acids and bases

Acids exhibit different colors according to their conjugate base forms.
Example: Phenolphthalein indicator that changes color when dissociated.
- Phenolphthalein (H_pH)
- Conjugate base: pH⁻

Demonstrated using common laboratory acid and base solutions:
  - 1 M Hydrochloric Acid (HCl)
  - 1 M Acetic Acid (diluted with deionized water)
  - Ammonia
  - 1 M Sodium Hydroxide (NaOH)
Phenolphthalein is added to these solutions:
  - In acidic solutions, the indicator remains colorless (approx. pH 0 to 2).
  - In weak acid (acetic acid), the color is also weak.
  - Ammonia (a weak base) has a significantly higher pH (approx. pH 12) shown by a strong color change to pink.

Color change: Transitions from clear to pink at pH 8, useful for titration endpoints.

Indicator used to demonstrate different pH levels, transitions:
  - Strong acid: thin orangey-pink color (pH 0 to 2).
  - Neutral solution: turns yellow (pH 7).
  - Strong bases show distinct changes in color compared to strong acids.

Soap is made from fats through a reaction with a strong base which breaks fats down into basic forms of carboxylic acids.
Mechanism of Soap:
- Soap molecules have hydrophobic tails that surround fat molecules, allowing them to mix with water despite being immiscible.

Hard water contains high levels of calcium and magnesium ions, making it difficult for soap to dissolve.
- Calcium and magnesium ions create insoluble compounds with soap, requiring more soap for effective cleaning.
- Residues left in kettles and other appliances can form calcium carbonate and magnesium carbonate, leading to build-up or scaling.

Acid-base behavior involves the transfer of protons (H⁺).
Example: Ammonia acts as a base by pulling a proton from water to form ammonium hydroxide.

Next chapter focuses on how chemical reactions generate electrical current:
- Introduction to oxidation-reduction (redox) reactions.

Definition: A reaction involving the transfer of electrons between reacting species.
Oxidation: Loss of electrons (increases charge).
Reduction: Gain of electrons (decreases charge).
Both processes occur simultaneously.

Half Reactions:
  - Reduction half-reaction: Where electrons are accepted, decreasing charge.
    - Example: Two hydrogen ions (H⁺) reduce to form a hydrogen molecule (H₂).
  - Oxidation half-reaction: Where electrons are lost, increasing charge.
    - Example: Magnesium (Mg) loses electrons to become magnesium ions (Mg²⁺).

Atom and charge must balance in half-reactions:
- Example: In magnesium oxidation, the half-reaction:
$Mg ightarrow Mg^{2+} + 2e^-$
Reduction:
$2H^+ + 2e^- ightarrow H_2$
The total must balance in terms of both mass and charge.

Real-world example of reactions:
- Burning magnesium in oxygen results in magnesium oxide, demonstrating a redox reaction:
$Mg + O_2 ightarrow MgO$

To produce electricity, redox reactions must be separated in space:
- Use of a salt bridge to allow ionic balance while separating oxidation and reduction sites.
Example setup:
- Zinc immersed in zinc sulfate solution
- Copper immersed in copper nitrate solution
Electrons flow through an external circuit due to the oxidation of zinc and reduction of copper ions, allowing work to be done:
 - Oxidation:
 $Zn ightarrow Zn^{2+} + 2e^-$
 - Reduction:
 $Cu^{2+} + 2e^- ightarrow Cu$