General Outcome B1 - Students will analyze the sources of acids and bases and their effects on the environment.

30–B1.1k define acids and bases in terms of proton donors and proton acceptors
30–B1.2k differentiate among acids, bases, neutral ionic compounds, neutral molecular compounds and strong and weak acids, based on appropriate diagnostic tests
30–B1.3k describe the relationship between pH and hydronium ion concentration
30–B1.4k explain, qualitatively, how buffers maintain a relatively constant pH when a small amount of acid or base is added to an aqueous system
30–B1.5k explain the importance of maintaining a relatively constant pH in a living system; e.g., the role of the hydrogen carbonate ion in maintaining the pH of blood, the evolution of the Arctic herb Artemisia tilesii in resisting acidic moisture by extracting calcium from the soil and pumping the calcium to its leaves
30–B1.6k trace the historical use of acid-base indicators; e.g., early Aboriginal methods of using extracts from natural substances
30–B1.7k explain what is meant by buffering capacity; e.g., soil or bedrock
30–B1.8k outline the chemical reactions (e.g., combustion reactions) that produce air pollutants (i.e., sulfur dioxide and nitrous oxides) that, when combined with water, ultimately result in acid deposition
30–B1.9k describe impacts on the biotic and abiotic components of the environment caused by acid deposition; e.g., lowered pH in water systems, accelerated corrosion, metal leaching from bedrock, the impact of leached metals on plants and the food chain.

1.1 Products of Combustion Reactions

Combustion Reactions and Their Products
Combustion is when a fuel (hydrocarbon) reacts with oxygen to form carbon dioxide and water.

fuel + oxygen	→	carbon dioxide + water
methane + oygen	→	carbon dioxide + water	word equation
CH4 + O2	→	CO2 + H2O	skeleton equation
CH4 + 2O2	→	CO2 + 2H2O	balanced equation

Example 1

C5H12 + 8O2	→	8CO2 + 6H2O	skeleton equation
C5H12 + O2	→	CO2 + H2O	balanced equation

Example 2

C4H8 + O2	→	CO2 + H2O	Enough oxygen
C4H8 + 8O2	→	4CO2 + 4H2O	Restricted oxygen

Example 3

C4H10 + O2	→	CO2 + H2O	skeleton equation
2C4H10 + 13O2	→	CO2 + H2O	balanced equation

Carbon Monoxide
When there isn’t enough oxygen in a combustion reaction, sometimes carbon monoxide is produced instead of carbon dioxide. This most often occurs when combustion occurs inside a building.

C8H16 + 12O2	→	8CO2 + 8H2O	Enough oxygen
C4H8 + 8O2	→	8CO + 8H2O	Restricted oxygen

When carbon dioxide mixes with water it undergoes a chemical reaction and forms carbonic acid. This falls as acid rain.

Equation 1:	CO2 (g) + H2O (l)	→	H2CO3 (aq)
Equation 2:	H2CO3	→	H+ + HCO- (aq)

Carbonic acid is a weak acid so not all of it dissociates.

Oxides of Sulfur

• Sulfur is often found in coal and oil.
• Dihydrogen sulfide H₂S is called sour gas and is found in oil and gas deposits.
• Sulfur can be removed from fuel during refining.
• When combusted they form sulfur oxide gas (sulfur dioxide & sulfur trioxide)

sulfur + oxygen	→	sulfur dioxide	word equation
S8 + O2	→	SO2	skeleton equation
S8 + 8O2	→	8SO2	balanced equation

When sulfur dioxide mixes with rain it forms sulfuric acid. This then falls as acid rain.

Equation 1:	SO2 (g) + O2 (g)	→	SO3 (g)
Equation 2:	SO3 (g) + H2O (l)	→	H2SO4 (aq)
Equation 3:	H2SO4 (aq)	→	H+ + HSO4- (aq)
Equation 4:	HSO4- (aq)	→	H+ + SO42- (aq)

Sulfuric acid is a strong acid which completely dissociates so it causes more damage.


