Lecture Presentation Chapter 16 Acids and Bases 2014

Stomach Acid and Heartburn • The cells that line your stomach produce hydrochloric acid. – To kill unwanted bacteria – To help break down food – To activate enzymes that break down food • If the stomach acid backs up into your esophagus, it irritates those tissues, resulting in heartburn. – Acid reflux © 2014 Pearson Education, Inc.

Curing Heartburn • Mild cases of heartburn can be cured by neutralizing the acid in the esophagus. – Swallowing saliva, which contains bicarbonate ion – Taking antacids that contain hydroxide ions and/or carbonate ions © 2014 Pearson Education, Inc.

Properties of Acids • Sour taste • Ability to dissolve many metals • Ability

Common Acids © 2014 Pearson Education, Inc.

Structures of Acids • Binary acids have acid hydrogens attached to a nonmetal atom.

Structure of Acids • Oxyacids have acid hydrogens attached to an oxygen atom. ü

Structure of Acids • Carboxylic acids have COOH group. ü HC 2 H 3

Properties of Bases • Taste bitter ü alkaloids = plant product that is alkaline Ø often poisonous • Feel slippery • Ability to turn red litmus paper blue • Ability to neutralize acids © 2014 Pearson Education, Inc.

Common Bases © 2014 Pearson Education, Inc.

Indicators • Indicators are chemicals that change color depending on the solution’s acidity or basicity. • Many vegetable dyes are indicators. – Anthocyanins • Litmus – From Spanish moss – Red in acid, blue in base • Phenolphthalein – Found in laxatives – Red in base, colorless in acid © 2014 Pearson Education, Inc.

Definitions of Acids and Bases • Arrhenius definition – Based on H+ and OH- • Brønsted–Lowry definition – Based on reactions in which H+ is transferred • Lewis definition – Based on reactions in which electrons are transferred © 2014 Pearson Education, Inc.

Arrhenius Theory • Acids: produce H+ ions in aqueous solution. HCl(aq) → H+(aq) +

Hydronium Ion • The H+ ions produced by the acid are so reactive they cannot exist in water. – H+ ions are protons! • Instead, they react with water molecules to produce complex ions, mainly hydronium ion, H 3 O+. H + + H 2 O H 3 O + – There also minor amounts of H+ with multiple water molecules, H(H 2 O)n+. © 2014 Pearson Education, Inc.

Arrhenius Theory • Bases: produce OH− ions in aqueous solution. Na. OH(aq) → Na+(aq)

Arrhenius Acid–Base Reactions • The H+ from the acid combines with the OH− from the base to make a molecule of H 2 O. – It is often helpful to think of H 2 O as H—OH. • The cation from the base combines with the anion from the acid to make a salt. acid + base → salt + water HCl(aq) + Na. OH(aq) → Na. Cl(aq) + H 2 O(l) © 2014 Pearson Education, Inc.

Problems with Arrhenius Theory • It does not explain why molecular substances, such as NH 3, dissolve in water to form basic solutions, even though they do not contain OH– ions. • It does not explain how some ionic compounds, such as Na 2 CO 3 or Na 2 O, dissolve in water to form basic solutions, even though they do not contain OH– ions. • It does not explain why molecular substances, such as CO 2, dissolve in water to form acidic solutions, even though they do not contain H+ ions. • It does not explain acid–base reactions that take place outside aqueous solution. © 2014 Pearson Education, Inc.

Brønsted–Lowry Acid–Base Theory • It defines acids and bases based on what happens in a reaction. • 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 or if there is OH− present. • All reactions that fit the Arrhenius definition also fit the Brønsted–Lowry definition. © 2014 Pearson Education, Inc.

Brønsted–Lowry Theory • The acid is an H+ donor. • The base is an H+ acceptor. – Base structure must contain an atom with an unshared pair of electrons. • In a Brønsted–Lowry acid–base reaction, the acid molecule donates an H+ to the base molecule. H–A + : B : A– + H–B+ © 2014 Pearson Education, Inc.

