4.1 A solution is a homogenous mixture of 2 or more substances The solute is(are) the substance(s) present in the smaller amount(s) The solvent is the substance present in the larger amount Soft drink ( l ) Air ( g ) Soft Solder ( s ) H 2 O N 2 Pb Sugar, CO 2 O 2 , Ar, CH 4 Sn Solution Solvent Solute
An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity. 4.1 nonelectrolyte weak electrolyte strong electrolyte
Strong Electrolyte – 100% dissociation Weak Electrolyte – not completely dissociated Conduct electricity in solution? Cations (+) and Anions (-) 4.1 NaCl ( s ) Na + ( aq ) + Cl - ( aq ) H 2 O CH 3 COOH CH 3 COO - ( aq ) + H + ( aq )
Ionization of acetic acid CH 3 COOH CH 3 COO - ( aq ) + H + ( aq ) 4.1 Acetic acid is a weak electrolyte because its ionization in water is incomplete. A reversible reaction. The reaction can occur in both directions.
Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner. 4.1 H 2 O
Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution 4.1 C 6 H 12 O 6 ( s ) C 6 H 12 O 6 ( aq ) H 2 O
Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Na + and NO 3 - are spectator ions 4.2 Pb 2+ + 2NO 3 - + 2Na + + 2I - PbI 2 ( s ) + 2Na + + 2NO 3 - PbI 2 Pb(NO 3 ) 2 ( aq ) + 2NaI ( aq ) PbI 2 ( s ) + 2NaNO 3 ( aq ) precipitate Pb 2+ + 2I - PbI 2 ( s )
Precipitation of Lead Iodide PbI 2 4.2 Pb 2+ + 2I - PbI 2 ( s )
4.2 Solubility is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.
Writing Net Ionic Equations <ul><li>Write the balanced molecular equation. </li></ul><ul><li>Write the ionic equation showing the strong electrolytes completely dissociated into cations and anions. </li></ul><ul><li>Cancel the spectator ions on both sides of the ionic equation </li></ul><ul><li>Check that charges and number of atoms are balanced in the net ionic equation </li></ul>4.2 AgNO 3 ( aq ) + NaCl ( aq ) AgCl ( s ) + NaNO 3 ( aq ) Ag + + NO 3 - + Na + + Cl - AgCl ( s ) + Na + + NO 3 - Ag + + Cl - AgCl ( s ) Write the net ionic equation for the reaction of silver nitrate with sodium chloride.
Chemistry In Action: An Undesirable Precipitation Reaction 4.2 CO 2 ( aq ) CO 2 ( g ) Ca 2+ ( aq ) + 2HCO 3 ( aq ) CaCO 3 ( s ) + CO 2 ( aq ) + H 2 O ( l ) -
Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas 4.3 Cause color changes in plant dyes. Aqueous acid solutions conduct electricity. 2HCl ( aq ) + Mg ( s ) MgCl 2 ( aq ) + H 2 ( g ) 2HCl ( aq ) + CaCO 3 ( s ) CaCl 2 ( aq ) + CO 2 ( g ) + H 2 O ( l )
Have a bitter taste. Feel slippery. Many soaps contain bases. Bases 4.3 Cause color changes in plant dyes. Aqueous base solutions conduct electricity.
Arrhenius acid is a substance that produces H + (H 3 O + ) in water Arrhenius base is a substance that produces OH - in water 4.3
A Br ø nsted acid is a proton donor A Br ø nsted base is a proton acceptor acid base acid base 4.3 A Br ø nsted acid must contain at least one ionizable proton!
