Ch. 4 - Silberberg, Chapter 4 The Major Classes of Chemical...

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Unformatted text preview: Silberberg, Chapter 4 The Major Classes of Chemical Reactions (Review) 4.1 The Role of Water as Solvent Oppositely charged ions from an ionic compound separate from each other when dissolved in water, become surrounded by water molecules, and spread randomly throughout a solution. The movement of the ions enables the solution to conduct electricity. • Electrolyte: a substance that conducts a current when dissolved in water. • Solvated: the process by which an ion is surrounded by solvent molecules. • Nonelectrolytes: covalent molecules that are polar in nature, dissolve in water, but do not dissociate. Water is a polar molecule: • The ionizing effect of water as a solvent is due to the distribution of it bonding electrons and the overall shape of the water molecule. 1. Sharing of electrons between nonidentical atoms in a covalent bond is unequal, with oxygen (O) attracting the bonding electrons with greater affinity than H. 2. This situation creates 2 poles of charge (+ for H, - for O). H-O is a polar bond because of this. 3. Non-bonding electron pairs on O of water distort molecule’s shape (O is group 6A, 2 bonding e-‘s, 4 nonbonding). 4. Water is a bent molecule to spatially accommodate the bonding H’s and 2 nonbonded pairs. • Non-bonding electron pairs on O act to attract (+) charged ions in solution. • Water separating the ionic compound greatly reduces the electrostatic force of attraction between the ion pair. 4.2 Aqueous Ionic Reactions Example: Molecular equation: Total ionic equation: Net ionic equation: 2AgNO3 (aq) + Na2CrO4(aq) Æ Ag2CrO4 (s) + 2NaNO3 (aq) + - - 2Ag (aq) + 2NO 3 (aq) + 2Na+(aq) + CrO4 - (aq) Æ Ag2CrO4(s) + 2Na+(aq) + 2NO3 (aq) 2Ag+(aq) + CrO4 - (aq) Æ Ag 2 CrO4 (s) - Spectator ions: Na+(aq), NO3 (aq) † • Molecular equation: shows all the reactants and products as if they were intact, † undissociated compounds. • Total ionic equation: shows all soluble ionic substances dissociated into ions. • Spectator ions: ions that appear on both side of the equation and not involved in the actual chemical change. • Net ionic equation: eliminates spectator ions and show the actual chemical change taking place. 4.3 Precipitation Reactions Two soluble ionic compounds react to form an insoluble product or precipitate (e.g. Ag2CrO4(s) from the previous example). Predicting whether a precipitation reaction will occur: 1. Note the ions present in the reactants. 2. Consider the possible combinations. 3. Decide whether any of the combinations are insoluble. See Table 4.1 pg. 139 for a list of precipitates. 4.4. Acid-Base Reactions • Acid: a substance that produces H+ ions when dissolved in water. 2O HX æHæÆ H+(aq) + X- (aq) æ • Base: a substance that produces OH- ions when dissolved in water. 2O MOH æHæÆ M+(aq) + OH- (aq) æ † • Strong acids and strong bases dissociate completely into ions in water. • Weak acids and bases dissociate very little in water and remain intact. † • Acids and bases react together in a neutralization reaction, to form water and salt. HX(aq) + MOH(aq) Æ MX (aq) + H2 O(l) acid base salt water • Titration: a basic (MOH) solution of known concentration (standard solution) is used to determine the concentration of an acidic solution (HX) that is unknown. † • The equivalence point in the titration is reached when all the moles of H+ ions in the solution have reacted with the measured amount of OH- ions added to the solution. • The end point of the titration occurs when a tiny excess of base is added, generally accepted to be the same as the equivalence point. Oxidation-Reduction Reactions (redox) • Redox reactions are the key chemical event for movement of electrons from one reactant to another. • Driving force in these reactions comes from the flow of electrons in a specific direction, from least electronegative atom to the most electronegative atom. X ætransfer oræ æelectronsÆY æ æ shift of ææ • X loses electrons • Y gains electrons • X is oxidized • Y is reduced † • X is the reducing agent • X increases in oxidation number • Y is the oxidizing agent • Y decreases in oxidation number Rules for assigning oxidation number (O.N.): 1. For an atom in its elemental form (Na, O2, Cl2, etc.) O.N.=0 2. For a monatomic ion, O.N. = ion charge 3. The sum of O.N. values for the atoms in a compound equals zero. The sum of O.N. values for the atoms in a polyatomic ion equals the ions charge. Balancing Redox Equations: 1. Assign oxidation numbers to all elements in the reaction. 2. From the changes in oxidation numbers, identify the oxidized and reduced species. 3. Compute the number of electrons lost in the oxidation and gained in the reduction, from the oxidation number changes. 4. Multiply one or both of these numbers by the appropriate factors to balance electron gain with electron loss. Use these factors are balancing coefficients. 5. Complete the balancing by inspection, adding states of matter (solid, liquid, gas). 4.6 Elemental Substances in Redox Reactions • Combination reactions: X + Y Æ Z • Decomposition reactions: Z Æ X + Y • Displacement reactions: X + YZ Æ XZ + Y † • Metathesis reactions (double displacement): WX + YZ Æ XZ +YW † † 4.7 Reversible Reactions: An Introduction to Chemical Equilibrium † See later chapters. ...
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This note was uploaded on 06/15/2010 for the course CHEM 124 taught by Professor Hascall during the Winter '08 term at Cal Poly.

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