Brønsted–Lowry Acids • Brønsted–Lowry acids are H+ donors. – Any material that has H can potentially be a Brønsted–Lowry acid. – Because of the molecular structure, often one H in the molecule is easier to transfer than others. • When HCl dissolves in water, the HCl is the acid because HCl transfers an H+ to H 2 O, forming H 3 O+ ions. – Water acts as base, accepting H+. HCl(aq) + H 2 O(l) → Cl–(aq) + H 3 O+(aq) acid © 2014 Pearson Education, Inc. base

Brønsted–Lowry Bases • Brønsted–Lowry bases are H+ acceptors. – Any material that has atoms with lone pairs can potentially be a Brønsted–Lowry base. – Because of the molecular structure, often one atom in the molecule is more willing to accept H+ transfer than others. • When NH 3 dissolves in water, the NH 3(aq) is the base because NH 3 accepts an H+ from H 2 O, forming OH–(aq). – Water acts as acid, donating H+. NH 3(aq) + H 2 O(l) NH 4+(aq) + OH–(aq) base acid © 2014 Pearson Education, Inc.

Amphoteric Substances • Amphoteric substances can act as either an acid or a base because they have both a transferable H and an atom with lone pair electrons. • Water acts as base, accepting H+ from HCl(aq) + H 2 O(l) → Cl–(aq) + H 3 O+(aq) • Water acts as acid, donating H+ to NH 3(aq) + H 2 O(l) NH 4+(aq) + OH–(aq) © 2014 Pearson Education, Inc.

Brønsted–Lowry Acid–Base Reactions • One of the advantages of Brønsted–Lowry theory is that it illustrates reversible reactions to be as follows: H–A + : B : A– + H–B+ • The original base has an extra H+ after the reaction, so it will act as an acid in the reverse process. • And the original acid has a lone pair of electrons after the reaction, so it will act as a base in the reverse process: : A– + H–B+ H–A + : B © 2014 Pearson Education, Inc.

Conjugate Acid–Base Pairs • In a Brønsted-–Lowry acid–base reaction, – the original base becomes an acid in the reverse reaction. – the original acid becomes a base in the reverse process. • Each reactant and the product it becomes is called a conjugate pair. © 2014 Pearson Education, Inc.

Conjugate Pairs A base accepts a proton and becomes a conjugate acid. An acid

Strong or Weak • A strong acid is a strong electrolyte. – Practically all the acid molecules ionize. • A strong base is a strong electrolyte. – Practically all the base molecules form OH– ions, either through dissociation or reaction with water. • A weak acid is a weak electrolyte. – Only a small percentage of the molecules ionize, • A weak base is a weak electrolyte – only a small percentage of the base molecules form OH– ions, either through dissociation or reaction with water, © 2014 Pearson Education, Inc.

Strong Acids • Strong acids donate practically all their H’s. ü 100% ionized in

$Weak Acids • Weak acids donate a small fraction of their H’s. ü Most$

Weak Acids • Weak acids donate a small fraction of their H’s. ü Most of the weak acid molecules do not donate H to water. ü Much less than 1% ionized in water • [H 3 O+] << [weak acid] © 2014 Pearson Education, Inc.

Examples of Acids © 2014 Pearson Education, Inc.

Strengths of Acids and Bases • Acid strength depends on the degree of attraction of the acid anion and the H+. • Commonly, acid or base strength is measured by determining the equilibrium constant of a substance’s reaction with water. Acid: HA + H 2 O A − + H 3 O + Base: + H 2 O HBase+ + OH− • The smaller ionization constant, K, the weaker the acid or base. • The position of equilibrium depends on the strength of attraction between the base form and the H+. – Stronger attraction means stronger base or weaker acid. © 2014 Pearson Education, Inc.

Ionic Attraction and Acid Strength © 2014 Pearson Education, Inc.

General Trends in Acidity • The stronger an acid is at donating H, the weaker the conjugate base is at accepting H. • Higher oxidation number = stronger oxyacid – H 2 SO 4 > H 2 SO 3; HNO 3 > HNO 2 • Cation stronger acid than neutral molecule; neutral stronger acid than anion – H 3 O+ > H 2 O > OH−; NH 4+ > NH 3 > NH 2− – Trend in base strength opposite © 2014 Pearson Education, Inc.

Autoionization of Water • Water is amphoteric; it can act as either an acid or a base. – Therefore, there must be a few ions present. • When water acts with itself as an acid and a base to form ions, this process is called autoionization. H 2 O (l) + H 2 O (l) H 3 O+ + OH– Can be written as H 2 O (l) H+ + OH–. • All aqueous solutions contain both H 3 O+ and OH–. – The concentration of H 3 O+ and OH– are equal in water. – [H 3 O+] = [OH–] = 10− 7 M at 25 °C © 2017 Pearson Education, Inc.