Monoprotic acids Strong electrolyte, strong acid Strong electrolyte, strong acid Weak electrolyte, weak acid Diprotic acids Strong electrolyte, strong acid Weak electrolyte, weak acid Triprotic acids Weak electrolyte, weak acid Weak electrolyte, weak acid Weak electrolyte, weak acid 4.3 HCl H + + Cl - HNO 3 H + + NO 3 - CH 3 COOH H + + CH 3 COO - H 2 SO 4 H + + HSO 4 - HSO 4 - H + + SO 4 2- H 3 PO 4 H + + H 2 PO 4 - H 2 PO 4 - H + + HPO 4 2- HPO 4 2- H + + PO 4 3-
Identify each of the following species as a Br ø nsted acid, base, or both. (a) HI, (b) CH 3 COO - , (c) H 2 PO 4 - Br ø nsted acid Br ø nsted base Br ø nsted acid Br ø nsted base 4.3 HI ( aq ) H + ( aq ) + I - ( aq ) CH 3 COO - ( aq ) + H + ( aq ) CH 3 COOH ( aq ) H 2 PO 4 - ( aq ) H + ( aq ) + HPO 4 2- ( aq ) H 2 PO 4 - ( aq ) + H + ( aq ) H 3 PO 4 ( aq )
Neutralization Reaction 4.3 acid + base salt + water HCl ( aq ) + NaOH ( aq ) NaCl ( aq ) + H 2 O H + + Cl - + Na + + OH - Na + + Cl - + H 2 O H + + OH - H 2 O
Oxidation-Reduction Reactions (electron transfer reactions) Oxidation half-reaction (lose e - ) Reduction half-reaction (gain e - ) 4.4 2Mg 2Mg 2+ + 4e - O 2 + 4e - 2O 2- 2Mg + O 2 + 4e - 2Mg 2+ + 2O 2- + 4e - 2Mg + O 2 2MgO
Zn is oxidized Cu 2+ is reduced Zn is the reducing agent Cu 2+ is the oxidizing agent 4.4 Ag + is reduced Ag + is the oxidizing agent Zn ( s ) + CuSO 4 ( aq ) ZnSO 4 ( aq ) + Cu ( s ) Zn Zn 2+ + 2e - Cu 2+ + 2e - Cu Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu ( s ) + 2AgNO 3 ( aq ) Cu(NO 3 ) 2 ( aq ) + 2Ag ( s ) Cu Cu 2+ + 2e - Ag + + 1e - Ag
Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. <ul><li>Free elements (uncombined state) have an oxidation number of zero. </li></ul>Na, Be, K, Pb, H 2 , O 2 , P 4 = 0 <ul><li>In monatomic ions, the oxidation number is equal to the charge on the ion. </li></ul>Li + , Li = +1 ; Fe 3+ , Fe = +3 ; O 2- , O = -2 <ul><li>The oxidation number of oxygen is usually –2 . In H 2 O 2 and O 2 2- it is –1 . </li></ul>4.4
<ul><li>The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1 . </li></ul>6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. <ul><li>Group IA metals are +1 , IIA metals are +2 and fluorine is always –1 . </li></ul>HCO 3 - O = -2 H = +1 3x( -2) + 1 + ? = -1 C = +4 4.4 7. Oxidation numbers do not have to be integers. Oxidation number of oxygen in the superoxide ion, O 2 - , is - ½ . Oxidation numbers of all the elements in HCO 3 - ?
The oxidation numbers of elements in their compounds 4.4
NaIO 3 Na = +1 O = -2 3x( -2 ) + 1 + ? = 0 I = +5 IF 7 F = -1 7x( -1 ) + ? = 0 I = +7 K 2 Cr 2 O 7 O = -2 K = +1 7x( -2 ) + 2x( +1 ) + 2x( ?) = 0 Cr = +6 4.4 Oxidation numbers of all the elements in the following ?
