Experiment III

Redox Reactions


In this experiment you will further explore a specific kind of chemical reactions, called Oxidation and Reduction reactions, or Redox. These reactions happen with a transfer of electrons from one atom to another, and they are responsible for a number of different things that happen in everyday real life. Rust on your car, the energy from batteries, the tarnishing of silver, and the reason the Statue of Liberty is the green color it is today instead of the shiny coppor color it used to be.


Important Terms for Before You Enter the Lab





Oxidizing Agent

Reducing Agent

Oxidizing Agent Strength

Reducing Agent Strength

Electron Transfer

Oxidation State

Half Reaction

Reverse Reaction







To understand redox chemisty, one very important characteristic of the elements is their electronegativity. Electronegativity is a numerical scale that relates to how much an element wants electrons. An element with a very high electronegativity value would want to GAIN electrons, while an element with a very low electronegativity value would want to LOSE electrons. The video below explains in further detail what electronegativity is.


Electronegativity Theory


The video explains that there is a general trend across the periodic table in relation to electronegativity. As you move to the right on the periodic table, the atomic number increases, meaning the number of protons increases. Protons are positively charged and with more of them, can attract more negatively charged electrons. Looking at the periodic table of electronegativity, you can see as you go across, the values increase.

However, when moving down the periodic table, the energy levels are increasing and it makes it harder for the protons to hold on to the higher energy electrons.


Watch the video below to see how this happens!


Electronegativity Trend on the Periodic Table 


Now that you've learned a little about electronegativity, try to answer some of the following questions


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Oxidation & Reduction


Oxidation = The losing of electrons

Reduction = The gaining of electrons


This also relates to electronegativity values, where elements with a high electronegativity (Want electrons) will undergo reduction very easily, and elements with low electronegativity (Do not want electrons) will undergo oxidation very easily


A redox (Reduction/Oxidation) reaction is the transfer of electrons from one atom to another.

Some atoms want to get rid of electrons, while other atoms want to take extra electrons, and this movement of electrons is a redox reaction.

In each redox reaction, there are two processes going on: Oxidation, and Reduction

The video below looks at oxidation and reduction across the periodic table.


Oxidation & Reduction on the Periodic Table


There are a few different fun ways to remember these defenitions





In both of these examples, the electrons are moving around, from one element to another, and since each electron has a charge of -1 the overall charge of an element changes during redox reactions.


Electron = -1 Charge

Oxidation = Loss of electron(s) = Loss of -1 for each electron, so... -(-1) = +1

During oxidation, the charge of the atom or ion INCREASES by the number of electrons lost


Reduction = Gain of electron(s) = Gain of -1 for each electron, so... +(-1) = -1

During reduction, the charge of the atom or ion DECREASES by the number of electrons gained


The following videos will help you understand exactly what is going on to an atom during oxidation and reduction.



Oxidation of Lithium



Reduction of Fluorine



 When you look at a reaction of only oxidation, or only reduction, you are looking at what is called a half reaction.


The half reaction for the oxidation of lithium would be:


Lithium lost an electron to become the Li+ ion


The half reaction for the reduction of fluorine would be:


The fluorine gas gained two electrons to become two fluoride ions


Try a few examples of half reactions and see if you can determine whether the reaction taking place is an oxidation reaction, or a reduction reaction


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Oxidation States


Each time an atom either gains or loses an electron, it has a different oxidation state. Some elements will only have a couple, such as the family I elements (Na(s), and Na+) Sodium has one electron ready to lose, so in addition to the neutral state Na(s) there is one additional oxidation state. Other metals, such as Manganese (Mn) can lose many electrons and therefore has many oxidation states (Mn(s), Mn+, Mn2+, Mn3+..... Mn7+)


Difference in Oxidation States



Oxidizing Agent and Reducing Agent


In an oxidation reaction, when the electron is removed from the atom, something must have caused this to happen. What causes the oxidation to occur is called the OXIDIZING AGENT. The Oxidizing agent is a chemical species that causes the electron to leave, and draws in the electron to itself.


In a reduction reaction, the electron being added to the atom needs to have come from somewhere, and that would be the REDUCING AGENT. The reducing agent is the chemical species that gives an electron to another atom.


When going through an oxidation/reduction equation, you need to determine what is the oxidizing agent, and what is the reducing agent. The video below shows how this is done.


What is the Difference Between Oxidizing & Reducing Agents? 




Now work through some equations and determine what are the oxidizing agents, and what are the reducing agents.



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 Critical Thinking Question:

You have a 1.0 Molar solution of Silver (I) Nitrate, and you toss a shiny new penny into the solution. The penny becomes a dull gray color with specs of shiny metalic color, while the solution transforms from a clear and colorless solution into a clear blue solution.

Did a reaction take place?

If so, what are the reacting species? Oxidizing Agent? Reducing Agent?

What are the products?

How could you confirm these things?



 Elements & Ions

When a metal has the ability to be oxidized, and has the ability to reduce something else (and is therefore a reducing agent), however when a metal such as sodium (Na) does this, the resulting cation now has the ability to gain an electron and be reduced, giving it the ability to oxidize something else (and therefore is an oxidizing agent) This inverse relationship is important with all of the metals (Reducing Agents) and their ions (Oxidizing Agents), as well as the non-metals (Oxidizing Agents) and their ions (Reducing Agents).


The video below explains in further detail, the differences between the elements and their ions, and which are reducing agents, and which are oxidizing agents.


