Chem 120 - Chapter 7 Study Guide PDF

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Chapter 7 - chemical reactions professor garima garg...

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Chemistry 120 (Intro) – Chapter 7- Chemical Equations/Reactions -

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Balancing Chemical Equations: o Never change the subscripts o Count the number of each atom on both sides of equation o Balance polyatomic ions as one unit o Balance metals before non-metals o Trial and error States/Symbols used o (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous o  = reaction is heated o hv = reaction requires light o time span (ex. 2h, 3d) = how long reaction carried out o temperature (ex. -25C) = temperature reaction carried out o chemical formula (ex. MnO4) = chemical that acts as a catalyst  catalyst = substance that speeds up rate of chemical reaction. Not consumed during reaction so its ignored when balancing. Double Displacement Reactions o Two ionic compounds trading anions. AX + BY  AY + BX o AX and BY are (aq) aqueous solutions. Charges of ions don’t change. o Reactions will proceed if one of the following is true:  Solid compound forms = Precipitation Reaction  Gas formed = Gas Evolution Reaction  Water or weak acid formed = Acid-Base Reaction or Neutralization Reaction  IF NONE OF THESE OCCUR THEN THERE IS NO REACTION o Solubility = degree which a compound will dissolve in a solvent (usually water)  Soluble = dissolve significantly  Insoluble = will not dissolve, remains solid in solution  Reaction that produces insoluble product is described as precipitation because the product falls out of the solution like rain precipitates  REFER TO SOLUBILITY RULES ON PAGE 83 IN LECTURE NOTES Complete Ionic Equations/Net Ionic Equations o Complete Ionic Equations show all “species” in solution for given reaction  Species = specific ions/compounds o Ex: NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq) Na+(aq) + Cl-(aq) + Ag+(aq) + NO3(aq)  AgCl(s) + Na+(aq) + NO3-(aq) o Spectator Ions are ions that “hang out” and do not participate in the reaction  Ions that are on both sides of a reaction are spectator ions  Doesn’t go under chemical change  Main purpose is to maintain constant charge in solution o Ionic Equations only show those chemicals which participate in the reaction  Frist write total ionic equation and then cancel anything that appears on both sides Ex- Total Equation: Na+(aq) + Cl-(aq) + Ag+(aq) + NO3(aq)  AgCl(s) + Na+(aq) + NO3-(aq) Net Ionic Equation: Ag+(aq) + Cl-(aq)  AgCl(s)

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Gas Evolution Reactions o When produced in a reaction, some compounds will immediately decompose into other products  Carbonic Acid (H2CO3) will decompose into CO2(g) and H2O(l)  Ammonium Hydroxide (NH4OH) will decompose into NH3(aq) and H2O(l)  Sulfurous Acid (H2SO3) will decompose into SO2(g) and H2O(l) o Each produced water and gas compound formed by atoms left after water is removed from starting formula o If any of these compounds appear in a reaction, cancel it and replace with decomposition product  Sulfides  H2S  Carbonates and Bicarbonates  CO2(g) and H2O(l)  Sulfites and Bisulfites  SO2(g) and H2O(l)  Ammonium  NH3(g) and H2O(l) Acids/Bases o Acids = compounds which produce H+ ions in a solution o Bases = compounds which produce OH- ions in solution o Strong acids are strong electrolytes; completely dissociates into ions in solution  Six strong acids that must be known: HCl/ HBr/ HI/ HNO3/ H2SO4/ HClO4 o Weak acids are weak electrolytes  Examples of weak acids: H3PO4/H2CO3/HC2H3O2/HF o Strong and weak bases; similar definitions apply  Only strong bases needed for this class include:  Group I Hydroxides: LiOH/NaOH/KOH/etc..  Some Group II Hydroxides: Ba(OH)2/Sr(OH)2/Ca(OH)2 Oxidation/Reduction o Metal is oxidized = loses electrons, Non-metal is reduced = gains electrons  “O.I.L R.I.G.” = Oxidation is less, Reduction is gain  “L.E.O” the lion says “G.E.R.” = Loss of electron is oxidation, Gain of electron is reduction o We can separate the main reaction into pair of half-reactions, one for oxidation and one for reduction  Ex: 2Na + Cl2  2NaCl Oxidation: 2Na  2Na+ + 2e- Reduction: Cl2 + 2e-  2ClCombustion Reactions o Chemical reacts with oxygen gas forming various products  For this class we will only consider combustion of organic compounds containing C,H,O o Compounds react with oxygen to form carbon dioxide (CO2) and water (H2O)  [Organic Compound] + O2(g)  CO2(g) + H2O (g)  Ex: Combustion of Benzene (C6H6) = 2C6H6(l) + 15O2(g)  12CO2(g) + 6H2O(g) Single Displacement Reactions o One element replaces another which is present as an ion in a compound. A + BX  AX + B  A and B are metals/metal ions and X is an anion  Two types: Displacement of metal with another/Displacement of halogen with another  Activity series (p. 93 lecture notes)

Many metals can displace H+ changing it to H2 transforming charged hydrogen to neutral Some metals can displace H+ from water leaving OH- behind. For these cases, think of water as HOH When does the reaction “go”? o If metal is more active than another, it will displace it.  Ex: Zn(s) + 2AgNO3(aq)  Zn(NO3)2(aq) + 2Ag(s)  Zinc is more active than silver (higher on activity series), zinc will displace silver  Ex: [Consider the reverse] Ag(s) + Zn(NO3)2(aq)  no reaction  Silver is less active than zinc, it can’t displace it, therefore there is no reaction Single Displacement Reactions: Displacement of Halogens o Anions derived from Group VIII can be displaced by a more active halogen  Activity of halogens decreases down the group F2 > Cl2 > Br2 > I2 [Most Active  Least Active] Combination Reactions o Two chemicals combine into one new chemical. A + B  C  Often difficult to predict products of combination reactions so we study a few general cases o Oxide formation: Metals/Non-metals react with oxygen to form metal/non-metal oxide  Ex: 4Li(s) + O2 (g)  2Li2O(s) and C(s) + O2(g)  CO2(g) o Reactions of Oxides  Metal oxides react with water to form metal hydroxides Ex: Na2O(s) + H2O(l)  2NaOH(aq)  Non-metals react with water to form oxyacids Ex: CO2 + H2O  H2CO3  Metal oxides and non-metal oxides combine to form a salt  Metal oxide + Carbon dioxide  metal carbonate Na2O(s) + CO2(g)  Na2CO3(s)  Metal oxide + sulfur trioxide  metal sulfate CaO(s) + SO3(g)  CaSO4 (s) Decomposition Reactions o Reverse of combination reactions. Z ∆ X + Y (Delta on top of arrow = heating)  

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Ex: 2HgO(s) ∆ →

2Hg(l) + O2(g)

o Metal chlorates decompose to form metal chlorides and oxygen gas  Ex: 2KClO3(s) ∆ 2KCl(s) + 3O2(g) →...