Balancing Oxidation-Reduction Reaction

There are ways in which we can balance chemical equation, simple equations can be balanced using inspection method .   Simple oxidation-reduction reaction can also be balanced by inspection method but complicated oxidation-reduction reaction cannot be balanced by this method.  In this post I will be discussing two methods of balancing Redox reaction, the oxidation number method and ion-electron method.

Balancing Redox Reaction using Oxidation Number Method

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In balancing redox reaction using oxidation method, there are guidelines to follow:
1.  Examine the reactants and products to determine if there is a change in the oxidation number by calculating the oxidation number or by inspection.  Use the technique used in my previous blog about how to determine the oxidation number.

a.  Write the oxidation number of the element above its symbol.  See to it that all elements in the reactant side and in the product side have their oxidation number.

b.  Diagram the number of electron lost by the oxidized element and the number of electron gained by the reduced element.

2.  Balance each element using a coefficient and remember that the electron lost and gain are equal.

a. Place a coefficient before the formula of the oxidized substance that corresponds to the number of electrons gained by the reduced substance.

b.  Place a coefficient before the formula of the reduced substance that corresponds to the number of electrons lost by the oxidized substance.

c.  Balance the remaining elements by inspection.

3.  After balancing the equation, check if the coefficient of the reactants and products are correct.

a. Place a check above the  symbol of the element to indicate that the coefficient are correct.  Meaning the number of atoms in the reactant side is equal to the number of atoms in the product side.

b.  For ionic equation, make sure also that the charges both in the reactant side and product side are equal.


Let us have an example:

1. Balance using oxidation number method the reaction between  Copper and Silver nitrate forming Copper (II) nitrate and solid silver, as shown in the unbalanced equation below:

Step 1.  Write the oxidation number of each element:

Step 2.  Diagram the number of electrons lost and gained.

Step 3.  Place a coefficient before the formula of the oxidized substance that corresponds to the number of electrons gained by the reduced substance.

Based from the figure above Cu loses 2e-, therefore to balance the electrons gain coefficient of two is placed in front of AgNO3  and Ag. Since silver gains only 1 e- therefore only 1 is the coefficient of Cu

Step 4.  Place a check above the  symbol of the element to indicate that the coefficient are correct.  Meaning the number of atoms in the reactant side is equal to the number of atoms in the product side.

Since number of atoms are balanced both in the reactant and product side, therefore the equation is already balanced.

Balancing Equation Using Ion-Electron Method in Acidic and Basic Solution

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Ion-electron method is also called half reaction method, used to balance ionic redox reactions in acidic and basic solution. A half reaction is  a part of redox reaction that shows the element that undergoes oxidation and reduction process. The ion-electron method balances the oxidation half-reaction with the reduction half-reaction separately.

 There are several steps to be followed in balancing ionic redox equation by ion-electron method.

1.  Write the half-reaction both the oxidation and reduction. Write the reactant and the product of the element that undergoes oxidation, Also write the reactant and product of the element that undergoes reduction.

2.  Balance the element in each  half-reaction by writing coefficient.

     a.  Balance all elements except oxygen and hydrogen.

     b.  Balance oxygen using H2O.

     c.  Balance hydrogen using H+.

Please take note:  For reaction in basic solution, add  OH- is to neutralize H+.

                              For example 2OH- neutralizes 2H+ to form 2H2O molecules.

     d,  Balance the charge by using e-.

3.  Multiply each half-reaction by a whole number to balance the number of electrons lost in oxidation process and the number of electrons gained in reduction process.

4.  Add the two half-reactions together and cancel similar ions, including electrons on each side of the equation.

5.  After balancing, check the number of atoms of all the elements if balanced.  Also calculate the net charged if balance both in the reactant side and in the product side.

Let us have an example:

Example 1.  Write a balanced equation on the reaction below in acidic solution:

Solution:

Step 1.   Write the half-reaction both the oxidation and reduction. 

