Objectives III. Reactions A. Reaction types 1. Acid-base reactions 2. Precipitation reactions 3. Oxidation-reduction reactions a. Oxidation number b. The role of the electron in oxidation- reduction

Objective (cont.) B. Stoichiometry 1. Ionic and molecular species present in chemical systems: net ionic equations

Single Replacement Reactions Element + Compound  Element + Compound If the element on the left is a metal, it must be more active than the metal (+) in the compound. The electromotive series (reduction potentials) tell you reactivity

Single Replacement Reactions If the element on the left is a halogen, it must replace a less active halogen (-) in the compound. The halogens reactivity is in order on the Periodic Table

Single Replacement Reactions Active metals replace less active metals in compounds in aqueous solutions Example: Magnesium turnings are added to an aqueous solution of iron(III) chloride Copper shavings are added to an aqueous solution of sodium chloride

Single Replacement Reactions Active metals replace hydrogen in acids. Mossy zinc is added to hydrochloric acid VERY active metals replace hydrogen in water Sodium metal is dropped into water Cadmium metal is subjected to steam

Single Replacement Reactions Active nonmetals (halogens) replace less active nonmetals (lower halogens) Ie, a halogen above replaces a halogen below it. Example: Chlorine gas is bubbled into a potassium iodide solution Solid iodine crystals are added to a solution of sodium chloride

Double Replacement (Metathesis) For a double replacement reaction (switching 2 positives), at least ONE of the following must occur: 1) a gas is formed, or 2) water is formed, or 3) a precipitant is formed.

Gases that form directly or are products of decomposition Common gases: H 2 S—any sulfide salt + acid  H 2 S + a salt NH 3 — ammonium salt + soluble hydroxide + heat  NH 4 OH as a product decomposes to NH 3 + H 2 O SO 2 —sulfite + acid  H 2 SO 3 + a salt  H 2 SO 3 as a product decomposes to SO 2 + H 2 O CO 2 —carbonate salt + acid  H 2 CO 3 + a salt  H 2 CO 3 as a product decomposes to CO 2 + H 2 O

Precipitates are not soluble in aqueous solutions SOLUBILITY RULES

All nitrates are soluble All ammoniums are soluble All Group I (alkali metal cations) are soluble All chlorates are soluble All acetates, exc Ag +  I, Ammonium, CAN SOLUBLE in WATER

Soluble w/ Exceptions All sulfates are soluble, EXCEPT, Ca,Sr,Ba,Ag,Pb,Hg All chlorides, bromides, iodides, are soluble, EXCEPT Ag, Pb,Hg Sulfates except caserba and Ag led Hg Halogens except Ag led Hg

INSOLUBLES except Group I & NH 4 + Carbonates are insoluble Phosphates are insoluble Sulfites are insoluble Chromates are insoluble, EXCEPT, Ca, Sr Hydroxides are insoluble, EXCEPT, Ca,Sr,Ba Sulfides are insoluble, EXCEPT, Group II

Double replacement AB + CD  CB + AD AB and CD must be water soluble Do not bring subscripts across the  Test each product for gas, water or insoluble If none of the above, NO REACTION

Net Ionic Reactions For each compound, determine if it is soluble. If it is, break it apart (+/-) Soluble: Na 2 SO 4 (aq)  2Na + + SO 4 2- Insoluble: doesn’t split Products H 2 CO 3, H 2 SO 3, NH 4 OH break apart as gases + water Solids (s) or (cr) and Gases (g) don’t split

Net Ionic Reaction Ammonium carbonate + nickel chloride in aqueous solution Step 1: write balanced formulas (NH 4 ) 2 CO 3 (aq) + NiCl 2 (aq) Step 2: complete the reaction and balance (NH 4 ) 2 CO 3 (aq) + NiCl 2 (aq)  NiCO 3 (aq) + 2NH 4 Cl (aq)

