AP Chemistry Study Guide 8 v Stomach acid and heartburn Ø The cells that line your stomach produce hydrochloric acid To kill unwanted bacteria To

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1 AP Chemistry Study Guide 8 v Stomach acid and heartburn Ø The cells that line your stomach produce hydrochloric acid To kill unwanted bacteria To help break down food To activate enzymes to break down food Ø If the stomach acid breaks up into your esophagus it irritates those tissues resulting in heartburn Acid reflux v Curing heartburn Ø Mild cases of heartburn can be cured by neutralizing the acid in the esophagus Swallowing saliva which contains bicarbonate ion Taking anti- acids that contain hydroxide ions and or carbonate ions v GERD Ø Chronic heartburn I a problem for some people Ø GERD (gastro- esophageal reflux disease) is chronic leaking of stomach acid into the esophagus Ø In people with GERD the muscles separating the stomach form the esophagus do not close tightly allowing stomach acid to leak into the esophagus v Properties of acids Ø Sour taste Ø Ability to dissolve many metals Ø Ability to neutralize bases Ø Change blue litmus paper to red v Six strong acids Ø Hydroxilic acid HCl Ø Hydrobromic acid HBr Ø Hyrdoiodic acid HI Ø Sulfuric acid H2SO4 Ø Nitric acid HNO3 Ø Perchloric acid HciO4 v Structures of acids Ø Binary acids have acid hydrogens attached to the nonmetal atom v Structure of acids Ø Oxyacids have acid hydrogens attached to an oxygen atom Ø Diprotic Ø Monoprotic Ø Carboxylic have COOH group Ø Only the first H in the formula is acidic v Properties of bases Ø Tastes bitter Alkaloids = plant product that is alkaline Often poisonous Ø Feel slippery Ø Ability to turn red litmus paper blue

2 Ø Ability to neutralize acids v Strong bases Ø Hydroxides/oxides of groups IA and IIA metals Excluding Be and Mg v Indicators Ø Indicators are chemicals that change color depending on the solutions acidity or basicity Ø Many vegetable dyes are indicators Anthocyanin Ø Litmus Form Spanish moss Red in acid blue in base Ø Phenolphthalein Found in laxatives Red in base colorless in acid v Definitions of aids and bases Ø Arrhenius definition Bases on H+ (acid) and OH- (base) Ø Bronsted- Lowry definition Based on reactions in which H+ is transferred Acid = H+ donor Base = H+ acceptor Ø Lewis definition All acid is electron poor (electron pair acceptor) Base electron rich (electron pair donor) v Arrhenius theory Ø Acids produce H+ ions in aqueous solution v Hydronium ion Ø The H+ ions produced by the acid ware so reactive they cannot exist in water H+ ions are protons Ø Instead they react with water molecules to produce complex ions mainly the hydronium ion H3O There are also minor amount of H+ with multiple water molecules v Arrhenius theory Ø Bases produce OH- ions in aqueous solution v Arrhenius acid- base reactions Ø The H+ from the acid combines with the OH- from the base to make molecules of water Ø The cation from the base combines with the anion from the acid to make salt Ø Acid + base à salt + water v Problems with Arrhenius theory Ø It does not explain why molecular substances such as NH3 dissolve in water to from basic solutions even though they do not contain OH- ions

3 Ø It does no explain how some ionic compounds such as Na2CO3 or Na2O dissolve in water to form basic solutions even though they do not contain OH- ions Ø It does not explain why molecular substances such as CO2 dissolve in water to form acidic solutions even though they do not contain H+ ions Ø It does not explain acid- base reactions that take place outside aqueous solution v Bronsted- Lowry acid base theory Ø It defines acids and bases based on what happens in a reaction Ø Any reaction involving H_ (proton) that transfers from one molecule to another is an acid- base reaction regardless of whether it occurs in aqueous solution of there is an OH- present Ø All reaction that fit the Arrhenius definition also fit the Bronsted- Lowry definition v The acid is an H+ donor v The base is an H+ acceptor Ø Base structure must contain an atom with an unshared pair of electrons Ø In Bronsted- Lowry acid base reaction the acid molecule donates an H+ to the base molecule v Bronsted- Lowry acids Ø Acids are donors Any material that has H can potentially be a Bronsted- Lowry acid Because of the molecular structure often one H in the molecules is easier to transfer than others Ø When HCl dissolves in water the HCl is the acid because Hl transfers an H+ to water forming H3O ions Water acts as base accepting H+ v Bronsted Lowry bases Ø H+ acceptors Any material that has atoms with lone pairs can potentially be a Bronsted- Lowry base Because of the molecular structure often one atom of the molecule is more willing to accept H+ transfer than others Ø When NH3 dissolves in water the NH3(aq) is the base because NH3 accepts an H+ from water forming OH- Water acts as acid donating H+ v Amphoteric substances Ø Can act as wither an acid or base because they have both transferrable H and an atom with lone pair electrons Ø Water v Bronsted- Lowry acid base reactions Ø One of the advantages of this theory is that it illustrates reversible reactions Ø The original base has an extra H+ after the reaction so it will act as an acid in the reversible process Ø And the original acid has a lone pair of electron after the reaction so it will act as a bas in the reverse process

