Oxidation reduction reactions
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Oxidation-Reduction Reactions - PowerPoint PPT Presentation

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Oxidation-Reduction Reactions. Oxidation and Reduction. Oxidation-reduction (redox) reactions involve transfer of electrons Oxidation – loss of electrons Reduction – gain of electrons Both half-reactions must happen at the same time

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Oxidation States electrons

  • Oxidation number assigned to element in molecule based on distribution of electrons in molecule

  • There are set rules for assigning oxidation numbers

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Oxidation electrons

  • Oxidation  reactions in which the atoms or ions of an element experience an increase in oxidation state

  • Ex. combustion of metallic sodium in atmosphere of chlorine gas

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Reduction to +1 (state of the ion)

  • Reduction  reactions in which the oxidation state of an element decreases

  • Ex. Chlorine in reaction with sodium

  • Each chlorine atom accepts e- and becomes chloride ion

  • Oxidation state decreases from 0 to -1

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Oxidation and Reduction as a Process to +1 (state of the ion)

  • Electrons are made in oxidation and acquired in reduction

  • For oxidation to happen during chemical reaction, reduction must happen as well

  • Number of electrons made in oxidation must equal number of electrons acquired in reduction

  • Conservation of mass

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Redox redox reaction Reactions and Covalent Bonds

  • Substances with covalent bonds also undergo redox reactions

  • Unlike ionic charge, oxidation number has no physical meaning

  • Oxidation number based on electronegativity relative to other atoms to which it is bonded in given molecule

  • NOT based on charge

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  • N hydrogen molecule is 0o electrons totally lost or gained

  • Hydrogen has donated a share of its bonding electron to chlorine

  • It has NOT completely transferred that electron

  • Assignment of oxidation numbers allows determination of partial transfer of e- in compounds that are not ionic

  • Increases/decreases in oxidation number can be seen in terms of completely OR partial loss or gain of e-

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  • Nitrate ion, NO monatomic ions and uncombined elements3-, is converted to nitrogen monoxide, NO

  • Nitrogen is reduced in this reaction

  • Instead of saying nitrogen atom is reduced, we say nitrate ion is reduced to nitrogen monoxide

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Balancing monatomic ions and uncombined elementsRedox Equations

Section 2

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Half-Reaction Method looking at them

  • Also called ion-electron method

  • Made of seven steps

  • Oxidation numbers assigned to all atoms and polyatomic ions to determine which species are part of redox process

  • Half-reactions balanced separately for mass and charge

  • Then added together

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  • To balance oxygen, H looking at them2O must be added to left side

  • This gives 10 extra hydrogen atoms on that side

  • So, 10 H atoms added to right side

  • In basic solution, OH- ions and water can be used to balance atoms

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  • H looking at them2O added to product side to balance oxygen atoms

  • 2 hydrogen ions added to reactant side to balance H atoms

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  • This ratio is already in lowest terms looking at them

  • If not, need to reduce

  • Multiply oxidation half-reaction by 1

  • Multiple reduction half-reaction by 8

  • Electrons lost = electrons gained

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Sample Problem looking at them

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Oxidizing and Reducing Agents form manganese (II) sulfate

Section 3

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  • Reducing agent form manganese (II) sulfate substance that has the potential to cause another substance to be reduced

  • They love electrons

  • Attain a positive oxidation state during redox reaction

  • Reducing agent is oxidized substance

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  • Oxidizing agent form manganese (II) sulfate substance that has the potential to cause another substance to be oxidized

  • Gain electrons

  • Attain a more negative oxidation state during redox reactions

  • Oxidizing agent is reduced substance

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Strength of Oxidizing and Reducing Agents form manganese (II) sulfate

  • Different substances compared and rated on relative potential as reducing/oxidizing agents

  • Ex. Activity series – related to each element’s tendency to lose electrons

  • Elements lose electrons to positively charged ions of any element below them in series

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  • Fluorine atom most highly EN atom electrons

  • Is also most active oxidizing agent

  • b/c of strong attraction for its own e-, fluoride ion is weakest reducing agent

  • Negative ion of strong oxidizing agent is weak reducing agent

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Autooxidation metals listed to displace other metals

  • Some substances can be both reduced and oxidized

  • Ex. Peroxide ions – O2-2 has relatively unstable covalent bond

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  • Hydrogen peroxide, H metals listed to displace other metals2O2, contains peroxide ion

  • Decomposes into water and oxygen as follows

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  • Hydrogen peroxide is both oxidized AND reduced metals listed to displace other metals

  • Oxygen atoms that become part of gaseous oxygen molecules are oxidized (-1  0)

  • Oxygen atoms that become part of water are reduced (-1  -2)

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  • Autooxidation  metals listed to displace other metalsa process in which a substance acts as both an oxidizing agent and a reducing agent

  • The substance is self-oxidizing and self-reducing

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Bombardier Beetle metals listed to displace other metals

  • Defends itself by spraying its enemies with an unpleasant hot chemical mixture

  • Catalyzed autooxidation of H2O2 produces hot oxygen gas

  • Gas gives insect ability to eject irritating chemical from abdomen