Acid-Base Chemistry Notes

Acids and Bases Overview

• Acids and bases are fundamental concepts in chemistry, studied from Grade 7 to Grade 12.

• Understanding of acids and bases evolves and becomes more accurate over time.

• Key Topics Covered:

  • Review and empirical definition of acids and bases

  • Arrhenius' definition of acids and bases

  • Bronsted-Lowry theory of acids and bases

  • The pH scale

  • Calculating pH

  • pH indicators

  • Titrations

  • Buffers

  • Environmental chemistry involving acids and bases

  • Diploma practice questions

Acids in Everyday Life

• Many foods are acidic.

  • Examples: Lactic acid, citric acid, ascorbic acid, butyric acid

• Essential Learning Outcome:

  • Describe the use of acids and bases in the everyday world and in industrial processes.

Effects of Acids on Teeth

• Soda pop contains carbon dioxide, phosphoric acid, and citric acid.

• These acids erode tooth enamel.

• Bacterial plaque can reach inner layers of teeth more easily, causing cavities.

Neutralizing Stomach Acid

• Antacids like AlkaSeltzer, Rolaids, or Tums are used to relieve upset stomachs.

• Baking soda (sodium hydrogen carbonate, $NaHCO_3$) can also neutralize stomach acid.

Industrial Uses of Acids and Bases

• Acids and bases are crucial in producing:

  • Fertilizers

  • Plastics

  • Petroleum products

  • Pigments

  • Synthetic fabrics

  • Dyes

  • Pesticides

  • Paper

  • Soap

• Acids or bases are among the top 10 chemicals produced industrially.

Empirical Definition of Acids and Bases

• Empirical definition: based on experience and observation.

• Ways to test for acids and bases:

  • Taste: Not recommended due to safety concerns.

  • Litmus paper: Bases turn red litmus blue, acids turn blue litmus red.

  • Electrical conductivity: Strong acids or bases are more conductive.

  • pH meter: Most accurate method.

• Essential Learning Outcome:

  • Differentiate between acids, bases, neutral ionic compounds, and neutral molecular compounds based on diagnostic tests.

Observations of Acids, Bases, and Neutral Compounds

• Acids:

  • Corrosive

  • Turn blue litmus red

  • React with metals to produce $H_2(g)$

  • Usually sour

  • Electrolytes (conduct electricity)

• Bases:

  • Corrosive

  • Turn red litmus blue

  • Usually bitter

  • Electrolytes (conduct electricity)

• Neutral Ionic Compounds:

  • No effect on litmus paper

  • No reaction with metals

• Neutral Molecular Compounds:

  • No effect on litmus paper

  • No reaction with metals

The pH Scale

• Acid: 0-6.9

• Neutral: 7

• Base: 7.1-14

• Examples:

  • Lemon Juice: Acidic

  • Cabbage Water: Neutral

  • Milk of Magnesia: Basic

  • Hydrochloric Acid: Acidic

  • Lye: Basic

Logarithmic Nature of the pH Scale

• The pH scale is logarithmic, similar to the seismic scale.

• A difference of 1 pH unit represents a tenfold change in acidity or basicity.

  • pH 3 is 10 times stronger than pH 4.

  • pH 2 is 100 times stronger than pH 4.

  • pH 13 is 10,000 times stronger than pH 9.

Neutralization of Acids and Bases

• Strong acids are neutralized with weak bases.

• Strong bases are neutralized with weak acids.

Arrhenius' Definition

• Dissociation/Ionization: Separation of a chemical substance into individual ions.

• Acids: Compounds that dissociate in solution to form hydrogen ions ($H^+$).

  • Example: $HCl(aq) → H^+(aq) + Cl^-(aq)$

• Bases: Compounds that dissociate in solution to form hydroxide ions ($OH^-$).

  • Example: $NaOH(aq) → Na^+(aq) + OH^-(aq)$

Strength vs. Concentration of Acids

• Concentration: Amount of substance in solution, measured in mol/L.

• Strength: Depends on the amount of dissociation.

• Strong acids dissociate completely, resulting in high hydronium ion concentration ($H_3O^+$) and lower pH.

• Weak acids dissociate partially, resulting in low hydronium ion concentration and higher pH.

Hydronium Ions

• Hydrogen ions ($H^+$) combine with water (H2O$) to form hydronium ($H3O).

• High concentration of $H_3O^+$ indicates an acid, while high concentration of $OH^-$ indicates a base.

Limitations of Arrhenius' Theory

• Ammonia ($NH_3$) acts as a base but does not contain $OH^-$ ions.

• Arrhenius' theory is considered false due to such exceptions.

Brønsted-Lowry Theory

• Proton: A hydrogen ion ($H^+$).

• Acid: Proton donor.

• Base: Proton acceptor.

• Essential Learning Outcome:

  • Identify acids and bases as proton donors and acceptors.

