1. UNIT 6 SOLUTION CHEMISTRY

2. KEY TERMS  Activity Series - A list of elements in order of chemical reactivity  Dispersion - Uniform spreading of matter within another form of matter  Dissociation - Separation of ions when an ionic solute is dissolved in a solvent  Reduced - The oxidation number of a reactant is reduced by gaining electrons  Oxidized - The oxidation number of a reactant is reduced by gaining electrons  Precipitation Reaction - The formation of a solid in a solution or inside another solid during a chemical reaction

3. KEY TERMS  Electrolyte - Solution that contains free ion and conducts electricity  Non-electrolyte - Does not dissociate into ions in solution and does not conduct electricity  Strong Electrolyte - Completely ionizes in solution and is a good conductor of electricity  Weak Electrolyte - Does not completely ionizes in solution and is a poor conductor of electricity

4. SOLUTION CHEMISTRY  There are two types of chemical reactions that are almost exclusively conducted within solutions:  Single Replacement Reactions: A + B Y → AY + B  Double Replacement Reactions: AX + BY → AY + BY  Some solutions will conduct electricity. These are called electrolytic solutions.

5. SINGLE REPLACEMENT REACTIONS  A + B Y → AY + B  Whether a reaction of this type can actually occur is dependent on two things:  The relative reactivity of the elements involved in the reaction  The ability to support the migration of electrons from one element to another (Reduction-Oxidation)

6. SINGLE REPLACEMENT REACTIONS ACTIVITY SERIES  The higher the metal is on the Activity Series, the greater the reactivity.  More active metals will replace less active metals.  The less active metal will precipitate out of the chemical reaction and settle to the bottom.  Iron is less reactive than sodium, therefore if iron ions are in the solution, they will come out of solution if sodium is added to the solution and precipitate as solid iron.  If copper was then added to the solution, nothing would happen to either the solution or the copper because copper is less reactive than the sodium ions now in the solution.

7. ACTIVITY SERIES OF HALOGENS Image used courtesy of” http://mschemedge.com/reactiontype3.htm

8. SINGLE REPLACEMENT REACTIONS REDUCTION-OXIDATION  Every Single Replacement reaction involves the transfer (migration) of electrons from one element to another.  The element that gains electrons becomes more negatively charged and therefore is said to be reduced (oxidation number is reduced).  The element that loses electrons becomes less negatively charged and therefore is said to be oxidized (oxidation number is increased).  Reactions where this transfer of electrons occurs are referred to as Reduction-Oxidation (REDOX) reactions.  All single replacement reactions are also REDOX reactions.

9. REDOX REACTIONS  All REDOX reactions result in the reduction of the oxidation number of one element as a reactant and the increase in the oxidation number in another reactant element when products are being formed.  Example:2 Al + 3 CuCl 2 → 2 AlCu 3 + 3 Cu All elements not combined with other elements are, by definition, neutral. They have an oxidation number of zero. The copper in the reactant compound is 2+ Copper went from 2+ to 0: REDUCED The aluminum in the product compound is 3+ Aluminum went from 0 to 3+: OXIDIZED

10. IDENTIFICATION OF OXIDATION NUMBERS  The oxidation number of elements is always zero.  Oxygen found in a compound almost always has an oxidation # of -2  Hydrogen found in compounds, except hydrides, always has an oxidation # of +1  The sum of all of the oxidation #’s for all of the elements in a compound must always equal zero.  Compounds are neutral regardless of how many elements are present.  The sum of all of the oxidation #’s in a polyatomic ion must equal the charge of the polyatomic ion.

11. OTHER REDOX REACTION TYPES  In addition to Single Replacement reactions, all of the reaction types are REDOX except Double Replacement reactions:  Synthesis : If one of the reactants is an element (oxidation # = zero), it will be oxidized or reduced when forming a compound.  Decomposition : If one of the products is an element (oxidation # = zero), it will be oxidized or reduced when formed from the breakdown of a compound.  Combustion : One of the reactants is the element oxygen with an oxidation # = zero, it will be reduced when forming both carbon dioxide and water.

12. INSOLUBILITY  There are many reasons why substances may not dissolve in a solvent  As discussed in our last unit, the nature of the solvent may not match the type of solvent with respect to polarity  But why can water dissolve certain ionic compounds and not others?  Some ionic bonds are too strong for water to separate because of the amount of charges holding the elements together in a compound.  The greater the difference in charge, the stronger the bonds  Some ionic bonds are too strong for water to separate because of how close together the elements are to each other  The shorter the distance between ionic nuclei, the stronger the bond. Coulomb’s Law: Bond strength is proportional to the product of the charges and inversely proportional to the distance between nuclei squared.

13. DOUBLE REPLACEMENT REACTIONS  Characterized by two reactants, both of which are ionic compounds.  The products will always be the new compounds that are formed by switching ions.  Whether the reaction will take place depends on whether or not the products are soluble in water.  If a Double Replacement reaction actually occurs, a precipitate will be formed.  A precipitate is an insoluble solid that is sometimes formed when two solutions of ionic compounds are mixed

14. STAAR CHART SOLUBILITY RULES Use the solubility chart to determine if the reaction will proceed If ONE product is NOT soluble, the reaction will go forward. If BOTH products ARE soluble in water, there will be NO REATION.

15. DOUBLE REPLACEMENT REACTION EXAMPLES Ca(NO 3 ) 2 + Na 3 PO 4  ? 3Ca(NO 3 ) 2 + 2 Na 3 PO 4  Ca 3 (PO 4 ) 2 + 6NaNO 3  The expected products of this reaction are Ca 3 (PO 4 ) 2 and NaNO 3  Checking the solubility chart, Ca 3 (PO 4 ) 2 is insoluble (will form a precipitate), so the reaction will proceed NH 4 I + K 2 SO 4  ? NH 4 I + K 2 SO 4  No Reaction (N.R.)  The expected products of this reaction, (NH 4 ) 2 SO 4 and KI, are both soluble in water  No reaction is expected

16. DISSOCIATION AND DISPERSION  Dissociation – When ionic compounds dissolve, they separate into the ions that were bonded together.  Example: Potassium sulfate will dissociate into potassium ions and sulfate ions. K 2 SO 4 → 2K 1+ + SO 4 2-  More examples of dissociation:  NaCl → Na 1+ + Cl 1-  MgCl 2 → Mg 2+ + 2 Cl 1-  Al 2 (SO 4 ) 3 → 2 Al 3+ + 3SO 4 2-  Dispersion – When complete dispersion occurs, the molecules of the solute are uniformly mixed with the molecules of the solvent.  This occurs with covalent compounds.

17. ELECTROLYTIC SOLUTIONS  Electrolytic solutions conduct an electrical current  Ions dissociate when dissolved in water  Ions, because they are charged particles, act as a path for the flow of electricity  The more ions present, the more electrolytic the solution is  Non-electrolytic solutions do not conduct an electrical current

18. ELECTROLYTIC SOLUTIONS Soluble Ionic Solute Non-Soluble Ionic Solute Non-Soluble Covalent Solute ElectrolyteYesNo

19. CONTROLLING ELECTROLYTIC SOLUTIONS  In order to increase the electrolytic properties of a solution you must increase the concentration of ions present in the solution  Reduce the volume of the solution while maintaining the amount of electrolytic solute  Increase the amount of electrolytic solute without modifying the volume  Using a different electrolytic solute which produces more ions when it dissociates  In order to decrease the electrolytic properties of a solution you must decrease the concentration of ions present in the solution  Dilute the electrolytic solution  Using a different electrolytic solute which produces less ions when it dissociates