Oxides of Nitrogen

• Nitrogen is 78% of our atmosphere. When air is sucked into an engine the heat and pressure can cause nitrogen to react with oxygen and form nitrogen oxides (NO_x: NO & NO₂).
• 𝑁₂ (𝑔) + 𝑂₂ (𝑔) → 2𝑁𝑂 (𝑔)
• 𝑁₂ (𝑔) + 2𝑂₂ (𝑔) → 2𝑁𝑂₂ (𝑔)
• These are removed from car exhaust by catalytic converters. The catalyst (metals) and heat in the converter break the NOx apart into nitrogen and oxygen gas.
• When nitrogen oxides get into the atmosphere they convert into nitric acid.

Equation 1:	NO (g) + O2 (g)	→	2NO2 (g)
Equation 2:	H2SO4 (aq)	→	H+ + HSO4- (aq)
Equation 3:	HNO3 (aq)	→	H+ + NO3- (aq)

• Nitric acid is a strong acid which completely dissociates so it causes more damage.

Monitoring Emissions in Alberta
Most oilfield and gas sites (drilling, refineries, pumps, compressors, etc…) have provincially monitored sensors.
If there is a gas release it is measured. Companies will be fined if they exceed provincial limits and heavily fined if they don’t report emissions. Dangerous gases are occasionally flared (burnt) to prevent releases that could harm people living nearby.

1.2 Chemistry of Acids and Bases

What Makes a Solution Acidic?
Arrhenius determined that:

• Acids are formed when hydrogen ions (H+) are released in water.
  HCL (aq) → H⁺ (aq) + Cl^- (aq)
• Acids are formed when hydrogen ions (H+) are released in water.
  NaOH (aq) → NA⁺ (aq) + OH⁺ (aq)
• Neutrals are formed when no hydrogen or hydroxide ions are released in water.
  NaCl (aq) → Na⁺ (aq) + Cl^- (aq)

Limitations to Arrhenius’s Theory

• Over time acidic substances without hydrogen were discovered.
• Basic substances without hydroxide were discovered.
• Water is polar (positive and negative side) which would attract loose positive hydrogen ions. This forms a hydronium ion that Arrhenius didn’t explain.

Exchange of Hydrogen Ions
Brønsted-Lowry formed a new theory for acids and bases.
According to this theory, a hydrogen ion is transferred from an acid (the donor) to a base (the acceptor) during acid-base reactions. The Brønsted-Lowry theory often refers to the hydrogen ion as a proton. The products of an acid-base reaction are a conjugate acid and a conjugate base.

Dissolving hydrogen chloride to form hydrochloric acid.
Arrhenius:	HCl (s) HCl (s) + H2O (l)	→ →	H+ (aq) + Cl- (aq) H+ (aq) + Cl- (aq) + H2O (l)
Brønsted-Lowry:	HCl (s) + H2O (l)	→	H3O+ (aq) + Cl- (aq)

Dissolving sodium hydroxide in water.
Arrhenius:	NaOH (s) NaOH (s) + H2O (l)	→ →	Na+ (aq) + OH- (aq) Na+ (aq) + OH- (aq) + H2O (l)
Brønsted-Lowry:	NaOH (s) + H3O+ (aq)	→	Na+ + 2H2O (aq)

Strong acids completely dissociate in water. They all breaks apart into ions.
HCl (aq) + H₂O (l) → H₃O⁺ (aq) + Cl^- (aq)
Weak acids partially dissociate in water. Some will break apart or go back together depending on concentrations.
H₂CO₃ (aq) + H₂O (l) ↔ H₃O⁺+ (aq) + HCO₃^- (aq)

Acid Deposition

• Acidic gases released into the environment react with water droplets in the atmosphere (clouds, rain & snow).
• These land on ground or waterways acidifying the land and water.

pH
The strength of an acid or base is measured as the pH of the acid.

