Science 30

Unit B: Chemistry and the environment Chapter 1: Acid Deposition

1.1- Products of Combustion Reactions

 Combustion reactions (eg. Cellular respiration, burning fossil fuels) are useful but produce emissions.

 Collisions between the methane and oxygen molecules form new molecules.

 If a hydrocarbon combusts, H 2 0 and CO 2 are formed; these are the waste products.

Chemical Equations

a) Oxides of Carbon

 Burning carbon compounds, or anything with biomass, or hydrocarbons results in carbon dioxide.

 Examples: C/R, volcanic eruptions  Carbon dioxide is important in the carbon cycle and as a greenhouse gas; contributes to climate change.

carbon cycle animation

Greenhouse effect

Greenhouse Effect - animated diagram

7

Hole in the ozone layer

Carbon monoxide

 Carbon monoxide (CO) is produced when the quantity of oxygen is limited in combustion.

 Caused by vehicle exhaust, furnaces (in poor condition); natural concentration of carbon monoxide in air is around 0.2 parts per million (ppm).

 Can bind to hemoglobin = decreased amount of oxygen to tissues.

CO

b) Oxides of sulfur

 Found when coal and crude oil or tar are burnt; found in natural gas as sour gas (hydrogen sulfide).

 Sour gas needs to have the hydrogen sulfide removed.

 The amount of SO 2 released depends on the sulfur content of coal, normally 0.7% to 2% by weight. High sulfur coal sometimes contains as much as 6% sulfur by weight.

Air pollution reduction

 Catalytic Converters  Catalytic Converters again  Energy efficient homes  Reduce recycle reuse Air Pollution

flaring

 When low quality natural gas is produced, it is burned off to produce SO 2 flaring.

and SO 3 =  Adds oxygen to hydrogen sulfide to form products.

c) Oxides of Nitrogen

 Whenever any fuel is combusted, nitrogen is present (includes breathing).

 When temperature reaches 650 °C, Nitrogen activates and forms NO and NO 2 .

 These are referred to as NO x compounds.

 Common sources:  Combustion of fuels in cars and furnaces.

 Fossil fuel power plants Fossil Fuels

d) Monitoring emissions

 The government creates standards to protect environment, organisms and support sustainability of resources.

 Monitored using specialized equipment and by an outside group.

 Cars are monitored using MAML labs.

1.2) Acids and Bases

 Acids, Bases and Neutral solutions have specific properties that are used to classify them.

 Acids: conducts a current (electrolyte), pH = 6 or less, corrosive, reacts with metals, tastes sour.

 Bases: Conducts current, corrosive, pH = 8 or more, feels slippery, bitter.

 Neutrals: pH = 7, can be electrolytic or not.

There are 4 different types of solutions

Neutral molecular Neutral ionic Acid base Solutions of compounds that are composed of non-metals only. Solutions of compounds that are composed of metals combined with non-metals Solutions of compounds that produce hydrogen ions. Solutions of compounds that produce hydroxide ions C 12 H 22 O 11(aq) NaCl(aq) HCl(aq) KOH(aq) CH 3 OH(aq) No effect on litmus CaBr 2 (aq) H 2 SO 4 (aq) Mg(OH) 2(aq No effect on litmus Turns litmus red Turns litmus blue

a) Types of deposition

 2 types of deposition:   Wet: emissions that contact precipitation and return as rain/snow.

Dry: gases and particles absorbed by the earth; deposited on any surface.

• Most deposition in Alberta is from dry deposition.

b) Acids

 Can be classified by properties (empirical) or by chemical composition.

 Classified as a molecular compound but behave like ionic compounds when dissolved in water (electrolytic solutions).

 The water molecules break the bonds in ionic compounds; these charges can now move in a direction = conduct electricity.

Electrostatic Attraction

 Force that pulls oppositely charged objects towards each other.

 Water pulls positive ions towards oxygen; creates a positive and negative charge.

 Dissociation occurs when 2 ions separate into different charges.

Arrhenius

 Svante Arrhenius formulated a theory in 1887 that all acids had a H+ ion and bases had an OH- ion.

 1.

2.

Problems: Not all acids and bases have an H+ or OH-.

H+ can not exist in water because it is so positively charged; actually forms H30 (hydronium ion) with a water particle.

