Chapter 4: TYPES OF CHEMICAL REACTIONS AND SOLUTION STOICHIOMETRY.
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Chapter 4: TYPES OF CHEMICAL REACTIONS AND SOLUTION STOICHIOMETRY
4–24–2 Aqueous Solutions Water is the dissolving medium, or solvent.
4–34–3 Some Properties of Water 1. Water is “bent” or V-shaped. 2. The O-H bonds are covalent. 3. Water is a polar molecule. 4. Hydration occurs when salts dissolve in water.
4–44–4 Figure 4.1: (Left) The water molecule is polar. (Right) A space-filling model of the water molecule.
4–54–5 Figure 4.2: Polar water molecules interact with the positive and negative ions of a salt assisting in the dissolving process.
4–64–6 Figure 4.3: (a) The ethanol molecule contains a polar O—H bond similar to those in the water molecule. (b) The polar water molecule interacts strongly with the polar O—H bond in ethanol. This is a case of "like dissolving like."
4–74–7 A Solute dissolves in water (or other “solvent”); changes phase (if different from the solvent); is present in lesser amount (if the same phase as the solvent). A Solvent retains its phase (if different from the solute); is present in greater amount (if the same phase as the solute).
4–84–8 Figure 4.4: Electrical conductivity of aqueous solutions.
4–94–9 Electrolytes Strong Electrolytes - conduct current efficiently NaCl, HNO 3 Weak Electrolytes - conduct only a small current vinegar, tap water Nonelectrolytes - no current flows pure water, sugar solution ※ Svante Arrhenius postulated: the extend to which a solution can conduct an electric current depends on the number of ions present.
4–10 Strong electrolytes: Substances that are completely ionized when they are dissolved in water. (1)soluble salts, (2) strong acids, (3) strong bases.
4–11 Figure 4.5: When solid NaCl dissolves, the Na + and Cl - ions are randomly dispersed in the water.
4–12 Figure 4.6: HCl(aq) is completely ionized. Arrhenius discoveries the nature of acids
4–13 ※ Arrhenius proposed that an acid is a substance that produces H +. HCl H + (aq) + OH - (aq) HNO 3 H + (aq) + NO 3 - (aq) H 2 SO 4 H + (aq) + HSO 4 - (aq)
4–14 Figure 4.7: An aqueous solution of sodium hydroxide. Strong bases - react completely with water to give OH - ions.
4–15 Figure 4.8: Acetic acid (HC 2 H 3 O 2 ) exists in water mostly as undissociated molecules. Only a small percentage of the molecules are ionized. Weak electrolytes:
4–16 Figure 4.9: The reaction of NH 3 in water. Weak bases - react only slightly with water to give OH - ions.
4–17 Molarity Molarity (M) = moles of solute per volume of solution in liters:
4–18 Common Terms of Solution Concentration Standard solution- concentration is exactly known. Stock solutions- routinely used solutions prepared in concentrated form. Concentrated solution- relatively large ratio of solute to solvent. (5.0 M NaCl) Diluted solution - relatively small ratio of solute to solvent. (0.01 M NaCl)
4–19 Figure 4.10: Steps involved in the preparation of a standard aqueous solution.
4–20 Figure 4.11: (a) A measuring pipet is graduated and can be used to measure various volumes of liquid accurately. (b) a volumetric (transfer) pipet is designed to measure one volume accurately.
4–21 Figure 4.12: Dilution Procedure (a) A measuring pipet is used to transfer 28.7mL of 17.4 M acetic acid solution to a volumetric flask. (b) Water is added to the flask to the calibration mark. (c) The resulting solution is 1.00 M acetic acid.
4–33 Types of Solution Reactions Precipitation reactions AgNO 3 (aq) + NaCl(aq) AgCl(s) + NaNO 3 (aq) Acid-base reactions NaOH(aq) + HCl(aq) NaCl(aq) + H 2 O(l) Oxidation-reduction reactions Fe 2 O 3 (s) + Al(s) Fe(l) + Al 2 O 3 (s)
4–34 Figure 4.13: When yellow aqueous potassium chromate is added to a colorless barium nitrate solution, yellow barium chromate precipitates.
4–35 Figure 4.14: Reactant Solutions: (a) Ba(NO 3 ) 2 (aq) and (b) K 2 CrO 4 (aq)
4–36 Figure 4.15a,b: The reaction of K 2 CrO 4 and Ba(NO 3 ) 2 (aq).
