Chapter 4 Chemical Formulas, Reactions, Redox and Solutions

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1 Terms to Know: Solubility Solute Solvent Solution Chapter 4 the amount of substance that dissolves in a given volume of solvent at a given temperature. a substance dissolved in a liquid to form a solution the dissolving medium in a solution a homogeneous mixture of a solute in a solvent Molarity, M Moles of solute per volume of solution in Liters (Expression of concentration ) M mole of solute Liter of solution Conductivity Electrolyte Nonelectrolyte Precipitation Reaction Spectator ions Reducing agent Oxidizing agent Reduction Oxidation Redox reaction the ability to conduct electricity in an aqueous solution Ionic compounds that conduct electricity in H 2 O Strong- excellent conductor, fully ionized Weak- poor conductor, partially ionized non-conductors, no ions present to conduct a reaction in which an insoluble substance forms and separates from the solution ions present in solution that do not participate directly in a reaction. Appear as both reacts and products. electron donor; a reactant that donates electrons to another substance to reduce the oxidation state of one of its atoms electron acceptor; a reactant that accepts electrons from another reactant a gain of electrons (a decrease in oxidation state) a loss of electrons (an increase in oxidation state) a reaction where electrons are gained (reduction) or lost (oxidation) Oxidation state (number) -the apparent charge of an atom Net-ionic equation Ionic equation Concepts I. Assigning Oxidation Numbers -a concept that provides a way to keep track of electrons in oxidation-reduction reactions according to certain rules an equation for a reaction in a solution, where strong electrolytes are written as ions, showing only those components that are directly involved in the chemical change. an equation for a reaction in solution, where strong electrolytes are written as ions Rules to calculate the oxidation number of any atom in any molecule. 1. In molecules group 1 and group 2 metals are +1 and +2 respectively. 2. Fluoride is always O is always -2 except in peroxides when it is -1 and OF 2 where it is H is always +1 with nonmetals and 1 with metals. 5. The sum of all the oxidation number of all the atoms equals the overall charge N 2 N = 0 Na 1+ Na = NH 4 H = +1 N = -3 SO 3 O = -2 S = +4

2 Chapter 4 II. Oxidation Reduction (Redox) (Single replacement reactions are redox rxns.) Involve a transfer of electron(s) o Oxidation- loss of electrons; the atom that is oxidized is the reducing agent o Reduction- gain of electrons; the atom that is reduced is the oxidizing agent Example: In the following rxn, identify the oxidized atom, reduced atom, oxidizing and reduction agents. Fe 3+ + Cu 1+ Cu 2+ + Fe 2+ Solution: First determine what is occurring with each atom. Then apply the definitions. Fe: Fe e - Fe 2+ reduction, oxidation agent Cu: Cu 1+ Cu e - oxidation, reduction agent Note: In half-rxns, when reduction occurs, e s are reactants and when oxidation occurs, e s are products. III. Balancing Redox Rxns {Note: Any time a question says a rxn occurs under certain conditions, it is most likely a redox & ½ reaction prob.} There are two methods that both arrive at the same answer. Your choice, choose the method that works for you. Method 1 Under acidic conditions 1 st Write the two half reactions. Then for each ½ reaction: i Balance the mass 1 Balance non- H and O atoms. 2 Add H 2 O to balance O s 3 Add H + s to balance H s ii Balance the charge by adding e - s. 2 nd Cross multiply to cancel e - s. 3 rd Add the two half-rxns to recreate the full rxn. Method 2 Under acidic conditions 1 st Write the two half reactions. Then for each ½ reaction: Balance non H and O atoms. Add e s to the appropriate side to match the change in the oxidation number. 2 nd Cross multiply to cancel e - s. 3 rd Add the two half-rxns to recreate the full rxn. 4 th Add H 2 O to balance O s 5 th Add H + s to balance H s Under basic conditions 1 st Balance ½ reactions as if acidic. 2 nd Add OH s to each side to cancel the H + s OH + H + H O. Cancel extra H 2 O s Example: Balance the following reaction: Fe 2+ + MnO 4 Mn 2+ + Fe 3+ Under Acidic Conditions: 1 st Fe 2+ Fe e - 8H + + MnO 4 + 5e - Mn H 2 O 2 nd 5( Fe 2+ Fe e - ) 1(8H + + MnO 4 + 5e - Mn H 2 O) 3 rd 5Fe H + + MnO 4 5Fe 3+ +Mn H 2 O Under basic conditions 1 st Balance ½ reactions as if acidic. 2 nd Add OH s to each side to cancel the H + s OH + H + H O. Cancel extra H 2 O s. Example: Balance the following reaction: Fe 2+ + MnO 4 Mn 2+ + Fe 3+ Under Acidic Conditions: 1 st Fe 2+ Fe e - [Fe: +2 to +3] MnO 4 + 5e - Mn 2+ [Mn: +7 to +2] 2 nd 5( Fe 2+ Fe e - ) 1(8H + + MnO 4 + 5e - Mn H 2 O) 3 rd 5Fe 2+ + MnO 4 5Fe 3+ +Mn 2+ 4 th 5Fe 2+ + MnO 4 5Fe 3+ +Mn H 2 O 5 th 5Fe H + + MnO 4 5Fe 3+ +Mn H 2 O Under basic conditions: acidic cond s 5Fe H + + MnO 4 5Fe 3+ +Mn H 2 O neuralize H + 8OH - + 5Fe H + + MnO 4 5Fe 3+ +Mn H 2 O + 8OH - 8H 2 O + 5Fe 2+ + MnO 4 5Fe 3+ +Mn H 2 O + 8OH - basic cond s 4H 2 O + 5Fe 2+ + MnO 4 5Fe 3+ +Mn OH -