1.

Hydronium ion is responsible for acidic properties.

c) Bronsted-Lowry Acid-base reactions

 Describes the actions of acids and bases during a chemical reaction.

 2 roles in the reaction:   Donor (acid) = gives H+ (proton) ion.

Acceptor (base) = accepts H+ (proton) ion.

 Product is a conjugate acid/base.

  Conjugate acid = formed when base accepts H+.

Conjugate base = formed when acid accepts H+.

Writing reactions

 Loss of a H+ ion by acid = conjugate base.

 Recognized by no H+ ion in formula.

 Gain of H+ ion by base = conjugate acid.

 Recognized by extra H+ ion in formula.

 Use the table of acids and bases (page 12 in data book) and follow 5 steps.

Steps in Bronsted-Lowry reactions

 1.

2.

3.

4.

5.

Follow these 5 steps to write reaction: Find the 2 solutions that are reacting.

Identify the acid and base.

1.

2.

Stronger acid is higher on the table- always choose the highest one if both are listed!

The base is the non-acid (or weak acid).

Write the reactants side of the equation.

Find conjugate form of acid and base.

Write conjugate forms on products side of the equation.

HCl acid-base

EXAMPLE - Conjugate Acids:

Write the formula for the conjugate acid of (a) F-, (b)NH3, (c) HSO4-, and (d) CrO4 2 .

Solution: In each case, the formula for the conjugate acid is derived by adding one H+ ion to the formulas above.

a.

HF

b.

NH4+

c.

H2SO4

d.

HCrO4 EXAMPLE - Conjugate Bases:

Write the formula for the conjugate base of (a) HClO3, (b)H2SO3, (c) H2O, and (d) HCO3-.

Solution: In each case, the formula for the conjugate base is derived by removing one H+ ion from the formulas above.

a. ClO3-

b.

HSO3-

c.

OH-

d.

CO3 2-

See page 174 Example problem 1.2/1.3

Proton hopping

 Confirmation of Bronsted-Lowry theory.

 Used lasers to do this; captured images of motion in chemical reactions.

 Helped predict outcomes of acid-base reactions.

d) Acid Deposition

 Emissions that are from human sources are

anthropogenic;

from combustion of energy sources.

 These emissions combine with water to form acid rain= acidic precipitation.

 Rain is acidic due to natural and human sources; the degree of acidity can be measured using pH.

e) pH and pH scale

 pH is measuring the amount of hydronium ions (H 3 0+) in a solution.

 The number of Hydronium ions influences:  Reactivity, amount of base needed to neutralize/to react.

 pH scale was developed in 1909 by Sorenson; designed to measure dilute acids.

 pH scale measures from 1-14 (1-6 = acid, 7= neutral, 8-14 = base).

pH

pH calculations

 1909 - S Ø ren Sørensen came up with “power of hydrogen” or pH  pH corresponds to the hydronium ion concentration in mol/L  [ H 3 O + (aq) ] = 10 -pH  pH = 5 then [ H 3 O + (aq) ] = 10  [ H 3 O + (aq) ] = 0.00001 mol/L -5 or 1.0 x 10 -5 mol/L 36

pH to mol/L and back again …



pH = - log [ H

3

O

+ (aq)

]



[ H

3

O

+ (aq)

] = 10

-pH 37

Converting pH to

H

3

O

+

concentration

 What is the concentration of hydrogen ions for an acid with a pH = 4.56?

 You must take the inverse of a log  That is probably the 10x button on your calculator  Type it in your calculator as  10 -4.56

= 0.000027542 mol/L  2.8 x 10 -5 mol/L (sig figs)  that is the concentration of hydrogen ions  be sure to put the negative sign in before the 4.56

38

Significant figures and pH/pOH

       Calculate the pH of a solution where the [H+] is 0.00100 M.

That's pretty easy, the answer is 3. After all 0.00100 is 10 ¯3 the negative log of 10 ¯3 is 3.

and But the pH is not written to reflect the number of significant figures in the concentration.

Notice that there are three sig figs in 0.00100 M.

So, our pH value should also reflect three significant figures.

Let's phrase that another way: in a pH (and a pOH), the only place where significant figures are contained is in the decimal portion.

