Ch 4: Types of Chemical Reactions and Solutions Stoichiometry Do Now: 1. True or false: a. Most ionic substances need an aqueous medium in order to react. b. Reactions occur at the level of atoms and molecules, therefore count of each reactant atoms/molecules/ions needs to be known in order to determine how much product can be produced. 2. If the above statements are true a. How can you determine the amount of particles involved in a reaction if you have 28 mL of NaCl(aq)
b. If you know that 50 g of CaS is dissolved in 250 mL of water, how would you determine the volume to use if you needed 0.21 mol of CaS for a reaction? 4.1 Water, the Common Solvent A. Structure of water 1. Oxygen's electronegativity is high (3.5) and hydrogen's is low (2.1) 2. Water is a bent molecule (shape) 3. Water is a polar molecule (charge differential) B. Hydration of Ionic Solutes 1. Positive ions attracted to the oxygen end of water 2. Negative ions attracted to the hydrogen end of water 3. Particle dissociates (or separate) into ions.
a. Level of dissociation depends on Ksp (solubility product constant) value. b. HIGH (+) Ksp value (soluble) Salt, CuCl 2, Ba(NO3)2 c. LOW Ksp value (insoluble) BaSO4, Ca(OH)2 C. Hydration of Polar Solute Molecules (Covalent) 1. Negative end of polar solute molecules are attracted to water's hydrogen 2. Positive end of polar solute molecules are attracted to water's oxygen 3. Particle stays together. Ex: sugar, alcohol D. "Like Dissolves Like" 1. Polar and ionic compounds dissolve in polar solvents like water
2. Nonpolar compounds like fats dissolve in nonpolar solvents like acetone NOTICE: DAILY QUIZ From Today Until the Time that Your Class Learns How to Behave and Be Considerate Towards Each Other, You Will Take a Daily Quiz Starting at the Bell Timed at 4-6 minutes Any talking or commenting during quiz will earn you an automatic ZERO No time extensions DO NOT START UNTIL TOLD TO DO SO.
4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes A. Definition of Electrolytes 1. A substance that when dissolved in water produces a solution that can conduct an electric current B. Strong electrolytes conduct current very efficiently 1. Completely ionized when dissolved in water a. Soluble Ionic compounds b. Strong acids (HNO3(aq), H2SO4(aq), HCl(aq)) c. Strong bases (KOH, NaOH)
C. Weak electrolytes conduct only a small current 1. Slightly ionized in solution a. Weak acids (organic acids - acetic, citric, butyric, malic) b. Weak bases (ammonia) D. Nonelectrolytes conduct no current 1. No ions present in solution a. alcohols, sugars 4.3 The Composition of Solutions A. Molarity 1. Moles of solute per liter of solution
M = molarity = moles of solute/liters of solution 2. Steps a. Convert Mass/volume/particles into moles b. Convert volume (may be based on mass and density) of solution into Liters. c. Divide by Liter of solution (NOT volume of water used) Example: 3. Calculate the molarity of a solution prepared by dissolving 11.5 g of solid NaOH in enough water to make a 1.50L of solution 4. Calculate the molarity of a solution prepared by dissolving 0.959 L of gaseous HCl in enough water to make 26.8 mL of solution.
B. Concentration of Ions in Solution 1. Ionic compounds dissociate in solution 2. multiplying the molarity by the number of ions present indicates concentration of ions** a. **Individuals IONS react, not complete formula units. Example 1. Give the concentration of each type of ion in the following examples a. 0.50 M Co(NO3)2 b. 1 M Fe(ClO4)3 C. Moles from Concentration 1. Liters of solution x molarity = moles of
solute Example: 1. Calculate the number of moles of Cl- ions in 1.75 L of 1.0 X 10-3 M ZnCl2 2. Typical blood serum is about 0.14 M NaCl. What volume of blood contains 1.0 mg NaCl D. Solutions of Known Concentration 1. Standard solution - a solution whose concentration is accurately known 2. Preparation of Standard solutions How much x How strong x What does it weigh? L x mol/L x g/mol = grams required to prepare the standard
3. Volumetric flask-used to prepare single, specific volume of known molarity. a. Determine mass of solute b. Mix with small amount of distilled water in beaker (or directly into volumetric flask **difficult) c. Transfer to volumetric flask (funnel) d. Rinse beaker multiple times with small amounts of distilled water and transfer to volumetric flask (ensures all particles of solute are transferred. (skip if solute added to flask directly) e. Rinse funnel into volumetric flask f. Add sufficient water to rise ALMOST to level of calibration mark (meniscus). g. Allow any water/solution clinging to neck to settle down. h. Add last several mL of distilled water dropwise until calibration mark (meniscus). Example:
1. To analyze the alcohol content of a certain wine, a chemist needs 1.00L of aqueous 0.200 M K2Cr2O7 solution. How much solid must be weighed out to make this solution? E. Dilution 1. Dilution of a volume of solution with water does not change the number of moles present 2. Solving dilution problems M1V1 = M2V2 Example: 1. What volume of 16M sulfuric acid must be used to prepare 1.5 L of a 0.10 M H2SO4 solution?
