Chapter 8 and 15 Notes - Lecture PDF

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Chapter 8 and 15: Chemical Reactions 

8.1: What is a chemical reaction? o Chemical reaction: one or more substances are converted into one or more new substances o Reactants: the starting material o Products: the newly created material o Chemical equation  Represents a chemical reaction  Reactants are on the lef  Products are on the right  Arrow symbol  means “react to give”  Reactants  products

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8.2-8.3: Balancing chemical equations o Takes energy to break a chemical bond o Balanced chemical equation  Shorthand way of representing a reaction  Number of atoms must be the same on both sides of an equation o Balancing equations  Balance by placing coefficients before reactants and products:  H2 + I2  2HI  Traditionally, chemists use the smallest whole numbers as coefficients

 Never balance equations using subscripts  Changing the subscript changes the compound o Guidelines for balancing equations  Write a skeletal equation by writing chemical formulas for each of the reactants and products  Balance atoms that occur in more complex substances first  Always balance atoms in compounds before atoms in pure elements  Balance atoms that occur as pure elements by adjusting on either side of the equation 

8.4: Types of reactions o Types:  Single-replacement reaction  AB + C  AC + B  One element replaces another in a compound  Double-displacement reaction  AB + CD  AD + BC  Two elements replace each other in two compounds  Decomposition reaction  AB  A + B  A compound breaks down into two or more substances  Combination reaction  A + B  AB  Two or more substances combine to make a compound  Combustion

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 Reacts with O2 (burns)  Fits into other classes  AB + CD  AD + BC 

15.1: Electrolytes and Nonelectrolytes o Electrolytes: medical and chemical term for ionic salt  Conduct electricity o Characteristic properties:  Acids: sour taste, litmus (indicator) turns red, eats away active metals  Bases: bitter taste, litmus turns blue, slippery when touching o Indicators: color change associates with the change of pH o Neutral substances: exhibits neither an acidic or basic behavior  pH = 7 o pH scale  Mathematical scale based on –log[H+]  pH less than 7 = acid  pH more than 7 = base  pH is 7 = neutral o Modern definitions of acids and bases:  Based on molecular definition of the compound  Based on electrolytes and nonelectrolytes  Electrolytes: conduct electricity  Nonelectrolytes: do not conduct electricity o How do you test whether or not a substance is an electrolyte or a nonelectrolyte?

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 Does it complete a circuit?  NaCl – strong electrolyte, conducted electricity and turned on the light  C12H22O11 – nonelectrolyte, does not turn on the light o Ions – mobile charge in a solution electricity = movement of electrons  Electrolytes break apart in aqueous solution allowing electrons to transfer between charged particles o Dissociation  Splitting of ionic compounds to individually charged components  Does not occur with molecular substances that dissolve without dissociation  Nonelectrolytes dissolve in water and will not conduct electricity  Electrolytes dissolve in water and dissociate into ions o Would you expect the following to be electrolytes or nonelectrolytes?  Al(NO3)3  Al3+ + 2NO3 Electrolyte 

15.2: Electrolytes Weak and Strong o Types of electrolytes  Strong electrolytes completely dissociate in water  100% dissociation = strong electrolyte  Weak electrolytes only partially dissociate in water

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15.3: Acid Weak and Strong o Arrhenius’ definition of acids

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 Any electrolyte that contains one or more hydrogen atom and produces H+ (actually H3O+) ions when dissolved in water  HCl + H2O  H3O+ + Clo Strong acids  Give a large number of H3O+ ions when placed in water  Strong electrolytes  Dissociate fully  Gives equal amounts of H+ as moles of acid dissolved o Weak acids  Partially dissociates when placed in water  Weak electrolyte  Small fraction of every mol dissociates to create H+  Weak o Number of H+ released doesn’t correlate with strength  Monoprotic acids  Produces only 1 H3O+ ion in solution  Diprotic acids  Produces 2 H3O+ ions in solution  Triprotic acids  Produces 3 H3O+ ions in solution 

15.4: Bases o Arrhenius’ definition of bases  Any electrolyte that contains a metal ion and hydroxide group (OH-) and produces hydroxide ions when dissolved in water

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o Weak and strong bases have similar definitions as weak and strong acids  Strong bases: readily dissociate to produce high numbers of OH Group 1 metals + OH Weak bases: only partially dissociate to produce OHo Acid-base neutralization  Occurs when the H3O+ of an acid is counteracted by the OH- of a base  HCl + NaOH  NaCl + H2O  Produce H2O + salt 

15.5: Another Definition for Acid and Base o Ammonia NH3 problems  Ammonia turns litmus blue but does not contain the OHion  Ammonia can be used to neutralize HCl solutions o Brønsted-Lowry definition of base  Any substance that removes an H+ ion from a solution o Brønsted-Lowry definition of acids and bases  Acids:  Anything that donates a proton (add an H3O+) to a solution  Bases:  Anything that accepts a proton (removes an H3O+) from a solution

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15.6: Weak Bases o Weak bases:

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 Slightly dissociates in water to produce a limited number of ions 

15.7: Is the Solution Acidic or Basic? o Defining solutions  Acidic if H3O+ is greater than OH Basic if OH- is greater than H3O+  Neutral if OH- = H3O+ o Aqueous solutions:  Ions and molecules disperse in water o Auto-dissociation:  The automatic separation and union of atoms in molecules when in a solvent such as water  The dissociation of water to form H3O+ and OHo Acidic  OH- < H3O+  1 x 10-7 < H3O+ o Basic  OH- > H3O+  OH- > 1 x 10-7 o Neutral  OH- = H3O+ o Important Equations  pOH = -log(OH-)

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8.5: Solubility and Precipitation Reactions o Precipitation reactions: between ionic species that results in the formation of a solid (precipitate)

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o Dissociation  Lattice breaks into its respective cation and anion o Spectator Ion: an ion that appears on both sides of the reaction  Do not precipitate and remain in the solution o Complete ionic equation: spectators included  Show all ions that take part in your reaction and spectator ions o Net ionic equation: spectators not included  Cancel spectator ions from two sides of the complete ionic equation 

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8.6: Introduction to Acid-Base Reactions...