Ion Product of Water • The product of the H 3 O+ and OH– concentrations is always the same number. • The number is called the ion product of water and has the symbol Kw. – Also know as the dissociation constant of water • [H 3 O+] × [OH–] = Kw = 1. 00 × 10− 14 at 25 °C – If you measure one of the concentrations, you can calculate the other. • As [H 3 O+] increases the [OH–] must decrease so the product stays constant. – Inversely proportional © 2014 Pearson Education, Inc.

Acid Ionization Constant, Ka • Acid strength is measured by the size of the equilibrium constant when it reacts with H 2 O. • The equilibrium constant for this reaction is called the acid ionization constant, Ka. – larger Ka = stronger acid © 2014 Pearson Education, Inc.

Acidic and Basic Solutions • All aqueous solutions contain both H 3 O+ and OH– ions. • Neutral solutions have equal [H 3 O+] and [OH–]. – [H 3 O+] = [OH–] = 1. 00 × 10− 7 • Acidic solutions have a larger [H 3 O+] than [OH–]. – [H 3 O+] > 1. 00 × 10− 7; [OH–] < 1. 00 × 10− 7 • Basic solutions have a larger [OH–] than [H 3 O+]. – [H 3 O+] < 1. 00 × 10− 7; [OH–] > 1. 00 × 10− 7 © 2014 Pearson Education, Inc.

Measuring Acidity: p. H • The acidity or basicity of a solution • is often expressed by p. H = −log[H 3 O+] – p. Hwater = −log[10− 7] = 7 – Need to know the [H 3 O+] concentration to find p. H • p. H < 7 is acidic; p. H > 7 is basic. • • p. H = 7 is neutral. [H 3 O+] = 10−p. H From Kw, [H 3 O+] [OH–] = 10− 14 © 2017 Pearson Education, Inc.

Sig. , Figs. , and Logs • When you take the log of a number written in scientific notation, the digits before the decimal point come from the exponent on 10, and the digits after the decimal point come from the decimal part of the number. log(2. 0 x 106) = log(106) + log(2. 0) = 6 + 0. 30303… = 6. 30303. . . • Because the part of the scientific notation number that determines the significant figures is the decimal part, the sig. figs. are the digits after the decimal point in the log(2. 0 × 106) = 6. 30 © 2014 Pearson Education, Inc.

What Does the p. H Number Imply? • The lower the p. H, the more acidic the solution; the higher the p. H, the more basic the solution. ü 1 p. H unit corresponds to a factor of 10 difference in acidity. • Normal range of p. H is 0 to 14. ü p. H 0 is [H 3 O+] = 1 M; p. H 14 is [OH–] = 1 M. ü p. H can be negative (very acidic) or larger than 14 (very alkaline). © 2014 Pearson Education, Inc.

p. OH • Another way of expressing the acidity/basicity of • a solution is p. OH = −log[OH ], [OH ] = 10−p. OH ü p. OHwater = −log[10− 7] = 7 ü Need to know the [OH ] concentration to find p. OH • p. OH < 7 is basic; p. OH > 7 is acidic; p. OH = 7 is • neutral. p. H + p. OH = 14. 0 © 2014 Pearson Education, Inc.

Relationship between p. H and p. OH • p. H + p. OH =

p. K • A way of expressing the strength of an acid or base is p. K. • p. Ka = −log(Ka), Ka = 10−p. Ka • p. Kb = −log(Kb), Kb = 10−p. Kb • The stronger the acid, the smaller the p. Ka. – Larger Ka = smaller p. Ka – Because it is the –log • The stronger the base, the smaller the p. Kb. – Larger Kb = smaller p. Kb © 2014 Pearson Education, Inc.

[H 3 O+] and [OH−] in a Strong Acid or Strong Base Solution • There are two sources of H 3 O+ in an aqueous solution of a strong acid—the acid and the water. • There are two sources of OH− in an aqueous solution of a strong acid—the base and the water. • For a strong acid or base, the contribution of the water to the total [H 3 O+] or [OH−] is negligible. – The [H 3 O+]acid shifts the Kw equilibrium so far that [H 3 O+]water is too small to be significant. • Except in very dilute solutions, generally < 1 × 10− 4 M – [H 3 O+]acid = [HA] © 2014 Pearson Education, Inc.

Finding p. H of a Strong Acid or Strong Base Solution • For a monoprotic strong acid [H 3 O+] = [HAcid]. – For polyprotic acids, the other ionizations can generally be ignored. • For H 2 SO 4, the second ionization cannot be ignored. – 0. 10 M HCl has [H 3 O+] = 0. 10 M and p. H = 1. 00 • 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. – 0. 10 M Ca(OH)2 has [OH−] = 0. 20 M and p. H = 13. 30. © 2014 Pearson Education, Inc.