4 v Conjugate acid base pairs Ø In Bronsted- Lowry acid base reaction The original base becomes an acid in the reverse reaction The original acid becomes a base in the revers process Ø Each reactant and the product it becomes is called a conjugate pair Ø A base accepts a proton and becomes a conjugate acid Ø An acid cones a proton and becomes a conjugate base v Arrow conventions Ø Chemists commonly use two kinds of arrows in reactions to indicate the degree of completion of the reactions Ø A single arrow indicated that all the reactant molecules are converted to product molecules in the end Ø A double arrow indicates the reaction stops when only some of the reactant molecules have been converted into products v Strong or weak Ø A strong acid is a strong electrolyte Practically all the acid molecules ionize Ø A strong base is a strong electrolyte Practically all these molecules form OH- ions either through dissociation or reaction with water Ø A weak acid is a weak electrolyte Only a small percentage of the molecules ionize ß à Ø A weak base is a weak electrolyte Only a small percentage of the base molecules with OH- ions either through dissociation or reaction with water ß à v Strong acids Ø Strong acids donate practically all of their hydrogens 100% ionized in water Strong electrolyte v Weak acids Ø Weak acids donate a small fraction of their hydrogens Most of the weak acid molecules do not donate H to water Much less than 1% ionized in water v Strengths of acids and bases Ø Commonly acid or base strength is measured by determining the equilibrium constant of a substance reaction with water Ø The farther the equilibrium position lies toward the products the stronger the acid or base Ø The position of equilibrium depends on the strength of attraction between the base and from the H+ Stronger attraction means stronger base or weaker acid v General trends in activity Ø The stronger an acid is at donating H the weaker the conjugate base is at accepting H Ø Higher oxidation number = stronger oxyacid Ø Carbon stronger acid than neutral molecules neutral stronger acid than anion

5 v Autoionization of water Ø Water is amphoteric it can act either as an acid to a base Therefore there must be a few ions present Ø About two out of every 1 billion water molecules form ions through a process called autoionization Ø All aqueous solutions contain both H3O and OH- Concentration of H3O and OH- are equal in water v Ion product of water Ø The product of the H3O and OH- concentration is always the same number Ø The number is called the ion product of water and has the symbol Kw Also known as the dissociation constant of water As the [H3O]+ increased the [OH- ] must decrease so the product stays constant Inversely proportional v Acid association constant Ka Ø Acid strength is measured by the size of the equilibrium constant when is reacts with water Ø The equilibrium constant for this reaction is called the ionization constant Ka The larger the Ka the stronger the acid v Acidic and basic solutions Ø All aqueous coltuions caontain both H3O and OH- ions Ø Natural solutions have equal [H3O]+ and [OH- ] Ø Acidic solutions have a larger [H3O]+ than [OH- ] Ø Basic solutions have a larger [OH- ] than [H3O]+ v Measuring acidity ph Ø The acidity or basicity of a solution is often expressed as ph Ø ph < 7 acidic Ø ph > 7 basic Ø ph = 7 neutral v Sigs figs and logs Ø When you take the log of a number written in scientific notation the digits before the decimal point some from the exponent on 10 and the digits after the decimal point from the decimal part of the number Ø Because the part of the scientific notation number that determines the significant figures is the decimal point the significant figures are the digits after the decimal point in the log v What does the ph number imply Ø The lower the ph the more acidic the solution the higher the ph the more basic the solution 1pH unit corresponds to the factor of 10 difference in acidity Ø Normal range of ph is 0-14 ph can be negative (very acidic) of larger than 14 (very alkaline) v poh Ø Another way of expressing the acidity/basicity of the solution is poh Ø Need to know the [OH- } concentration to find poh Ø poh < 7 is basic