Baseball Analogy

• If a proton is a baseball, an acid is the pitcher (donor) and a base is the catcher (acceptor).

Conjugate Acids and Bases

• Acids donate $H^+$ to become conjugate bases.

• Bases accept $H^+$ to become conjugate acids.

• Example: HCl(g) + H2O(l) → H3O+(aq) + Cl(aq)- In this reaction, HCl donates a proton to water, forming hydronium ions (H3O+) as the conjugate acid and chloride ions (Cl-) as the conjugate base.

  • HCl is the acid (proton donor).

  • $H_2O$ is the base (proton acceptor).

  • $H_3O^+$ is the conjugate acid of the base.

  • $Cl^-$ is the conjugate base of the acid.

Relative Strengths of Acids and Bases

• Strong acids:

  • Hydrochloric acid ($HCl$)

  • Sulfuric acid (H2SO4)

  • Nitric acid ($HNO_3$)

  • Hydronium ion ($H_3O^+$)

• Weak acids (examples):

  • Oxalic acid ($HOOCCOOH$)

  • Sulfurous acid (H2SO3)

  • Hydrogen sulfate ion ($HSO_4^-$)

Predicting Changes in Charge

• Acids donate (lose) a proton ($H^+$), becoming more negatively charged.

• Bases accept (gain) a proton ($H^+$), becoming more positively charged.

Brønsted-Lowry Theory Summary

• Acids donate protons ($H^+$) to bases.

• Bases accept protons ($H^+$) from acids.

• When an acid donates a proton, it forms a conjugate base.

• When a base accepts a proton, it forms a conjugate acid.

• Strong acids are at the top left of the table; strong bases are at the bottom right.

The pH Scale and Acidity

• The pH scale measures the degree of acidity or hydronium ($H_3O^+$) concentration.

• Hydronium ion concentration can approximate pH.

• Neutral solutions have equal concentrations of hydronium and hydroxide ions.

pH Scale Details

• Acidic: 0-6.9, high hydronium, low hydroxide

• Neutral: 7, equal hydronium and hydroxide

• Basic: 7.1-14, low hydronium, high hydroxide

Logarithmic Relationship of the pH Scale

• Each step on the pH scale represents a tenfold change in hydronium ion concentration.

Comparing Acidity

• For two solutions, the one with the lower pH is more acidic.

Tenfold Difference on the pH Scale

• A one-number change on the pH scale corresponds to a 10-fold difference in hydronium ion concentration (logarithmic).

Calculating pH

• Given hydronium ion concentration: pH = -log{10}[H3O^+]

• Given pH: $[H_3O^+] = 10^{-pH}$

• These formulas are in the data booklet.

• Essential Learning Outcomes:

  • Describe the relationship between pH and hydronium ion concentration.

  • Calculate pH given a hydronium ion concentration.

  • Calculate hydronium ion concentration given pH.

Acid-Base Indicators

• Substances that change color depending on the pH of the solution.

• Litmus paper turns red in acid and blue in base.

• Indicators were originally used as dye by Indigenous people.

• Essential Learning Outcomes:

  • Explain everyday uses of acid-base indicators.

  • Determine the pH of an unknown solution given indicator colors.

  • Determine the color an indicator will be given the pH of a solution.

Indicator Color Changes

• Methyl red changes from red to yellow between pH 4.8 and 6.0.

• If a solution is red after adding methyl red, pH is less than 4.8.

• If a solution is yellow after adding methyl red, pH is greater than 6.0.

• If a solution is orange after adding methyl red, pH is between 4.8 and 6.0.

Indicator Practice Questions

• The more indicators used, the more accurate the estimate of pH.

• Example: Alizarin yellow in acid is yellow (pH less than 7).

Titrations

• Titrations determine the unknown concentration of a solution by reacting it with measured quantities of a solution with known concentration (standardized solution) until the endpoint is reached.

• Usually involves adding a base to an acid until neutralization occurs.

• Indicators show when neutralization has occurred.

• Essential Learning Outcome:

  • Identify the equipment used in a titration experiment.

Standardized Solution

• The solution with known concentration used in a titration.

• The volume of standardized solution used to neutralize the unknown solution allows calculation of the unknown concentration, using the formula C1V1 = C2V2

Titration Equipment

• Pipette: Transfers the unknown solution.

• Burette: Contains the standardized solution.

• Erlenmeyer flask: Contains the unknown solution and indicator.

• Indicator shows the color change at the endpoint.

Steps in Performing a Titration

• Step 1: Add the standardized solution to the burette and record the initial volume.

• Step 2: Pipette the unknown solution into the Erlenmeyer flask and add 2-3 drops of indicator.

• Step 3: Slowly titrate the standardized solution into the flask until neutralization (endpoint). Observe the color change.

• Step 4: Record the final volume of the burette and calculate the volume used.

Titration Calculations

• When the mole-to-mole ratio is 1:1, use CaVa = CbVb to calculate the unknown value.

