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19 Cards in this Set
- Front
- Back
Oxidation
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Loss of electrons
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Reduction
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Gain of electrons
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Oxidizing Agent
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Causes another atom to undergo oxidation and is itself reduced (gains electrons)
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Reducing Agent
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Causes another atom to undergo reduction and is itself oxidized (losses electron)
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Common Oxidizing Agents
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Contain oxygen or a similarly electronegative element
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Common Reducing Agents
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Contain metal ions or hydrides (H-)
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Oxidation Number- Free Element
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Number of zero
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Oxidation Number- Monatomic Ion
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Equal to the charge on the ion
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Oxidation Number- Metals in Compounds
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Group IA: +1, Group IIA: +2
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Oxidation Number- Group VIIA in Compounds
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Number of -1 (unless combined with an element of higher electronegativity)
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Oxidation Number- Hydrogen
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+1 unless paired with a less electronegative element in which case it is -1
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Oxidation Number- Oxygen
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+2, expect in peroxides (-1) or in compounds with more electronegative elements
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Half-Reaction Method
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1. Separate two half-reactions
2. Balance the atoms by adding H and O (acidic: water and H+, basic: water and OH-) 3. Balance the charges by adding e- 4. Multiply half reaction to even out number of electrons 5. Add half reactions and cancel out terms on both sides |
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Complete Ionic Equation
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Accounts for all of the ions present in a reaction. To write split all aqueous compounds into their relative ions
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Net Ionic Equations
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Ignore spectator ions (ions appearing on both sides of the reaction)
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Specific Net Ionic Equations
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Reaction with no aqueous salt: same as overall balanced reaction. Double displacement without solid salt formation: no net ionic equation
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Disproportionation Reactions
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Redox reaction in which one element is both oxidized and reduced forming at least two elements with different oxidation states
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Oxidation-Reduction Titration
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Titrations that follow the transfer of charge with indicators that change colors based on voltages of the solution
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Potentiometric Titration
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Redox titration where a voltmeter or external cell measures the electromotive force of a solution
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