1. Aqueous Reactions and Solution Stoichiometry

2. ► Aqueous Solutions – a solution which water is the solvent ► Solution = Solute + Solvent ► Solute = smaller part being dissolved ► Solute = larger fraction doing the dissolving ► Molarity is used to express conc.

3. Molarity ► M = moles of solute/ liters of solution ► A 0.38g sample of sodium nitrate is placed in a 50.0 ml volumetric flask and is filled with water to the mark. What is the molarity of the solution?

4. Molarity ► How many grams of potassium permanganate are needed to prepare 2.0 liters of a 0.2 molar solution of the salt?

5. Dilution Problems ► M1V1= M2V2 ► How many liters of 12 molar sulfuric acid is necessary to make 2 liters of a 6 molar conc. of the acid?

6. Specific Gravity or Density Problems ► Commercially available conc. HCl is in a aq sol cont 38% HCl by mass. D = 1.19g/ml a) what is the molarity of the solution? b) how many ml of conc HCl are required to make 1.00 l of a 0.10 molar solution?

7. Electrolytes ► Solutes that separate into tow or more ions when dissolved in water are called electrolytes. ► Solutes that remain uncharged molecules are called non- electrolytes. ► Either may be soluble in water.

8. Electrolytes ► NaCl (s)  Na + + Cl- ionic solids dissociate into ions. ► Electrical Conductivity Strong electrolytes conduct electricity very Strong electrolytes conduct electricity very well because ions totally dissociate well because ions totally dissociate Weak electrolytes only weakly conduct Weak electrolytes only weakly conduct because ions only partially dissociate because ions only partially dissociate Nearly all ionic cmps are strong and molecular cmps are usually non-electrolytes Nearly all ionic cmps are strong and molecular cmps are usually non-electrolytes

9. Acids, Bases, And Salts ► Acids – are solutes that can ionize to produce hydrogen ions H+(aq) ► Strong acids dissociate and exist almost totally as ions ► Ex HCl (aq)  H + + Cl -

10. Weak Acids ► HC 2 H 3 O 2  H + + C 2 H 3 O 2-  Weak acids only partially dissociate the double arrows indicate that the reaction occurs in significant amounts in both directions Acedic Acid or Vinegar is the classic example.

11. Acids Cont. ► H+ is simply a proton – acids are often called proton donors ► Monoprotic acids – contribute one H+ per molecule of acid ex. HCl or HNO 3. ► Diprotic acids contribute two H+ per moleclue ex. H 2 SO 4 ► Polyprotic acids contribute greater than one H+ ex. H 2 SO 4 or H 3 PO 4.

12. Bases ► Are substances that accept H+ ions or a proton. ► Ex H +(aq) + OH- (aq)  H 2 O (l) ► Ex H +(aq) + NH 3-((aq)  NH 4(aq) ► OH- & NH 3- are the bases. OH- is the most common base in aqueous solution. Sometimes bases are defined as substances with increased OH- ions in solution.

13. Note ► Acids and bases are electrolytes – if a strong acid or base they are strong electrolytes if a weak acid or base a weak electrolyte. ► Strong Acids HCl (aq) Hydrochloric H 2 SO 4 Sulfuric HNO 3 (aq) Nitric HBr Hydrobromic HClO 4 (aq) Prechloric HI Hyrooiodic HClO 3 (aq) chloric

14. ► Weak Acids HF Hydrofluoric H 2 C 2 H 3 O 2 Acedic ► Strong bases Grp I Hydroxides LiOH, NaOH, KOH, RbOH, and CsOH Grp II Hydroxides Ca(OH) 2, Sr(OH) 2, Ba(OH) 2 Weak Base NH 3 Ammonia

15. Salts ► Are ionic compounds formed from acids by the replacement of one or more H+ ions with some other cation. ► For example replacing the H+ ion in HCl(aq) with Na+ ion yields NaCl.

