1 / 33

Reduction-Oxidation Reactions

Reduction-Oxidation Reactions. ( Redox ). Redox -defined. Oxidation state (oxidation number) A system for tracking electrons in redox reactions An atom in a pure element has no charge and is assigned an oxidation state of zero. Redox -defined. REDuction-OXidation

sbeede
Download Presentation

Reduction-Oxidation Reactions

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  1. Reduction-Oxidation Reactions (Redox)

  2. Redox-defined Oxidation state (oxidation number) • A system for tracking electrons in redox reactions • An atom in a pure element has no charge and is assigned an oxidation state of zero

  3. Redox-defined • REDuction-OXidation • chemical reactions in which atoms have their oxidation number (oxidation state) changed. • This can be either a simple or a complex process

  4. Redox-defined • REDuction-OXidation • Oxidation is the loss of electrons or an increase in oxidation state • Reduction is the gain of electrons or a decrease in oxidation state • “OIL RIG” (oxidation is loss, reduction is gain)

  5. Redox-defined • Electrons will go to the molecule, atom, or ion with the greater electronegativity • The oxidant(oxidizing agent) removes electrons from another substance i.e. it oxidizes other substances, and is thus itself reduced • The reductant(reducing agent) transfers electrons to another substance i.e. it reduces others, and is thus itself oxidized

  6. Redox-defined

  7. Redox-defined

  8. Redox-examples

  9. Trends in oxidation state

  10. Oxidation-Reduction Between Nonmetals

  11. Oxidation-Reduction Between Nonmetals

  12. Rules for Assigning Oxidation Numbers • The oxidation number of a free element is always 0. • The atoms in He and N2, for example, have oxidation numbers of 0. • The oxidation number of a monatomic ion equals the charge of the ion. • For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3.

  13. Rules for Assigning Oxidation Numbers • The usual oxidation number of hydrogen is +1. • The oxidation number of hydrogen is -1 in compounds containing elements that are less electronegative than hydrogen • The oxidation number of oxygen in compounds is usually -2. • Exceptions include OF2, since F is more electronegative than O, and BaO2, due to the structure of the peroxide ion, which is [O-O]2-.

  14. Rules for Assigning Oxidation Numbers • The oxidation number of a Group IA element in a compound is +1. • The oxidation number of a Group IIA element in a compound is +2. • The oxidation number of a Group VIIA (17) element in a compound is -1, except when that element is combined with one having a higher electronegativity. • The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl.

  15. Rules for Assigning Oxidation Numbers • The sum of the oxidation numbers of all of the atoms in a neutral compound is 0. • The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion. • For example, the sum of the oxidation numbers for SO42- is -2.

  16. Balancing Oxidation-Reduction Reactions by the Half-reaction Method • Consider the reaction of aluminum with oxygen to produce aluminum oxide 4Al(s) + 3O2(g) 2Al2O3 • This reaction can be separated into a half-reaction for the substance oxidized 4Al(s) 4Al3+(s)+ 12e- • And a half-reaction for the substance reduced 3O2(g) + 12e- 6O2-(s)

  17. Balancing Oxidation-Reduction Reactions by the Half-reaction Method 4Al(s) + 3O2(g) 2Al2O3 4Al(s) 4Al3+(s)+ 12e- 3O2(g) + 12e- 6O2-(s) The number of electrons lost (oxidized) must equal the number of electrons gained (reduced)

  18. Electrochemistry • A galvanic cell, or voltaic cell, named after Luigi Galvani, or Alessandro Volta respectively, is an electrochemical cell that derives electrical energy from spontaneous redox reactions taking place within the cell. • It generally consists of two different metals connected by a salt bridge, or individual half-cells separated by a porous membrane.

  19. Electrochemistry

  20. Electrochemistry • A galvanic cell consists of two half-cells, • the electrode of one half-cell is composed of metal A, and the electrode of the other half-cell is composed of metal B; the redox reactions for the two separate half-cells are thus: • An+ + ne- A • Bm+ + me- B

  21. Electrochemistry • these two metals can react with each other: • the metal atoms of one half-cell are able to induce reduction of the metal cations of the other half-cell

  22. Electrochemistry • The standard electrical potential of a cell can be determined by use of a standard potential table for the two half cells. • identify the two metals reacting in the cell. Then look up the standard electrode potential, in volts, for each of the two half reactions.

  23. Electrochemistry

  24. Electrochemistry • the reaction of zinc metal and copper ion • To figure the potential or voltage for a redox reaction, add the voltages for the two half-reactions. One of the reactions will have to be “flipped” to make it an oxidation reaction. • Cu2+ + 2e- → Cu 0.34v • Zn2+ + 2e- → Zn -0.76v • Zn → Zn2+ + 2e- 0.76v • Zn + Cu2+ → Zn2+ + Cu 1.10v

  25. Electrochemistry • What is the potential difference (voltage) created in a redox reaction between zinc metal and iron (II) ions? • Zn2+ + 2e- → Zn -0.76v • Fe2+ + 2e- → Fe -0.45v • Zn + Fe2+ → Zn2+ + Fe 0.31v

  26. Electrochemistry • What is the potential difference (voltage) created in a redox reaction between liquid bromine and iodine ions • Br2 + 2I- → 2Br- + I2 0.53v

  27. https://www.youtube.com/watch?v=eJ1e2Lj_tsQ

  28. Relative Strength of Oxidizing and Reducing Agents Reducing Agents Oxidizing Agents S Li Li+ W T K K+ E R Ca Ca2+ A O Na Na+ K N Mg Mg2+ E G Al Al3+ R E Zn Zn2+ R Cr Cr3+ Fe Fe2+ Ni Ni2+ Sn Sn2+ Pb Pb2+ H2 H3O+ H2S S Cu Cu2+ I– I2 MnO42– MnO4– Fe2+ Fe3+ Hg Hg22+ Ag Ag+ NO2– NO3– S Br– Br2 T W Mn2+ MnO2 R E SO2 H2SO4 (conc.) O A Cr3+ Cr2O72– N K Cl– Cl2 G E Mn2+ MnO4– E R F– F2 R

  29. Redox-examples • H2 + F2 → 2 HF (synthesis) • Fe + CuSO4 → FeSO4 + Cu (replacement) • 4 Fe + 3 O2 → 2 Fe2O3 (synthesis, “rusting”) • CH4 + 2 O2 → CO2 + 2 H2O (replacement, “combustion”)

  30. Redox –applications • Cleaning products, fertilizer • Electrochemical cells (batteries) • Electroplating • Production of compact discs • Cellular respiration • Photosynthesis • etc…

More Related