Ever wondered how chemical reactions actually work in real life?...
Chemistry160 views·Updated 14 Jul 2026·10 pagesDynamic Equilibrium and Acids Study Notes
Marianne @mdesprsdavies_wcgo
1 / 10
1
of 10
Dynamic Equilibrium Fundamentals
Think of chemical reactions like a busy motorway - traffic flows both ways, but sometimes it reaches a point where the flow balances out perfectly. That's exactly what happens in dynamic equilibrium.
In a closed system (where nothing can escape or be added), reversible reactions eventually reach a state where the forward and backward reactions happen at equal rates. The key word here is dynamic - the reaction never actually stops, but the overall amounts of each substance stay constant.
Dynamic equilibrium only occurs in closed systems with reversible reactions. It's like a perfectly balanced seesaw that's still moving - the position stays the same, but the motion continues underneath.
Quick Check: Remember that dynamic equilibrium means the reaction is still happening, just at equal rates in both directions!
2
of 10
Le Chatelier's Principle
Here's where chemistry gets predictable - Le Chatelier's Principle is like a universal rule that helps you predict what happens when you mess with an equilibrium reaction. If you disturb the balance, the reaction will shift to counteract your change.
The position of equilibrium tells you which side of the reaction is favoured. Think of it like a seesaw - if there's more product than reactant, the equilibrium "lies to the right" (towards products).
When you add a catalyst, don't expect the equilibrium position to change. Catalysts speed up both forward and backward reactions equally, so they just help you reach equilibrium faster without affecting the final balance.
Pro Tip: Le Chatelier's Principle is your best friend for predicting reaction behaviour - the system always tries to undo whatever change you make!
3
of 10
Concentration and Pressure Changes
Want to push a reaction in a particular direction? Changing concentration is your go-to method. If you increase the concentration of a reactant, the equilibrium shifts right to use up that excess. Decrease it, and the equilibrium shifts left to make more.
For gaseous reactions, pressure changes create interesting effects. Higher pressure favours the side with fewer gas molecules - it's the reaction's way of reducing pressure by making fewer particles bash against the container walls.
The magic formula: count the molecules on each side of the equation. If you've got 3 molecules on the left and 2 on the right, higher pressure will push the equilibrium towards the right side.
Remember: Pressure changes only matter for gases, and only when the number of molecules differs on each side!
4
of 10
Temperature Effects
Temperature changes are where things get interesting because they actually change which direction is favoured. Exothermic reactions release heat, whilst endothermic reactions absorb it.
When you increase temperature, the equilibrium shifts to favour the endothermic direction (the one that absorbs heat). It's like the reaction is trying to cool things down by absorbing your extra heat energy.
If you want more products from an exothermic forward reaction, decrease the temperature. The equilibrium will shift right to produce more heat and counteract the cooling. This is why many industrial processes use carefully controlled temperatures.
Key Insight: Temperature is the only factor that actually changes the equilibrium constant - everything else just shifts the position!
5
of 10
The Equilibrium Constant (Kc)
The equilibrium constant (Kc) is chemistry's way of putting a number on how far a reaction proceeds. It's calculated using the concentrations of products and reactants at equilibrium, and it stays constant at a given temperature.
For any reaction aA + bB ⇌ cC + dD, the expression is: Kc = [C]^c[D]^d / [A]^a[B]^b. Products go on top, reactants on bottom, and the powers come from the balanced equation.
Homogeneous equilibria involve everything in the same phase (all gases or all in solution). These are the most straightforward to work with because you include all species in your Kc expression.
Top Tip: The equilibrium constant only changes with temperature - pressure and concentration changes don't affect its value!
6
of 10
Kc Examples and Heterogeneous Systems
Let's see Kc in action with real industrial processes. For the contact process (2SO₂ + O₂ ⇌ 2SO₃), the expression becomes: Kc = [SO₃]² / [SO₂]²[O₂]. Notice how the powers match the coefficients in the balanced equation.
