Transcript Redox and Electrochemistry

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+ = +1
CaCl2
4. For a polyatomic ion, the sum of the
oxidation states must equal the overall
charge
– Example: SO42-
Rules for assigning Oxidation Numbers
5. Metals of group 1 always have an
oxidation number of +1
6. Metals of groups 2 always have an
oxidation number of +2
7. Fluorine is always -1, other halogens are
usually -1
8. Aluminum is always +3
Rules for assigning Oxidation Numbers
9. Oxygen is usually -2
Exceptions:
– When paired with F (OF2), oxygen will be +2
– Peroxides (H2O2), oxygen will be -1
10. 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. SO323. 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+ + 3eCl-  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
Examples
1. H2 + Cl2  2HCl
2. Fe + ZnO  Zn + FeO
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
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
Balancing Equations
1. Assign oxidation numbers to all
substances in the equation
2. Identify the substances that are being
oxidized and reduced (there will be 2
substances that have a change oxidation
numbers)
3. Write the oxidation and reduction ½
reactions
Balancing Equations
4. Balance (cancel out) the electrons in the
½ reaction
5. Balance the rest of the equation
6. Check
Examples
1. Fe + Cl2 → FeCl3
2. Fe + CuSO4 → FeSO4 + Cu
3. KMnO4 + HCl → KCl + MnCl2 + H2O + Cl2
Activity Series
Reference Table J
Metals
• The most reactive metals are listed at the
top
• An uncombined metal will replace (in a
chemical compound) any other metal ion
on the table that is below that uncombined
metal
• Reactive metals lose electrons easily (low
ionization energy)
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)
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?