Chapter 4 Student Notes PHW

Chapter 4 Student - Chapter 4 Objectives 4.1 Discuss the idea of a chemical solution and the components of a solution Emphasize that water is the

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Unformatted text preview: Chapter 4: Objectives 4.1. Discuss the idea of a chemical solution and the components of a solution. Emphasize that water is the most common solvent in a solution. 4.2. Introduce molarity (M) as a calculation of a way describing the the atomic/molecular composition of a solution. 4.3. Emphasize acid-base, precipitation, and redox reactions as common acidsolution chemistry reactions. 4.4. Include examples of solution stoichiometry calculations that will be will emphasized throughout the semester in both lecture and laboratory. laboratory. 8/31/2009 Zumdahl Chapter 4 1 Water, the Most Common Solvent Hydration or Dissolution: Two (or more) substances spread out, or disperse, into each other at the level of individual atoms, molecules, or ions. Solution: Solvent: Solute(s): Solute(s): In general, the solute and solvent can be any combination of solid (s), liquid (l), and solid (s (l gaseous (g) phases. However, we will be primarily interested in liquid solutions. (g solutions. 8/31/2009 Zumdahl Chapter 4 2 1 Ionic Compounds in Water All ionic compounds are solids (salts) in the range of temperature in which water salts) is a liquid. They have rigid lattices in which strong (Coulomb) forces hold the constituent ions in place. The high melting points of ionic compounds indicate that a good deal of energy must be supplied to destroy the lattice and produce a liquid (molten) form in which (molten) the ions move more freely. Water is a polar molecule. When water interacts with an ionic compound such as NH4NO3, the ionic compound dissociates into a positive ion (NH4+, a cation) and a negative ion (NO3−, an anion). cation) anion). 8/31/2009 Zumdahl Chapter 4 3 Ionic Compounds in Water NH4NO3(s) + H2O(l) → NH4+(aq) + NO3–(aq) 8/31/2009 Solvation of the ions by water molecules provides the Zumdahl Chapter 4 energy gain needed to break apart the crystal lattice. 4 2 Dissolution of One Liquid in Another Liquid Pairs of liquids that mix in any proportion are termed miscible. miscible. Liquids that do not mix in any proportion are termed immiscible. immiscible. Liquids that are partially miscible mix in some proportions but not in others. "Like Dissolves Like" Substances with similar intermolecular attractive forces tend to be soluble in one Another. ▬► Polar liquids tend to be miscible in other polar liquids & non-polar nonliquids tend to be miscible in other non-polar liquids. nonExample: Ethanol: CH3CH2OH 8/31/2009 n-octane: CH3CH2CH2CH2CH2CH2CH2CH3 Zumdahl Chapter 4 5 Electrolytes and Non-Electrolytes Mainly ionic compounds such as potassium sulfate (K2SO4) and sodium chloride (NaCl) which, as a solute, increase the (NaCl) electrical conductance of water when they dissolve. Do not increase the conductance of water when they dissolve. ~100% dissociated in aqueous solution. Increase conductivity much more than weak electrolytes. electrolytes. 8/31/2009 Zumdahl Chapter 4 6 3 The Composition of Solutions The actual amount of solute in any given volume of solution depends on how concentrated or dilute the solution happens to be. This is expressed by a ratio: The concentration (c) of a solute in a solution equals the chemical amount (i.e., no. of moles) of the solute (n) divided by the volume (V) (V of the entire solution. csolute = molarity 8/31/2009 ≡ ≡ M Zumdahl Chapter 4 7 Calculate the molarity of a solution made by dissolving 10.0 g of Al(NO3)3 in enough water to make 250.0 mL of solution. nAl(NO3)3 = = 0.0469 moles of Al(NO3)3 Vsolution = = 0.250 L = 0.188 mol L-1 cAl(NO3)3 = 8/31/2009 Zumdahl Chapter 4 8 4 Types of Chemical Reactions 8/31/2009 1. Precipitation Reactions (Chapter 13) – Ionic Equations and Net Ionic Equations – Predicting Precipitation Reactions 2. Acids and Bases (Chapters 12 & 13) – Arrhenius Acids and Bases Theory – Strong and Weak acid – Naming Acids – Weak Bases – Modifying the Arrhenius model – Acid-Base Titrations Acid3. Oxidation−Reduction (Redox) Reactions Oxidation− Redox) – Oxidizing and Reducing Agents – Oxidation Number – Types of Redox Reactions » Combination and Decomposition » Oxygenation » Hydrogenation » Displacement Reactions Zumdahl Chapter 4 9 » Disproportionation Approximate Solubility Designations • Soluble • More than of solute per 100 g water • Slightly Soluble • of solute per 100 g water • Insoluble • Less than 8/31/2009 of solute per 100 g water Zumdahl Chapter 4 10 5 Predicting Dissolution and Precipitation Reactions CHEM1310 McKelvy Lecture 8/31/2009 Zumdahl Chapter 4 11 Ionic Equations and Net Ionic Equations BaCl2(aq) + K2SO4(aq) → BaSO4(s) + 2 KCl(aq) aq) aq) KCl(aq) Ba2+(aq) + 2 Cl−(aq) + 2 K+(aq) + SO42−(aq) → aq) (aq) aq) aq) Ionic equation BaSO (s) + 2 K+(aq) + 2 Cl−(aq) (aq) 4 Ba2+(aq) + 2 Cl−(aq) + 2 K+(aq) + SO42−(aq) → aq) (aq) aq) aq) Spectator ions BaSO4(s) + 2 K+(aq) + 2 Cl−(aq) (aq) Net ionic equation 8/31/2009 Zumdahl Chapter 4 12 6 8/31/2009 Zumdahl Chapter 4 13 Write a net ionic equation to represent the formation of the precipitate observed when aqueous solutions of CaCl2 and NaF are mixed. Identify the spectator ions in this process. CaCl2(aq) ≈ Ca2+(aq) + 2 Cl−(aq) aq) aq) aq) NaF(aq) ≈ Na+(aq) + F−(aq) NaF(aq) aq) aq) CaCl2(aq) + NaF(aq) → aq) NaF(aq) Ca2+(aq) + 2 Cl−(aq) + Na+(aq) + F−(aq) → aq) aq) aq) aq) Spectator ions CaF2(s) + Na+(aq) + 2 Cl−(aq) aq) aq) Net ionic equation 8/31/2009 Zumdahl Chapter 4 14 7 Acids and Bases and Their Reactions (more in Chapter 12) Acid/Base Definitions 1. Arrhenius Acids and Bases • Acids are H+ donors • Bases are OH− donors 2. Arrhenius-Broadened Definition Arrhenius• Acids increase H+ concentration • Bases increase OH− concentration 3. Brønsted-Lowry Acids and Bases Brø nsted• Acids donate H+ • Bases accept H+ 8/31/2009 Zumdahl Chapter 4 Arrhenius 1903 Nobel Prize 15 The Person Behind the Science Svante Arrhenius (1859-1927) Highlights – Specialized in the conductivities of electrolytes. Postulated that in forming a solution, a salt dissociates into charged particles. – Formulated the concept of activation energy. The Arrhenius equation gives the quantitative basis of the relationship between the activation energy and the rate at which a reaction proceeds (Chapter 17). 17). – In 1884 Arrhenius proposed definitions for acids and bases. Acids produce hydrogen ions in solution and bases produce hydroxide ions in solution. Moments in a Life – Awarded 1903 Nobel Prize in chemistry – The lunar crater Arrhenius is named after him 8/31/2009 Zumdahl Chapter 4 16 8 Acid/Base Neutralization Reactions H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq) → aq) aq) aq) aq) H2O(l) + Na+(aq) + Cl−(aq) O(l aq) aq) (net ionic equation) Neutralization Concept (at the equivalence point) point) = = 8/31/2009 Zumdahl Chapter 4 17 Titration of an Acid with a Base End point Base (Titrant): Standard Solution 8/31/2009 Indicator Equivalence point Stoichiometric point Zumdahl Chapter 4 18 9 Strong Acids and Bases Strong acids and bases are strong electrolytes which dissociate essentially completely in water. Example: n HCl(g) → n H+(aq) + n Cl−(aq) HCl(g aq) aq) 7 Common Strong acids: Hydrochloric acid acids: Perchloric acid Chloric acid Hydrobromic acid Sulfuric acid Hydroiodic acid Nitric acid Example: n NaOH(s) → n Na+(aq) + n OH−(aq) NaOH(s aq) aq) hydroxide salts are strong bases bases 8/31/2009 Zumdahl Chapter 4 19 Hydrofluoric acid (HF) Many Weak Acids Oxalic acid (H2C2O4) Acetic acid (CH3COOH) Formic acid (HCOOH) Phosphoric acid (H3PO4) CH3COOH(aq) → CH3COO−(aq) + H+(aq) (dissociation of acetic acid) COOH(aq) aq) aq) Neutralization Reaction CH3COOH(aq) + NaOH(aq) → NaCH3COO(aq) + H2O(l) (overall equation) CH3COOH(aq) + Na+(aq) + OH−(aq) → COOH(aq) aq) aq) Na+(aq) + CH3COO−(aq) + H2O(l) (aq) O(l (ionic equation) (net ionic equation) 8/31/2009 Zumdahl Chapter 4 20 10 Compute the molarity of a solution of sodium hydroxide if 25.64 mL of solution must be added to a solution containing 0.5333 g of KHC8H4O4 (potassium hydrogen phthalate) to reach the equivalence point . nKHP = = 2.611 x 10-3 mol KHP nNaOH = (at the equivalence point) point) [NaOH] = Concept = 0.1018 mol L-1 nbase = nacid nNaOH = nKHP 8/31/2009 (at the equivalence point) point) Zumdahl Chapter 4 21 Oxidation − Reduction Reactions Redox: an extensive and important class of reactions that is Redox: characterized by the transfer of electrons. 2 Mg(s) + O2(g) → 2 MgO(s) MgO(s Mg(s ↓loss 2x 2e2e ↑gain 2e = 1 x 2 x 2e- Magnesium is : it gives up electrons as the charge on its atoms increases from zero to +2. Oxygen is : it gains electrons as the charge on its atoms decreases from zero to -2 (i.e., becomes more negative). 