Finding the p. H of a Weak Acid • There also two sources of H 3 O+ in an aqueous solution of a weak acid—the acid and the water. • However, finding the [H 3 O+] is complicated by the fact that the acid undergoes only partial ionization. • Calculating the [H 3 O+] requires solving an equilibrium problem for the reaction that defines the acidity of the acid. HA + H 2 O © 2017 Pearson Education, Inc. A + H 3 O +

Finding the p. H of a Weak Acid 1. Write the equation for the acid, and use it as a guide to prepare the ICE table. 2. Complete the ICE table based on the given information. Keep in mind the reaction stoichiometry for the change line. 3. Substitute the expressions for equilibrium concentrations into the K expression. 4. Determine [H 3 O+] from the calculated value of x, and calculate the p. H as necessary. 5. Check the answer. © 2017 Pearson Education, Inc.

Percent Ionization • 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. – Concentration of ionized acid = [H 3 O+] © 2017 Pearson Education, Inc.

Relationship between [H 3 O+]equilibrium and [HA]initial • Increasing the initial concentration of acid results in increased [H 3 O+] at equilibrium. • Increasing the initial concentration of acid results in decreased percent ionization. • This means that the increase in [H 3 O+] concentration is slower than the increase in acid concentration. © 2017 Pearson Education, Inc.

Why Doesn’t the Increase in H 3 O+ Keep Up with the Increase in HA? • The reaction for ionization of a weak acid is as follows: HA(aq) + H 2 O(l) A−(aq) + H 3 O+(aq) • According to Le Châtelier’s 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. • The result will be a larger [H 3 O+] in the dilute solution compared to the initial acid concentration. • This will result in a larger percent ionization. © 2014 Pearson Education, Inc.

Finding the p. H of Mixtures of Acids • For a mixture of a strong acid with a weak acid, the complete ionization of the strong acid provides more than enough [H 3 O+] to shift the weak acid equilibrium to the left so far that the weak acid’s added [H 3 O+] is negligible. • For mixtures of weak acids, you generally need to consider only the stronger for the same reasons, as long as one is significantly stronger than the other and their concentrations are similar. © 2014 Pearson Education, Inc.

Strong Bases • The stronger the base, the more willing it is to accept H. ü Use water as the standard acid. • For ionic bases, practically all units are dissociated into OH– or accept H’s. ü Strong electrolyte ü Multi-OH strong bases completely dissociated © 2014 Pearson Education, Inc.

$Weak Bases • In weak bases, only a small fraction of molecules accept H’s.$

Weak Bases • In weak bases, only a small fraction of molecules accept H’s. ü Weak electrolyte ü Most of the weak base molecules do not take H from water. ü Much less than 1% ionization in water • [HO–] << [weak base] • Finding the p. H of a weak base solution is similar to finding the p. H of a weak acid. © 2014 Pearson Education, Inc.

Base Ionization Constant, Kb • Base strength is measured by the size of the equilibrium constant when it reacts with H 2 O. B + H 2 O OH− + BH+ • The equilibrium constant is called the base ionization constant, Kb. – Larger Kb = stronger base © 2017 Pearson Education, Inc.

Common Weak Bases © 2017 Pearson Education, Inc.

Structure of Amines © 2014 Pearson Education, Inc.

Acid–Base Properties of Ions and Salts • Salts are ionic compounds. • Soluble salts form ions in solution. • Cation of a strong base and anion of a strong acid do not ionize water; hence, it doesn’t affect solution p. H. • Cation of a weak base and anion of a weak acid will ionize water and affect the p. H of solution. – Na. HCO 3 solutions are basic. • Na+ is the cation of the strong base Na. OH. • HCO 3− is the conjugate base of the weak acid H 2 CO 3. • Anions tend to form basic or neutral solutions. • Cations tend to form acidic or neutral solutions. © 2014 Pearson Education, Inc.

Anions as Weak Bases • Every anion can be thought of as the conjugate base of an acid. • Therefore, every anion can potentially be a base. – A−(aq) + H 2 O(l) HA(aq) + OH−(aq) • The stronger the acid, the weaker the conjugate base. • An anion that is the conjugate base of a strong acid is p. H neutral. Cl−(aq) + H 2 O(l) HCl(aq) + OH−(aq) • An anion that is the conjugate base of a weak acid is basic. F−(aq) + H 2 O(l) © 2017 Pearson Education, Inc. HF(aq) + OH−(aq)

Strength of Conjugate Acid–Base Pairs © 2017 Pearson Education, Inc.