4–37 Figure 4.15c: The reaction of K 2 CrO 4 and Ba(NO 3 ) 2 (aq). (cont'd)
4–38 Figure 4.16: Precipitation of silver chloride by mixing solutions of silver nitrate and potassium chloride. The K+ and NO 3 - ions remain in solution.
4–39 Figure 4.17: The reaction of KCl(aq) with AgNO 3 to form AgCl(s).
4–40 Table 4.1 Simple Rules for Solubility of Salts in Water 1.Most nitrate (NO 3 ) salts are soluble. 2.Most alkali (group 1A) salts and NH 4 + are soluble. 3.Most Cl , Br , and I salts are soluble (NOT Ag +, Pb 2+, Hg 2 2+ ) 4.Most sulfate salts are soluble (NOT BaSO 4, PbSO 4, HgSO 4, CaSO 4 ) 5.Most OH salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH) 2, Ca(OH) 2 are marginally soluble) 6.Most S 2 , CO 3 2 , CrO 4 2 , PO 4 3 salts are only slightly soluble.
4–41 Describing Reactions in Solution 1. Molecular equation (reactants and products as compounds) AgNO 3 (aq) + NaCl(aq) AgCl(s) + NaNO 3 (aq) 2. Complete ionic equation (all strong electrolytes shown as ions) Ag + (aq) + NO 3 (aq) + Na + (aq) + Cl (aq) AgCl(s) + Na + (aq) + NO 3 (aq) 3. Net ionic equation (show only components that actually react) Ag + (aq) + Cl (aq) AgCl(s) Na + and NO 3 are spectator ions.
4–42 Stoichiometry Steps for reactions in solution.
4–43 Performing calculations for acid-base reactions.
4–44 Key Titration Terms Titrant - solution of known concentration used in titration Analyte - substance being analyzed Equivalence point - enough titrant added to react exactly with the analyte (Stoichiometric point) Endpoint - the indicator changes color so you can tell the equivalence point has been reached. Indicator- a color substance with its color change to mark the endpoint of titration.
4–45 Figure 4.18: The titration of an acid with a base.
4–46 Figure 4.19: The reaction of solid sodium and gaseous chlorine to form solid sodium chloride.
4–47 Sum of oxidation states = 0 in compounds Sum of oxidation states = charge of the ion
4–48 Figure 4.20: A summary of an oxidation-reduction process, in which M is oxidized and X is reduced.
4–49 Balancing by Half-Reaction Method 1. Write separate reduction, oxidation reactions. 2.For each half-reaction: Balance elements (except H, O) Balance O using H 2 O Balance H using H + Balance charge using electrons 3. If necessary, multiply by integer to equalize electron count. 4.Add half-reactions. 5.Check that elements and charges are balanced.
4–50 Balancing Oxidation-Reduction Reactions Cr 2 O 7 2- (aq) + SO 3 - (aq) Cr 3+ (aq) + SO 4 2- (aq) How can we balance this equation?
4–51 Method of Half Reactions Cr 2 O 7 2- (aq) 2Cr 3+ (aq) SO 3 - (aq) SO 4 2- (aq) How many electrons are involved in each half reaction?
4–52 6e - + Cr 2 O 7 2- (aq) 2Cr 3+ (aq) SO 3 - (aq) SO 4 2- (aq) + 2e - How can we balance the oxygen atoms? Method of Half Reactions (cont.)
4–53 6e - + Cr 2 O 7 2- (aq) Cr 3+ (aq) + 7H 2 O H 2 O + SO 3 - (aq) SO 4 2- (aq) + 2e - How can we balance the hydrogen atoms? Method of Half Reactions (cont.)
4–54 This reaction occurs in an acidic solution. 14H + + 6e - + Cr 2 O 7 2- (aq) Cr 3+ (aq) + 7H 2 O H 2 O +SO 3 - (aq) SO 4 2- (aq) + 2e - + 2H + How can we balance the electrons? Method of Half Reactions (cont.)
4–55 14H + + 6e - + Cr 2 O 7 2- (aq) Cr 3+ (aq) + 7H 2 O 3[H 2 O +SO 3 - (aq) SO 4 2- (aq) + 2e - + 2H + ] Cr 2 O 7 2- (aq) + 3SO 3 - (aq) + 8H + (aq) 2Cr 3+ (aq) + 3SO 4 2- (aq) + 4H 2 O(l) Method of Half Reactions (cont.)
4–56 Half-Reaction Method - Balancing in Base 1.Balance as in acid. 2.Add OH that equals H + ions (both sides!) 3.Form water by combining H +, OH . 4.Check elements and charges for balance.