Chapter 6: Neutralizing the Threat of Acid Rain

Is normal rain acidic?

Is acid rain worsein some parts of thecountry?

Is there a way to “neutralize” acidrain?

What does the word acid mean to you?

6.1

There are several ways to define an acid.

HCl(g) H+(aq) + Cl–(aq)  H2O

Consider hydrogen chloride gas, dissolved in water:

The most common definition for a acid is:

Any substance that causes hydrogen ions to be formed in an aqueous solution.

(a proton)

6.1

A H+ ion (a proton) is very reactive and does not actually exist in water. Instead, when the ion is formed it reacts with a water molecule to form a hydronium ion, H

3O+.

H+ (aq) + H2O → H

3O+ (aq)

The chloride ion (Cl-) remains unchanged, but is surrounded by water molecules to keep it dissolved.

The hydronium ion is often abbreviated:H

3O+ (aq) = H+ (aq)

and in many cases, aqueous is assumed,so the (aq) is often not added. O

H

H H+

hydronium ion

or:

HCl(g) + H2O(l) H3O+(aq) + Cl–(aq)

HCl(g) H+(aq) + Cl–(aq)  H2O

(a proton)

Therefore, the same reaction between gaseous HCl and water could be written in two different ways:

The most common definition for a base is:

Any substance that causes hydroxide ions (OH-) to be produced in aqueous solution.

NaOH (s) + H2O → Na+ (aq) + OH- (aq)

Notice that when written in this manner, the equation cannot be balanced. In this case water is the solvent, not a reactant, and more properly written above the arrow to indicate reaction conditions:

NaOH(s) Na+(aq) + OH–(aq)Sodium hydroxide

H2O

Sodium ion Hydroxide ion

6.2

Ca(OH)2(s) Ca2+(aq) + 2OH–(aq)H2O

What about ammonia (NH3)? It is a base, but has no OH group.

NH3(aq) + H2O(l) NH4OH(aq) ammonium hydroxide

NH4OH(aq) NH4+(aq) + OH–(aq) H2O

Calcium hydroxide produces 2 equivalents of OH– :

Acid/Base Strength

The chemistry definition of strength refers only to how easily the acid or base forms ions in water:

Strong acids/bases completely form ions

Weak acids/bases only partially form ions, so stay as the initial compound

Strength has nothing to do with concentration or reactivity

Strong acids: fully ionize to form H+ ions:

HCl (aq) → H+1 (aq) + Cl-1 (aq) 99.9% of HCl molecules split apart

Weak acids: only a small percentage of the molecules form hydronium ions:

HF (aq) ↔ H+1 (aq) + F-1 (aq)

less than 1% of HF molecules split apart

Soluble ionic compounds completely split apart into individual ions in water.

Therefore, soluble metal hydroxides fully split apart and are strong bases:

NaOH (aq) → Na+1 (aq) + OH-1 (aq)

For 'insoluble' metal hydroxides, only a very small amount of ions are formed:

FeOH (s) ↔ Fe+1 (aq) + OH-1 (aq)

most of the compound remains as a solid

Strong/weak – refers to % ionization

Concentrated/dilute – refers to the amount of acid/base per liter of solution

Reactivity – depends on the acid or base, and with what it is reacting

HCl

Strong acid

Concentrated in your stomach

Does not react readily with stomach lining

Does react to digest food particles

HF

Weak acid

Extremely reactive with most substances

Even dilute solutions eat through glass

HC2H

3O

2 – acetic acid

Weak acid

Dilute in vinegar

used to make pickles, salad dressing...

Concentrated solutions would be hazardous to ingest.

When acids and bases react with each other, we call this a neutralization reaction.

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

acid + base salt + water

6.3

In neutralization reactions, hydrogen ions from an acid combine with the hydroxide ions from a base to form molecules of water (hydrogen hydroxide).

The other product is a salt (an ionic compound).

There are three ways chemical equations are typically written:

Chemical equation (standard form): shows reactants and products properly balanced.

Ionic equation: shows ions that are present in solution.

Net ionic equation: shows only ions and compounds that participate in the reaction.

Spectator ions – appear on both sides of the ionic equation, and not shown in net ionic equations.

6.3

2 HBr(aq) + Ba(OH)2(aq) BaBr2(aq) + 2 H2O(l)

Consider the reaction of hydrobromic acid with barium hydroxide.

This reaction may be represented with a molecular, ionic, or net ionic equation:

Molecular (standard chemical equation):

Ionic: 2 H+(aq) + 2 Br

–(aq) + Ba2+(aq) + 2 OH–(aq) Ba2+(aq) + 2 Br

–(aq) + 2 H2O(l)

Net Ionic: 2 H+(aq) + 2 OH–(aq) 2 H2O(l)

or by dividing both sides of the equation by 2 to simplify it: H+(aq) + OH–(aq) H2O(l)

6.3

2 HBr(aq) + Ba(OH)2(aq) BaBr2(aq) + 2 H2O(l)

Ionic:2 H+(aq) + 2 Br–(aq) + Ba2+(aq) + 2 OH–(aq) Ba2+(aq) + 2 Br–(aq) +2H2O(l)

Remove the species that appear unchanged on both sides of the reaction – these are called spectator ions.

