Chemical Formulas, Reactions, Redox and Solutions

Chapter 4 Chemical Formulas, Reactions, Redox and Solutions

Terms to Know: Solubility the amount of substance that dissolves in a given volume of solvent at a given temperature. Solute a substance dissolved in a liquid to form a solution Solvent the dissolving medium in a solution Solution a homogeneous mixture of a solute in a solvent Molarity, M Moles of solute per volume of solution in Liters (Expression of concentration )

mole of solute M  Liter of solution

Conductivity the ability to conduct electricity in an aqueous solution

Electrolyte Ionic compounds that conduct electricity in H2O Strong- excellent conductor, fully ionized Weak- poor conductor, partially ionized Nonelectrolyte non-conductors, no ions present to conduct Precipitation Reaction a reaction in which an insoluble substance forms and separates from the solution Spectator ions ions present in solution that do not participate directly in a reaction. Appear as both reacts and products. Reducing agent electron donor; a reactant that donates electrons to another substance to reduce the oxidation state of one of its atoms Oxidizing agent electron acceptor; a reactant that accepts electrons from another reactant Reduction a gain of electrons (a decrease in oxidation state) Oxidation a loss of electrons (an increase in oxidation state) Redox reaction a reaction where electrons are gained (reduction) or lost (oxidation) Oxidation state (number) -the apparent charge of an atom -a concept that provides a way to keep track of electrons in oxidation-reduction reactions according to certain rules Net-ionic equation 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. Ionic equation an equation for a reaction in solution, where strong electrolytes are written as ions

Concepts

I. Oxidation – Reduction (Redox) (Single replacement reactions are redox rxns.)  Involve a transfer of electron(s) o Oxidation- loss of electrons o Reduction- gain of electrons o Oxidizing agent- the atom that is reduced o Reducing agent- the atom that is oxidized Chapter 4 Chemical Formulas, Reactions, Redox and Solutions Example: In the following rxn, identify the oxidized atom, reduced atom, oxidizing and reduction agents. Fe3+ + Cu1+  Cu2+ + Fe2+

Solution: First determine what is occurring with each atom. Then apply the definitions. Fe: +3  +2 Fe3+ + 1e-  Fe2+ reduction, oxidation agent Cu: +1  +2 Cu1+  Cu2+ + 1e- oxidation, reduction agent

Note: In half-rxns, when reduction occurs, e’s are reactants and when oxidation occurs, e’s are products.

II. Balancing Redox Rxns {Note: Any time a question says a rxn occurs under certain conditions, it is most likely a redox & ½ reaction prob.}

Under acidic conditions 1 Write the two half reactions. Then for each ½ reaction, i. Balance the mass 1. Balance non- H and O atoms. 2. Add H2O to balance O’s 3. Add H+’s to balance H’s ii. Balance the charge by adding e1-’s. 2 Cross multiply to cancel e1-’s. 3 Add the two half-rxns to recreate the full rxn.

Under basic conditions 1 Balance as if acidic. 2 Add OH1-’s to each side to cancels the H+’s 3 Cross multiply to cancel e1-’s. 4 Add the two half-rxns to recreate the full rxn.

Example: Balance the following reaction: 2+ 1- 2+ 3+ Fe + MnO4  Mn + Fe

 Under Acidic: 5( Fe2+  Fe3+ + 1e-) + 1- - 2+ 1(8H + MnO4 + 5e  Mn + 4H2O) 2+ + 1- 3+ 2+ 5Fe +8H + MnO4  5Fe + Mn + 4H2O

 Under Basic: Fe2+  Fe3+ + 1e-

- + 1- - 2+ - acidic cond’s 8OH + 8H + MnO4 +5e  Mn + 4H2O + 8OH + neuralize H 8H2 O 1- - 2+ - basic cond’s 4H2O +MnO4 + 5e  Mn +8OH Cross-multiply and add as before = same answer

III. Assigning Oxidation Numbers

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 –1. 3. O is always -2 except in peroxides when it is -1 and OF2 where it is +2. 4. 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

N2 N = 0 Chapter 4 Chemical Formulas, Reactions, Redox and Solutions Na1+ Na = +1 1+ NH4 H = +1 N = -3 2- SO3 O = -2 S = +4

IV. Ionic Solutions

1. Ionic compounds dissociate when dissolved in aqueous media (aka water). a+ b- Ex. AbBa(s)  A (aq) + B (aq)

2. Insoluble ionic compounds precipitate when formed in H2O a+ b- Ex. A (aq) + B (aq)  AbBa(s)

3. Soluble ionic compounds “redissociate” when formed in water leaving no net reaction. Ex. Aa+(aq) + Bb- (aq)  Aa+(aq) + Bb-(aq) No reaction

V. Ionic Reactions  a+ b- c+ d- These reactions involve 2 solid ionic compounds dissolved in H2O, A B + C D  When two opposite ions attract they will form either, 1. A soluble compound that will then re-dissociate into ions, or 2. An insoluble compound that precipitates (we only care about this one)  Ions that appear as both reacts 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.

You will be given questions that involve the mixing of two ionic solutions and you will be asked to predict if a chemical rxn occurs and if it does to write the balanced chemical equation for the rxn.

Example: A solution of MgCl2 (aq) is mixed with a solution of AgNO3. Write the chemical equation for the reaction.

