Chapter 10 Acids and Bases in Our Environment Lecture Presentation Chapter 10 Acids and Bases in Our Environment Bradley Sieve Northern Kentucky University Highland Heights, KY

10.1 Acids Donate Protons and Bases Accept Them Taste sour Cause of taste in vinegar and citric fruits Essential to the chemical industry Sulfuric acid is the number one chemical produced in the United States

10.1 Acids Donate Protons and Bases Accept Them Taste bitter Feel slippery React with skin oils to form slippery soaps Heavily used in industry, with sodium hydroxide being the most common

10.1 Acids Donate Protons and Bases Accept Them Brønsted–Lowry Defines acids and bases based on hydrogen ions (H+) Acids donate an H+ Bases accept H+

10.1 Acids Donate Protons and Bases Accept Them HCl reacting with water HCl donates a hydrogen ion and therefore is an acid Water accepts the hydrogen ion and therefore is a base

10.1 Acids Donate Protons and Bases Accept Them Ammonia reacting with water Water donates a hydrogen ion and therefore is an acid Ammonia accepts the hydrogen ion and therefore is a base

10.1 Acids Donate Protons and Bases Accept Them Some chemicals can behave as either an acid or a base Depends on what it is paired with Water is the main example of this behavior

Concept Check Identify the behavior as an acid or a base of each participant in the reaction H2PO4– + H3O+  NH3PO4 + H2O

Concept Check In the forward reaction (left to right), H2PO4– gains a hydrogen ion to become H3PO4. In accepting the hydrogen ion, H2PO4– is behaving as a base. It gets the hydrogen ion from the H3O+, which is behaving as an acid. In the reverse direction, H3PO4 loses a hydrogen ion to become H2PO4– and is thus behaving as an acid. The recipient of the hydrogen ion is the H2O, which is behaving as a base as it transforms to H3O+.

10.1 Acids Donate Protons and Bases Accept Them Lewis definition More general than the Brønsted–Lowry definition Focuses on lone electron pairs Acids

10.1 Acids Donate Protons and Bases Accept Them Bases donate a lone pair of electrons Arrows show the movement of electrons Water is donating a pair of electrons to form a bond with the H atom

10.1 Acids Donate Protons and Bases Accept Them Acids accept a lone pair of electrons HCl is accepting the lone pair from the water

Concept Check How is it possible for carbon dioxide, CO2, to behave as an acid when it has no hydrogen ions to donate?

Concept Check Carbon dioxide behaves as an acid when it accepts the lone pair on the oxygen atom of a water molecule.

10.1 Acids Donate Protons and Bases Accept Them Salts Are ionic products of acid-base reactions NaCl is a common salt but not the only type HCl + NaOH  NaCl + H2O HCl + KOH  KCl + H2O

10.1 Acids Donate Protons and Bases Accept Them Neutralization A reaction between an acid and a base Forms salt and water [insert table 10.1]

Concept Check Is a neutralization reaction best described as a physical change or a chemical change?

Concept Check New chemicals are formed during a neutralization reaction, meaning the reaction is a chemical change.

10.2 Some Acids and Bases Are Stronger Than Others Strength of an acid or a base is the measure of how much remains after being added to water A strong acid or base will have little of the original chemical left A weak acid or base will have a lot of the original chemical left

10.2 Some Acids and Bases Are Stronger Than Others Hydrochloric acid (HCl) is a strong acid All of it is converted to ions

10.2 Some Acids and Bases Are Stronger Than Others Acetic acid (C2H4O2) is a weak acid Much of the original chemical remains when added to water

10.2 Some Acids and Bases Are Stronger Than Others The stronger the acid or base, the better it conducts electricity Pure Water HCl Acetic Acid

10.2 Some Acids and Bases Are Stronger Than Others Corrosive ability is based on concentration, not the strength of an acid or a base Highly concentrated solutions tend to be more corrosive Less concentrated solutions tend to be less corrosive

10.3 Solutions Can Be Acidic, Basic, or Neutral Amphoteric A substance that can behave as an acid or a base Water is a good example

10.3 Solutions Can Be Acidic, Basic, or Neutral Pure water contains roughly 1 × 10–7 M H+ and OH– ions The concentration of H3O+ ([H3O+]) multiplied by the concentration of OH– ([OH–]) always equals a constant Kw Kw = 1 × 10–14

Concept Check In pure water, the hydroxide-ion concentration is 1.0 × 10–7 M. What is the hydronium-ion concentration?