How did we go from Ionic to Net Ionic?

6.3

2 HBr(aq) + Ba(OH)2(aq) BaBr2(aq) + 2 H2O(l)

Ionic:2 H+(aq) + 2 Br–(aq) + Ba2+(aq) + 2 OH–(aq) Ba2+(aq) + 2 Br–(aq) + 2 H2O(l)

spectator ions

How did we go from Ionic to Net Ionic?

Net Ionic: H+(aq) + OH–(aq) H2O(l)

6.3

The hydroxide and hydronium concentrations in water are related by the expression:

Kw = [H+][OH–] = 1 x 10–14 (at 25 oC)

where Kw is the ion-product constant for water, and remember that [x] means molarity (moles/liters) of x.

Since Kw is a constant, as the hydronium concentration is increased,

the hydroxide concentration decreases, and vice versa.

Knowing the hydroxide ion concentration, we can calculate the [H+] and vice versa.

6.3

Kw = [H+][OH–] = 1 x 10–14 (at 25 oC)

This expression shows that both ions, H+ and OH-, are always present in water.

For acidic solutions, the hydronium concentration is greater.

For basic solutions, the hydroxide concentration is greater.

For neutral solutions, the two concentrations are equal (10-7 M), not zero.

6.4

pH is used to compare the concentration of the H+ ions present in aqueous solutions (easier to compare than using exponential values).

The mathematical expression for pH is a log-based scale and is represented as:

pH = –log[H+]

So for a solution with a

[H+] = 1.0 x 10–3 M, the pH = –log(1.0 x 10–3), or –(–3.0) = 3.0

Since pH is a log scale based on 10, a pH change of 1 unit represents a power of 10 change in [H+].

That is, a solution with a pH of 2 has a [H+] ten times that of a solution with a pH of 3.

6.6

Measuring pH with a pH meter

Or with litmus paper Base applied to red litmus

Acid applied to blue litmus

6.4

To measure the pH in basic solutions, we make use of the expression below to calculate [H+] from [OH–].

Kw = [H+][OH–] = 1 x 10–14 (at 25 oC)

The three possible aqueous solution situations are:

[H+] = [OH–] a neutral solution (pH = 7)[H+] > [OH–] an acidic solution (pH < 7)[H+] < [OH–] a basic solution (pH > 7)

6.4

Common substances and their pH values

Note that “normal” rain is slightly acidic.

Rain water is naturally slightly acidic due to dissolved carbon dioxide:

CO2 (g) ↔ CO

2 (aq)

CO2 (aq) + H

2O (l) ↔ H

2CO

3 (aq)

H2CO

3 (aq) ↔ H+1 (aq) + HCO

3-1 (aq)

In fact any pure water left exposed to air will become slightly acidic due to small amounts of dissolved carbon dioxide

6.5

Rainwater is naturally acidic, but ocean water is basic.

Three chemical species responsible for maintaining ocean pH:

C

O

-O O-

carbonate ion CO3

2- (aq)

C

O

HO O-

bicarbonate ion HCO3

- (aq)

C

O

HO OH

carbonic acid H2CO3 (aq)

Base Base Acid

Many ocean creatures, coral, mollusks, urchins, have shells made out of CaCO

3

They get the CO3

-2 to make their shells from

dissolved ions in the ocean

The shells from these dead animals are buried in the ocean floor, but help keep the CO

3-2 (aq)

concentration constant and neutralize acid from dissolved CO

2

Over the past 200 years, the amount of carbon dioxide in the atmosphere has increased, so more carbon dioxide is dissolving in the ocean and forming carbonic acid.

Ocean acidification – the lowering of ocean pH due to increased atmospheric carbon dioxide

The H + produced from the dissociation of carbonic acid reacts with carbonate ion in seawater to form the bicarbonate ion:

H+(aq) + CO32–(aq) → HCO3

–(aq)

This reduces the concentration of carbonate ion in seawater, so the calcium carbonate in the shells of sea creatures begins to dissolve to maintain the concentration of carbonate ions in seawater:

CaCO3(s) → Ca2+(aq) + CO32–(aq)

The buffering from the buried shells is too slow of a process to compensate for the carbon dioxide changes over the last 200 years

6.5

Chemistry of Carbon Dioxide in the Ocean

6.6

Why is rain naturally acidic?

Carbon dioxide in the atmosphere dissolves to a slight extent in water and reacts with it to produce a slightly acidic solution of carbonic acid:

CO2(g) + H2O(l) H2CO3(aq) carbonic acid

H2CO3(aq) H+(aq) + HCO3–(aq)

The carbonic acid dissociates slightly leading to rain with a pH around 5.3.

6.6

6.6

But acid rain can have pH levels lower than 4.3 (10 times the normal acidity) – where is the extra acidity coming from?

