Ch8: Galvanic Cells and Batteries

Overview: Batteries & the “Bigger Picture”

Batteries = devices that convert chemical energy $o$ electrical energy.
Direct illustration of the First Law of Thermodynamics: energy is neither created nor destroyed, only transformed. This means the total energy of the isolated system (the battery) remains constant, but its form changes.
In chemical vocabulary, “battery chemistry” $\approx$ “galvanic-cell chemistry” because most everyday batteries are galvanic (a.k.a. voltaic) cells. These cells harness the energy from spontaneous redox reactions.
Practical remarks
- Many individual galvanic cells can be wired together to form one large battery (e.g., a car’s lead–acid battery contains many heavy cells to supply sufficient power, typically $12V$ ).

Common Commercial Battery Types & Everyday Uses

Nickel–cadmium (NiCd)
- Toys, older portable electronics. Known for high discharge rates and ability to be recharged many times.
- Being phased out due to toxicity of cadmium and memory effect.
Nickel–metal hydride (NiMH)
- Replaces NiCd in many devices; safer & more energy-dense (higher capacity). Less prone to memory effect than NiCd.
Alkaline
- Flashlights, calculators, small appliances, clocks, etc. These are primary (non-rechargeable) batteries, known for good performance and low cost.
Primary lithium (non-rechargeable)
- LED lighting, smoke alarms; notable for long shelf life and high energy density due to lithium's low atomic weight.
Lead–acid
- Vehicle starter batteries; large, heavy, many cells in series (typically six $2V$ cells in a car battery). Known for their ability to deliver high surge currents.
Lithium-ion / lithium-polymer
- Laptops, tablets, phones (connected to Chap-1 discussion on portable electronics). High energy density, lightweight, and rechargeable, making them ideal for portable electronics.

Fundamental Chemistry of Galvanic (Voltaic) Cells

Galvanic reaction = a spontaneous redox reaction that simultaneously
- Generates electrical current (electrons in an external wire).
- Produces heat (exothermic component often present).
- Proceeds “rapidly/quickly,” historically tied to the word “galvanic.” The spontaneity is driven by a negative Gibbs Free Energy change (\Delta G < 0).
Essential requirement: physical separation of oxidation & reduction half-reactions so that electrons are forced to move through an external circuit, thus producing usable electrical work.

Redox Basics Refresher

Oxidation: increase in oxidation number; electrons are released (LEO: Lose Electrons, Oxidation).
Reduction: decrease in oxidation number; electrons are gained (GER: Gain Electrons, Reduction).
Must occur simultaneously (“red-ox”)—electron donor (reducing agent) and acceptor (oxidizing agent) paired.
Simple classroom demo: copper wire in blue $\text{AgNO}_3$ forms a silver coating—visual proof of a spontaneous redox event (electrons flow Cu $o$ Ag $^+$ ).

Anatomy of a Daniel (Zn/Cu) Cell – Canonical Example

Historical name: “Daniel” or “Daniell” (John Daniell, British chemist) cell.
Physical setup
- Left half-cell (anode)
- Zinc metal electrode immersed in $\text{ZnSO}_4(aq)$ . Zinc is more easily oxidized than copper.
- Oxidation: $\text{Zn}(s) \; o\; \text{Zn}^{2+}(aq) + 2e^-$ ( $\uparrow$ oxidation number $0 \; o\; +2$ ). Electrons are produced here.
- Electrode labelled anode (–): site of e⁻ generation in a galvanic cell, where the negative charge builds up due to electron release.
- Right half-cell (cathode)
- Copper strip in $\text{CuSO}_4(aq)$ . Copper ions are more easily reduced than zinc ions.
- Reduction: $\text{Cu}^{2+}(aq) + 2e^- \; o\; \text{Cu}(s)$ ( $\downarrow$ oxidation number $+2 \; o\; 0$ ). Electrons are consumed here.
- Electrode labelled cathode (+): site of e⁻ consumption in a galvanic cell, where a relatively positive potential attracts electrons.
- Electron pathway: anode $o$ external wire $o$ cathode ( $\text{e}^-$ flow). Electrons cannot flow through the solution.
- Overall balanced redox equation:
  $\text{Zn}(s) + \text{Cu}^{2+}(aq) \;\longrightarrow\; \text{Zn}^{2+}(aq) + \text{Cu}(s)$

