AFNS Chemistry chapter Electrochemistry

Good morning, class. Please settle down and open your notebooks to Chapter 6: Solutions and Electrochemistry. For those of you preparing for the AFNS Initial Test, this chapter is vital as it bridges the gap between liquid chemistry and the electrical principles used in medical equipment and biological systems.

I. Solutions and Concentration Units

A solution is a homogeneous mixture where a solute is dissolved in a larger amount of solvent. While we often think of solids in liquids, such as saline (NaCl in water), they can exist in any state—all mixtures of gases, like air, are considered solutions.

Solubility Factors:

• “Like dissolves like” ⭐: Polar solvents dissolve polar solutes, and non-polar solvents dissolve non-polar ones.
• Temperature: Most solids become more soluble as temperature increases, but gases become LESS soluble. This is why oxygen levels in blood change with body temperature.
• Pressure: Affects only gases (Henry’s Law).

Concentration Units:

• Molarity (M): Moles of solute per litre of solution. This is the standard for lab work and medicine ⭐.
• Molality (m): Moles per kg of solvent, used specifically for colligative property changes.
• ppm (parts per million): Used for trace analysis and water quality.

II. Colligative Properties and Medical Significance

Colligative properties depend only on the number of solute particles, not their identity. This is a critical concept for understanding IV fluids.

1. Vapour Pressure Lowering: Adding a non-volatile solute decreases the vapour pressure.
2. Boiling Point Elevation: Solutions boil at higher temperatures than pure solvents (e.g., adding salt to cooking water or antifreeze in a radiator) ⭐.
3. Freezing Point Depression: Solutions freeze at lower temperatures (e.g., salt on icy roads or antifreeze in engines) ⭐.
4. Osmotic Pressure: The pressure required to stop osmosis. IV fluids must be isotonic to blood plasma (~308 mOsm/L) to prevent cell damage.

The van’t Hoff factor (i) accounts for electrolytes: non-electrolytes like glucose have i = 1, but NaCl dissociates into two ions, so i = 2. ⭐ REMEMBER FOR AFNS:

• Isotonic: 0.9% NaCl (normal saline) and 5% glucose (dextrose) ⭐.
• Hypotonic: Water enters cells, causing them to burst (haemolysis).
• Hypertonic: Water leaves cells, causing them to shrink (crenation).

III. Redox Reactions: OIL RIG

Redox reactions involve the transfer of electrons. We use the most important memory aid: OIL RIG (Oxidation Is Loss, Reduction Is Gain) ⭐. Alternatively, remember “LEO the lion says GER” (Loss of Electrons = Oxidation, Gain = Reduction).

• The species that loses electrons is oxidised and acts as the reducing agent ⭐.
• The species that gains electrons is reduced and acts as the oxidising agent ⭐.
• In the reaction Zn + CuSO₄ → ZnSO₄ + Cu, Zn is the reducing agent because it increases its oxidation state from 0 to +2.

IV. Electrochemical and Daniell Cells

We distinguish between two types of cells:

1. Galvanic (Voltaic): Converts chemical energy to electrical energy via a spontaneous reaction ⭐ (e.g., a battery).
2. Electrolytic: Uses electrical energy to drive a non-spontaneous reaction ⭐ (e.g., electroplating).

The Universal Rule:

• Anode = Oxidation (AN OX) ⭐.
• Cathode = Reduction (RED CAT) ⭐.
• In a Galvanic cell, the Anode is negative (−) and the Cathode is positive (+).

The Daniell Cell: A classic galvanic cell using Zinc and Copper. The Zinc electrode dissolves (oxidation at anode) while Copper deposits on the electrode (reduction at cathode). The Salt Bridge is essential to maintain electrical neutrality. The cell EMF is calculated as E°cathode − E°anode ⭐.

V. Electrolysis and Faraday’s Laws

Electrolysis involves the decomposition of a substance using electricity.

• Water: Produces H₂ at the cathode and O₂ at the anode in a 2:1 volume ratio ⭐.
• Brine (NaCl aq): This “Chlor-alkali process” produces H₂ gas, Cl₂ gas, and NaOH solution ⭐.

Faraday’s Laws ⭐:

• First Law: Mass deposited (m) is directly proportional to the charge passed (Q=It).
• Second Law: Masses are proportional to equivalent masses.
• One Faraday (F) is the charge of 1 mole of electrons, approximately 96,500 C ⭐.

VI. Electrode Potentials and Industrial Uses

The Standard Hydrogen Electrode (SHE) is our reference point, assigned a potential of 0.00 V ⭐.

• A more positive E° indicates a stronger oxidising agent (more easily reduced) ⭐.
• A more negative E° indicates a stronger reducing agent (more easily oxidised) ⭐.
• Lithium is the strongest reducing agent, while Fluorine is the strongest oxidising agent.

Industrial Applications:

• Electroplating: The object to be plated is always the cathode.
• Aluminium Extraction: The Hall-Héroult process uses molten cryolite to lower the melting point of Al₂O₃. Aluminium is the 3rd most abundant element ⭐.
• Fuel Cells: Used in space shuttles; they burn H₂ and O₂ to generate electricity and water directly.

Study these spontaneous reaction criteria and the isotonic values carefully, class. We will proceed to the practice lab in our next session. Class dismissed.