Step 2.  Balance all the elements and the charge.

Step 3.  Multiply the two half reaction with a whole number to balance the number of e-.  Since the first half-reaction has only one e- and the second half reaction has 5e-  then the first half-reaction must be multiplied by 5.

Step 4.  Add the two half-reactions together and cancel 5e- both in the reactant side and product side,

Step 5.  Verify the equation if balance by checking the number of atoms and charges.

               

   

Example 2.  Write a balance equation of the equation below in basic solution.

Solution:

Step 1.  Write the half-reaction the oxidation and reduction.

Step 2.   Balance all the elements, balance the O by adding H2O, and to balance H add H+.  Balance charges by adding e-. Since this is for basic solution neutralize H+  with OH-.

In basic solution H+ should be neutralize with the same number of OH-.

Step 3.  Balance the number of e- by multiplying whole number to the half reaction. Since the first half reaction lost 2e- and the second half-reaction gained 3e-,  to balance the lost and gained of electrons multiply the first half-reaction with 3 to make it 6e- and the second half-reaction 2 to balance the e-.

Step 4.  Add the two half-reactions, and cancel all similar ions including e-.

Step 5.  Check if the number of atoms and charges are equal.

TRY THIS:

Balance the following:

Oxidation Number

Oxidation number is also called as oxidation states, it refers to the electrons lost or gain during chemical reaction.  To determine whether an atom undergoes oxidation or reduction in a chemical reaction, oxidation number is determined before and after reaction. Let us see for example the reaction of H2 and Cl2 in the formation of HCl:

In the above equation, in the reactant side H2 has an oxidation number of 0, Cl2 has also oxidation number 0 while in the product side H has +1 oxidation number and Cl has -1 oxidation number.
H showed an increase of oxidation number therefore it is the one that undergoes oxidation or loses an electron while Cl showed a decrease of oxidation number therefore it undergoes reduction or gains an electron.

Now in order to identify if an atom undergoes oxidation or reduction, we need to assign oxidation numbers to the different atoms in the reactant side and in the product side. There are rules to follow in assigning oxidation numbers.  Below are the rules:

1.  In free elements, each atom has an oxidation number equal to zero.  Example Na, K, Mg, Cu,  etc. Diatomic molecules also have zero oxidation number like O2, Cl2, F2, N2, Cl2, Br2, H2,  including other molecules like F4, S8.

2.  For monoatomic ions there oxidation number is equal to the charge on their atoms.  You can use periodic table in determining the oxidation  number or the different list of monoatomic ions in my previous blog.  Example elements in Group 1A ions  have +1 oxidation number, group 2A ions have +2 oxidation number.

3.  Oxidation number of oxygen is always -2 but in peroxide compounds oxygen has a charge of -1. Example in H2O (water), H has a charge of +1 and O has a charge of -2 and H2O2 (hydrogen peroxide), H has +1 charge and O as a charge of -1.

There are two atoms of oxygen having -1 charge, that is why peroxide has -2 oxidation number.

4,  Oxidation number of hydrogen is always +1 when paired with other nonmetals or polyatomic anions except for hydride. Example in HCl, hydrogen has an oxidation number of +1.  Hydride has -1 oxidation number, example CaH2, hydrogen here has an oxidation number of -1.

5. Flourine has -1 oxidation number in all compounds.  Other atoms in the halogen group ( Cl, Br, I) also has -1 oxidation number except when they are bonded with oxygen forming polyatomic ions, their positive oxidation numbers are used.  Example :

6.  In a neutral molecule, the total charge of the atom must be equal to zero.  In polyatomic ions the total charge is equal to the charge of the ion.

 Example of neutral molecule is MgCl2,  

      Mg has a charge of +2 and Cl has a charge of -1 x 2 = -2.  Therefore +2 - 2  = 0 or neutral.   

Example of polyatomic ion,

What is the charge of Sulfur?

