May 21 - Chapter 17 textbook OXIDATION-REDUCTION. Objective :To determine O.N. for atoms in elements and compounds. HW : Complete worksheet. STUDY PAGE 606-607 from textbook. DO NOW OBSERVATION SKILLS!.
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Objective :To determine O.N. for atoms in elements and compounds.
HW : Complete worksheet. STUDY PAGE 606-607 from textbook
Whenever an atom loses an electron another atom has to gain one. Both reactions are simultaneous.
Mg + O2 MgO
Magnesium lost 2 electrons because oxygen took them. The metal LOST electrons, the NON METAL GAINED electrons
Chemist use the O.N. to determine how many electrons are either gained or lost by an atom or ion in a chemical reaction.
O.N. is the charge or partial charge of an atom in a compound or an ion.
1. For all uncombined elements O.N. = O (FREE ELEMENTS)
2. For monoatomic ions the charge equals O.N.
3. Metals of group 1 in compounds O.N.= +1.
Metals of group 2 in compounds O.N.= +2
4. Fluorine in compounds is always – 1.
Other halogens -1 in binary compounds with metals.
5. Hydrogen + 1 except in metal hydrides (combined with metals of group 1 or 2)
6. Oxygen is -2 except when combined with F (is =2) or in peroxides (-1).
7. THE SUM OF THE OXIDATION NUMBERS IN ALL COMPOUNDS MUST BE ZERO
8. FOR POLYATOMIC IONS THE SUM OF THE O.N. IS EQUAL TO THE CHARGE OF THE ION
Objective: How to keep track of electron transfers in chemical reactions?
HW: finish worksheet and read page 604 to 605.
Answer question 1 from page 611
The specie that gets oxidized loses electrons and its oxidation number increases.
The specie that gets reduced gains electrons and its oxidation number decreases.
A redox reaction can always be broken down as 2 half reactions that show the atom or ion that is being oxidized and the one that is being reduced.
MASS AND CHARGE has to be conserved in a half reaction
1. find the o.n. of each element in the reaction. Determine which is being reduced an which is being oxidized
2. Balance the masses first
3. Complete each half reaction with electrons.( LEO GER)
4. Verify that masses and charges are balanced.
R.A. gets oxidized.
Its O.N. increases
Active metals are good RA
Look for changes in the oxidation number or the atoms. If one element changed the O.N. then for SURE is redox reaction.
All single replacement, synthesis, decomposition and combustion reactions are REDOX.
Double replacement reactions are not redox.
Objetive: Spontaneous Redox Reaction
To use table J to predict if a single replacement reaction will occurr.
Spontaneous Reactions : happen without external help.
CuSO4 + Zn Zn SO4 + Cu
In a single replacement reaction the most active element replaces the other element from a compound. (TABLE J)
The metal above gets oxidized the one below will get reduced.
For non metals the one above gets reduced the one below gets oxidized.
F2 + NaCl
I2 + NaCl
Cl2 + Na I
The higher the metal is in table J, the most active it is, the more tendency to became oxidized (lose electrons)
On Top of table J best reducing agents
Towards the bottom metals tend to gain electrons then they became reduced and are good reducing agents
F2 has the greatest tendency to gain electrons ( became reduced) is the BEST OXIDIZING AGENT.
Ag (NO3) + Cu
Zn+2 + Co
MgCl2 + Ni
K + FeCl3
Li + Mg 2+
Ag (NO3) + Cu
Zn+2 + Co
MgCl2 + Ni
K + FeCl3
Li + Mg 2+
What SPONTANEOUS reaction would occur if we have
Cu, Cu2+, Zn and Zn2+ together.
Hint USE TABLE J
Where the electrons flow? Which loses which gains?
In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released.
If we can place the two metals in two different containers and connect them with a wire the electrons will flow from the Zn to the Cu and we will have an electric current – ELECTRICITY
But it does not work if we do not close the circuit – USE A SALT BRIDGE
*Anode: where the oxidation occurs.
*Cathode: where the reduction occurs.
Voltaic or galvanic cell: produce ELECTRICITY from an spontaneous chemical reaction.
ELECTRON FLOW. FROM THE ONE THAT GETS OXIDIZED TO THE ONE THAT GETS REDUCED.
ANODE – NEGATIVE source of electrons
A typical voltaic cell
REDuction at the cathode
Cathode: the electrode at which reduction occurs
Anode: the electrode at which oxidation occurs
Salt bridge: a tube containing strong electrolyte, a pathway to allow the ions to move from one side to another. PERMIT THE MIGRATION OF IONS
1. A cell uses the reaction Mn + Ni2+ Ni + Mn2+ to produce electricity.
Write the half-reaction that occurs at the anode.
b) Write the half-reaction that occurs at the cathode.
c) Which species in this cell loses electrons?
d) As the cell produces electricity, which ion increases in concentration?
Mn Mn2+ + 2e-
Ni2+ + 2e- Ni
1. Voltaic Cell
A type of electrochemical cell that converts chemical energy to electrical energy by a spontaneous redox reaction.
In 1800, Volta built the voltaic pile and discovered the first practical method of generating electricity. Constructed of alternating discs of zinc and copper, with pieces of cardboard soaked in salt water b/w the metals, the voltaic pile produced electrical current.
An apparatus that uses a redox reaction to produce electrical energy (voltaic cell) or uses electrical energy to cause a chemical reaction (electrolytic cell).
Uses electricity to force a reaction that is not spontaneous to occur.
NOTICE THAT IN ELECTROLYTIC CELLS THERE IS A BATTERY OR A POWER SOURCE PRESENT
Electrolysis: to decompose a substance using electricity
Is used to cover an object with metal using electricity.
THE OBJECT TO BE COVERED GOES IN THE CATHODE
POLARITIES : ANODE – POSITIVE
CATHODE- NEGATIVE SOURCE OF ELECTRONS – CONNECTED TO NEGATIVE