Introduction to Acids and Bases
Welcome! In this chapter, we are going to explore the world of acids and bases. You have probably encountered these in everyday life—like the citric acid in lemons or the alkaline nature of soap. In Chemistry 9701, we use the Brønsted–Lowry theory to understand exactly how these substances behave. By the end of these notes, you’ll be able to identify them, predict how they react, and even draw the graphs that describe their "personality" during a reaction!
1. Common Acids and Alkalis
Before we dive into the theory, you need to be on a first-name basis with the most common characters in this chapter. You should memorize these formulas:
Common Acids
• Hydrochloric acid: \( HCl \)
• Sulfuric acid: \( H_{2}SO_{4} \)
• Nitric acid: \( HNO_{3} \)
• Ethanoic acid: \( CH_{3}COOH \) (This is the acid found in vinegar!)
Common Alkalis (Bases)
• Sodium hydroxide: \( NaOH \)
• Potassium hydroxide: \( KOH \)
• Ammonia: \( NH_{3} \)
Quick Review: Acids usually have an \( H \) at the start (or end, in the case of organic acids like ethanoic acid), and bases often have an \( OH \) (hydroxide) or consist of Ammonia.
2. The Brønsted–Lowry Theory
This theory is actually very simple if you think of it as a game of "catch" using a proton (which is just a hydrogen ion, \( H^{+} \)).
• Brønsted–Lowry Acid: A species that donates a proton (\( H^{+} \)).
• Brønsted–Lowry Base: A species that accepts a proton (\( H^{+} \)).
Analogy: Imagine a game of soccer. The Acid is the player who kicks the ball (the proton), and the Base is the player who receives it.
Did you know? A Hydrogen atom is just one proton and one electron. If you remove the electron to make \( H^{+} \), all that is left is a single proton. That’s why we use the terms "proton" and "\( H^{+} \) ion" interchangeably!
Key Takeaway: Acids give away \( H^{+} \); Bases take in \( H^{+} \).
3. Strong vs. Weak: The Concept of Dissociation
Students often get "strong" confused with "concentrated." In Chemistry, strength refers to how easily the substance splits apart (dissociates) in water.
Strong Acids and Bases
These are the "all-in" players. They fully dissociate (split up) into ions when dissolved in aqueous solution.
Example: \( HCl(aq) \rightarrow H^{+}(aq) + Cl^{-}(aq) \)
In this solution, there are almost no \( HCl \) molecules left; they have all turned into ions.
Weak Acids and Bases
These are more "hesitant." They only partially dissociate in aqueous solution. Most of the molecules stay stuck together.
Example: \( CH_{3}COOH(aq) \rightleftharpoons CH_{3}COO^{-}(aq) + H^{+}(aq) \)
Notice the reversible reaction arrow! This shows that the reaction happens in both directions, and a state of equilibrium is reached where only a small percentage of ions are formed.
Memory Aid: Strong = Single arrow (complete). Weak = Wobbly arrow (reversible).
Key Takeaway: Strong = 100% split into ions. Weak = only a tiny bit split into ions.
4. Comparing Behavior: How can we tell them apart?
If you have a bottle of a strong acid and a bottle of a weak acid of the same concentration, how can you tell which is which? Don't worry if this seems tricky; just look at the ions!
• pH Value: A strong acid has a higher concentration of \( H^{+} \) ions, so it will have a lower pH than a weak acid.
• Electrical Conductivity: Since strong acids have more ions, they conduct electricity better than weak acids.
• Reaction with Reactive Metals: Both will react with a metal (like Magnesium) to produce Hydrogen gas. However, the strong acid will fizz much more vigorously because it has more \( H^{+} \) ions ready to react immediately.
5. The pH Scale and Water
The pH scale measures how acidic or alkaline a solution is.
• Acidic solutions: pH is below 7.
• Neutral solutions (Pure Water): pH is exactly 7.
• Alkaline solutions: pH is above 7.
Common Mistake: Thinking pH 1 is "weak" because the number is small. Remember: Small number = Strong Acid!
6. Neutralisation and Salt Formation
When an acid meets a base, they cancel each other out. This is called neutralisation.
The Ionic Equation
For any neutralisation between an aqueous acid and an alkali, the "real" action is always the same:
\( H^{+}(aq) + OH^{-}(aq) \rightarrow H_{2}O(l) \)
What is a Salt?
A salt is formed during neutralisation when the \( H^{+} \) ion of an acid is replaced by a metal ion (or an ammonium ion).
Example: \( HCl + NaOH \rightarrow NaCl + H_{2}O \)
Here, Sodium Chloride (\( NaCl \)) is the salt.
7. pH Titration Curves
A titration curve is a graph showing how the pH changes as you add an acid to a base (or vice versa). There are four main shapes you need to recognize:
1. Strong Acid + Strong Base: Starts at pH 1, ends at pH 13. Has a very long, vertical section at pH 7.
2. Strong Acid + Weak Base: Starts at pH 1, ends around pH 9. The vertical section is in the acidic region (below 7).
3. Weak Acid + Strong Base: Starts at pH 3-4, ends at pH 13. The vertical section is in the alkaline region (above 7).
4. Weak Acid + Weak Base: No clear vertical section. This is difficult to perform a titration for!
Step-by-Step for sketching:
1. Identify the starting pH (Acid in flask? Start low. Base in flask? Start high).
2. Identify the finishing pH (based on the strength of what you are adding).
3. Draw the "equivalence point" (the vertical bit) where the two neutralize each other.
8. Indicators
An indicator is a substance that changes color at a specific pH. To choose the right one for a titration, the indicator must change color within the vertical section of your titration curve.
• Methyl Orange: Changes color at low pH (Acidic). Great for Strong Acid + Weak Base titrations.
• Phenolphthalein: Changes color at high pH (Alkaline). Great for Weak Acid + Strong Base titrations.
Quick Review Box:
• Acid: Proton donor.
• Base: Proton acceptor.
• Strong: Fully dissociated.
• Neutralisation: \( H^{+} + OH^{-} \rightarrow H_{2}O \).
• Indicator: Must match the vertical range of the graph.