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! In your notebook record the experiment in words and then describe the experiment like a chemist (with a CHEMICAL REACTION! )

Indicate the metal and the non metal How metals react? How non metals react? Review : draw the dot diagram for each element, and for the compound formed How many electrons and protons in each reactant? How many electrons and protons in each element in the product?

REDOX REACTIONS (electron transfer reactions) 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

What happens to a car when the paint chips off?

LEO GER Losing Electrons is Oxidation. Gaining Reduction

OIL RIG Oxidation Is Loss. Reduction Gain.

REDOX REACTIONS REDuction – OXidation reactions Electrons are transferred from the element that is being oxidized to the one that is being reduced.

Oxidation Number (O.N.) 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.

RULES FOR ASSIGNING O.N. 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

Group work Get in your groups and practice what you just have learnt and complete handout. Finish the rest for homework.

Answer to finding the Oxidation state +7 +5 +4 +6 +2 +3 +5 +6

Review Find the O.N. for each element Cl2 K2S2O3 NO3- PO43- CaCr2O7 KNO2 OF2 H2O2 KH

May 22 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

DO NOW OBSERVATION SKILLS! In your notebook record the experiment in words and then describe the experiment like a chemist (with a CHEMICAL REACTION! )

PRACTICE HCl + Mg -> MgCl2 + H2 Indicate the oxidation state of each element in the reaction. READ HANDOUT “ANALYZING OXIDATION-REDUCTION REACTIONS”

VOCABULARY SIMULTANEOUS = at the same time OXIDATION = to lose electrons REDUCTION = to gain electrons OXIDATION NUMBER = charge or partial charge over an element.

To recognize redox reactions… 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.

PRACTICE – your turn! Get in your group and work with the reactions in the handout. Determine the O.N. for each element and decide which element got oxidized and which got reduced (TIP ALWAYS IN THE REACTANTS SIDE!!!)

May 24 How to recognize oxidizing and reducing agents? How to separate a redox reaction into half reactions?

AGENTS They are always found in the reactants side.

REDUCING AGENT When a substance is oxidized it LOSES electrons. Its O.N. increases. It is being oxidized and is making the other substance in the reaction reduced. Then the one that gets oxidized is the REDUCING AGENT. R.A. gets oxidized. Its O.N. increases Active metals are good RA

OXIDIZING AGENTS When a substance is reduced it GAINS electrons . Its O.N. decreases. It is being reduced, it takes electrons from the other specie, it makes the other specie to get oxidize. The one that gets reduced is the OXIDIZING AGENT Oxidizing Agent : Gets reduced Its ON decreases.

2Mg + O2  MgO Mg: oxidation number changes from 0 to +2 . It increased. Mg gets oxidized. Is the reducing agent. Oxygen: Oxidation number changes from 0 to -2 , it decreases is the Oxidizing agent.

Half reactions Hw p 164 q 32 to 36 And finish worksheet #3

Half reactions 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.

May 28 Objetive: Spontaneous Redox Reaction To use table J to predict if a single replacement reaction will occurr.

Activity series 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)

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 

METALS IN TABLE J 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

Non Metals F2 has the greatest tendency to gain electrons ( became reduced) is the BEST OXIDIZING AGENT.

Predict if the reaction will occur Ag (NO3) + Cu  Zn+2 + Co  MgCl2 + Ni  K + FeCl3  Li + Mg 2+

May 29 ELECTROCHEMISTRY OBJECTIVE: To distinguish between electrochemical and electrolytic cells To identify the different parts of a cell and explain their purpose.

Do now 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?

Voltaic Cells In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released.

IDEA! 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

Voltaic Cells A typical cell looks like this. The oxidation occurs at the anode. The reduction occurs at the cathode.

ELECTRODES Where the reduction or oxidation takes place. They are usually metals or they can be made of graphite. RED CAT REDuction at the cathode

ELECTROCHEMISTRY *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. POLARITIES ANODE – NEGATIVE source of electrons CATHODE- POSITIVE

REDuction at the cathode A typical voltaic cell RED CAT 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

Problems 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 Mn Mn2+

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.

Electrochemical cell An apparatus that uses a redox reaction to produce electrical energy (voltaic cell) or uses electrical energy to cause a chemical reaction (electrolytic cell).

Electrolysis: to decompose a substance using electricity 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

Electroplating 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