Homework Page 180-181 Questions (9, 12, 15, 17,) (19, 21, 23 a, b, 25 ) DO NOW Take out HW, then do the practice problems below. Write the dissociation equation for the dissolution of Ammonium phosphate in water. If 543.0 g of the above compound is dissolved to form 345 mL of solution
determine the molarity of the solution. Determine the concentration of each type of ion dissolved in solution. Topic: Types of Chemical Reactions AIM: To differentiate, predict products and solve stoichiometry problems for three different categories of reactions. Consider a 1.00L sample each of two solutions of 1.00 M Potassium chromate and 2.00 M Barium nitrate. 1. Write the dissociation equation for each. 2. Draw a particulate diagram for each solution (separate) 3. Write a balanced chemical equation for the reaction between Potassium chromate and Barium nitrate.
4. Which, if any, is the limiting reactant? Justify 5. Draw a particulate diagram for the final reaction mixture. HW: Textbook Page 182 # 29, 31, 33, 37, 40, 42, 43 4.4: Type of Chemical Reactions (Rxn) Millions of chemical reactions Can be categorized into 3 primary groups Precipitation reactions Acid-Base reactions Oxidation-Reduction (Redox) reactions
Figure 4.14: Reactant Solutions: (a) Ba(NO3)2(aq) and (b) K2CrO4(aq) 4.5: Precipitation Reactions A. Precipitation Reactions 1. A reaction in which a solid (precipitate) is formed. B. Process 2. Dissociation a. Ionic compounds dissolve in water and the ions separate and move independently Ex 1: AgNO3(aq) + NaCl(aq) products 3. Determination of Products
a. Recombination of ions b. Elimination of reactants as products c. Identifying the precipitate 4. Determine whether reaction takes place (solubility rules). a. Dependent on ability to separate products successfully b. If there are NO insoluble products, reaction has NOT taken place. c. If both products are insoluble, reaction has NOT taken place. 3a. Solubility Rules Practice Set 1: Justify whether each of these reactions will occur. 1.
Ag2SO4 + CaCl2 6. KCl + AgNO3 4.6 Describing Reactions in Solution The Molecular Equation A. 1. Gives the overall reaction stoichiometry, not necessarily the actual forms of reactants and products in solution. Example: Na2(CO3) + Ca(NO3)2 ????
i. B. Dont EVER give this as your answer. The Complete Ionic Equation 1. Represents as ions all reactants and products that are strong electrolytes C. The Net Ionic Equation *** 1. Includes only those components that take part in the chemical change
2. Does not indicate spectator ions. i. D. ***ONLY valid form of representing reaction on AP exam. ALL species MUST be written AS their prescribed behavior in waterions or formula units/molecules. Practice: Write the COMPLETE and the NET ionic reaction 1. 2.
Aqueous potassium chloride is added to aqueous silver nitrate to form a silver chloride precipitate and aqueous potassium nitrate Aqueous potassium hydroxide is mixed with aqueous iron (III) nitrate to form a precipitate of iron (III) hydroxide and aqueous potassium nitrate. 3. Lithium iodide solution is mixed with ammonium sulfate solution. 4. Calcium carbonate and Barium hydroxide are added to water. 4.7: Stoichiometry of
Precipitation Reactions Practice 1. Calculate the mass of PbSO4(s) formed when 1.25 L of 0.0500 M Pb(NO3)2 and 2.00 L of 0.0250 M Na2SO4 are mixed. a) Based on question 2, determine the moles of excess reactant left over. b) Determine the moles of each ion in solution AFTER the reaction is complete. c) Determine the concentration of each ion at the end of the reaction.
Strategy Is this a limiting reactant problem? How do you know? What conversions are required? Do we know how to make the conversion? Will ALL Limiting Reactant ions be used up? Why or why not? What portion of Excess Reactant ions will be left over? Both ions in equal concentrations? Why or why not? What does concentration (Molarity) depend on?
