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Redox and Electrochemistry. Redox Reactions. Reduction – Oxidation reactions Involve the transfer of electrons from one substance to another. +. The oxidation numbers of the atoms will change…. one goes up (oxidation) and one goes down (reduction). Oxidation Number (Oxidation State). - PowerPoint PPT Presentation
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Redox and
Electrochemistry
Redox Reactions• Reduction – Oxidation reactions• Involve the transfer of electrons from one
substance to another
The oxidation numbers of the atoms will change…. one goes up (oxidation) and one goes down (reduction)
+
Oxidation Number (Oxidation State)• Used to keep track of the transfer of
electrons
• Number is assigned to every atom in a chemical formula, in accordance with certain rules
• NOT an ionic charge, but is often the same as the ionic charge– Possible oxidation states are given on the
periodic table (upper right hand corner)
Rules for assigning Oxidation Numbers
1. For a neutral compound, the sum of the oxidation states must be zero
2. The oxidation state of any atom in an uncombined element is zero
• Element not in chemical combination with another element
• Examples: Na, Mg, H2, Cl2
Rules for assigning Oxidation Numbers
3. The oxidation state of a monatomic ion is equal to its charge
– Examples: Na+ =
4. In an ionic salt, the oxidation number of each ion is equal to its charge
– Examples: CaCl25. For a polyatomic ion, the sum of the oxidation
states must equal the overall charge– Example: SO4
2-
Rules for assigning Oxidation Numbers
6. Metals of group 1 always have an oxidation number of +1
7. Metals of groups 2 always have an oxidation number of +2
8. Fluorine is always -1, other halogens are usually -1
9. Aluminum is always +3
Rules for assigning Oxidation Numbers
10. Oxygen is usually -2
Exceptions: – When paired with F (OF2), oxygen will be +2
– Peroxides (H2O2), oxygen will be -1
11. Hydrogen is usually +1
Exceptions: – Metal hydrides (Group 1 or 2 metals paired
with hydrogen), LiH, CaH2, hydrogen will be -1
Examples
Assign an oxidation state to each element in the following:
1. H2SO4
2. SO32-
3. K2CrO4
4. CrCl3
• Reduction– Reduction of charge by gaining electrons
Na+ + e- Na
O + 2e- → O2-
• Oxidation– Increase in charge by loss of electrons
Fe Fe3+ + 3e-
Cl- Cl + e-
LEO the lion says GER
Losing Electrons Oxidation
Gaining Electrons Reduction
Conservation of Matter/Conservation of Charge
• Mass must be conserved – Mass on both sides must be the same
(balanced)
• Charge must be conserved– Net charge on both sides must be the same
(balanced) – add electrons to the higher side
• Reduction and Oxidation reactions must occur together (REDOX reactions)
Half Reactions
• Every Redox reaction consists of a reduction and oxidation reaction
• Each reaction is called a ½ reaction
• A separate equation can be written for each ½ reaction
Half Reactions
• Net charge and mass must be the same on both sides of the equation
• The number of electrons must balance out, electrons do not appear in the net equation
• One ½ reaction is reduction and the other is oxidation
Spectator Ion
• Does not change oxidation states in the reaction, same oxidation state on both sides of the equation
• Not every species in an equation is oxidized or reduced, some are spectator ions
1. Assign oxidation states to each element in the reaction
2. Identify the 2 substances that are changing oxidation states
3. Write the half reactions• Balance the mass• Balance the charge (add electrons to the
higher side)
Examples
1. H2 + Cl2 2HCl
2. Fe + ZnO Zn + FeO
Reducing Agent
• Substance which is oxidized– Serves as a source of electrons to make
the reduction reaction occur
– Good reducing agents are substances that lose (donate) electrons easily – elements with low ionization energies
Examples: group 1 and 2 metals
Oxidizing Agent
• Substance which is reduced– Accepts (gains electrons)
– Good oxidizing agents are substances that gain electrons (highly electronegative elements)
Examples: Group 17 elements
Activity Series
Reference Table J
Metals• The most reactive metals are listed at the
top• A reaction will occur spontaneously if the
metal is higher than the metal ion that it is trying to replace
• Reactive metals lose electrons easily (low ionization energy)
• Higher on the table = More likely to be oxidized
Examples
Ba + ZnCl2 → Zn + BaCl2• Ba will replace Zn because Ba is
above Zn– Ba is more reactive than Zn
• More reactive means that it loses electrons easier
Nonmetals• For the halogen nonmetals listed in Table
J, the most reactive ones are at the top• For nonmetals, high reactivity means that
they are likely to gain electrons (high electronegativity)
• Higher on the table = More likely to be reduced
Example: F2 will replace any other halogen (it is the most
reactive)
Examples
1. Which metal is most reactive?
a. Fe b. Zn Cu
2. Will Ba react with Mn2+?
3. Will Na+ react with Cr?
4. Will this reaction occur spontaneously? Mg + Co(NO3)2 →
5. If this reaction does occur, what products would be made?
Balancing Equations
1. Assign oxidation numbers to all substances in the equation
2. Write the oxidation and reduction ½ reactions
3. Balance (cancel out) the electrons in the ½ reaction
4. Balance the rest of the equation
5. Check
Examples
1. Fe + Cl2 → FeCl3
2. Fe + Cu2+ Cu + Fe3+
3. KMnO4 + HCl → KCl + MnCl2 + H2O + Cl2