     S  (-2x4)  =  -2

     S  (-8)  =  -2

     S  =  -2  +  8

     S  =  +6

Therefore the oxidation number of S in  is + 6.

TRY THIS:

Identify the oxidation number of each element of the following compounds and ions:

CLICK HERE FOR THE ANSWER:

Oxidation-Reduction Reactions

Oxidation-Reduction reaction is also called redox reaction, which involves transfer of electrons between two species.  The oxidation number of atoms or ions change by either gaining or losing electron. Oxidation-reduction are so vital that it occurs in the different processes from our environment, plants and animals. Example of which is  the reaction of Magnesium and oxygen

2Mg(s)   +   O₂(g)    à    2MgO(s)

The example above involves the reaction of magnesium atom with the oxygen atom producing magnesium oxide. Magnesium oxide is made up of  Mg⁺² ion and O^-2 ion.  We have to analyze what happens to magnesium and oxygen in the reaction by trying to identify if they undergo oxidation or reduction.

             2Mg    à     2 Mg⁺²     +     4e^-

O₂    +    4e^‑    à     2O^-2

The above equations are the two half-reactions that occur in the reaction of magnesium and oxygen. Half-reaction shows the electrons involved in a redox reaction..  Two magnesium atoms gives up 4 electrons to oxygen atoms and oxygen atoms gain 4 electrons from magnesium atom. The sum of the two half-reactions give the overall reaction which is

2Mg   +   O₂   +  4e^-     à     2Mg⁺²     +     2O^-2     +    4e^-     

If we cancel electrons on both sides of the equations, it will look like this:

2Mg    +     O₂      à     2Mg⁺²     +     2O^-2

Checking the equation if balance, Mg has 2 atoms, O has also 2 atoms and both charges are equal to zero in the reactant side and in the product side.

Understanding about the two half-reactions, oxidation reaction, refers to the half-reaction that involves the loss of electron while a reduction reaction refers to the half-reaction that involves the gain of electrons.  In the above reaction, Mg is oxidized and O is reduced.  Mg is called the reducing agent, since it is the one responsible for the O to be reduced.  On the other hand O is the oxidizing agent since it is the one responsible for the Mg to be oxidized.  Therefore, reducing agent is the substance that donates electrons and is oxidized while oxidizing agent is the substance that accepts electrons and is reduced.

TRY THIS PROBLEM:

Write the half-reaction and identify the reducing and oxidizing agent of the following redox reaction:

1.       4Fe  +   3O₂     à      2Fe₂O₃

2.        Cl₂    +    2NaBr     à    2NaCl   +   Br₂ 

1.    3.        Si   +   2F₂     à    SiF₄

1.    4.        H₂   +   Cl₂    à    2HCl

Acid-Base Reaction

In my previous post, I discussed the different types of chemical reactions, the combination reaction, decomposition, single replacement and double replacement reaction. Acids and bases undergo reaction.  Acid-Base Neutralization is a kind of double replacement reaction producing salt and water.  Another reaction between acid and base  resulting to the production of gas.

Acid-Base Neutralization

Acid-Base Neutralization reaction is an acid and base reaction producing salt and water.   It is called neutralization since the acidity and basicity of acids and bases are neutralized. See the the format below:

When acid and base react it produce salt and water. It is called neutralization since the H+ ion concentration is being neutralized with OH- ion making it neutral.  The acidity and basicity of the two reactants are canceled out.

For example, when a solution of HCl is mixed with NaOH, neutralization occur producing water and salt.

Since both HCl and NaOH are strong electrolytes, both are completely ionized in water.  The ionic equation is shown below:

Therefore, the reaction can be represented by the net ionic equation as shown below:

Na+ and Cl-  are spectator ions.