3 IV. Ionic Solutions Chapter 4 1. Soluble ionic compounds dissociate when dissolved in water and are always written as ions. ex. A b B a (s) A a+ (aq) + B b- (aq) 2. Soluble ionic compounds redissociate when formed in water leaving no net reaction because all ions cancel. ex. A a+ (aq) + B b- (aq) A a+ (aq) + B b- (aq) 3. Insoluble ionic compounds precipitate when formed in H 2 O and do not dissolve in H 2 O. ex. A a+ (aq) + B b- (aq) A b B a (s) V. Ionic Reactions When solutions of ionic compounds are mixed, a chemical reaction may or may not occur. A reaction occurs if anything remains after cancelling the spectator ions. You will be asked to predict if a chemical rxn occurs and if it does to write the balanced chemical equation for the rxn. Things you should be aware of These reactions involve 2 solid ionic compounds dissolved in H 2 O, A a+ B b- + C c+ D d- When two opposite ions attract they will form either a soluble or insoluble compound (apply the rules). A soluble compound re-dissociates into ions so soluble compounds are written as ions An insoluble compound precipitates and is written as an empirical formula. Ions that appear as both reactants and products are spectator ions and are removed from the ionic chemical equation. Once the spectator ions are removed, we write the net ionic equation. Example: A solution of MgCl 2 (aq) is mixed with a solution of AgNO 3. Predict the reaction that occurs. Solution: To solve: 1 st in the question, connect cation to anion to predict the potential products 2 nd apply solubility rules soluble are ions; and insoluble are formulas 3 rd write the equation cancelling the spectator ions. Potential products are MgNO 3 (soluble) and AgCl (insoluble). Mg 2+ + Cl 1 + Ag 1+ + NO 3 1 Mg 2+ + NO AgCl So, Ag 1+ + Cl AgCl VI. Conducting Electrolytes ionic compounds that conduct electricity in H 2 O or ionized covalent compounds. Strong- excellent conductor, typically because they are 100% ionized Weak- poor conductors, typically because they are partially ionized Non-Electrolytes nonconductors, cause no ions present to conduct, typically covalent compounds. Example: Which Conduct More? 1. Sugar vs. NaCl 2. Acetic acid vs. NaCl 3. 1 M NaCl vs. 0.5 M NaCl (1M has more ions) 4. 1 M BaCl 2 vs. 1 M NaCl (1M with 3 ions vs 1M with 2 ions)

4 Chapter 4 VII. Review of Formula Writing and Ionic Equation Writing TABLE 1 IONS FORMULAS To write Neutral Binary Formulas from Ions 1 Write the two ions side-by-side, cation first, without the charges. Example A a+ + B b- 1 A a+ + B b- A B 2 Determine how many of each ion are needed to create a neutral formula, write that number as a subscript. [Typically this is done by criss-crossing and reducing the charges.] 3 Note that if a polyatomic ion is used, the subscript is written outside of ( ) s 2 A a+ + B b- A b B a 3 A a+ b- + BO x A b (BO x ) a Al 3+ + OH Al 1 (OH) 3 note that you do not have to write the 1 s, but it helps TABLE 2 FORMULAS IONS Dissociating Neutral Binary Formulas Into Ions 1 Split the formula into two halves. Often this is done by writing the first symbol as the first half an the remaining symbols as the second half. Do NO write any subscripts unless they are inside ( ) s or follow two capital letters (see example) Example A b B a + or A b (BO x ) a + 1 A b B a A + B or A b (BO x ) a A + BO x 2 The first ion is positive the second is negative. 2 A b B a A + + B - or A b (BO x ) a A + + BO x 3 Now determine the magnitude of the charge. Do the opposite of Ions to formulas by criss-crossin the subscripts into the position of charges. 4 To make sure you are right (and often you are not) you need to check your charges with the periodic table or with the charges on the polyatomic ions y were to memorize. Trend For Charges For simple compounds, any element within the vertical family headed by the following s and p block elements will have the corresponding charge. H 1+, Be 2+, B 3+, C 4+/4-, N 3-, O, F, He no charge 3/4 A b B a A a+ + B b- or A b (BO x ) a A a+ + BO x 4 Examples of Checking Charges Cu 2 O Cu 1+ + O Note the charge for O matches the periodic table. So we assume the charges are correct. K 2 SO 4 K 1+ + SO 4 Note the charge for K and SO 4 matches the periodic table and the memorized charge, so we assume the charges are correct. Also note, the subscript 4 was not criss-crossed. The invisible 1 outside the invisible ( ) s was criss-crossed. The 4 follows two capital letters and stays put. For more complex compounds (AP Chem), the rules for determining oxidation numbers must be followed.. CuO Cu 1+ + O Note that the charge for O should be a from the periodic table. So, we must double all charges to get it from to. So... CuO Cu 2+ + O TABLE 3 NAMES FORMULAS To write Neutral Binary Formulas from Names 1 Write the symbols for the corresponding cation and anion 2 Write the formula using Table 1. Example: Write the formula of sodium sulfate 1 sodium = Na 1+ sulfate = SO 4 2 Na 1+ + SO 4 Na 2 SO 4