• Low pH values are acidic, less than 6
• High pH values are basic, more than 8
• pH is a measurement of the concentration of hydronium
• pH = -log₁₀[H₃O⁺ (aq)]
• [H₃O⁺ (aq)] = 10^-pH

Ex 1. What is the pH of a solution with a hydronium ion concentration of 3.7x10^-2 mol/l?


Ex 2. What is the pH of a solution with a hydronium ion concentration of 1.6x10-10 mol/l?

You can estimate the pH of a solution as the exponent on the 10.

[H3O+]	Estimate	Actual
3.5 × 10-2	2	1.5
8.7 × 10-10	10	9.1
3.5 × 10-8	8	7.3

Ex 3. What is the hydronium ion concentration if the pH of a solution is 3.4?

Ex 4. What is the hydronium ion concentration if the pH of a solution is 8.1?

Using Indicators to Estimate pH

• Indicators are chemicals that change colors at different pH’s.

Ex. An unknown chemical is tested with orange IV which turns yellow. Then phenolphthalein turns pink. Then indigo carmine turned blue. Estimate the pH of the solution.

Orange IV turning yellow pH >2.8
Phenolphthalein turning pink pH >10
Indigo Carmine turning blue pH < 11.4
The pH is between 10 and 11.4

Ex. An unknown chemical is tested with methyl red and turns red. Methyl orange turns yellow.

Methyl red turns red pH < 4.8
Methyl orange turns yellow pH > 4.4
Bromcresol green turns green 3.8 < pH < 5.4

Using a pH meter

• Acids and bases are electrolytes so they conduct electricity.
• A pH meter measures how easily the liquid conducts electricity to determine the concentration of hydronium ions and the pH.

1.3 Impact of Acid Deposition on Ecosystems

Effects of Acid Deposition on the Environment and Ecosystems

• Acid deposition acidifies land and waterways. Some plants and animals can’t live in a more acidic environment so they die.
• Plants and animals that can survive in more acidic environments (very few) can expand their habitat.
• Some areas are so acidic nothing can survive so the lakes are dead.
• Some lakes on the Canadian shield are metamorphic or intrusive rocks that contain acid buffers (limestone). The hydronium reacts with the limestone on the bottom of the lake and neutralize the acid.
• Calcium carbonate CaCO3 (limestone) reacts with acid and neutralizes it. The limestone only reacts with acids and doesn’t affect the water in normal conditions.
• Some lakes have been “limed” when ground limestone is added to lakes to prevent acidification.

Reaction of Hydronium Ion and carbonate Ion

H3O+ (aq)	+	CO32- (aq)	→	H2O (l)	+	HCO3- (aq)
acid		base		conjugate base		conjugate acid

Limestone (calcium carbonate ) dissolves in water and dissociates.
CaCO₃ (s) → Ca²⁺ (aq) + CO₃^2- (aq)
This forms carbonate ions. As carbonate is used up by acid, more limestone dissolves releasing more carbonate into the lake.
CO₃^2- (aq) + H₃O⁺ (aq) ↔ HCO₃^- (aq) + H₂O (l)
If there is a lot of hydronium, the carbonate will accept a proton (H⁺) to become hydrogen carbonate to decrease the hydronium concentration, making it less acidic.
If the lake water becomes more basic (less hydronium) the hydrogen carbonate releases a hydrogen ion (proton) to form more hydronium.
HCO₃^- (aq) + H₂O (l) ↔ CO₃^2- (aq) + H₃O⁺ (aq)
Changes in hydronium ion concentration (pH) cause the reaction to go one way and release hydronium, or the other way and absorb hydronium. This is due to partial dissociation.

Plant Nutrients, Metal Leaching, and pH

• When acids flow through soil they can react with metals and release metal ions.
• Some metals ions like lead, mercury and aluminum are toxic with different effects.
• Some metals ions like calcium and magnesium are beneficial for people and plants.

