Chemistry Chapter: Self-Ionization of Water
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Questions and Answers

What is the self-ionisation equation of water?

The self-ionisation equation of water is H₂O ↔ H⁺ + OH⁻.

What is the ionic product of water (Kw) at 25°C?

At 25°C, the ionic product of water (Kw) is 1 × 10⁻¹⁴.

How do you calculate pH from hydrogen ion concentration?

pH is calculated using the formula pH = -log₁₀[H⁺].

What is the pH of a solution with a hydrogen ion concentration of 7.1 x 10^-7 mol/L?

<p>The pH of that solution is approximately 6.15.</p> Signup and view all the answers

Define a neutral solution in terms of pH.

<p>A neutral solution has a pH equal to 7.</p> Signup and view all the answers

What is the equation to calculate pH from hydrogen ion concentration?

<p>pH = -log<del>10</del>[H^+^]</p> Signup and view all the answers

Calculate the hydrogen ion concentration when the pH of a solution is 3.7.

<p>[H^+^] = 1.995 x 10^-4^ mol/L</p> Signup and view all the answers

What differentiates strong acids from weak acids using the dissociation constant Ka?

<p>Strong acids have large K<del>a</del> values, while weak acids have small K<del>a</del> values.</p> Signup and view all the answers

Provide an example of a strong base and explain its characteristic.

<p>Sodium hydroxide (NaOH) is a strong base because it readily accepts protons.</p> Signup and view all the answers

Explain the relationship between strong acids and their conjugate bases.

<p>A strong acid has a weak conjugate base, such as HCl and Cl^-.</p> Signup and view all the answers

How does adding OH^- affect the equilibrium of the indicator reaction?

<p>Adding OH^- shifts the equilibrium to the left, favoring the reverse reaction and resulting in Colour 1 being seen.</p> Signup and view all the answers

What is the significance of the pH range for an indicator?

<p>The pH range indicates the interval over which the indicator changes color, allowing for effective monitoring of pH changes during titrations.</p> Signup and view all the answers

Describe the setup of a pH titration experiment.

<p>The setup includes a burette filled with a base, a beaker containing an acid and a magnetic stir bar, and a pH sensor connected to a computer to record pH changes.</p> Signup and view all the answers

What would you observe if H^+ is added to the indicator solution?

<p>Adding H^+ would react with OH^-, decreasing its concentration and shifting the equilibrium right, resulting in Colour 2 being seen.</p> Signup and view all the answers

How do indicators like methyl orange and phenolphthalein differ in terms of pH range and color change?

<p>Methyl orange has a pH range of 3-5, changing from red to yellow, while phenolphthalein has a range of 8-10, changing from colorless to pink.</p> Signup and view all the answers

What is the pH of a 0.04 M H2SO~4 solution?

<p>The pH is 0.097.</p> Signup and view all the answers

Calculate the pH of a 0.15 M NaOH solution.

<p>The pH is 13.18.</p> Signup and view all the answers

Identify the limitations of the pH scale.

<p>The pH scale is limited to 0.1-14, only applicable at 25°C, and does not work at extremely low concentrations.</p> Signup and view all the answers

How do you calculate the pH of a weak acid or weak base?

<p>You must know the dissociation constant (K<del>a</del> or K<del>b</del>) for the substance.</p> Signup and view all the answers

What is the formula for calculating the concentration of H^+ ions from a weak acid?

<p>[H^+] = \sqrt{K_a \times M_{acid}}.</p> Signup and view all the answers

Calculate the pH of a 0.1 M solution of methanoic acid with Ka of 2.1 x 10^-4.

<p>The pH is 2.34.</p> Signup and view all the answers

Describe the action of an acid-base indicator in a solution.

<p>An acid-base indicator changes color based on the pH of the solution.</p> Signup and view all the answers

What happens to an acid-base indicator when H^+ ions are added to the solution?

<p>Adding H^+ ions favors the reverse reaction, resulting in Colour 1 being seen.</p> Signup and view all the answers

Explain the role of indicators in acid-base titrations and provide an example.

<p>Indicators change color at a specific pH range, signaling the endpoint of a titration. For example, methyl orange is suitable for titrating a strong acid against a weak base.</p> Signup and view all the answers

What characterizes a strong acid when titrated against a strong base?

<p>A strong acid dissociates almost completely in water, resulting in a sharp increase in pH at the equivalence point of the titration.</p> Signup and view all the answers

Describe the pH change observed during the titration of a weak acid and a strong base.