Relationship between Ka of an Acid and Kb of Its Conjugate Base • Many reference books give tables only of Ka values because Kb values can be found from them. When you add equations, you multiply the Ks. © 2017 Pearson Education, Inc.

Cations as Weak Acids • The stronger the base, the weaker the conjugate acid. • A cation that is the counterion of a strong base is p. H neutral. • A cation that is the conjugate acid of a weak base is weakly acidic. NH 4+(aq) + H 2 O(l) NH 3(aq) + H 3 O+(aq) • A cation that is a small, highly charged metal ion forms weakly acidic solution. © 2017 Pearson Education, Inc.

Metal Cations as Weak Acids • Cations of small, highly charged metals are weakly acidic. – Alkali metal cations and alkali earth metal cations are p. H neutral. – Cations are hydrated. Al(H 2 O)63+(aq) + H 2 O(l) Al(H 2 O)5(OH)2+ (aq) + H 3 O+(aq) © 2014 Pearson Education, Inc.

Classifying Salt Solutions as Acidic, Basic, or Neutral • If the salt cation is the counterion of a strong base and the anion is the conjugate base of a strong acid, it will form a neutral solution. – Na. Cl Ca(NO 3)2 KBr • If the salt cation is the counterion of a strong base and the anion is the conjugate base of a weak acid, it will form a basic solution. – Na. F © 2014 Pearson Education, Inc. Ca(C 2 H 3 O 2)2 KNO 2

Classifying Salt Solutions as Acidic, Basic, or Neutral • If the salt cation is the conjugate acid of a weak base and the anion is the conjugate base of a strong acid, it will form an acidic solution. – NH 4 Cl • If the salt cation is a highly charged metal ion and the anion is the conjugate base of a strong acid, it will form an acidic solution. – Al(NO 3)3 © 2014 Pearson Education, Inc.

Classifying Salt Solutions as Acidic, Basic, or Neutral • If the salt cation is the conjugate acid of a weak base and the anion is the conjugate base of a weak acid, the p. H of the solution depends on the relative strengths of the acid and base. – NH 4 F because HF is a stronger acid than NH 4+, Ka of NH 4+ is larger than Kb of the F−; therefore, the solution will be acidic. © 2014 Pearson Education, Inc.

Ionization in Polyprotic Acids • Because polyprotic acids ionize in steps, each H has a separate Ka. • Ka 1 > Ka 2 > Ka 3 • Generally, the difference in Ka values is great enough so that the second ionization does not happen to a large enough extent to affect the p. H. – Most p. H problems just do first ionization. – Except H 2 SO 4 uses [H 2 SO 4] as the [H 3 O+] for the second ionization. • [A 2−] = Ka 2 as long as the second ionization is negligible. © 2014 Pearson Education, Inc.

Ionization in H 2 SO 4 • The ionization constants for H 2 SO 4 are as follows: H 2 SO 4 + H 2 O HSO 4 + H 3 O+ strong HSO 4 + H 2 O SO 42 + H 3 O+ Ka 2 = 1. 2 × 10− 2 • For most sulfuric acid solutions, the second ionization is significant and must be accounted for. • Because the first ionization is complete, use the given [H 2 SO 4] = [HSO 4−]initial = [H 3 O+]initial. © 2017 Pearson Education, Inc.

Dissociation of a Polyprotic Acid © 2017 Pearson Education, Inc.

Strengths of Binary Acids • The more d+ H─X d− polarized the bond, the more acidic the bond. • The stronger the H─X bond, the weaker the acid. • Binary acid strength increases to the right across a period. – Acidity: H─C < H─N < H─O < H─F • Binary acid strength increases down the column. – Acidity: H─F < H─Cl < H─Br < H─I © 2014 Pearson Education, Inc.

Relationship between Bond Strength and Acidity © 2014 Pearson Education, Inc.

Strengths of Oxyacids, H–O–Y • The more electronegative the Y atom, the stronger the oxyacid. – HCl. O > HIO – Acidity of oxyacids decreases down a group. • Same trend as binary acids – Helps weaken the H–O bond. • The larger the oxidation number of the central atom, the stronger the oxyacid. – H 2 CO 3 > H 3 BO 3 – Acidity of oxyacids increases to the right across a period. • Opposite trend of binary acids • The more oxygens attached to Y, the stronger the oxyacid. – Further weakens and polarizes the H–O bond – HCl. O 3 > HCl. O 2 © 2014 Pearson Education, Inc.