Charge Balance & the Salt Bridge

Oxidation makes left beaker more positive (extra $\text{Zn}^{2+}$ ions accumulate in solution).
Reduction depletes $\text{Cu}^{2+}$ so right beaker becomes negative (excess $\text{SO}_4^{2-}$ remains in solution).
Salt bridge (often a U-tube of $\text{Na}<em>2\text{SO}</em>4(aq)$ gel or a porous disk): essential component that maintains electrical neutrality by allowing ion flow between the half-cells.
- Anions ( $\text{SO}_4^{2-}$ ) drift left (towards the anode) to neutralize $\text{Zn}^{2+}$ surplus.
- Cations ( $\text{Na}^+$ ) drift right (towards the cathode) to neutralize $\text{SO}_4^{2-}$ surplus (which is now in excess due to $\text{Cu}^{2+}$ depletion).
- Maintains electrical neutrality $o$ continuous electron flow; without it, the reaction would quickly stop due to charge buildup.
Student Q&A clarification:
- The bridge is pre-filled with salt solution; ions are already dissociated and simply migrate as the reaction proceeds, preventing charge imbalance.

Cell-Diagram (Line Notation) Conventions

Compact “shorthand” for an electrochemical cell; example for the Daniel cell: $\text{Zn}(s) \ | \ \text{Zn}^{2+}(aq) \ || \ \text{Cu}^{2+}(aq) \ | \ \text{Cu}(s)$
Left side = anode half-cell (oxidation); right side = cathode half-cell (reduction).
Single vertical line $|$ separates phases (e.g., solid electrode $\text{Zn}(s)$ from aqueous solution $\text{Zn}^{2+}(aq)$ ).
Double line $||$ denotes the salt bridge, indicating a separation between the two half-cells.
Reading the diagram instantly tells you:
- Which species undergo oxidation/reduction according to their position.
- Expected sign of each electrode (anode on left is negative, cathode on right is positive in a galvanic cell).

Key Terminology & Sign Conventions (must memorize)

Anode – where oxidation occurs; in galvanic cells it is negative (–) because it's the source of electrons.
Cathode – where reduction occurs; in galvanic cells it is positive (+) because it attracts electrons.
Electron flow: anode $o$ cathode (alphabetical order helps: A $o$ C).
Current direction (I): conventional current is opposite electron flow (cathode $o$ anode), defined as the direction of positive charge movement.

Sample Concept-Check (Class Activity) Statement set: identify the incorrect one

Salt bridge completes the circuit – True.
Electrons flow anode $o$ cathode – True.
Cathode is positive in a galvanic cell – True.
Oxidation occurs at the cathode – False oldsymbol{\to} oxidation occurs at the anode.
Correction #1: “Oxidation occurs at the anode.”
OR correction #2: keep location, change process – “Reduction occurs at the cathode.”

Broader Connections & Implications

Demonstrates how controlling redox separation lets chemists harness energy instead of releasing it solely as heat. This principle is fundamental to electrochemistry.
Same principles underlie corrosion, metal plating (electrolysis, reverse of galvanic cell), fuel cells (where reactants are continuously supplied), and biological electron-transport chains (e.g., in cellular respiration).
Ethical / environmental side:
- Choice of battery chemistry impacts toxicity (NiCd, Pb), recyclability, and resource scarcity (Li, Co). Responsible disposal and recycling are crucial.
- Engineering safer, longer-life chemistries (e.g., solid-state batteries, flow batteries) is a critical sustainability goal to reduce environmental impact and improve performance.

Formula & Numerical Reminders

Electron bookkeeping: half-reactions must be balanced for both charge and mass (multiply by coefficients if needed to equalize electron transfer).
Cell potential (not yet derived here) adds via $E<em>{cell}=E</em>{cathode}-E_{anode}$ once standard potentials are known.

Minimum must-know list for exams

Definitions of anode, cathode, oxidation, reduction.
Ability to write half-reactions & overall redox equation for a galvanic cell.
Role &