So, the correct answer to the above problem is 3.000. Three sig figs and they are all in the decimal portion, NOT (I repeat NOT) in the whole number portion.

Example ...



What is the pH of a solution with a [H

3

O

+

] of: 1.89 x 10

-4

mol/L

Answer



pH = - log [1.89 x 10

-4

]



pH = 3.723 this is an acidic solution because the pH is less than 7

1 a) 1 b) 1 c) 1 d) 1 e) pH = 7 pH = 11 pH = 2 pH = 4 pH = 14

Practice

= 10 -7 = 1 x 10 -7 mol/L = 10 -11 = 1 x 10 -11 mol/L = 10 -2 = 1 x 10 -2 mol/L = 10 -4 = 1 x 10 -4 mol/L = 10 -14 = 1 x 10 -14 mol/L 41

Practice

2 a) [H 3 O + (aq) ] = 10 -3 mol/L 2 b) [H 3 O + (aq) ] = 10 -5 mol/L 2 c) [H 3 O + (aq) ] = 10 -7 mol/L 2 d) [H 3 O + (aq) ] = 10 -10 mol/L pH = - log [ 10 -3 ] =

3

pH = - log [ 10 -5 ] =

5

pH = - log [ 10 -7 ] =

7

pH = - log [ 10 -10 ] =

10

42

f) Indicators

 First nations used natural acids to adjust the color of the dyes made from leaves, berries and bark.

 An indicator is anything that changes color in response to a change in pH.

 Common indicators are shown in the table on page 12 of your booklet.

 Used to measure the pH of a substance.

 pH meter is more accurate and gives exact measure of pH.

The Rainbow Connection Demonstration

Examples:

According to the acid-base indicator table, what is the

color

of each of the following indicators in the solutions of given pH?

(a)

Phenolphthalein

in a solution with a

pH = 12.7

.

RED

(b)

Bromothymol blue

in a solution of

pH = 2.8

YELLOW

(c)

Methyl orange

in a solution of

pH = 3. RED

(d)

Thymol blue

in a

pH = 5.0

solution

YELLOW

(e)

Litmus

in a solution with a

pH of 8.2

BLUE

Example Problem:



Separate samples of an unknown solution turned both

methyl orange

and

bromothymol blue

to

yellow

, and turned

bromocresol green

to

blue

. The

pH

of the unknown solution is likely __________

Example Problem:



methyl orange

= yellow = 4.4+



bromothymol blue

= yellow = 6.0-



bromocresol green

=

blue

=

5.4+



The

pH

of the unknown solution is likely

between 5.4 and 6.0.

1.3) Impact of Acid Deposition

 Higher levels of sulphates and nitrates in rainwater = higher concentration of hydronium ions and lower pH in water.

 Wind patterns affect the deposition; provides way to trace pollution.

 Alberta soil is slightly basic (alkaline) from the carbonate caused by erosion of limestone; neutralizes the acid deposition.

a) Buffering

 A buffer is used to resist pH change in soil or lake water.

 The buffering capacity is the ability of a substance to resistance pH change when an acid/base is added.

 Specific plants need to be used in acidic soils; most are not able to grow due to lack of nutrients.

 Nutrients are deposited in soil from biogeochemical cycles and through neutralization reactions.

b) pH and plants

 Plants absorb nutrients into the roots from the soil, up to the leaves.

 Acid deposition changes nutrients so they are insoluble; plants can not use them.

 Nutrient deficiencies causes plants to die or become diseased.

 Chlorosis = calcium deficiency causes decreased chlorophyll = yellow leaves.

 A neutral pH is best for plant growth; too acidic or basic results in death.

c) Leaching

 acid in soil makes metal ions (aluminum and mercury) available to plants.

 Plants take those nutrients into their roots and they dissolve in the soil or groundwater = leaching.

 Affects ecosystem by:  Decreased root growth.

    Prevents absorption of calcium.

Reduce decomposing soil bacteria.

Mercury causes damage to gills.

Methyl mercury traps in tissues = bioaccumulation

Biomagnification

 A pollutant increases in concentration up a food chain.

 Causes disease and death.

Measured in ppm (10 6 ) or ppb (10 9 ) or ppt (10 12 )

d) Effect on Biotic factors

 The abiotic (non-living) factors in an ecosystem affect the biotic (living).