Relationship between Electronegativity and Acidity © 2014 Pearson Education, Inc.

Relationship between Number of Oxygens on the Central Atom and Acidity © 2014 Pearson

Lewis Acid–Base Theory • Lewis acid–base theory focuses on transferring an electron pair. – Lone pair bond – Bond lone pair • Does NOT require H atoms • The electron donor is called the Lewis base. – Electron rich; therefore nucleophile • The electron acceptor is called the Lewis acid. – Electron deficient; therefore electrophile © 2014 Pearson Education, Inc.

Lewis Bases • The Lewis base has electrons it is willing to give away to or share with another atom. • The Lewis base must have a lone pair of electrons on it that it can donate. • Anions are better Lewis bases than neutral atoms or molecules. – N: < N: − • Generally, the more electronegative an atom, the less willing it is to be a Lewis base. – O: < S: © 2014 Pearson Education, Inc.

Lewis Acids • They are electron deficient, either from being attached to electronegative atom(s) or not having an octet. • They must have an empty orbital willing to accept the electron pair. • H+ has empty 1 s orbital. • B in BF 3 has empty 2 p orbital and an incomplete octet. • Many small, highly charged metal cations have empty orbitals they can use to accept electrons. • Atoms that are attached to highly electronegative atoms and have multiple bonds can be Lewis acids. © 2014 Pearson Education, Inc.

Lewis Acid–Base Reactions • The base donates a pair of electrons to the acid. • It generally results in a covalent bond forming H 3 N: + BF 3 H 3 N─BF 3 • The product that forms is called an adduct. • Arrhenius and Brønsted–Lowry acid–base reactions are also Lewis. © 2014 Pearson Education, Inc.

Examples of Lewis Acid–Base Reactions © 2014 Pearson Education, Inc.

Examples of Lewis Acid–Base Reactions Ag+(aq) + 2 : NH 3(aq) Ag(NH 3)2+(aq) Lewis

U. S. Fuel Consumption • Over 85% of the energy use in the United States comes from the combustion of fossil fuels. – Oil, natural gas, coal • Combustion of fossil fuels produces CO 2. CH 4(g) + 2 O 2(g) → CO 2(g) + 2 H 2 O(g) • Natural fossil fuels also contain small amounts of S that burn to produce SO 2(g). S(s) + O 2(g) → SO 2(g) • The high temperatures of combustion allow N 2(g) in the air to combine with O 2(g) to form oxides of nitrogen. N 2(g) + 2 O 2(g) → 2 NO 2(g) © 2014 Pearson Education, Inc.

What Is Acid Rain? • Natural rain water has a p. H of 5. 6. – Naturally slightly acidic due mainly to CO 2 • Rain water with a p. H lower than 5. 6 is called acid rain. • Acid rain is linked to damage in ecosystems and structures. © 2014 Pearson Education, Inc.

What Causes Acid Rain? • Many natural and pollutant gases dissolved in the air are nonmetal oxides. – CO 2, SO 2, NO 2 • Nonmetal oxides are acidic. CO 2(g) + H 2 O(l) H 2 CO 3(aq) 2 SO 2(g) + 2 H 2 O(l) 2 H 2 SO 4(aq) 4 NO 2(g) + 2 H 2 O(l) 4 HNO 3(aq) • Processes that produce nonmetal oxide gases as waste increase the acidity of the rain. – Natural – volcanoes and some bacterial action – Man made – combustion of fuel © 2014 Pearson Education, Inc.

p. H of Rain in Different Regions Figure 15. 16 pg 740 © 2014

Weather Patterns • The prevailing winds in the United States travel west to east. • Weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced. • Much of the Northeast United States has rain of very low p. H, even though it has very low sulfur emissions, due in part to the general weather patterns. © 2014 Pearson Education, Inc.

Damage from Acid Rain • Acids react with metals and materials that contain carbonates. • Acid rain damages bridges, cars, and other metallic structures. • Acid rain damages buildings and other structures made of limestone or cement. • Acidifying lakes affects aquatic life. • Soil acidity causes more dissolving of minerals and leaching more minerals from soil, making it difficult for trees. © 2014 Pearson Education, Inc.

Damage from Acid Rain © 2014 Pearson Education, Inc.

Acid Rain Legislation • 1990 Clean Air Act attacks acid rain. – Forces utilities to reduce SO 2 • The result is acid rain in the Northeast stabilized and beginning to be reduced. © 2014 Pearson Education, Inc.