1.4 Quantifying Acid Deposition and Monitoring Its Effects

Titration—Quantifying Acid in a Solution
A titration uses a known concentration to neutralize and change the pH of an unknown concentration solution.

H3O+ (aq)	+	OH- (aq)	→	H2O (l)	+	H2O (l)
acid		base		conjugate base		conjugate acid

C = n/V
n_i = n_f C_iV_i = C_fV_f
Ex.Ex. An environmental scientist takes a sample of acidic lake water. The scientist performs a titration using 25 mL of 0.010 mol/L NaOH to neutralize the 10 mL of acidic lake water. What is the concentration of the acid?
C₁V₁ = C₂V₂
25 ml × 0.010 mol/l = c₂ × 10 ml
c₂ = 0.025 mol/l
Limestone in water releases carbonate ions. The carbonate ions are the buffer that absorb hydrogen decreasing the hydronium ion concentration.
However once the limestone is all dissolved there is no more carbonate to stop the pH from changing because the buffer (carbonate ions) are all used up.

1.5 Learning from Acid Deposition

Acidic smoke particles can be removed from smoke stacks before it is released into the environment:

• Electrostatic precipitators use electrostatic attraction to collect the particles before they leave the chimney.
• Scrubbers spray limewater through smoke stacks to neutralize acidic particles before they are released. Catalytic converters add oxygen to convert CO to CO₂ and NO_x to N₂ and O₂.

Scrubbers

• Limewater with sulfur dioxide
• Calcium carbonate and sulfur dioxide from calcium sulfite and carbon dioxide
  CaCO₃ (aq) + SO₂ (g) → CaSO₃ (s) + CO₂ (g)
• The calcium sulphite is a solid. When it gets to the liquid at the bottom of the scrubber tank it settles at the bottom of the tank and can be collected.
• The liquid limewater can then be reused (sprayed again). Over time more calcium carbonate must be added to replace what was converted to calcium sulfite.

Catalytic Convertor

• The catalyst metals help these reactions to occur.
  

  2CO (g) + O+ (g)	→	2 CO2 (g)
  2 NO (g) + O2 (g)	→	N2 (g) + 2 O 2 (g)
  2 NO2 (g)	→	N2 (g) + 2 O 2 (g)

• Also any unburned hydrogen or carbons or hydrocarbons will also react with extra oxygen to form carbon dioxide and water.

Air pollution is most often experienced as smog, a yellow/orange-ish fog. Depending on weather and geography some locations have more smog than others. Temperature inversions and mountains (significant elevation changes) can trap pollution causing more smog.

Photochemical Smog Reaction 1
Within car engine	2 CO (g) + O2 (g)	→	2 CO2 (g)
In atmosphere	2 NO (g) + O2 (g)	→	2 NO2 (g)
Photochemical Smog Reaction 2
Powered by sunlight	NO2 (g)	→	NO (g) + O (g)
Photochemical Smog Reaction 3
	O (g) + O2 (g)	→	O3 (g)
	Note: This reaction involves another substance, like N2 (g) present in the atmosphere to act as a catalyst.

Thinking Smart—Not Creating New Problems
Hydrogen fuel cells combine hydrogen gas (from fuel tank) with oxygen (from air) in a special plastic layer to form water and release electricity. It doesn’t produce any pollution and isn’t hot enough to generate NOx. The problem is where to get hydrogen from. Electric cars are all powered from one electric power generation station. That makes one location when the exhaust can be scrubbed, electrically precipitated and/or catalytic converted releasing very little pollution for the amount of power produced. Sulfur can be removed from fuel (gasoline, diesel) before it used. North America has changed from 500 ppm in 2000 to 15 ppm in 2012 European fuel is currently 10 ppm, their sulfur free fuel is 3-5 ppm.

Chapter 1 Summary

Chapter 1 Review Questions