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10 Cards in this Set
- Front
- Back
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Oxidising agents |
1. Get reduced 2. Types: - KMnO4 - K2Cr2O7 - H2O2 - Cl2(Halogens tend to gain electrons due to high electronegativity) |
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KMnO4 |
1. Observation: Purple→colourless 2. Ionic equation: MnO4-(aq) + 8H+(aq) + 5e-→ Mn2+(aq) +4H2O(l) 3. Use: Test for reducing agents |
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K2Cr2O7 |
1. Observation: Orange→Green 2. Ionic equation: K2Cr2O7 2-(aq) + 14H+(aq) +6e-= 2Cr3+ +7H2O(l)3. Use: Test for reducing agents |
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H2O2 |
1. Observation: Orange→Green 2. Ionic equation: H2O2(l) + 2H(aq) + 2e-(aq) → 2H2O(l) 3. both oxidising and reducing |
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Chlorine |
1. Observation: Solution turns darker in color when reacted with bromide or iodide 2. Ionic equation: Cl2(g) + 2e-(aq) → 2Cl-(aq) 3. Use: Oxidises 3Br- to Br2 and 2I- to I2 |
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Reducing agents |
1. Get oxidised 2. Types: - Potassium Iodide(aq) - Reactive metals - Hydrogen |
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KI |
1. Observation: Solution turns brown 2. Ionic equation: 2I-(aq) → I2 + 2e- (reacts with Cl2) 3. Use: Test for oxidising agents |
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Reactive Metals, e.g. Al, Zn |
1. Observation: Less reactive metals are produced 2. Use: Displacement of less reactive metals |
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Hydrogen |
1. Observation (Reacts with CuO): Reddish-brown solid(Cu) is formed 2. Use: Reduces copper(II) oxide to copper |
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Disproportionation Reaction (spelling!) |
1. Examples - 2H2O2(l) → 2H2O(l) +O2(g) (H2O2, O) - Cu2O(s) + H2SO4(aq) → Cu(s) +CuSO4(aq) + H2O(l) (Cu2O, Cu) 2. Definition: A disproportionation reaction occurs when a substance is both oxidised and reduced simultaneously to form two different products. |