6 Ø poh > 7 is acidic Ø poh = 7 is neutral v Relationship between ph and poh Ø ph + poh at 25 degrees C Ø you can used poh to find the ph f a solution v pk Ø A way of expressing the strength of an acid or base is pk Ø The stronger the acid the smaller the pka Larger Ka = smaller pka Ø The stronger the base the smaller the pkb Larger Kb = smaller pkb v [H3O]+ and [OH- ] in a strong acid or strong base solution Ø There are two sources of H3O+ in an aqueous solution of a strong acid the acid and the water Ø There are two sources of OH- in an aqueous solution of a strong acid the base and the water Ø For a strong acid or base the contribution of the water to the total [H3O]+ and [OH- ] is negligible v Finding ph of a strong acid or strong base solution Ø For a monoprotic strong acid [H3O+] =[HAcid] For polyprotic acids the other ionizations can generally be ignored For a strong ionic base [OH- ] = (number OH- ) *[base] For molecular bases with multiple lone pairs available only one lone pair accepts an H- the other reactions can generally be ignored v Other notes and equations that you need to know Ø Kw = [H3O]+[OH- ] = 1 *10-14 Ø ph = - log[h3o+] Ø [H3O+] = 10 - ph Ø poh = - log[oh- ] Ø [OH- ] = 10 - poh Ø ph + poh = 14 v Finding ph of a strong acid or strong base solution Ø For a monoprotic strong acid [H3)] = [HAcid] For polyprotic acids the other ionizations can generally be ignored Not for H2SO4 Ø For a strong ionic base [OH- ]=(number OH- 0*([base]) For molecular bases with multiple lone pairs available only one lone pair accepts an H the other reactions can generally be ignored v Finding the ph of a weak acid Ø There are also two sources of H3O in an aqueous solution of a weak acid the acid and the water Ø However finding the [H3O+] is complicated by the fact that the acid only undergoes partial dissociation Ø Calculating [H3O+] requires solving an equilibrium problem for the reaction that defines the acidity of the acid

7 v Percent ionization Ø Another way to measure the strength of an acid is to determine the percentage of acid molecules that ionize when dissolved in water this is called the percent ionization The higher the percent ionization the stronger the acid Percent ionization = molarity of the ionized acid/initial molarity of the acid * 100 Or: [H3O+]/[HA] * 100 v Relationship between [H3O+] equilibrium and [HA]initial Ø Increasing the initial concentration of acid results in increased [H3O+] at equilibrium Ø Increasing the initial concentration of acid results in decreased percent ionization Ø This means that the increase in [H3O+] concentration is slower than the increase in acid concentration v Why doesn t the increase in [H3O+] keep up with the increase in [HA] Ø The reaction for the ionization of a weak acid is HAaq + H2Ol ß à A- aq + H3O+aq Ø According to Le Chateliers principle if we reduce the concentrations of all the (aq) components the equilibrium should shift to the right to increase the total number of dissolved particles We can reduce the (aq) concentrations by using a more dilute initial acid concentration Ø The result will be a larger [H3O+] v Finding the ph of mixtures of acids Ø Generally you can ignore the contribution of the weaker acid to the [H3O+]eq Ø For a mixture of a strong acid with a weak acid the complete ionization of the strong acid provides more than enough [H3O+] to shift the weak acid equilibrium to the left so fat that the weak acid s added [H3O+] is negligible Ø For mixtures of weak acids you generally only need to consider the stronger for the same reasons as long as one is significantly stronger than the other and their concentrations are similar v Strong bases Ø LiOH Ø NaOH Ø KOH Ø Sr(OH)2 Limited solubility Ø Ba(OH)2 Limited solubility Ø The stronger the base the more willing it is so accept H Use water as the standard acid Ø For ionic bases particularly all units are dissociated into OH- or accept H s Strong electrolyte Multi- OH strong bases completely dissociated v Weak bases

8 Ø In weak bases only a small fraction of molecules accept H s Weak electrolyte Most of the weak base molecules do not take H from water Much less than 1% ionization in water Ø [HO} << [weak base] Ø Finding the ph of a weak base solution is similar to finding the ph of a weak acid v Base ionization constant Kb Ø Base strength is measured by the size of the equilibrium constant when it reacts with H2O Base + water ß à OH- +Hbase Ø The equilibrium constant is called the base ionization constant kb Larger Kb Kb = [OH- ]*[H:base]/[:base] v Equilibrium Constant Ø Even though the concentrations of reactants and products are not equal at equilibrium there is a relationship between them Ø The relationship between the chemical equation and the concentrations of reactants and products is called the law of mass action Ø For the general equation aa + bb ß à cc + dd the law of mass action gives the relationship below The lowercase letters represent the coefficients of the balanced chemical equation Always products over reactants Ø K is called the equilibrium constant Unitless v What does the value of Keq imply Ø When the value of Keq >> 1 when the reaction reaches equilibrium there will be many more product molecules present than reactant molecules Ø The position of equilibrium favors products Ø When the value of Keq << 1 when the reaction reached equilibrium there will be many ore reactant molecules present than the product molecules Ø The position of equilibrium favors reactants

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