Elements vs. Ions


Balancing Redox Reactions

When looking at a reduction/oxidation (redox) reaction, there will always be one compound undergoing oxidation, and one compound undergoing reduction. This is a transfer of electrons, where the electron(s) leave the compound undergoing oxidation and the compound undergoing reuction gains the electron(s). The half reactions are very useful for balancing the equation.

Let's work through an example


You can first separate the equation into the ionic equation

Now identify the spectators, and the Oxidizing Agent and Reducing Agent


NO3- is always identified as a spectator, and therefore will not participate in the reaction

I2 has the ability to gain electrons and become I- so therefore is an Oxidizing Agent

Cu+ has the ability to lose electrons and become Cu2+ so therefore is a Reducing Agent


Now you can look at the half reactions

Reduction of Cu+


Oxidation of I2


Start by balancing the electrons. The oxidation reaction requires 2 electrons, but the oxidation reaction only generates one electrons. This means that in order to generate enough electrons to oxidize one molecule of I2, you must reduce two ions of Cu+ to Cu2+


After which, you can add the reactions together combining everything on the right side of the reaction and everything on the left side of the reaction

And now the electrons can cancel out on each side to give the fully balanced and correct net ionic redox reaction


In this way, both the elements are balanced as well as the overall charge on each side of the equation. Try balancing the following examples.




Metals & Metal Ions

Metals themselves are solid materials, some hard, other soft that will act as reducing agents. The metals have unfilled valance electron shells, and and relatively low electronegativity. This means that the metal atoms would want to lose their outer electrons and become a positively charged metal ion.

The metal ions are the oxidized form of the metal atoms and act as oxidizing agents They have already lost electrons and therefore have a positive charge, and still have a relatively low electronegativity. This means that the metal atoms have the ability to gain the electrons back and become the metal atom again.


If you look at the periodic table, you can notice the trend of electronegativity, and can then see the reactivity of the metals. Take a look at the videos below explaining the appearance and reactivity of the alkali metals in family 1.


Appearance of Alkali Metals




Reactivity of Alkali Metals





Questions to ponder.....

If a particular metal is a very strong reducing agent (very reactive), would its accompanying ion be a strong or weak oxidzing agent?

Will a metal with a low electronegativity (such as Na) be a better or worse reducing agent than a metal with a high electronegativity (such as Al)?

Will a metal ion with a low electronegativity (such as Cs+) be a better or worse oxidizing agent than a metal ion with a high electronegativity (such as Ga3+)?



 Try some questions below to see if you can figure out whether the following pairs are oxidizing agents or reducing agents, and predict which one is the stronger based on their position in the periodic table.

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Pre Lab Experiment

Watch the videos below and see if you can guess the relative reactivity of the following metals before the answer comes up. Potassium (K), Calcium (Ca) Magnesium (Mg), and Sodium (Na).  


Theoretical Metal Reaction Strength


Reaction of the Metals & Formation of Hydroxide



After viewing the above videos, what is the reactivity order of Potassium (K), Calcium (Ca) Magnesium (Mg), and Sodium (Na)?


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Analyzing a Redox Reaction of Metals and Metal Ions


In the Lab!!


In the lab you will be reacting metals (reducing agents) with metal ions (oxidizing agents). From both your observations in seeing which combinations form a spontaneous reaction, as well as using the information about the metals and their electronegativity and placement on the pereodic table, you will be able to determine the relative strerengths of the reducing agents (metals) and oxidizing agents (metal ions).


Preparing a Redox Reaction Assay



 Determining a Spontaneous Reaction


What are the ways you know a reaction has taken place?

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Analysis of a Spontaneous Reaction





You can analyze and predict the reactions of metals and metal ions based on other redox reactions. In the lab you will conduct redox experiments and determine which pairs of oxidizing agents and reducing agents will react in a given set.



Important Note: When preforming these redox reactions please allow ~30min to pass to allow slow reactions to take place. Some reactions may be subtle, such as bubbles forming on the metal surface


The picture above shows how the analysis is set up, where there are mixtures of oxidizing agents and reducing agents in each well of a well plate.


Based on the appearance, you can visually see if a reaction has occured between the two species. You can see that in the well with Zn2+ and Mg, there is a reaction because there are multiple bubbles that formed. In the well with Cu and Zn2+ there is a reaction because there was a color change. Answer whether there was a reaction or not in the following combinations. Keep in mind that Na+ is a spectator ion. This is what the metals should look like. Also, a combination of a metal and its metal ion is a reference as well, where there is no reaction.


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From this information you can determine the relative strengths of the oxidizing agents and reducing agents used in the experiment.


Based on the picture, please rank the oxidizing agents and reducing agents in order of how many reactions took place


Oxidizing Agents (Mg2+, Cu2+, Zn2+, Ag+)

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Reducing Agents (Mg, Cu, Zn)

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 These rankings are the relative strengths of the oxidizing agents and reducing agents that were used in the experiment!


Oxidizing Agent Strength


Reducing Agent Strength


Question to consider.....

- In the above picture, you see that there is a reaction between the silver ion (Ag+) and the zinc metal (Zn). Do you think a reaction will occur between the silver metal (Ag) and zinc ion (Zn2+)?

- What are some ways you can tell if a reaction has occured?



Additional redox chemistry can happen without metals, and can occur greatly with the halogens and halides as you will see next!