Heterogeneous equilibria involve different phases (like solids with gases). Here's the crucial rule: don't include solids in your Kc expression. For H₂O + C ⇌ H₂ + CO, you get Kc = [H₂][CO] / [H₂O].
The Haber process gives us Kc = [NH₃]² / [N₂][H₂]³. These industrial examples show how equilibrium principles apply to real-world chemistry.
Remember: Solids don't appear in Kc expressions because their concentrations don't change significantly!
7
of 10
Acid-Base Theory Essentials
The Brønsted-Lowry theory makes acid-base chemistry crystal clear: acids donate protons (H⁺), whilst bases accept them. This definition works for way more reactions than you might think.
Strong acids like HCl completely dissociate in water: HCl → H⁺ + Cl⁻. Weak acids like ethanoic acid only partially dissociate: CH₃COOH ⇌ H⁺ + CH₃COO⁻. Notice the equilibrium arrow for weak acids!
Don't confuse strong/weak with concentrated/dilute. Concentration tells you how much acid is dissolved, whilst strength tells you how completely it breaks apart. You can have a dilute solution of a strong acid!
Key Point: All neutralisation reactions are fundamentally the same - H⁺ + OH⁻ → H₂O - which is why they all release similar amounts of energy!
8
of 10
The pH Scale and Calculations
The pH scale is a logarithmic way of expressing H⁺ concentration: pH = -log[H⁺]. Each pH unit represents a 10-fold change in acidity - pH 1 is ten times more acidic than pH 2.
Converting from H⁺ concentration to pH is straightforward: just stick the concentration into your calculator and use the -log function. For example, [H⁺] = 0.0100 mol dm⁻³ gives pH = 2.00.
Going the other way (pH to concentration) requires the inverse operation. For pH = 3.42, you calculate 10^ = 3.80 × 10⁻⁴ mol dm⁻³. Most scientific calculators have a "10^x" function for this.
Calculator Tip: Practice these conversions - they're fundamental for any quantitative acid-base work and frequently appear in exams!
9
of 10
More pH Conversions and Titration Intro
Getting comfortable with pH calculations takes practice, but the pattern is always the same. Whether you're going from concentration to pH (use -log) or pH to concentration (use 10^), the key is staying organised with your powers of ten.
Volumetric analysis (titrations) is chemistry's way of finding unknown concentrations. You know one concentration and measure two volumes, then use stoichiometry to find the unknown concentration.
The beauty of titrations lies in their simplicity - you're just using the fundamental relationship n = cV (moles = concentration × volume) combined with the balanced chemical equation to work out ratios.
Success Strategy: Always convert cm³ to dm³ by dividing by 1000 - this is the most common mistake in titration calculations!
10
of 10
Titration Calculations Made Simple
Here's the step-by-step approach that works every time: write the balanced equation, find moles in the known solution, use stoichiometry to find moles of the unknown, then calculate concentration.
For the example: HCl + NaOH → NaCl + H₂O. With 25.0 cm³ of 0.10 mol dm⁻³ NaOH neutralising 15.0 cm³ of HCl, first find moles of NaOH: n = 0.10 × 0.025 = 2.50 × 10⁻³ mol.
Since the reaction is 1:1, you've got the same moles of HCl. Finally, concentration = moles ÷ volume = 2.50 × 10⁻³ ÷ 0.015 = 0.167 mol dm⁻³. The systematic approach makes these calculations manageable every time.
Exam Success: Always show your working clearly - even if you make an arithmetic error, you'll get method marks for the correct approach!
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Chemistry160 views·Updated 14 Jul 2026·10 pagesDynamic Equilibrium and Acids Study Notes
Marianne @mdesprsdavies_wcgo
Ever wondered how chemical reactions actually work in real life? This guide breaks down the fascinating world of chemical equilibrium and acid-base chemistry - essential concepts that explain everything from industrial processes to the pH of your drinks.