. Electrochemistry studies the relationship between:

1. A) Heat and pressure
2. B) Chemical reactions and electrical energy
3. C) Light and chemical reactions
4. D) Mass and volume

✔ Answer: B) Chemical reactions and electrical energy

2. Oxidation is defined as:
3. A) Gain of electrons
4. B) Gain of protons
5. C) Loss of electrons
6. D) Loss of neutrons

✔ Answer: C) Loss of electrons

3. Reduction is defined as:
4. A) Loss of electrons
5. B) Gain of electrons
6. C) Loss of protons
7. D) Gain of neutrons

✔ Answer: B) Gain of electrons

4. OIL RIG stands for:
5. A) Oxidation Involves Loss, Reduction Involves Gain (of electrons)
6. B) Oxygen In Liquid, Reduction In Gas
7. C) Oxidation Is Light, Reduction Is Gas
8. D) None of the above

✔ Answer: A) Oxidation Involves Loss, Reduction Involves Gain (of electrons)

5. In an electrochemical cell, oxidation occurs at the:
6. A) Cathode
7. B) Salt bridge
8. C) Anode
9. D) Electrolyte

✔ Answer: C) Anode

6. In an electrochemical cell, reduction occurs at the:
7. A) Anode
8. B) Cathode
9. C) Salt bridge
10. D) Electrolyte

✔ Answer: B) Cathode

7. The unit of conductance is:
8. A) Ohm (Ω)
9. B) Siemens (S) or mho
10. C) Volt (V)
11. D) Ampere (A)

✔ Answer: B) Siemens (S) or mho

8. An electrolytic cell converts:
9. A) Chemical energy to electrical energy
10. B) Electrical energy to chemical energy
11. C) Heat to electrical energy
12. D) Light to chemical energy

✔ Answer: B) Electrical energy to chemical energy

9. A galvanic (voltaic) cell converts:
10. A) Electrical energy to chemical energy
11. B) Heat to electrical energy
12. C) Chemical energy to electrical energy
13. D) Light to chemical energy

✔ Answer: C) Chemical energy to electrical energy

10. The standard hydrogen electrode (SHE) has an EMF of:
11. A) +1.0 V
12. B) −1.0 V
13. C) 0.00 V
14. D) +0.5 V

✔ Answer: C) 0.00 V

11. Faraday’s First Law states that the mass of substance deposited is proportional to:
12. A) Voltage applied
13. B) Quantity of electricity (charge) passed
14. C) Temperature
15. D) Concentration of electrolyte

✔ Answer: B) Quantity of electricity (charge) passed

12. One Faraday (F) equals:
13. A) 96,500 coulombs
14. B) 9,650 coulombs
15. C) 6.022 × 10²³ coulombs
16. D) 1 ampere

✔ Answer: A) 96,500 coulombs

13. Electrolysis of water produces:
14. A) H₂ at cathode and O₂ at anode
15. B) O₂ at cathode and H₂ at anode
16. C) H₂ at both electrodes
17. D) O₂ at both electrodes

✔ Answer: A) H₂ at cathode and O₂ at anode

14. Electroplating uses which type of cell?
15. A) Galvanic cell
16. B) Fuel cell
17. C) Electrolytic cell
18. D) Daniel cell

✔ Answer: C) Electrolytic cell

15. In the Daniell cell, the negative electrode is:
16. A) Copper
17. B) Zinc
18. C) Iron
19. D) Silver

✔ Answer: B) Zinc

16. EMF of a cell = E°(cathode) − E°(anode). A cell has E°(cathode) = +0.34V and E°(anode) = −0.76V. EMF =
17. A) 0.42 V
18. B) 1.10 V
19. C) −0.42 V
20. D) −1.10 V

✔ Answer: B) 1.10 V

17. The salt bridge in an electrochemical cell:
18. A) Provides electrons
19. B) Maintains electrical neutrality by allowing ion flow
20. C) Increases voltage
21. D) Acts as anode

✔ Answer: B) Maintains electrical neutrality by allowing ion flow

18. Corrosion of iron (rusting) is an example of:
19. A) Reduction only
20. B) Electrochemical oxidation
21. C) Condensation
22. D) Electrolysis

✔ Answer: B) Electrochemical oxidation

19. Which metal is used to prevent rusting of iron by cathodic protection?
20. A) Copper
21. B) Gold
22. C) Zinc
23. D) Lead

✔ Answer: C) Zinc

20. Resistance (R) and conductance (G) are related by:
21. A) G = R
22. B) G = 1/R
23. C) G = R²
24. D) G = R + 1

✔ Answer: B) G = 1/R

21. Strong electrolytes completely dissociate in water. An example is:
22. A) Acetic acid
23. B) Glucose
24. C) NaCl
25. D) Ethanol

✔ Answer: C) NaCl

22. Weak electrolytes partially dissociate in water. An example is:
23. A) HCl
24. B) NaOH
25. C) CH₃COOH (acetic acid)
26. D) NaCl

✔ Answer: C) CH₃COOH (acetic acid)

23. Non-electrolytes do NOT conduct electricity because they:
24. A) Have too many ions
25. B) Do not ionise in solution
26. C) Are solids
27. D) Are acidic

✔ Answer: B) Do not ionise in solution

24. The Nernst equation accounts for the effect of _____ on cell EMF:
25. A) Temperature and concentration
26. B) Mass only
27. C) Pressure alone
28. D) Surface area

✔ Answer: A) Temperature and concentration

25. During electrolysis of molten NaCl, Na is deposited at the:
26. A) Anode
27. B) Salt bridge
28. C) Cathode
29. D) Electrolyte

✔ Answer: C) Cathode