5 Writing, Dissociating and Naming Ionic Formulas TABLE 4 NAMING IONS and BINARY IONIC COMPOUNDS Naming Binary Ionic Compounds Example You never name formulas. Name ONLY ions! So If you are given a formula to name you must first dissociate it into ions and then name the ions. or If you are given a name and asked to write a formula, you must first write the ions and then write the formula. Naming Cations if s- or p-block elements The name is the elemental name. if d-block elements The name is the elemental name followed by the charge as a Roman numeral in parentheses. Naming Anions if mono-atomic (only one elemental symbol) Then the name is the elemental name, but with the ending changed to ide. if polyatomic (more than one elemental symbol) Then the name is the memorized name. Naming Cations if s- or p-block elements Li 1+ lithium. Al 3+ aluminum if d-block elements Ni 1+ nickel (I). Ni 3+ nickel (III) Naming Anions if monatomic: F fluoride (not fluorine) O oxide (not oxygen) As 3- arsenide (not arsenic) if polyatomic: These are the names to memorize C 2 H 3 O 2 acetate CO 3 carbonate CN cyanide OH hydroxide NO 3 nitrate NO 2 nitrite SO 4 sulfate SO 3 sulfite 3- PO 4 phosphate 3- PO 3 phosphite Note that there exists one polyatomic cation that should be memorized. NH 1+ 4 is ammonium All other cations will be one elemental symbol. To name a formula 1 Dissociate into ions, then 2 Name the cation followed by the anion. To name a formula ex. Write the name of CuCO 3 1 CuCO 3 Cu 2+ + CO 3 2 Cu 2+ = copper (II) CO 3 = carbonate so, the name is copper (II) carbonate

6 VIII. Solubility Rules The following rules must be memorized. These rules are applied to a solution of an ionic compound and tell you whether that compound is soluble and so must be written as ions. OR If the compound is insoluble and so must be written as an empirical formula. Aside from the very first rule, the rules are applied by looking at the anion and then checking for exceptions. Example: What are the solubilities of each compound? Solution: MgNO 3 Soluble All nitrates are soluble, no exceptions so you would write Mg 2+ + NO 3 AgCl Insoluble Most chlorides are soluble, this is an exception, so you would write AgCl K 2 SO4 Soluble Most sulfates are soluble or first rule, so you would write K 1+ + SO 4 K 2 C 2 O 4 Soluble Most oxalates are insoluble this is an exception or first rule, so write K 1+ + C 2 O 4 Solubility Rules Must Be Memorized! TABLE 1 SOLUBLE ANIONS Soluble Compounds FIRST RULE Most salts containing alkali metal ions and the ammonium ion are soluble. Exceptions Salts of nitrate, NO 3 chlorate, ClO 3 perchlorate, ClO 4 acetate, CH 3 CO 2 Most salts of Cl, Br, and I Halides of Ag 1+, Hg 2+ 2, and Pb 2+ Compound containing fluoride, F Fluorides of Mg 2+, Ca 2+, Sr 2+, Ba 2+, and Pb 2+ Salts of sulfate, SO 4 Sulfates of Ca 2+, Sr 2+, Ba 2+, and Pb 2+ TABLE 2 INSOLUBLE ANIONS Insoluble Compounds All salts of carbonate, CO 3 phosphate, PO 3-4 oxalate, C 2 O 4 chromate, CrO 4 Most metal sulfides, S most metal hydroxides, OH, and oxides, O Exceptions Salts of NH 1+ 4 and the alkali metal cations. Salts of NH 1+ 4 and the alkali metal cations. Salts of NH 1+ 4 and the alkali metal cations.

AP Chemistry Unit #4. Types of Chemical Reactions & Solution Stoichiometry

AP Chemistry Unit #4. Types of Chemical Reactions & Solution Stoichiometry AP Chemistry Unit #4 Chapter 4 Zumdahl & Zumdahl Types of Chemical Reactions & Solution Stoichiometry Students should be able to: Predict to some extent whether a substance will be a strong electrolyte,