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Dynamic Equilibrium Fundamentals
Think of chemical reactions like a busy motorway - traffic flows both ways, but sometimes it reaches a point where the flow balances out perfectly. That's exactly what happens in dynamic equilibrium.
In a closed system (where nothing can escape or be added), reversible reactions eventually reach a state where the forward and backward reactions happen at equal rates. The key word here is dynamic - the reaction never actually stops, but the overall amounts of each substance stay constant.
Dynamic equilibrium only occurs in closed systems with reversible reactions. It's like a perfectly balanced seesaw that's still moving - the position stays the same, but the motion continues underneath.
Quick Check: Remember that dynamic equilibrium means the reaction is still happening, just at equal rates in both directions!
2
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By signing up you accept Terms of Service and Privacy Policy
Le Chatelier's Principle
Here's where chemistry gets predictable - Le Chatelier's Principle is like a universal rule that helps you predict what happens when you mess with an equilibrium reaction. If you disturb the balance, the reaction will shift to counteract your change.
The position of equilibrium tells you which side of the reaction is favoured. Think of it like a seesaw - if there's more product than reactant, the equilibrium "lies to the right" (towards products).
When you add a catalyst, don't expect the equilibrium position to change. Catalysts speed up both forward and backward reactions equally, so they just help you reach equilibrium faster without affecting the final balance.
Pro Tip: Le Chatelier's Principle is your best friend for predicting reaction behaviour - the system always tries to undo whatever change you make!
3
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Concentration and Pressure Changes
Want to push a reaction in a particular direction? Changing concentration is your go-to method. If you increase the concentration of a reactant, the equilibrium shifts right to use up that excess. Decrease it, and the equilibrium shifts left to make more.
For gaseous reactions, pressure changes create interesting effects. Higher pressure favours the side with fewer gas molecules - it's the reaction's way of reducing pressure by making fewer particles bash against the container walls.
The magic formula: count the molecules on each side of the equation. If you've got 3 molecules on the left and 2 on the right, higher pressure will push the equilibrium towards the right side.
Remember: Pressure changes only matter for gases, and only when the number of molecules differs on each side!
4
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Temperature Effects
Temperature changes are where things get interesting because they actually change which direction is favoured. Exothermic reactions release heat, whilst endothermic reactions absorb it.
When you increase temperature, the equilibrium shifts to favour the endothermic direction (the one that absorbs heat). It's like the reaction is trying to cool things down by absorbing your extra heat energy.
If you want more products from an exothermic forward reaction, decrease the temperature. The equilibrium will shift right to produce more heat and counteract the cooling. This is why many industrial processes use carefully controlled temperatures.
Key Insight: Temperature is the only factor that actually changes the equilibrium constant - everything else just shifts the position!
5
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The Equilibrium Constant (Kc)
The equilibrium constant (Kc) is chemistry's way of putting a number on how far a reaction proceeds. It's calculated using the concentrations of products and reactants at equilibrium, and it stays constant at a given temperature.
For any reaction aA + bB ⇌ cC + dD, the expression is: Kc = [C]^c[D]^d / [A]^a[B]^b. Products go on top, reactants on bottom, and the powers come from the balanced equation.
Homogeneous equilibria involve everything in the same phase (all gases or all in solution). These are the most straightforward to work with because you include all species in your Kc expression.
Top Tip: The equilibrium constant only changes with temperature - pressure and concentration changes don't affect its value!
6
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Kc Examples and Heterogeneous Systems
Let's see Kc in action with real industrial processes. For the contact process (2SO₂ + O₂ ⇌ 2SO₃), the expression becomes: Kc = [SO₃]² / [SO₂]²[O₂]. Notice how the powers match the coefficients in the balanced equation.
Heterogeneous equilibria involve different phases (like solids with gases). Here's the crucial rule: don't include solids in your Kc expression. For H₂O + C ⇌ H₂ + CO, you get Kc = [H₂][CO] / [H₂O].