Practice 2. Determine the mass of the precipitate formed with 50. mL of 1.75 M Na2(CO3) is mixed with 75 mL of 2.5 M Ca(NO3)2 a. Based on question 2, determine the moles of excess reactant left over. b. Determine the moles of each ion in solution AFTER the reaction is complete. c. Determine the concentration of each ion at the end of the reaction. 4.8 Acid-Base Reactions AIM: 1. To understand the different concepts of an ACID vs. BASE
2. To understand the net ionic reaction for acidbase reactions 3. To learn how to titrate to achieve neutralization reaction. DO NOW: 4. How do you recognize an acid from formula? A base from formula? 5. Identify the products of the following reaction. Determine whether reaction will take place. Sulfuric acid reacts with Sodium hydroxide A. Defining Acids vs. Bases 2 Definitions 1. Arrhenius
a. Acids: substance that produces H+ ions in solution i. Strong acids: virtually every molecule ionizes -Example: Hydrochloric, Nitric, Sulfuric, Perchloric Acid b. Base: substance that produces OH- ions in solution i. 2. Strong bases: Sodium hydroxide, Lithium hydroxide
Bronsted-Lowry: a. Acid: substance that is a proton donor b. Base: substance that is a proton acceptor (discuss in detail in ch. 14) List of Strong acids and Bases Strong acidscompletely ionized in water to give one or more H+ (protons) per acid molecule HI HBr
HClO4 HCl HClO3 H2SO4 H+(aq) + HSO4-(aq) (HSO4- is a weak acid that contributes additional protons) Strong Basescompletely ionized in water to give one or more OH- ions per base molecule NaOH KOH
LiOH RbOH CsOH Group 2 hydroxidesnot fully soluble, but forms strong bases. Ca(OH)2 Ba(OH)2 Sr(OH)2 HNO3 BONUS QUESTION: that Mg(OH)2 is missing from the list of Everything elseNotice
is a WEAK strong Yet,dissociate Mg(OH)2 isfully) the active ingredient of many acidbases. (doesnt antacids? Is this a good choice for an antacid? Why or why not? How is it able to neutralize stomach acid (HCl)? B. Acid-Base Reactions Net ionic equation for STRONG acid-STRONG base reactions 1. H+(aq) + OH-(aq) H2O(l) The hydroxide ion can be assumed to completely
react with even a weak acid in solution Ex: OH- (aq) + HC2H3O2(aq) H2O(l) + C2H3O2-(aq) 2. a) i. b) i. Notice Acetic Acid shown as molecule, not ions. Why? Weak Acids and Bases must be written as molecular formulas because they DO NOT dissociate fully. **write species as their
prescribed behavior in water. Also true for weak base with hydrogen ion. Practice: write neutralization reaction between ammonium hydroxide and hydrochloric acid. Often called NEUTRALIZATION reaction 3. a) Why? Practice Determine whether the following reactions are Acid-Base
neutralization? If yes, predict product. 1. HCl + NaOH 2. HNO3 + CaSO4 3. Ca(OH)2 + C6H12O6 4. H3PO4 + KOH C. Stoichiometry Calculations for
Acid-Base Reactions 1. List the species present in the combined solution before any reaction occurs; decide what reaction will occur 2. Write the balanced net ionic equation for this reaction 3. Calculate the moles or reactants a. For reactions in solution, use volumes of the original solutions and their molarities 4. Determine the limiting reactant where appropriate 5. Calculate the moles of the required reactant or product 6. Convert to grams or volume of solution as Practice Problem 1
What volume of a 0.100 M HCl solution is needed to neutralize 25.0 mL of 0.350 M NaOH What volume of a 0.100 M HCl solution is needed to neutralize 25.0 mL of 0.350 M NaOH Strategy? Determine Determine Determine Determine Molarity
moles of OH- present. moles of H+ needed. H+ to Acid ratio volume based on given Practice Problem 2 In a certain experiment, 28.0 mL of a 0.250 M HNO3 and 53.0 mL of 0.320 M KOH are mixed. a. Calculate the volume of water formed in the resulting reaction. b. What is the concentration of H+ or OH- ions in excess after the reaction goes to completion. c. Is the resulting solution acidic, basic or neutral. Explain. d. What is the molarity of the solution in terms of the
i. resulting salt? ii. Excess H+ or OH- ions? In a certain experiment, 28.0 mL of a 0.250 M HNO3 and 53.0 mL of 0.320 M KOH are mixed. a. Calculate the volume of water formed in the resulting reaction. b. What is the moles of H+ or OH- ions in excess after the reaction goes to completion. c. Is the resulting solution acidic, basic or neutral.
