1. Chapter 4 Chemical Reactions and Solution Stoichiometry

2.  water  single most important substance on earth  life evolved around water  essential for chemical Rx’s, weather, transportation, life  molecular polarity attributed to bent shape

3.  hydrates ionic solids  Video(click here) Video(click here)

4.  solubilty  dissolve-ability  quantity of solute that can be dissolved in a quantity of solvent ▪solute – substance being dissolved ▪solvent – substance doing the dissolving  general rule – “like dissolves like”  solution – homogeneous mixture

5.  electrolyte  aqueous solution that conducts an electric current ▪any solution that contains ions in dissolved or molten state ▪ions carry e - from negative electrode to positive electrode

6.  strong electrolyte  substance that completely ionizes/dissociates in water ▪soluble salts ▪example – NaCl (s)  Na + (aq) + Cl - (aq) ▪strong acids ▪example – HCl (L)  H + (aq) + Cl - (aq) ▪strong bases ▪example – NaOH (s)  Na + (aq) + OH - (aq)

7.  weak electrolytes  substances that only partially dissociate/ionize  100 particles added to solution; only a couple ionize/dissociate, rest remain unreacted  partially soluble salts ▪100 AgCl (s)  Ag + (aq) + Cl - (aq) + 99AgCl (s)  weak acids ▪100 HC 2 H 3 O 2(L)  5 H + (aq) + 5 C 2 H 3 O 2 - (aq) + 95 HC 2 H 3 O 2(aq)  weak bases ▪100 NH 3(g) + 100 H 2 O (L)  5 NH 4 + (aq) + 5 OH - (aq) + 95 NH 3(aq)

8.  nonelectrolyte  substance that does not form ions when dissolved  example – C 12 H 22 O 11(s)  C 12 H 22 O 11(aq)  molecules(excluding acids and bases)

9.  concentration of solutions  comparison of amount of solute to amount of solvent or solution  concentrate / dilute – qualitative terms  quantitative concentrations ▪molarity(M) = moles solute ▪exampleexample liter solution

10.  standard solution  solution of accurate known concentration  stock solution  concentrated solutions prepared or purchased for dilution  dilution  process of adding H 2 O to a concentrated or a stock solution  mole solute after dilution = mole solute before dilution  vol solution after dilution > vol solution before dilution

11.  dilution is an inverse relationship  as volume of solution increases, concentration decreases  M 1 V 1 = M 2 V 2

12.  precipitation reactions  double displacement reaction in which one of the products forms a precipitate ▪precipitate – insoluble solid ▪Solubility rules(table 4.1)  ionic equation – all soluble substance dissociate NaCl(aq) + AgNO 3 (aq)  NaNO 3 (aq) + AgCl(s) Na + (aq) + Cl - (aq) + Ag + (aq) + NO 3 - (aq)  Na + (aq) + NO 3 - (aq) + AgCl(s)

13. Na + (aq) + + NO 3 - (aq) Na + (aq) + NO 3 - (aq) + net ionic equation – only reacting materials - no spectator ions Cl - (aq)+ Ag + (aq) AgCl(s) 

14.  stoichiometry of precip reactions 1) write net ionic equation for Rx 2) calc. mol of each reactant 3) determine limiting reactant 4) calc. mass of product example

15.  acid-base reactions  acid – contain ionizable hydrogen ▪proton or H + donor  base – contain OH - ▪proton or H + acceptor  neutralization reaction ▪acid + base  water + salt ▪H + + OH -  H 2 O H 2 SO 4 (aq) + Ca(OH) 2 (aq)  2 H 2 O(l) + CaSO 4 (aq) 2 H + + SO 4 2- + Ca 2+ + 2 OH -  2 H 2 O(l) + Ca 2+ + SO 4 2-

16.  titration – controlled addition of a solution of known concentration to a solution of unknown concentration  video(click here)click here

17.  titrant – solution of known conc.  analyte – solution being analyzed  endpoint – point at which the indicator changes color  equivalence point – point at which enough titrant has been added to completely react with the analyte  example example

18.  oxidation – reduction reactions  reactions in which one or more e - ‘s are transferred 2 Na(s) + Cl 2 (g)  2 NaCl(s)  Na loses e - to Cl in the reaction ▪Na goes thru oxidation ▪Cl goes thru reduction  oxidation states or oxidation numbers ▪an arbitrary way of keeping track of e - ▪similar to ion charges ▪covalently bonded atoms have oxidation numbers ▪Imaginary charges

19.  rules for assigning oxidation numbers 1) an atom of a free element = 0, compounds = 0 a.Na = 0, Cl 2 = 0, H 2 O = 0 2) a monatomic ion = charge a.Na + = +1, N 3- = -3, 3) fluorine in compounds = -1 4) oxygen in compounds = -2, except in peroxides = -1 or with F then = +2 a.more electronegative = -, less electronegative = + 5) hydrogen in compounds = +1 6) the sum of oxidation numbers in a polyatomic ion = charge a.SO 4 2- = -2 example

20.  redox reactions characteristics  e - transfer ▪literally w/ atoms and ions ▪figuratively w/ atoms and molecules  oxidation and reduction occur simultaneously and equally ▪substance being oxidized = reducing agent ▪e - donor ▪substance being reduced = oxidizing agent ▪e - acceptor

21.  balancing Redox Rx’s(half reaction method) 1) write separate half-reactions for oxidation and reduction 2) balance each half-reaction A. first atomically(except H and O) I.balance oxygen atoms by adding H 2 O II.balance hydrogen by adding H + III.if reaction occurs in alkaline conditions add OH - for each H + to both sides of equation a)combine H + and OH - on same side to form water B. then electrically 3) equalize e - ‘s in oxidation with e - ‘s in reduction 4) add half-reactions and cancel identical species 5) check for atomic and electrical balance - exampleexample