When magnesium burns in the presence of oxygen the octet rule explains why magnesium & oxygen come toghether. Mg has 2 valence and O has 6….

When magnesium reacts with chlorine the octet rule also allows us to understand why the atoms come together. Determine the formula for Magnesium Chloride.

What did these have in common? In both of these reactions the Magnesium gave up electrons and its partner (the O or the Cl) gained electrons. When an atom gains or loses an electron we have names to describe it.

Reduction When something gains an electron it has been reduced. This makes sense because electrons are ________charged. So gaining a negative makes it reduced.

Oxidation When something loses an electron it is oxidation.

Which one of these gained and which on lost electrons. Which one was oxidized and which one was reduced

They always go together One atom cannot gain an electron (reduced) unless another one can lose an electron (oxidized) Since they always go together we call the reaction a Redox reaction (reduction & oxidation)

Oxidation numbers For some reactions it is easy to find the charges and see the oxidation & reduction. For other reactions it is not so easy. In those reactions we can use the oxidation numbers at the top right of the reference table.

Oxidation numbers Each element has different possible numbers of electrons it could gain or lose. If something gains 1 electrons (reduced) it will be -1, if it is reduced by 2 electrons it is -2 If something gained 1 or 2 electrons it will be +1 or + 2

Oxidation number rules These are 5 main rules of oxidation numbers.

Oxidation rules Uncombined elements have an oxidation number of zero (Na, Mg, Cl2 or O2) For group 1 & 2 elements when combined they always have a charge of +1 or +2 (look at the reference table) Oxygen is almost always -2 (except when with fluorine then it’s a +2 --the regents wont ask on this--)

Oxidation rules 4) Fluorine is always a -1 (the other halogens are also almost always -1). 5) Hydrogen is a -1 when with a metal but +1 when with a non metal 6) The sum of all oxidation numbers in a compound need to equal zero. (there will be an exception soon)

Oxidation These rules may sound complicated but once you get the hang of them it becomes very simple.

Examples What are the oxidation numbers of the atoms in HNO3 --------------------------

What are the oxidation numbers of the atoms in HNO3 O is (almost) always -2 (rule 3) H is with a non metal so it is -1 (rule 5) The N then must have an oxidation of ______ to make everything even.

Examples What is the oxidation number of platinum in K2PtCl6 What is the oxidation number of Sulfur in H2SO4

How to spot a redox reaction Not all reactions are redox reactions 2NaOH + H2SO4 Na2SO4 + 2H2O This is not a redox Mg +Br2  MgBr2 This is a redox The way to spot a redox is seeing if oxidation numbers have changed!

Example CuO + H2  Cu + H2O Find the oxidation numbers before and after.

Spot the redox CaCO3 + 2HCl  CaCl2 +H2O + CO2 Or Zn + 2AgNO3  2Ag + Zn(NO3)2

Hint for redox’s When ever you see a reaction that has an element by itself on one side… and bonded on the other….. That tells you the oxidation number MUST have changed. Meaning that reaction is always the redox.

Half reactions A chemical equation shows the formulas of the reactents & products, but doesn’t show how electrons are exchanged. A half reaction shows the electrons being gained or lost.

Reduction A reduction half reaction shows an atom gaining electrons…. & the oxidation charge goes ______ Fe3+ + 3e-  Fe

Oxidation In an Oxidation half reaction we show an atom losing electrons and the oxidation charge goes _______ Fe  Fe3+ + 3e- Notice how charge is conserved!!

Example Write the half reaction showing the reduction of Fe(II) Write the half reaction showing the oxidation of Mg

Example Write the oxidation half reaction for 1) K 2) Ca Write the reduction half reaction for 1) Ag+ 2) F

Which was oxidized & reduced Cu + 2AgNO3  Cu(NO3)2 + 2Ag Then write the half reactions

Electro chemical cells In a redox reaction there is an exchange of electrons… In electro chemical cells we can harness that electron exchange for our own purposes…

Battery

Electro chemical cells In an electro chemical cell there are two sides called electrodes The electrode that oxidation occurs at (loses electrons) in the ANODE The electrode that gains electrons is the CATHODE

Reduction = cathode Oxidation = Anode RED CAT AN OX

Voltaic cells A voltaic cell is where this electron exchange happens spontaneously (A battery)

Voltaic cells In this example the Zn Anode lost electrons, which traveled thru the wire to the copper cathode. There Cu2+ from the solution left the liquid and went into the cathode to accept the e-’s

Salt bridge There needs to be another connection called a salt bridge between the beakers so that the ions can stabilize (not as important for the regents)

How can you tell the anode & cathode?

Look at the overall reaction and determine which gained or lost electrons… The one that was reduced (lost) is the _______ The one that was oxidized (gained) is the _______

Describe the direction of electron flow in the external circuit in this operating cell. [1]

State the purpose of the salt bridge in this voltaic cell?

Allow 1 credit. Acceptable responses include, but are not limited to: The salt bridge allows for the migration of ions between the half-cells

A simpler way to find anode & cathode Look at table J on your reference. The metal that is higher on the activity chart will be oxidised (anode) The metal is where the reduction will occuer (cathode)

Electrolytic cells So far we have learnt about voltaic cells that are spontaneous, (happen without us forcing them to) The reverse can also happen, where the electrons are pushed in a non natural way.

When electricity is used to force a chemical reaction to occur it is called electrolysis. In this type of reaction the oxidized metal is lower down on the chart Electrolysis is used to plate a metal

The electrons leave the anode, and travel thru the wire. The metal ions leave the anode. And travel to the cathode (the spoon) to meet up with the electrons.