Explain. d. What is the molarity of the solution in terms of the i. resulting salt? ii. Excess H+ or OH- ions? Strategy Is it a Limiting Reactant problem? Determine moles of one reactant needed with respect to the other. Determine moles water formed Determine mass, then volume (DH2O = 1.0 g/mL)
Not a gas, cannot use 22.4 L/mol Determine excess ion amount used and remaining. Determine acid/base/neutral depending on excess ion remaining Determine molarity of salt formed and H+ or OH- ion. Dont forget to ADD the volume of water formed from reaction to total volume from
aqueous solutions. Is the volume of water formed significant? D. Acid-Base Titrations 1. Vocabulary a. Titrant - Solution of known concentration b. Analyte - Solution of unknown concentration c. Equivalence point - Point at which the amount of titrant added to analyte results in perfect neutralization d. Indicator - a substance added at the beginning of the titration that changes color at the equivalence point e. Endpoint - the point at which the indicator changes color
2. Requirements for a successful titration a. the exact reaction between titrant and analyte must be known b. the reaction must proceed rapidly c. the equivalence point must be marked accurately (select the appropriate indicator) d. the volume of titrant required to reach the equivalence point must be known accurately e. for acid-base titrations, the titrant should be a strong acid or a strong base Video 1 Video 2
4.9-4.10: Redox Reactions AIM: 1. To differentiate between reduction and oxidation of elements 2. To learn how to balance Redox reactions. DO NOW: 3. Identify the oxidation state of oxygen, Aluminum, Carbon, Hydrogen and Nitrogen. ---in elemental form? ---in compound form? 4.9: Oxidation-Reduction Reactions 1. Reactions in which one or more electrons
are transferred are known as oxidationreduction or REDOX reactions a. Example: 2Na(S) + Cl2(g) 2NaCl(s) b. Most reactions used for ENERGY PRODUCTION (photosynthesis, combustion, cellular respiration) are redox reactions c. Apparent when products or reactants are ionic OR both elements and compounds as reactants / products d. Covalent compounds can also be represented by assumed or actual electron transfers. 2. Oxidation States a. b.
c. AKA oxidation numbers Helps keep track of electrons in redox reactions ** especially with covalent compounds History/Rules i. in covalent compounds, oxidation numbers for elements are based on arbitrary assignment given to certain elements. ii. In partial sharing of electrons, element with stronger affinity is assumed to have all the electrons. d. e.
Ionic compounds: oxidation number is equal to charge of ion. In electrically neutral compounds, sum of oxidation numbers must equal ZERO. In ions, the sum must equal the charge on the ion. Table 4.2: Rules for Assigning Oxidation States ** MEMORIZE 3. Examples of oxidation numbers Determine oxidation number for each atom in the following 1. CO2 2.
SF6 3. Fe3O4 4. NO3- 5. PO4-3 6.
NH4+1 4. Characteristics of Redox Reactions Characterized by TRANSFER of electrons b. Ionic compoundsliteral transfer Example: 2Na(s) + Cl2(g) NaCl(s) c. Covalent compoundspresumed transfer Example: CH4(g) + 2O2(g) CO2(g) + 2H2O(g) a.
5. Oxidation VS. Reduction a. Oxidation: an increase in the oxidation state (a loss of electrons) Example: 2Na 2Na+1 b. Reduction: A decrease in oxidation state (a gain of electrons) Example Cl2 2Cl-1 Practice: Determine oxidation states of each element in the following reactions. Determine which elements are oxidized and which are reduced 1. 2Al(s) + 3I2(s) 2AlI3(s) 2.
2PbS(s) + 3O2(g) 2PbO(s) + 2SO2(g) 3. PbO(s) + CO(g) Pb(s) + CO2(g) 4.10: Balancing REDOX Reactions Useful to separate into 2 half reactions For oxidation and reduction Ex: Ce4+ + Sn2+ Ce3+ + Sn4+ Balance
each half reaction separately Equalize number of electrons transferred. **differs for acidic vs. basic media Add up half reactions to obtain full balanced reaction 2 + Acidic Media 2 + Add OH- to both sides of equation
(equal to H+) Form H2O on the side containing H+ and OH- ions Eliminate number of H2O appearing on both sides Check that elements and charges are balanced BASIC MEDIA
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