<p>The pH rises gradually, becoming almost vertical around the equivalence point, indicating a significant change but not a sudden jump.</p> Signup and view all the answers

What is the significance of the conjugate acid-base pair?

<p>A conjugate pair consists of an acid and its corresponding base that differ by one hydrogen ion, indicating the reversible nature of acid-base reactions.</p> Signup and view all the answers

How does the strong dissociation of sulfuric acid differ from that of a weak monobasic acid?

<p>Sulfuric acid almost completely dissociates in water to produce more hydrogen ions, while a weak acid only partially dissociates.</p> Signup and view all the answers

Calculate the pH of a 0.5 M NaOH solution and explain your steps.

<p>pOH = -log10[OH⁻] = -log10(0.5) ≈ 0.3, thus pH = 14 - pOH = 14 - 0.3 = 13.7.</p> Signup and view all the answers

Why is it not possible to detect the endpoint in a titration of a weak acid with a weak base?

<p>The gradual pH rise does not yield a sudden change, making it difficult to pinpoint the endpoint with common indicators.</p> Signup and view all the answers

State the observed color change when a weak acid indicator is added to a strong base.

<p>The solution turns purple as hydroxide ions remove hydrogen ions from the indicator, shifting the equilibrium.</p> Signup and view all the answers

Flashcards

Self-ionization of water

The process where water molecules react with each other to produce H+ and OH- ions.

Ionic product of water (Kw)

The constant that expresses the product of the hydrogen ion and hydroxide ion concentrations in water at a specific temperature.

pH

A measure of the hydrogen ion concentration in a solution, calculated as the negative base-10 logarithm of the hydrogen ion concentration (mol/L).

pH of an acidic solution

A solution is acidic if its pH is less than 7.

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pH of a basic(alkaline) solution

A solution is basic(alkaline) if its pH is greater than 7

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pH of a solution

A measure of the hydrogen ion concentration in a solution, calculated as -log₁₀[H⁺].

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Strong Acid

A chemical compound that readily donates a proton(H⁺) in water.

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Weak Acid

A chemical compound that is poor at donating a proton(H⁺) in water.

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Conjugate Acid-Base Pair

Two chemical species related by the gain or loss of a single proton (H⁺).

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Calculating [H⁺] from pH

To find the hydrogen ion concentration ([H⁺]) from the pH, use the formula: [H⁺] = 10⁻ᵖʰ.

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Indicator Range

The pH interval where a specific indicator visibly changes colour. It's the range where the indicator's molecular structure alters due to pH changes, causing a colour shift.

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Le Chatelier's Principle (Indicators)

Adding base (OH-) shifts the indicator equilibrium towards the acidic form (colour 1), while adding acid (H+) shifts it towards the basic form (colour 2). This principle explains why indicators change colour based on the presence of acid or base.

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pH Titration

A controlled experiment measuring pH changes during the addition of a base to an acid. It helps determine the equivalence point, where the acid and base completely neutralize each other.

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pH Sensor in Titration

A device measuring pH changes during the addition of a base to an acid. It is connected to a computer to record the pH readings over time, enabling plotting of a pH versus time graph.

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What is the purpose of a pH titration?

To determine the equivalence point of a reaction between an acid and a base, where the amount of added base exactly neutralizes the acid.

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pH Titration Curve

A graph that shows the change in pH of a solution as a strong base (or weak base) is added to a strong acid (or weak acid) solution.

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End Point

The point in a titration where the indicator changes color, signifying that the reaction is complete.

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Strong Acid-Strong Base Titration

The titration of a strong acid with a strong base, resulting in a steep pH jump at the equivalence point.

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Weak Acid-Strong Base Titration

The titration of a weak acid with a strong base, resulting in a gradual pH change.

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Suitable Indicator

An indicator that changes color close to the equivalence point, highlighting the completion of the reaction.

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Brønsted-Lowry Acid

A substance that donates a proton (H+) in a chemical reaction.

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Strong acid pH calculation

To calculate the pH of a strong acid solution, you first need to determine the concentration of hydrogen ions (H+) produced by the acid. Then, use the formula: pH = -log10[H+].

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Strong base pH calculation

To calculate the pH of a strong base solution, you first need to determine the concentration of hydroxide ions (OH-) produced by the base. Then, use the formula: pOH = -log10[OH-]. Finally, calculate pH using the relationship: pH + pOH = 14.

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pH scale limitations

The pH scale, while useful, has limitations. Its range is typically restricted from 0 to 14, but concentrations outside this range are theoretically possible. It's also specific to dilute aqueous solutions at 25°C. Additionally, it doesn't apply at extremely low concentrations.