Concept Check 1.0 × 10–7 M, because in pure water, [H3O+] = [OH–].

Concept Check What is the concentration of hydronium ions in a solution if the concentration of hydroxide ions is 1.0 × 10–3 M?

Concept Check 1.0 × 10–11 M, because [H3O+][OH–] must equal 1.0 × 10–14 = Kw.

10.3 Solutions Can Be Acidic, Basic, or Neutral Types of Solutions Neutral is when the solution has equal number of H3O+ and OH– ions Acidic is when the solution has more H3O+ ions than OH– ions Basic is when the solution has fewer H3O+ ions than OH– ions

10.3 Solutions Can Be Acidic, Basic, or Neutral

Concept Check How does adding ammonia, NH3, to water make a basic solution when there are no hydroxide ions in the formula for ammonia?

10.3 Solutions Can Be Acidic, Basic, or Neutral Ammonia indirectly increases the hydroxide-ion concentration by reacting with water: NH3 + H2O  NH4+ + OH– This reaction raises the hydroxide-ion concentration, which has the effect of lowering the hydronium-ion concentration. With the hydroxide-ion concentration now higher than the hydronium-ion concentration, the solution is basic.

10.3 Solutions Can Be Acidic, Basic, or Neutral pH Scale Used to describe acidity of a solution pH = –log [H3O+] Acidic pH < 7 Neutral pH = 7 Basic pH > 7

10.3 Solutions Can Be Acidic, Basic, or Neutral

10.3 Solutions Can Be Acidic, Basic, or Neutral Two common ways to determine pH

10.4 Buffer Solutions Resist Changes in pH Resist large changes in pH Contain two components, each neutralizing an acid or a base Often prepared by mixing a weak acid and a salt of that weak acid

10.4 Buffer Solutions Resist Changes in pH A strong acid added to the buffer is neutralized by the base component

10.4 Buffer Solutions Resist Changes in pH A strong base added to the buffer is neutralized by the acid component

Concept Check Why do most buffer solutions contain two dissolved components?

Concept Check The weak base component neutralizes any incoming acid, and the weak acid component neutralizes any incoming base.

10.4 Buffer Solutions Resist Changes in pH Blood contains several buffer systems Maintains pH between 7.35 and 7.45 Outside this range, proteins become denatured One buffer system in the blood is based on the carbonic acid/sodium bicarbonate system

10.4 Buffer Solutions Resist Changes in pH

10.5 Rainwater Is Acidic Rainwater is naturally acidic due to CO2 CO2 + H2O  H2CO3 Normal pH falls around 5.6 but ranges from 5 to 7 Acidity can lead to erosion and the formation of caves

Acid rain is rain with a pH lower than 5 10.5 Rainwater Is Acidic Acid rain is rain with a pH lower than 5 Occurs when moisture absorbs pollutants Formation of sulfuric acid in rainwater 2SO2(g) + O2(g)  2SO3(g) SO3(g) + H2O(l)  H2SO4(aq) Corrodes metal, paint, and other exposed substances

10.5 Rainwater Is Acidic (a), (b) These two photographs show the same obelisk in New York City’s Central Park before and after the effects of acid rain. (c) Many forests downwind from heavily industrialized areas, such as in the northeastern United States and Europe, have been noticeably hard-hit by acid rain.

10.5 Rainwater Is Acidic

Concept Check When sulfuric acid, H2SO4, is added to water, what makes the resulting aqueous solution corrosive?

Concept Check Because H2SO4 is a strong acid, it readily forms hydronium ions when dissolved in water. Hydronium ions are responsible for the corrosive action.

10.6 Carbon Dioxide Acidifies the Oceans Atmospheric CO2 dissolves in water to form carbonic acid (H2CO3) Is quickly neutralized in ocean waters by dissolved alkaline minerals CO2 is not released back into the environment

10.6 Carbon Dioxide Acidifies the Oceans This trapping of CO2 in the ocean removes basic carbonate compounds from the water These basic components are used for survival of aquatic life, including coral and shelled organisms Lowers the pH of the ocean water 0.1 pH over the past 100 years Has not occurred naturally for 56 million years

10.6 Carbon Dioxide Acidifies the Oceans