Step 3: Determine solubility and split 2NH 4 + (aq) + CO 3 2- (aq) + Ni 2+ (aq) + 2Cl - (aq)  NiCO 3 (s) + 2NH 4 + (aq) + 2Cl - (aq) Step 4: Cross off those matches from left to right. The net ionic is the leftovers. CO 3 2- (aq) + Ni 2+ (aq)  NiCO 3 (s) “final answer”

Another example Na 2 CO 3 (aq) + HCl (aq)  Na 2 CO 3 (aq) + 2HCl (aq)  H 2 CO 3(aq) + 2NaCl (aq) 2Na + + CO H + + 2Cl -  H 2 O + CO 2(g) + 2Na + + 2Cl - Final answer CO 3 2- (aq) + 2H + (aq)  H 2 O + CO 2(g)

Another example Carbon dioxide gas is bubbled into an aqueous solution of magnesium hydroxide When carbon dioxide is bubbled into water, it forms carbonic acid. Carbonic acid can be a reactant but not a product. H 2 CO 3(aq) + Mg(OH) 2(aq) 

H 2 CO 3(aq) + Mg(OH) 2(aq)  MgCO 3 (s) + 2H 2 O H 2 CO 3(aq) + Mg 2+ (aq) + 2OH - (aq)  MgCO 3 (s) + 2H 2 O Since nothing cancels, that is the net ionic reaction.

Review equations Solid calcium sulfite + acetic acid Liquid bromine is added to a container of sodium iodide crystals. An aluminum strip is immersed in a solution of silver nitrate Solid ferric carbonate is added to hydrochloric acid A solution of ammonium nitrate is added to aqueous sodium hydroxide

Stoichiometry with Solutions Molarity = mol/Liter of solution M*L = moles {M is Molarity} G/MM = moles {MM is molar mass} PV/(RT) = moles

Acid-Base gas formation What volume of gas is generated at 755 torr and 35 o C when collected over water from the complete reaction of 7.6 g sodium sulfide with sufficient hydrochloric acid solution? (remember this will be collected over water)

Precipitant formation How many grams of precipitate can be collected when 14.3 g of magnesium chloride reacts with 40.0 mL of 2.5 M silver nitrate solution?

Assay/ Percent purity A sample of g of an unknown compound containing barium ions is dissolved in water and treated with excess sodium sulfate. If the mass of precipitate formed is g, what is the % Ba in the original unknown compound?

Assay/Purity Titration A mL sample of a sulfuric acid solution used as a battery acid (d = g/mL) is diluted to mL with water. A mL sample of the diluted acid requires mL of M KOH to the equivalence point. What is the M and mass % sulfuric acid in the battery?

Oxidation-Reduction (REDOX) Oxidation-reduction reactions occur when there is a transfer of electrons One element will gain electrons(  -, therefore, reduced in oxidation #) One element must lose an equivalent # of electrons(  +, therefore, oxidized) OXIDATION = REDUCTION

REDOX The element reduced is the OXIDIZING agent The element oxidized is the REDUCING agent

REDOX Balancing 1. Determine the oxidation number of every element {H = +1, O = -2, elements = 0} 2. Determine the 2 elements that changed oxidation numbers 3. Write the two ½ reactions 4. Balance that element 5. Balance oxidation number by adding electrons 6. Balance charge by adding H + ions in acidic, OH - ions in basic 7. Balance hydrogen by adding water 8. Add the two ½ reactions together.

IO Mn 2+  I - + MnO 2 In acidic solution. The sum of the oxidation of a polyatomic ion is the charge. The sum of the oxidation number of a compound is zero.

IO Mn 2+  I - + MnO 2 In acidic…..

SO 2 + I 2  SO 3 + I - In basic solution

H 2 O 2 + ClO 4 -  ClO O 2 In basic solution (O anion exception for peroxides)

Assay/Purity P. 99, #65, 68

Assay and Purity w/ Redox An iron ore sample weighing g is dissolved in HCl and becomes Fe +2. The solution is then titrated with 29.43mL of M KMnO 4. What is the % iron in the sample?