 Acid deposition decreases the biodiversity of the system.

 Acid deposition affects the ecosystem in:  Decreasing soil bacteria.

 Destroying waxy coating on plants.

 Damaging aquatic ecosystems.

1.4) Monitoring effects of acid deposition

 Classification of acids is done in 2 ways:  Quantitatively = involves measurement (pH, titration)  Qualitative = involves properties, characteristics, attributes (color, observations)

a) Titrations

 Used to determine the amount of acid/base present.

 Uses an acid, base and an indicator.

 When the “end point” is reached, the indicator will change color and the amount of acid/base needed to neutralize the solution is known.

Calculating the Concentration of Acids and Bases 

Calculating concentrations

of an acid solution.

: Calculate the concentration, in moles per litre, of 250 mL of a solution containing 0.243 mol of potassium hydroxide, KOH (s), used to analyze the concentration  Given: n = .243 mol V= 250 mL (0.250 L)  Formula: C = n/V = 0.243 mol/0.250 L = 0.972 mol/L  Answer: The KOH has a concentration of 0.972 mol/L  The

CONCENTRATION

of an acid can easily be changed by adding more solvent (water).

 What would the concentration of KOH be if an additional 150 mL of water was added to the solution?

Technique

    Solution #1: “ standard solution ” Will know the concentration and the volume used.

Solution #2: “ unknown ” Will know volume used but NOT the concentration Acid / Base titrations will involve a

NEUTRALIZATION

Reaction Neutralization Example: acid + base → salt + water HCl (aq) + NaOH (aq) → NaCl (aq) + H 2 O (l)

Titration Set-up

TITRANT

 solution with known concentration  goes in the

buret SAMPLE

 solution with

un

known concentration  goes in

Erlenmeyer Flask

Titration Process

 The titrant will be added to sample drop by drop until they have reacted fully (there is an equal amount of acid and base) and the

ENDPOINT

is reached.



How do you know you have reached the endpoint?

The Erlenmeyer flask will also contain an endpoint.

INDICATOR

that will change color when the reaction has reached the  The indicator changes colour at the endpoint because the [H 3 O + ] and therefore pH will have changed sufficiently.

How do you determine the volume of titrant used?

 Read the volume of titrant in the buret before the titration begins AND once the endpoint has been reached  Subtract the volumes  REMEMBER - read the

bottom

meniscus!!

of the

   

b) Acid Base Stoichiometry

KOH (aq) + HCl (aq)  KCl (aq) + HOH (aq) Calculate ‘ n ’ for the given substance Use a molar ratio to calculate ‘ n ’ for the required substance Make the appropriate calculation to answer the question you are asked.

Calculate: Find the concentration of a solution of potassium hydroxide, KOH(aq), if it requires 8.32 mL of a 0.100 mol/L standard solution of hydrochloric acid to neutralize 10.0 mL of the potassium hydroxide solution.

Strong and Weak Acids and Bases

 Strength

≠

Concentration  Strength refers to the % of the acid or base that dissociates in water  You can change the concentration but you can NOT change strength 

Strong

acids and strong bases

dissociate completely

in aqueous solutions  

Weak acids

- only a portion of the acid molecules release protons

Weak base

– only a portion of the base molecules accept protons

Examples of Strong and Weak Acids / Bases

 Example:  NaOH is a strong base   NaOH dissociates completely in H 2 O NaOH  Na + and OH  Example:  Acetic Acid is a weak acid  Most acetic acid molecules DO NOT release protons but instead remain in undissociated form   CH 3 COOH + H 2 O  CH 3 COO + H 3 O + For every weak acid there is a conjugate base (

See page 12 of Data Booklet

) What is the conjugate base of acetic acid?

c) Buffers

 A Buffer is a weak acid / weak base conjugate pair  Most effective when there are equal amount of acid and base  Example: H 2 CO 3 in Blood  CO 2 acid) and H 2 O react to make H 2 CO 3 (carbonic  Carbonic acid resists changes in blood pH H 2 CO 3 + H 2 O → H 3 O + + HCO 3 -

Sources of buffering

 Natural sources include hydrogen carbonate in your blood, levels of carbonate in the soil and limestone deposits.

 Buffers do not work forever however; Eastern Canada has reached the capacity to buffer soil and has acidic deposition currently.