To solve: 1st in the question, “connect” cation to anion to predict the potential products 2nd apply solubility rules to the products and identify the compound that will be precipated. 3rd write the equation between the two ions to form the insoluble product. Example Answer: Potential products are MgNO3 (soluble) and AgCl (insoluble). So, Ag1+ + Cl1-  AgCl

VI. Conducting

Electrolytes- ionic compounds that conduct electricity in H2O  Strong- excellent conductor, cause fully ionized  Weak- poor conductors, cause partially ionized

Non-Electrolytes- nonconductors, cause no ions present to conduct.

Questions. 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 BaCl2 vs. 1 M NaCl (1M with 3 ions vs 1M with 2 ions) Chapter 4 Chemical Formulas, Reactions, Redox and Solutions

VII. Review of Formula Writing and Ionic Equation Writing Ions to Formulas Example 1st Write the two ions side-by-side, positive first, Aa+ + Bb- → ______without the charges 1st Aa+ + Bb- → A B

nd nd a+ b- 2 Determine how many of each ion are needed to 2 A + B → AbBa create a neutral formula. Typically this is done by criss-crossing and reducing the charges.

rd rd a+ b- 3 Note that if a polyatomic ion is used, the subscript 3 A + BOx → Ab(BOx)a is written outside of ( )’s 3+ 1- Al + OH → Al1(OH)3 note that you do not have to write the 1’s, but it helps

Formulas to ions Example st 1 Split the formula into two “halves”. This is done AbBa → ___ + ___ or Ab(BOx)a → __ + __ by writing the symbol as the first half and the st remaining symbols as the second half. Do NOT 1 AbBa → A + B or Ab(BOx)a → A + BOx write any subscripts unless they are inside ( )’s or follow two capital letters (see example)

nd nd + - + - 2 The first ion is positive the second is negative. 2 AbBa → A + B or Ab(BOx)a → A + BOx

3rd Now determine the magnitude of the charge. Do the opposite of “Ions to formulas” by criss-crossing a+ b- a+ b- the subscripts into the position of charges. 3/4 AbBa → A + B or Ab(BOx)a → A + BOx

4th Note that to make sure you are right you need to check your charges with the periodic table or with 1+ 2- the charges on the polyatomic ions you were to Cu2O → Cu + O memorize. Note the charge for O matches the periodic table. So we assume the charges are correct. Trend For Charges 1+ 2- Any element within the vertical family headed by these K2SO4 → K + SO4 s and p block elements will have the corresponding charge. Note the charge for K and SO4 matches the periodic table and the memorized charge, so we assume the charges are correct.

1+ 2+ 3+ 4+/4-, 3- 2- 1- Also note, the subscript “4” was not criss-crossed. The invisible H , Be , B , C N , O , F , He no charge “1” outside the invisible “( )’s” was criss-crossed. The “4” follows two capital letters and stays put. Note that this is only a trend and does not apply in many situations. CuO → Cu1+ + O1- Note that the charge for O should be a 2- from the periodic table. So, we must double all charges to get it from 1- to 2-. So... CuO → Cu2+ + O2- Naming Ionic Compounds You never name formulas. Name 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 Example formula, you must first write the ions and then write the formula.

Naming Cations Naming Cations if s or p block elements if s or p block elements The name is the elemental name. Li1+ lithium. if d-block elements Al3+ aluminum The name is the elemental name followed by the if d-block elements charge as a Roman numeral in parentheses. Ni1+ nickel (I). Ni3+ nickel (III)

Naming Anions Naming Anions if mono-atomic (only one elemental symbol) if monatomic: Then the name is the elemental name, but with the F1- fluoride (not fluorine) ending changed to –ide. O2- oxide (not oxygen) As3- arsenide (not arsenic) if polyatomic (more than one elemental symbol) Then the name is the memorized name. if polyatomic: 2- CO3 carbonate To name a formula CN1- cyanide Dissociate into ions, then name the cation followed by the OH1- hydroxide 1- anion. NO3 nitrate 1- NO2 nitrite 2- To write a formula from a name SO4 sulfate 2- From the name write the formulas for the cation and the SO3 sulfite 3- anion, then write the formula. PO4 phosphate 3- PO3 phosphite

Note that there exists one polyatomic cation that should be memorized. 1+ NH4 is ammonium All other cations will be one elemental symbol. Solubility Rules That Must Be Memorized

Soluble Compounds Exceptions Most salts containing alkali metal ions and the ammonium ion are soluble. 1- Salts of nitrate, NO3 1- chlorate, ClO3 1- perchlorate, ClO4 1- acetate, CH3CO2 1- 1- 1- 1+ 2+ 2+ Most salts of Cl , Br , and I Halides of Ag , Hg2 , and Pb Compound containing fluoride, F1- Fluorides of Mg2+, Ca2+, Sr2+, Ba2+, and Pb2+ 2- 2+ 2+ 2+ 2+ Salts of sulfate, SO4 Sulfates of Ca , Sr , Ba , and Pb

Insoluble Compounds Exceptions 2- 1+ All salts of carbonate, CO3 Salts of NH4 and the alkali metal cations. 3- phosphate, PO4 2- oxalate, C2O4 2- chromate, CrO4 2- 1+ Most metal sulfides, S Salts of NH4 and the alkali metal cations. 1- 2- 1+ most metal hydroxides, OH , and oxides, O Salts of NH4 and the alkali metal cations.