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Weak acid pH calculation

Calculating the pH of a weak acid solution requires considering its dissociation constant (Ka). Use the formula: [H+] = √(Ka * [acid]), where [acid] is the acid's concentration in moles per liter.

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Weak base pH calculation

Similar to weak acids, calculating the pH of a weak base solution involves its dissociation constant (Kb) and the formula: [OH-] = √(Kb * [base]), where [base] is the base's concentration in moles per liter.

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Acid-base indicator

An acid-base indicator is a substance that changes color depending on the pH of the solution it's in. This color change is due to the indicator's structure shifting as it acts as a weak acid or base.

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Indicator color change: Le Chatelier's Principle

The color change in an indicator is explained by Le Chatelier's Principle. Adding acid favors the reverse reaction, showing the acid form's color. Adding base removes H+ ions, favoring the forward reaction and showing the base form's color.

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How does an indicator work?

Indicators change color because of the shift in equilibrium of the chemical reaction within the indicator molecule. The indicator exists in two forms, one of which is acidic and one of which is basic. The acidic form has a different color than the basic form, and when the pH changes, the equilibrium shifts to favor one form or the other, changing the solution's color.

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Study Notes

Self-Ionization of Water

  • Water conducts electricity when it contains dissolved ions.
  • Pure water has a very small current due to self-ionization.
  • Water self-ionizes as follows: H₂O → H⁺ + OH⁻
  • The concentration of H⁺ and OH⁻ ions in pure water is very small.
  • The equilibrium in the self-ionization reaction is strongly to the left.
  • The concentration of H₂O remains relatively constant.
  • Equilibrium constant (K) can be written as: K = [H⁺][OH⁻]/[H₂O]
  • Since [H₂O] is effectively constant, the expression can be simplified to: Kw = [H⁺][OH⁻]
  • At 25°C, Kw = 1 x 10⁻¹⁴
  • To find [H⁺] given Kw, use the fact [H⁺]=[OH⁻]: [H⁺] = √(1 x 10⁻¹⁴) = 1 x 10⁻⁷ mol/L

pH Scale

  • pH is defined as: pH = -log₁₀[H⁺]
  • Square brackets indicate concentration in moles per litre.
  • pH is the negative logarithm to the base 10 of hydrogen ion concentration.
  • pH of a solution is measured in moles per litre.
  • pH values range from 0 to 14.
  • pH values below 7 are acidic; above 7 are alkaline; 7 is neutral.

Strengths of Acids and Bases

  • Strong acid: A good proton donor (e.g., HCl, H₂SO₄)
  • Weak acid: A poor proton donor (e.g., ethanoic acid)
  • Strong base: A good proton acceptor (e.g., NaOH, KOH)
  • Weak base: A poor proton acceptor (e.g., NH₃)
  • Acid dissociation constants (Ka) measure acid strength.
  • Larger Ka values indicate stronger acids.
  • Base dissociation constants (Kb) measure base strength.
  • Larger Kb values indicate stronger bases.

Calculating pH of Strong Acids and Bases

  • Strong acids and bases are fully dissociated in water.
  • The pH of a strong acid solution can be calculated by using the concentration of H+ ions in solution.
  • The pH of a strong base solution can be calculated by using the concentration of OH-ions in solution.

Calculating pH of Weak Acids and Bases

  • Calculating the pH of a weak acid solution requires consideration of the dissociation constant (Ka).
  • The approach uses the equilibrium constant (Ka) and the initial concentration of the weak acid to find the concentration of H+.
  • To calculate the pH of a weak acid solution, we must know its dissociation constant.

pH Indicators

  • Acid-base indicators: Substances that change color depending on the pH of the solution.
  • Indicators have a specific pH range where their color changes.
  • Choosing the correct indicator depends on the pH range of the solution being tested.
  • Indicator color changes are affected by adding H⁺ or OH⁻ ions according to Le Chatelier's principle.

pH Titration

  • A pH titration is an experiment to track pH changes during a titration.
  • The process involves adding a base (from a burette) to an acid (in a beaker).
  • pH is monitored continuously.
  • The pH data is typically plotted as a graph.
  • The graph (titration curve) shows how pH changes as the base is added.
  • Titration curves for strong acid-strong base, weak acid-strong base and weak acid-weak base reactions have typical shapes.

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Explore the self-ionization of water and understand the pH scale through this quiz. You'll learn about the equilibrium constants, the formation of H⁺ and OH⁻ ions, and how pH values indicate acidity or alkalinity. Test your knowledge and see how well you grasp these fundamental concepts in chemistry.

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