Types of Oxidation-Reduction Reactions Combination Reaction Decomposition Reaction 0 0 +3 -1 +1 +5 -2 +1 -1 0 4.4 A + B C 2Al + 3Br 2 2AlBr 3 2KClO 3 2KCl + 3O 2 C A + B
Types of Oxidation-Reduction Reactions Combustion Reaction 0 0 +4 -2 4.4 0 0 +2 -2 A + O 2 B S + O 2 SO 2 2Mg + O 2 2MgO
Types of Oxidation-Reduction Reactions Displacement Reaction Hydrogen Displacement Metal Displacement Halogen Displacement 4.4 0 +1 +2 0 0 +4 0 +2 0 -1 -1 0 A + BC AC + B Sr + 2H 2 O Sr(OH) 2 + H 2 TiCl 4 + 2Mg Ti + 2MgCl 2 Cl 2 + 2KBr 2KCl + Br 2
The Activity Series for Metals Hydrogen Displacement Reaction M is metal BC is acid or H 2 O B is H 2 4.4 M + BC AC + B Ca + 2H 2 O Ca(OH) 2 + H 2 Pb + 2H 2 O Pb(OH) 2 + H 2
The Activity Series for Halogens Halogen Displacement Reaction 4.4 0 -1 -1 0 F 2 > Cl 2 > Br 2 > I 2 Cl 2 + 2KBr 2KCl + Br 2 I 2 + 2KBr 2KI + Br 2
Types of Oxidation-Reduction Reactions Disproportionation Reaction Element is simultaneously oxidized and reduced. 0 +1 -1 4.4 Cl 2 + 2OH - ClO - + Cl - + H 2 O Chlorine Chemistry
Precipitation Acid-Base Redox (H 2 Displacement) Redox (Combination) 4.4 Ca 2+ + CO 3 2- CaCO 3 NH 3 + H + NH 4 + Zn + 2HCl ZnCl 2 + H 2 Ca + F 2 CaF 2 Classify the following reactions.
Chemistry in Action: Breath Analyzer 4.4 3CH 3 COOH + 2Cr 2 (SO 4 ) 3 + 2K 2 SO 4 + 11H 2 O +6 +3 3CH 3 CH 2 OH + 2K 2 Cr 2 O 7 + 8H 2 SO 4
Solution Stoichiometry The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. M KI M KI 500. mL = 232 g KI 4.5 M = molarity = moles of solute liters of solution What mass of KI is required to make 500. mL of a 2.80 M KI solution? volume of KI solution moles KI grams KI 166 g KI 1 mol KI x 2.80 mol KI 1 L soln x 1 L 1000 mL x
Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution. 4.5 Dilution Add Solvent Moles of solute before dilution (i) Moles of solute after dilution (f) = M i V i M f V f =
M i V i = M f V f M i = 4.00 M f = 0.200 V f = 0.06 L V i = ? L 4.5 = 0.003 L = 3 mL 3 mL of acid + 57 mL of water = 60 mL of solution How would you prepare 60.0 mL of 0.200 M HNO 3 from a stock solution of 4.00 M HNO 3 ? V i = M f V f M i = 0.200 x 0.06 4.00
Gravimetric Analysis 4.6 <ul><li>Dissolve unknown substance in water </li></ul><ul><li>React unknown with known substance to form a precipitate </li></ul><ul><li>Filter and dry precipitate </li></ul><ul><li>Weigh precipitate </li></ul><ul><li>Use chemical formula and mass of precipitate to determine amount of unknown ion </li></ul>
Titrations In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point – the point at which the reaction is complete Indicator – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL the indicator changes color 4.7
4.7 25.00 mL = 158 mL What volume of a 1.420 M NaOH solution is Required to titrate 25.00 mL of a 4.50 M H 2 SO 4 solution? WRITE THE CHEMICAL EQUATION! volume acid moles acid moles base volume base H 2 SO 4 + 2NaOH 2H 2 O + Na 2 SO 4 4.50 mol H 2 SO 4 1000 mL soln x 2 mol NaOH 1 mol H 2 SO 4 x 1000 ml soln 1.420 mol NaOH x M acid rx coef. M base
Chemistry in Action: Metals from the Sea CaCO 3 ( s ) CaO ( s ) + CO 2 ( g ) Mg(OH) 2 ( s ) + 2HCl ( aq ) MgCl 2 ( aq ) + 2H 2 O ( l ) CaO ( s ) + H 2 O ( l ) Ca 2+ ( aq ) + 2OH ( aq ) - Mg 2+ ( aq ) + 2OH ( aq ) Mg(OH) 2 (s ) - Mg 2+ + 2e - Mg 2Cl - Cl 2 + 2e - MgCl 2 ( aq ) Mg ( s ) + Cl 2 ( g )