16. Identifying strong and weak electrolytes ► most salts are strong electrolytes ► most acids are weak electrolytes except for the strong acids ► most strong bases hydroxides of metals are strong electrolytes except NH 3 which is a weak base and electrolyte

17. ► Neutralization reactions are those between an acid and a base they yield salt and water ► Ionic Equations a. varying amounts of detail can be included in chem eq depending on what information is relevant to the problem at hand a. varying amounts of detail can be included in chem eq depending on what information is relevant to the problem at hand b. molecular equations - have all the species written as associated molecules even though they may be disassociated b. molecular equations - have all the species written as associated molecules even though they may be disassociated c. complete ionic eq – strong electrolytes are written as dissociated ions; molecules, solids, and weak electrolytes are written as associated ions c. complete ionic eq – strong electrolytes are written as dissociated ions; molecules, solids, and weak electrolytes are written as associated ions

18. d. net ionic equations – an ionic eq in which the spectator ions have been cancelled. d. net ionic equations – an ionic eq in which the spectator ions have been cancelled. ► Rules for Converting Molecular eq to Ionic 1. Ionic Sub indicated as dissolved in solution such as NaCl (aq) are normally written as ions 1. Ionic Sub indicated as dissolved in solution such as NaCl (aq) are normally written as ions 2. Ionic substances that are insoluble (do not dissolve readily) either as reactants or products (precipitates) are rep as formulas of the compounds 2. Ionic substances that are insoluble (do not dissolve readily) either as reactants or products (precipitates) are rep as formulas of the compounds

19. d. Molecular substances that are strong electrolytes such as strong acids are written as ions. e. Molecular substances that are weak electrolytes are represented as their molecular formulas f. Spectator ions – ions in an equation that do not take part in the equation are cancelled from both sides of the equation

20. Examples ► Ba(C 2 H 3 O 2 ) 2 + K 2 CrO 4  ► Ba(C 2 H 3 O 2 ) 2 + K 2 CrO 4  BaCrO 4 + 2KC 2 H 3 O 2 Ba + 2C 2 H 3 O 2 + 2K + CrO 4  BaCrO 4 + 2K Ba + 2C 2 H 3 O 2 + 2K + CrO 4  BaCrO 4 + 2K +2 C 2 H 3 O 2 +2 C 2 H 3 O 2 Ba + CrO 4  BaCrO 4 Ba + CrO 4  BaCrO 4

21. Examples ► Ca(C 2 H 3 O 2 ) 2 + H 2 C 2 O 4  ► Ca(C 2 H 3 O 2 ) 2 + H 2 C 2 O 4  CaC 2 O 4 + HC 2 H 3 O 2 ► Ca+ 2C 2 H 3 O 2 + H 2 C 2 O 4  CaC 2 O 4 + HC 2 H 3 O 2 ► Weak acids do not split

22. Examples ► NH 4 NO 3 + NaOH  ► NH 4 NO3 + NaOH  NH 4 OH + Na + NO 3 ► NH 4 + NO 3 + Na + OH  NH 3 + H 2 O + Na + NO 3 + NO 3 ► NH 4 + OH  NH 3 + H 2 O

23. Metathesis Reactions ► Reactions in which two ionic reactants exchange ion partners ► AX + BY  AY + BX ► Precipitation reactions are those metathesis reactions in which an insoluble solid product or precipitate forms A. Precipitate –insoluble solid product A. Precipitate –insoluble solid product B. Solubility – the amount of a substance that can dissolve in a given amount of solvent B. Solubility – the amount of a substance that can dissolve in a given amount of solvent

24. ► (expressed in g/100 ml or in liters/mol) A substance is considered insoluble if its solubility is less that 0.01 mol/l. ► Predicting solubility is a matter of know the solubility rules (Table 4.3 or what I should know for AP list) ► Some metathesis reactions from water or another weak electrolyte or non-electrolyte in ionic and net ionic equations (appear as associated species) ► Some methesis reactions form gases these are usually binary compounds of non-metals exam ples H 2 S, CO 2, HCN.

25. Redox It’s electro-magic ► I. Oxidization Numbers A. Oxidization – loss of electrons. LEO B. Reduction – gain of electrons – GER C. When one reactant loses e’s another reactant must gain them

26. Intro to electrochemistry ► D. reactions in which electrons are trans are called oxidization – reduction or redox reactions ► E. oxidization number (ON) rules 1. atoms in the elemental form – oxidization # = 0 1. atoms in the elemental form – oxidization # = 0 2. monatomic ions – the oxidization # = the charge of the ion ex Na + = 1 S -2 = -2 3. non-metals usually have neg ON although they can sometimes be positive a> Oxygen usually = -2 except peroxides O 2 -2 = each O = -1

27. b) hygrogen usually = +1 when bonded to non-metals = -1 when bonded to metals c) Fluorine = -1 in all cmps – halogens = -1 in most binary cmps but when combine with oxygen as in oxyanions they have a + number 4. The sum of oxidizations #’s a) in a neutral cmp = 0 a) in a neutral cmp = 0 b) in a polyatomic ion = charge of ion b) in a polyatomic ion = charge of ion Work sample exercise 4.6