The most acidic rain falls in the eastern third of the United States, with the region of lowest pH being roughly the states along the Ohio River valley.

The extra acidity must be originating somewhere in this heavily industrialized part of the country.

6.6

But acid rain can have pH levels lower than 4.3 (10 times the normal acidity)

Analysis of rain for specific compounds confirms that the chief culprits are the oxides of sulfur and nitrogen:

sulfur dioxide (SO2), sulfur trioxide (SO3), nitrogen monoxide (NO), and nitrogen dioxide (NO2).

These compounds are collectively designated SOx and NOx and are often referred to as “Sox and Nox.”

6.6

Oxides of sulfur and nitrogen are acid anhydrides, literally “acids without water.”

SO2(g) + H2O(l) H2SO3(aq) sulfurous acid

SO3(g) + H2O(l) H2SO4(aq) sulfuric acid

And then: H2SO4(aq) 2 H+(aq) + SO42–(aq)

SOx react with water to form acids:

6.6

4 NO2(g) + 2 H2O(l) + O2(g) 4 HNO3(aq) nitric acid

Like sulfuric acid, nitric acid also dissociates to release the H+ ion:

HNO3(aq) H+(aq) + NO3–(aq)

What about the NOx?

6.7

Sulfur dioxide emissions are highest in regions with many coal-fired electric power plants, steel mills, and other heavy industries that rely on coal.

Allegheny County, in western Pennsylvania, is just such an area, and in 1990 it led the United States in atmospheric SO2 concentration.

6.7

How does the sulfur get into the atmosphere?

The burning of coal.Coal contains 1–3% sulfur and coal burning power plants usually burn about 1 million metric tons of coal a year!

S(s) + O2(g) SO2(g)

2 SO2(g) + O2(g) 2 SO3(g)

Once in the air, the SO2 can react with oxygen molecules to form sulfur trioxide, which acts in the formation of aerosols.

Burning of sulfur with oxygen produces sulfur dioxide gas, which is poisonous.

The highest NOx emissions are generally found in states with large urban areas, high population density, and heavy automobile traffic.

Therefore, it is not surprising that the highest levels of atmospheric NO2 are measured over Los Angeles County, the car capital of the country.

6.8

Nitrogen dioxide gas in the atmosphere reacts with the hydroxyl radical to formnitric acid.

NO2(g) + · OH(g) → HNO3(l)

6.10

Emissions of NOx and SO2 in the U.S.

These SOx and NO

x compounds react with

airborne moisture to acidify rain.

But if no water vapor is present, they may be trapped in particulates on dry land.

When rain does come, they then form acids.

The term 'acid rain' is used to identify both processes of acidifying rain water.

Some effects of acid rain are:

Erosion of buildings

Limestone and marble are mainly CaCO

3(a base), and react with acids just

as coral shells do.

Rusting of metal in buildings and cars

6.11

Effects of acid rain: damage to marble

1944 At present

These statues are made of marble, a form of limestone composed mainly of calcium carbonate, CaCO3.

Limestone and marble slowly dissolve in the presence of H+ ions:

CaCO3(s) + 2 H+(aq) Ca2+(aq) + CO2(g) + H2O(l)

6.11

Effects of acid rain: rusting metal

Iron metal dissolves when exposed to hydrogen ions:

4 Fe(s) + 2 O2(g) + 8 H+(aq) → 4 Fe2+(aq) + 4 H2O(l)

Now the aqueous Fe2+ ions react with oxygen to form rust (Fe2O3):

4 Fe2+(aq) + O2(g) + 4 H2O(l) → 2 Fe2O3(s) + 8 H+(aq)

Billions of dollars are spent annually to protect bridges, cars, buildings, and ships from the reactions above.

Human health

Asthma

Bronchitis

Even death in some cases

Visibility

Causes haze or smog

6.12

“They stop at Overlooks, idling their engines as they read

signs describing Shenandoah’s ruggedly scenic terrain:

Hogback, Little Devil Stairs, Old Rag, Hawksbill.

Jammed along the narrow road, the cars and motor homes

add to the miasmic summer haze that cloaks the hills.”

Ned Burks,

Environmental writer describing Skyline Drive in Virginia

Harms surface fresh water

More acidic water can kill some or all plants and animals in lakes and rivers

Some regions have natural limestone bedrock which helps neutralize the acid

Regions with more granite bedrock do not have this buffering ability

6.13

Effects of acid rain: damage to lakes and streams

Harms forests

Stunts growth

Weakens/kills plants

Acid rain removes nutrients from the soil

Harms entire ecosystem

6.9

Plants need nitrogen in a chemical form that reacts more easily, such as the ammonium ion, ammonia, or the nitrate ion, in order to grow.

Simplified Nitrogen Cycle

6.9

Bacteria do not provide enough reactive nitrogen to supply the world with food, so synthetic fertilizers must be added.

Haber-Bosch process for large scale fertilizer production: N2(g) + H2(g) → 2 NH3(g)