposition Reaction • A decomposition reaction is when a single compound decomposes into two or more other substances. AB A + B 2KClO3(s) 2KCl(s) + 3O2(g) 2NO2(g) 2NO(g) + O2(g)

  10. 0 Displacement Reaction • A displacement (single replacement) reaction is a reaction where one element displaces another element. A + BC B + AC Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g) Cu(s) + 2AgNO3(aq) Cu(NO3)2(aq) + 2Ag(s)

  11. 0 Metathesis Reaction • A metathesis (double replacement) reactionis a reaction where two compounds switch cations to form two new compounds. A+B- + C+D- A+D- + C+B- CaCl2(aq) + Na2CO3(aq) CaCO3(s) + 2NaCl(aq) AgNO3(aq) + KCl(aq) ???

  12. 0 SOLUBILITY • Solubility – the amount of a substance that can be dissolved in a given quantity of solvent (like water) at a specific temperature • Unsaturated – amount of substance less than saturated • Saturated – the exact amount at solubility • Supersaturated – excess amount of substance

  13. 0 How Solubility Influences Rxn • When a substance is soluble in water, it will appear with a subscript of (aq) meaning that the substance is broken up into it’s ions incorporated into the water lattice. • When a substance is insoluble in water, it will be written with a subscript of (s), (l), or (g) and will precipitate out of solution.

  14. Solubility of Ionic Compounds in Water Soluble CompoundsExceptions

  15. Insolubility of Ionic Compounds in Water Insoluble CompoundsExceptions

  16. 0 A Look at Metathesis Again • Looking back at slide 10 to the first reaction: when the cations rearranged, the CaCO3 being insoluble by our definition is recorded as CaCO3 (s). The CaCO3 would precipitate out of solution as a solid. • Looking at AgNO3(aq) + KCl(aq) ?, we can rearrange the cations and conclude that the AgCl is a solid and will precipitate out of solution

  17. 0 Types of Metathesis Reactions • Three classifications of metathesis reactions • Precipitation reaction - formation of a solid Pb(NO3)2(aq) + Na2CO3(aq) PbCO3(s) + 2NaNO3(aq) • Neutralization reaction - formation of water HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) • Gas formation reaction - CO2, H2S, SOx, & NOx are typically formed

  18. 0 Precipitation Reactions • A solid precipitate is produced in the rearrangement of cations as follows: Pb(NO3)2(aq) + Na2CO3(aq) PbCO3(s) + 2NaNO3(aq) • The Ionic Equation is expressed as: Pb+2 + 2NO3- + 2Na+ + CO3-2 PbCO3(s) + 2Na+ + 2NO3- • After neglecting the spectator ions, the net ionic equation will look like: Pb+2(aq) + CO3-2(aq) PbCO3(s)

  19. 0 Reviewing Ionic Compounds Ca+2 + 2Cl- CaCl2 Each ion comes together based on charge to form an overall neutral ionic compound. 3Ca+2 + 2PO4-3 Ca3(PO4)2 The cation and the polyatomic ion come together based on charge to form an overall neutral ionic compound.

  20. 0 Net Ionic Equations (NIE) • If you were given the reactants Ca(NO3)2 and Na3PO4 you should be able to predict the precipitate and write a balance equation, the ionic equation, and the net ionic equation (NIE) for this reaction. • The NIE for these reactants is as follows: 3Ca+2(aq) + 2PO4-3(aq) Ca3(PO4)2(s)

  21. carbonate ion CO3-2 sulfate ion SO4-2 sulfite ion SO3-2 hydroxide OH- phosphate PO4-3 permanganate MnO4- chromate CrO4-2 dichromate Cr2O7-2 ammonium NH4+ oxalate C2O4-2 bicarbonate HCO3- cyanide ion CN- acetate C2H3O3- 0 Common Polyatomic Ions

  22. 0 Neutralization Reaction • A neutralization reaction is a reaction that occurs between an acid and a base with the production of a salt and water. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) acid base salt water

  23. 0 Gas Formation Reaction • A gas formationreaction is a metathesis reaction that generates a gas as a product. • Metal carbonates or bicarbonates + acid • Metal sulfides + acid • Metal sulfites + acid • Ammonium salts and strong base

  24. 0 Metal Carbonates • Metal carbonates or bicarbonates when combined with an acid form salt, water and carbon dioxide gas. Na2CO3(aq)+ 2HCl(aq) 2NaCl(aq)+ H2O(l)+ CO2(g) • Where CO2 gas is given off

  25. 0 Metal Sulfides • Metal sulfides when combined with an acid form salt and hydrogen sulfide gas. Na2S(aq)+ 2HCl(aq) 2NaCl(aq)+ H2S(g) • Where H2S gas is given off

  26. 0 Metal Sulfites • Metal sulfites when combined with an acid form salt, water, and sulfur dioxide gas. Na2SO3(aq)+ 2HCl(aq) 2NaCl(aq)+ H2O(l)+ SO2(g) • Where SO2 gas is given off