• This formula works only when the ratio is 1:1. Do not use this method in Chemistry 30.

• The number of moles of acid equals the number of moles of base at neutralization.

Buffers

• A buffer is a substance that resists changes in pH when exposed to an acid or a base.

• Example: Tums and Rolaids contain carbonate ions ($CO_3^{2-}$), which neutralize acids and maintain pH.

• Blood contains buffers to minimize pH changes, which can be fatal.

• Essential Learning Outcomes:

  • Explain how buffers maintain a relatively constant pH when a small amount of acid or base is added.

  • Explain the importance of maintaining a relatively constant pH in living systems.

  • Explain what is meant by buffering capacity.

Environmental Chemistry: Acid Rain

• Acid deposition includes both dry and wet deposition with a pH less than 5.6.

• The acidity arises from sulfur oxides (SOx$) and nitrogen oxides ($NOx).

• Dry deposition: SOx(g)$and$NOx(g) transported by wind.

• Wet deposition: SOx(g)$and$NOx(g) reacting with water in the atmosphere to form rain, snow, or hail.

Sources of Acid Deposition

• Natural: Volcanic eruptions, forest fires.

• Anthropogenic: Burning fossil fuels, refining crude oil, automobile combustion.

• Sulfur oxides ($SO_x$ ) from industry.

• Nitrogen oxides ($NO_x$) from combustion in cars

Chemical Reactions Leading to Acid Deposition

Sulfur Oxides:

• SO2(g) + H2O(l) → H2SO3(aq)

• Hydrogen sulfide (H2S) is also involved in the formation of sulfuric acid (H2SO4) through further oxidation.

• S8(s) + O2(g) → SO2(g) + SO3(g)

• SO3(g) + H2O(l) → H2SO4(aq)

Nitrogen Monoxides:

• N2(g) + O2(g) → NO_2(aq)

• NO2(g) + H2O(l) → HNO2(aq) + HNO3(aq)

Effects of Acid Deposition

Ecological impacts:

• Death of aquatic populations.

• Defoliation of plants and trees.

• Changes in soil pH.

• Washing away soil nutrients (Ca2+, Mg2+).

• Leaching of fertilizers and metal ions (nickel, zinc, lead) into water.

Societal Impacts:

• Increased rates of asthma and bronchitis.

• Corrosion of buildings, monuments, and statues.

Soil Nutrients and pH

• Iron is most available in acidic soil.

• Potassium, sulfur, boron, and molybdenum are most available in alkaline (basic) soil.

Limestone and Granite

• Limestone ($CaCO_3$) releases carbonate ions that neutralize acids (buffering).

• Granite contains silicon compounds that do not react with acids (no buffering).

Ways to Reduce Acid Deposition

• Electrostatic precipitators: Collect fly ash from emissions before release.

• Scrubbers: Remove sulfur dioxide from combustion by reacting $SO_2$ with a weak base.

• Catalytic converters: Reduce nitrogen oxides, carbon monoxide, and hydrocarbons by converting them to nitrogen, carbon dioxide & $H_2O$

• Liming: Adding calcium carbonate (base) to lakes to neutralize acidified water

• Using fuel cell cars (turns $H_2$ into electricity)

• Carpooling, walking, biking

• Alternate energy sources (wind and water)

• Turning off lights and appliances when not in use

• Recycling

• Improved insulation

Photochemical Smog

• Combination of smoke and fog, with emissions like NOx$, particulates,$H2S(g), and tropospheric ozone (ground-level ozone).

• Irritates eyes, nose, throat, and can cause asthma.

• Photochemical smog emissions increase as $NO_x$ levels increase during morning rush hours.

Ozone

• Ozone is O3(g)$and is formed when:$O(g) + O2(g) entoque O3(g)

• Stratospheric ozone protects us from damaging UV radiation

• Tropospheric ozone (ground level ozone, close to the ground) is "BAD"

Acid and Bases Summary

*   Empirical definitions are based on observable characteristics.
*   Acids: Taste sour, conduct electricity in solution, react with some metals to produce hydrogen
*   Bases: Taste bitter, conduct electricity in solution, feel slippery turn, neutralize acids
*   A change of 2 on a pH scale is a change of 10 x 10  in acid strength
*   Dissolving happens to molecular compounds, while dissociation -  breaking into ions - happens to ionic compounds.
*   Arrhenius: acids are compounds that give $H^+$ when dissolved in water, bases are substances that give $OH^-$ when dissolved in water.
*   Bronsted - Lowry Theory: Acids donate a proton to water to make hydronium, bases accept a proton from water to make hydroxide.
* Titrations uses solutions of known concentration to find solutions of unknown concentration. acid base titrations are neutralization reactions that use an indicator to find the endpoint, Be able to identify the equipment used (diploma exam question) Understand the experimental procedure, make titration calculations
*   Understand how buffers work, and buffer curves.   Environmental chemistry involving acids and bases