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Gen Chem (12)
| Question | Answer |
|---|---|
| What is oxidation? What is reduction? | Oxidation is the LOSS of electrons. Reduction is the GAIN of electrons. OIL RIG |
| What is a reducing agent? What is an oxidizing agent? | A reducing agent reduces the other species — it gets oxidized itself (loses electrons). An oxidizing agent oxidizes the other species — it gets reduced itself (gains electrons). |
| Name common oxidizing agents and common reducing agents. | Oxidizing agents: O₂, O₃, H₂SO₄, halogens. Reducing agents: H₂, Fe, Zn, alkali metals. |
| What is oxidation state/number? | A positive or negative number assigned to an atom representing its degree of oxidation or reduction. Changes in oxidation number represent electrons gained or lost when a bond is broken or formed. |
| What are the 5 rules for determining oxidation number? | Elemental state = 0 Monatomic ions = their charge F = always −1; O = −2 (−1 in peroxides, +1 bonded to F); H = usually +1 (except −1 in metal hydrides) Alkali metals = +1; alkaline earth metals = +2 Sum of all oxidation numbers = overall charge |
| What are half-reactions? | The individual oxidation and reduction components of a redox reaction. They must be coupled together for a redox reaction to occur. When added together, electrons cancel out to give the overall balanced reaction. |
| What are the 7 steps to balance a redox reaction in ACIDIC conditions? | Write the half-reactions Balance all atoms except O and H Balance O by adding H₂O Balance H by adding H⁺ Balance charges by adding e⁻ Multiply equations so the number of e⁻ is equal Add half-reactions; cancel e- and common terms |
| How do you balance a redox reaction in BASIC conditions differently from acidic? | Follow the same 7 steps as acidic, but after adding H⁺ (step 4): for every H⁺ added, add the same number of OH⁻ to BOTH sides. On the side where H⁺ was added, combine H⁺ + OH⁻ to form H₂O. |
| What are electrochemical cells? What do anode and cathode do? | Electrochemical cells (batteries) convert between chemical and electrical energy via redox reactions. Oxidation always occurs at the anode; reduction always occurs at the cathode. Electrons always flow from anode to cathode. |
| What is a galvanic (voltaic) cell? | A spontaneous electrochemical cell (ΔG < 0, E°cell > 0). It converts chemical energy into electrical energy. Does not require an external power source. Anode is negative (−), cathode is positive (+). Electrons flow anode → cathode. |
| What is an electrolytic cell? | A non-spontaneous electrochemical cell (ΔG > 0, E°cell < 0). It converts electrical energy into chemical energy. Requires an external power source. Anode is positive (+), cathode is negative (−). Electrons are forced to flow non-spontaneously. |
| How do galvanic and electrolytic cells compare? | Both have ox at the anode and red at the cat. In galvanic cells, anode is (−) and cathode is (+); in electrolytic cells, anode is (+) and cathode is (−). Galvanic is spontaneous; electrolytic is non-spontaneous. |
| What is a salt bridge and what does it do? | A tube filled with nonreactive electrolytes (NaNO₃ or KNO₃). It balances charge between the two half-cells by allowing ions to flow, maintaining electrical neutrality. Anions flow toward the anode; cations flow toward the cathode. |
| What is molten electrolysis? | A type of electrolytic setup where an ionic compound is melted at high temperature to form a conductive liquid electrolyte. Anions flow to the anode and are oxidized; cations flow to the cathode and are reduced. |
| What is electroplating? | An electrolytic process that deposits a metal onto a conductive object using electrical current. The object being plated is the cathode (−); the metal source is the anode (+). |
| What is electrochemical cell notation? | Shorthand notation: anode is always on the LEFT, cathode on the RIGHT. A single vertical line ( | ) separates phase changes; a double vertical line ( ‖ ) separates the anode from the cathode. Example: Cu(s) | Cu²⁺(aq) ‖ Zn²⁺(aq) | Zn(s) |
| What is standard cell potential (E°cell) and how is it calculated? | E°cell is the sum of the oxidation and reduction potentials, measured in volts. Formula: E°cell = E°reduction + E°oxidation. Standard conditions: 298 K, 1 atm, 1.0 M concentration. |
| What does E°cell tell you about spontaneity? | E°cell > 0 → spontaneous reaction (ΔG < 0) → galvanic cell. E°cell < 0 → non-spontaneous reaction (ΔG > 0) → electrolytic cell. |
| What are the 4 steps to calculate standard cell potential? | Write the ox and red half-reactions Look up E°reduction values Balance half-reactions; reverse the sign of the oxidation half-reaction (E°oxidation = −E°reduction). Add the potentials: E°cell = E°reduction + E°oxidation |
| What is the relationship between standard cell potential and Gibbs free energy? | ΔG° = −nFE°cell, where n = moles of electrons transferred, F = Faraday's constant (96,500 C/mol e⁻), E°cell = standard cell potential. |
| What is Faraday's constant and how is it used in electrolysis calculations? | Faraday's constant (F) = 96,500 coulombs per mole of electrons. Formula for moles of product deposited by electroplating: (I × tₛ) / (n × F) = moles of product, where I = current (amps), tₛ = time (seconds), n = moles of electrons transferred. |
| What is the relationship between reduction potential and oxidizing/reducing strength? | More positive reduction potential → stronger oxidizing agent (wants to gain electrons). More negative reduction potential → stronger reducing agent (wants to lose electrons). |
Created by:
smurtab