  27. 0 Ammonium Salts • Ammonium salts when combined with a base produce salt, water and ammonia. NH4Cl (aq)+ NaOH(aq) NaCl(aq)+ H2O(l)+ NH3(g) • Where ammonia gas is given off

  28. 0 Combustion Reactions • A combustion reaction is a reaction with molecular oxygen to form products in which all elements are combined with oxygen. CH4 + 2O2  CO2 + 2H2O

  29. 0 Limiting Reactants • One of the reactants is in limited supply and thus restricts the amount of product formed. • Think of it as: If you wanted to bake a batch of peanut butter cookies and the recipe calls for 1 cup of peanut butter and all you have is ½ a cup, even though you have all the other ingredients, you can at most make ½ a batch of cookies.

  30. 0 Limiting Reactants (cont.) • Consider the combustion reaction: CH4 + 2O2 CO2 + 2H2O • How much CO2 can be produced if you have 0.13g of methane and 0.45g of O2?

  31. 0 Percent Yield • The maximum amount of product that can be obtained from a chemical reaction is the theoretical yield. • The actual amount produced in a chemical process is the actual yield. • The percent yield is equal to the actual yield divided by the theoretical yield times 100%.

  32. 0 Redox Reactions • Oxidation of an element takes place when electrons are lost from the valence shell of the element. • Reduction of an element takes place when electrons are added to the valence shell of the element. • Redox reactions show the transfer of electrons that takes place during oxidation and reduction.

  33. 0 Redox Reactions (cont.) • All oxidation and reduction reactions involve transfer of electrons between substances. • View CD-ROM screen 5.12 • Ag+ accepts electrons for Cu and is reduced to Ag and Cu loses electrons to Ag+ and is oxidized to Cu+2 in the following redox rxn: 2Ag+(aq) + Cu(s) 2Ag(s) + Cu+2(aq)

  34. 0 Redox Reactions (cont.) • The oxidation half reaction is : Cu(s) – 2e- Cu+2(aq) • The reduction half reaction is: 2Ag+(aq) + 2e- 2Ag(s) • Cu is called the reducing agent because it caused Ag+ to be reduced; and Ag+ is called the oxidizing agent because it caused Cu to be oxidized.

  35. 0 Determining Oxidation Numbers • Each atom in a pure element has an oxidation number of zero. • For monoatomic ions, the ox. number is equal to it’s ionic charge. • F is always –1, other halogens are –1 as well except with oxygen or fluorine. • The ox. number for H is +1 except with hydrides (CaH2).and O is –2 except with peroxides (H2O2). • The  ox.# must = 0 for a compound or = to the overall charge of polyatomic ion being considered.

  36. 0 Balancing Redox Reactions • We can use the balance of electrons transferred in a redox reaction to help us balance the overall equation. • Consider the unbalanced equation: Zn(s) + HCl(aq) ZnCl2(aq) + H2(g) • The balanced equation takes into consideration the oxidation of the Zn and the reduction of the H+. • Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

  37. 0 Molarity Molarity = Moles of Solute Liters of Total Solution • Symbol for molarity is M • Units are moles/Liter

  38. Solution Preparation • To prepare a 1.0M solution of NaCl, you would determine how many grams of NaCl is contained in 1.0 moles of NaCl and then dissolve that amount in a 1.0L volumetric flask. You would then qs with distilled H2O. • 1.0M NaCl = 1 mole (or 58.44g) NaCl 1.0L of solution • How much NaCl would you use to make a 0.1M solution of NaCl? As 1/10 of a mole = 5.844g NaCl, you would dissolve 0.1mole (5.844g) of NaCl in 1.0L of solution.

  39. Acids An acid is defined as follows: • Arrhenius – releases H+ when dissolved in H2O • Bronsted-Lowrey – a substance that can donate a proton to another substance • Lewis – a substance that can accept a pair of electrons from another atom to form a new bond

  40. Bases A base is defined as follows: • Arrhenius – releases OH- when dissolved in H2O • Bronsted-Lowrey – a substance that can accept a proton from another substance • Lewis – a substance that can donate a pair of electrons to another atom to form a new bond

  41. pH and Concentrations of Acids and Bases pH = -log [H+] 1 – acidic – 7 – basic – 14 When dealing with [H+] less than 0.1M (pH=1), we use activity coefficients instead of pH.

  42. pH of Household Items • pH of vinegar = 2.80 • pH of soda = 2.90 • pH of orange juice = 3.80 • pH of pure water = 7.00 • pH of blood = 7.40 • pH of ammonia = 11.00 • pH of oven cleaner = 11.7

  43. Titration • A method for quantitative analysis of a substance by essentially complete reaction in solution with a measured quantity of a reagent of known concentration. • Often used in redox reactions • Many redox reactions go rapidly to completion in aqueous media to determine the equivalency point. • Typically used for neutralization reactions. • Acid is titrated with a base using an indicator to determine the equivalency point of the neutralization reactions.

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