8/31/2009 Zumdahl Chapter 4 22 11 8/31/2009 Zumdahl Chapter 4 23 Oxidation Numbers (also called oxidation states) are determined states) for the atoms in covalently bonded compounds by applying the following set of simple rules: rules: 1. In the free elements, each atom has an oxidation number of Zero elements, no matter how complicated the element. i.e., Na (s), Fe(s), K(s), Mg(s) H2, O2, N2, Cl2 all have zero oxidation states Fe(s), K(s), Mg(s) 2. The oxidation number of the atoms in a neutral molecule must add up to zero; those in an ion must add up to the charge on the ion. 3. Alkali metal (Group I) atoms have oxidation number +1, and alkaline earth (Group II) atoms have oxidation number +2 in their compounds; atoms of Group III elements usually have oxidation number +3 in their compounds. 8/31/2009 Zumdahl Chapter 4 24 12 Oxidation Numbers (con’t) (con’ 4. Fluorine always has an oxidation number of -1 in its compounds. The other halogens have oxidation number -1 in their compounds, except in compounds with oxygen and with other halogens, in which they can have positive oxidation numbers. 5. Hydrogen is assigned an oxidation number of +1 in its compounds, except in metal hydrides such as LiH, in which rule 3 take precedence and hydrogen has an oxidation number of -1. 6. Oxygen is assigned an oxidation number of -2 in compounds. There are two exceptions: in compounds with fluorine, rule 4 takes takes precedence, and in compounds that contain O—O bonds, rules 3 − 5 O— take precedence. Thus, the oxidation number of oxygen in OF2 is +2; in OF peroxides (e.g., HO-OH, FO-OF and NaO-ONa), its oxidation number is HOFONaO- ONa), ±1, and in superoxides (e.g., KO-O, FO-O), its oxidation number is ±½. KO- FO8/31/2009 Zumdahl Chapter 4 25 Oxidation and Reduction An atom is oxidized (_____ electrons) if its oxidation number increases in a chemical reaction Na · → Na+ + e- Loss of a valence electron An atom is reduced ( electrons) if its oxidation number decreases. 8/31/2009 - e - + Cl → Cl− Chapter 4 Gain of a valence electron Zumdahl 26 13 Oxidation State (Oxidation Number) – Can assign for Ionic and Covalent bonded compounds – Not formal electric charges, rather what the charge would be if the compound were ionic – Range from -3 to +7 Examples 4. VCl4 1. CrO3 0 = (Cr * 1) + (O * 3) 5. Mn2O7 2. TlCl3 Net Charge = ∑ (Oxid State x N) where " Oxid State" is the oxidation number 3. Mn3N2 of an element and N is the number of atoms of that element 8/31/2009 Zumdahl Chapter 4 27 Examples (con’t) MnO4− -1 = (Mn * 1) + (O * 4) -1 = (Mn *1 ) + (-2 * 4) (- Net Charge = ∑ (Oxid State x N) where " Oxid State" is the oxidation number of an element and N is the number of atoms of that element 8/31/2009 Zumdahl Chapter 4 28 14 Oxidation Electron Number Change Change Term Oxidation Increase Loss of Electrons Reduction Decrease Gain of Electrons Oxidizing Agent, does the oxidizing Decrease Picks Up electrons Reducing Agent, does the reducing Increase Supplies Electrons Substance Oxidized Increase Loses Electrons Substance Reduced Decrease Gains Electrons 8/31/2009 Zumdahl Chapter 4 29 Oxidation-Reduction reactions Oxidation Number Change Electron Change Oxidizing Agent, does the oxidizing Decrease Picks Up electrons Reducing Agent, does the reducing Increase Supplies Electrons Substance Oxidized Increase Loses Electrons Substance Reduced Decrease Gains Electrons Term ½ O2 H2 → H2O ½ Cl2 + Na → NaCl 2 H+ 8/31/2009 + + Mg → Mg++ + H2 Zumdahl Chapter 4 30 15 A summary of an oxidation-reduction oxidationprocess, in which M is oxidized and X is reduced. Li F Li + 8/31/2009 e- + e- F Zumdahl Chapter 4 31 Chapter 4 Types of Chemical Reactions and Solution Stoichiometry 4.1 4.2 Water, the Common Solvent The Nature of Aqueous Solutions: Strong and Weak Electrolytes 4.3 The Composition of Solutions 4.4 Types of Chemical Reactions 4.5 Precipitation Reactions 4.6 Describing Reactions in Solution 4.7 Selective Precipitation 4.8 Stoichiometry of Precipitation Reactions 4.9 Acid-Base Reactions Acid4.10 Oxidation-Reduction Reactions Oxidation4.11 Balancing Oxidation-Reduction Equations (skip) Oxidation4.12 Simple Oxidation-Reduction Titrations (skip) Oxidation8/31/2009 Zumdahl Chapter 4 32 16 ...
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This note was uploaded on 10/06/2009 for the course CHEM 1101 taught by Professor Bottomley during the Fall '08 term at Georgia Institute of Technology.

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