The Haber process gives us Kc = [NH₃]² / [N₂][H₂]³. These industrial examples show how equilibrium principles apply to real-world chemistry.
Remember: Solids don't appear in Kc expressions because their concentrations don't change significantly!
7
of 10Sign up to see the content. It's free!
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Acid-Base Theory Essentials
The Brønsted-Lowry theory makes acid-base chemistry crystal clear: acids donate protons (H⁺), whilst bases accept them. This definition works for way more reactions than you might think.
Strong acids like HCl completely dissociate in water: HCl → H⁺ + Cl⁻. Weak acids like ethanoic acid only partially dissociate: CH₃COOH ⇌ H⁺ + CH₃COO⁻. Notice the equilibrium arrow for weak acids!
Don't confuse strong/weak with concentrated/dilute. Concentration tells you how much acid is dissolved, whilst strength tells you how completely it breaks apart. You can have a dilute solution of a strong acid!
Key Point: All neutralisation reactions are fundamentally the same - H⁺ + OH⁻ → H₂O - which is why they all release similar amounts of energy!
8
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The pH Scale and Calculations
The pH scale is a logarithmic way of expressing H⁺ concentration: pH = -log[H⁺]. Each pH unit represents a 10-fold change in acidity - pH 1 is ten times more acidic than pH 2.
Converting from H⁺ concentration to pH is straightforward: just stick the concentration into your calculator and use the -log function. For example, [H⁺] = 0.0100 mol dm⁻³ gives pH = 2.00.
Going the other way (pH to concentration) requires the inverse operation. For pH = 3.42, you calculate 10^ = 3.80 × 10⁻⁴ mol dm⁻³. Most scientific calculators have a "10^x" function for this.
Calculator Tip: Practice these conversions - they're fundamental for any quantitative acid-base work and frequently appear in exams!
9
of 10Sign up to see the content. It's free!
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By signing up you accept Terms of Service and Privacy Policy
More pH Conversions and Titration Intro
Getting comfortable with pH calculations takes practice, but the pattern is always the same. Whether you're going from concentration to pH (use -log) or pH to concentration (use 10^), the key is staying organised with your powers of ten.
Volumetric analysis (titrations) is chemistry's way of finding unknown concentrations. You know one concentration and measure two volumes, then use stoichiometry to find the unknown concentration.
The beauty of titrations lies in their simplicity - you're just using the fundamental relationship n = cV (moles = concentration × volume) combined with the balanced chemical equation to work out ratios.
Success Strategy: Always convert cm³ to dm³ by dividing by 1000 - this is the most common mistake in titration calculations!
10
of 10Sign up to see the content. It's free!
- Access to all documents
- Improve your grades
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By signing up you accept Terms of Service and Privacy Policy
Titration Calculations Made Simple
Here's the step-by-step approach that works every time: write the balanced equation, find moles in the known solution, use stoichiometry to find moles of the unknown, then calculate concentration.
For the example: HCl + NaOH → NaCl + H₂O. With 25.0 cm³ of 0.10 mol dm⁻³ NaOH neutralising 15.0 cm³ of HCl, first find moles of NaOH: n = 0.10 × 0.025 = 2.50 × 10⁻³ mol.
Since the reaction is 1:1, you've got the same moles of HCl. Finally, concentration = moles ÷ volume = 2.50 × 10⁻³ ÷ 0.015 = 0.167 mol dm⁻³. The systematic approach makes these calculations manageable every time.
Exam Success: Always show your working clearly - even if you make an arithmetic error, you'll get method marks for the correct approach!
We thought you’d never ask...
Our AI Companion is a student-focused AI tool that offers more than just answers. Built on millions of Knowunity resources, it provides relevant information, personalised study plans, quizzes, and content directly in the chat, adapting to your individual learning journey.
You can download the app from Google Play Store and Apple App Store.
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