Presentation on theme: "Reactions in Aqueous Solutions Chapter 4. GENERAL PROPERTIES."— Presentation transcript:
Reactions in Aqueous Solutions Chapter 4
HCl ( l ) H + (aq) + Cl − (aq) CH 3 COOH (aq) H + (aq) + CH 3 COO − (aq)
Review of Concepts The diagrams here show three compounds AB 2 (a), AC 2 (b), and AD 2 (c) dissolved in water. Which is the strongest electrolyte and which is the weakest? (For simplicity, water molecules are not shown.)
Example K 3 PO 4 (aq) + Ca(NO 3 ) 2 (aq) Example 4.2 page 125 in textbook
Review of Concepts Which of the diagrams here acuratly describes the reaction between Ca(NO 3 ) 2 (aq) and Na 2 CO 3 (aq) ? For simplicity, only the Ca 2+ (yellow) and CO 3 2− (blue) ions are shown.
ACID Sour taste Color changes in plant dyes React with metals to produce H 2 gas React with carbonates and bicarbonates to produce CO 2 gas Aqueous acid solutions conduct electricity Taste bitter Feel slippery Color changes in plant dyes Aqueous base solutions conduct electricity BASE General Properties
ACID Proton donor Monoprotic Diprotic Triprotic Proton acceptor BASE Brønsted Acid and Bases
STRONG ACIDS HI HBr HClO 4 HCl H 2 SO 4 HNO 3 NaOH KOH LiOH RbOH CsOH Ca(OH) 2 Ba(OH) 2 Sr(OH) 2 STRONG BASES Strong acids/bases are strong electrolytes and will completely dissociate in water.
Review of Concepts Which of the following diagrams best represents a weak acid? Very weak acid? Strong acid? The proton exists in water as the hydronium ion. All acids are monoprotic. (For simplicity, water molecules are not shown.)
Reaction between an acid and a base Generally aqueous solutions result in water and a salt Ex: HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O ( l ) *this is a strong acid and strong base so they completely dissociate and the net ionic equation is H + (aq) + OH − (aq) H 2 O( l ) Ex: HCN(aq) + NaOH(aq) NaCN(aq) + H 2 O( l ) *this is a weak acid and strong base so the acid does not completely ionize in water. When writing the ionic and net ionic equations you cannot break the weak acid apart! The net ionic equation is HCN (aq) + OH − (aq) CN − (aq) + H 2 O ( l ) Acid-Base Neutralization
Gas formation Certain salts react with acids to produce gaseous products HNO 3 breaks down into H 2 O ( l ) + NO 2 (g) + NO (g) H 2 CO 3 breaks down into H 2 O ( l ) + CO 2 (g) H 2 SO 3 breaks down into H 2 O ( l ) + SO 2 (g) NH 4 OH breaks down into H 2 O ( l ) + NH 3 (g) H 2 S (g) CO 2 (g) H 2 (g) If you get one of these as a product in your molecular equation, they immediately breakdown as above Gasses do not ionize
Double Replacement Rxns Review Driving ForceHow do you recognize it? Precipitate You must memorize the solubility rules. Any compound formed from two ions can be recognized as soluble (written as separate ions) or as a precipitate (written as a molecule). Gas formed You must memorize the combinations that decompose into gases (there are 4). You must also memorize the gases that form. For example, when you H 2 SO 3 as a product, you must know it decomposes into H 2 O and SO 2 gas. Weak electrolyte You must memorize the short list of strong acids and strong bases so you will recognize all the weak acids and bases that dissolve, but do not dissociate into ions. The weak base ammonia, NH 3, is in this category. It exits in water as NH 3 (aq) and only slightly forms the ions NH OH −
OXIDATION REACTION Reaction that involves the loss of electrons Contains reducing agent- donates electrons Involves the gain of electrons Contains oxidizing agent- accepts electrons REDUCTION REACTION Half-reaction OIL RIG
Oxidation Number Charge of the atom would have in a molecule if electrons were transferred completely Rules Uncombined elements = 0 Neutral compounds sum = 0 Ion = ion charge (polyatomic ions sum to charge) Exceptions Hydrogen +1 w/ nonmetals, −1 w/ metals Oxygen −2 except w/ fluorine (+2), in peroxides (−1) Fluorine ALWAYS −1
More common oxidation numbers are in red.
Example 2 Mg (s) + O 2 (g) 2 MgO (s)
Types of Redox Reactions Combination 2 Al (s) + 3 Br 2 ( l ) 2 AlBr 3 (s) Decomposition 2 NaH (s) 2 Na (s) + H 2 (g) Combustion C 2 H 8 (g) + 5 O 2 (g) 3 CO 2 (g) + 4H 2 O ( l ) Displacement Hydrogen Ca (s) + 2 H 2 O ( l ) Ca(OH) 2 (s) + H 2 (g) Metal Zn (s) + 2 HCl (aq) ZnCl 2 (aq) + H 2 (g) Halogen Cl 2 (g) +2 KBr (aq) 2 KCl (aq) + Br 2 ( l ) Disproportionation 2 H 2 O 2 (aq) 2H 2 O ( l ) + O 2 (g)
Activity Series For Halogens: F 2 > Cl 2 > Br 2 > I 2
Elements most likely to undergo disproportionation
Concentration Molarity = moles of solute liters of solution nVnV M = KCl (s) K + (aq) + Cl − (aq) H2OH2O Ex: 1M KCl solution Ex: 1M Ba(NO 3 ) 2 solution Ba(NO 3 ) 2 (s) Ba 2+ (aq) + 2 NO 3 − (aq) H2OH2O
Example How many grams of potassium dichromate (K 2 Cr 2 O 7 ) are required to prepare a 250 mL solution whose concentration is 2.16M?
Dilutions M i V i = M f V f
Example Describe how you would prepare 5.00x10 2 mL of a 1.75M H 2 SO 4 solution, starting with an 8.16M stock solution of H 2 SO 4.
Review of Concepts What is the final concentration of a 0.6M NaCl solution if its volume is doubled and the number of moles of solute is tripled?
Example A g sample of an ionic compound containing chloride ions and an unknown metal is dissolved in water and treated with and excess of AgNO3. if g of AgCl precipitate forms, what is the percent by mass of Cl in the original compound?
Example How many mL of a 0.610M NaOH solution are needed to neutralize 20.0 mL of a 0.245M H 2 SO 4 solution?
Example A mL volume of M KMnO 4 solution is needed to oxidize mL of a FeSO 4 solution in an acidic medium. What is the concentration of the FeSO 4 solution in molarity? The net ionic equation is 5Fe 2+ + MnO 4 − + 8H + Mn Fe H 2 O