Lets consider a reaction between a weak acid and a strong base,

The ionic equation of this reaction is:

HCN being a weak base will not be ionized in water.  And the net ionic equation of this reaction :

Other reactions of acids and bases are the following:

Acid-Base Reactions Leading to Gas Formation

Some salts that are basic when it reacts with an acid produced gas.  These salts are the following:

Examples of reaction between this salts and acids are shown below:

Definition of Acids and Bases and their Properties

How do you identify an acid and a base?  You may say that acid has sour taste. Yes your are right!  There are several examples of acids that you are familiar, like for example ascorbic acid (vitamin c), citric acid, acetic acid in vinegar and lactic acid in milk.  But take note not all acids can be tasted.  You need to use acid-base indicators to identify if a substance is an acid or base. Base, on the other hand is said to be bitter in taste and slippery to touch like detergents.  Below are definitions of acids and bases from different chemists and some of the properties of acid and bases that you can use to identify the difference between the two.

Definition of Acids and Bases

Arrhenius Definition of Acids and Bases

Swedish chemist Svante Arrhenius define acids as substances that produce H+ ions in aqueous solution while bases are substances that produces OH- ions in aqueous solutions.

Let us look at the example below:

The equation above hydrochloric acid when dissolved in water is ionized to form H+ ions and Cl- ions.

The above equation shows that NaOH when dissolved in water is ionized to produce Na+ ions and OH- ions.

Its the H+ ions that makes a substance an acid while  OH- ions make the substance a base. 

Bronsted-Lowry Acids and Bases Definition

Danish chemist Johannes Bronsted  and English chemist Thomas Martin Lowry , define acids as substances that donate protons (hydrogen ions) while bases are substances that accept protons (hydrogen ions).  Their definition is quite broad compared to Arrhenius definition, Arrhenius acids and bases are only limited to aqueous solution, whereas in Bronsted-Lowry definition of acids and bases includes substances capable of donating proton while acids includes any substances capable of accepting proton.

The above equation shows the ionization of hydrochloric acid in water.  But this is not only what happens in the solution the H+ ions or the proton is being attracted to the oxygen of water molecule forming a hydrated proton H3O+, called hydronium ion. Therefore, the ionization of hydrochloric acid in water look like this:

The above equation shows that HCl is the Bronsted-Lowry acid, since it donated proton or H+ ion to H2O, while H2O is the Bronsted-Lowry base since it accepted protons from HCl.  

Let us have an additional example.  Identify if Bronsted-Lowry acid or base:

Given a) HBr and b) NO2-.  

a) HBr is an example of Bronsted-Lowry acid since its capable of donating proton in once dissolve in water.

b) NO2- is an example of Bronsted-Lowry base since it has capability to accepts H+ ion or proton to form HNO2.

General Properties of Acids and Bases

Acids

• Acids have sour taste.
• Acids change blue litmus paper to red.
• Acids react with other metals (higher in position in the activity series) to produce hydrogen gas.
• Acids also react with carbonates and bicarbonates (like NaCO3, NaHCO3)  to produce carbon dioxide gas.
• Aqueous acids are electrical conductor.

Bases

• Bases have bitter taste.
• Bases are slippery to touch.
• Bases change red litmus paper to blue.
• Bases are electrical conductor

Molecular Equation, Ionic Equation, Net Ionic Equation

In my previous post we learn about chemical equation and the types of chemical reactions.  All cited equations are example of molecular equation.  Molecular equations show the reaction of molecules where all species are present which exist as whole unit or molecules, as shown in the example below:

The equation above is an example of molecular equation and a double replacement reaction, It shows the reaction of lead nitrate molecule in reaction with potassium iodide forming lead  iodide precipitate and potassium nitrate molecules.  Although the equation gives all the required information of the reactants and products, it does not show what happens to the substances in solution.

Ionic equation on the other hand is a kind of equation which shows what happens to the substances in solution.  It shows dissolve species as free ions. See the example below:

This shows what happens to the reaction of lead nitrate  and potassium iodide in solution.  If you will notice all reactants and products are dissociated in water except lead iodide.  This is because lead iodide is not soluble in water.  Now in predicting if precipitate will be formed in the reaction, refer to the solubility rules.  Once the solute is soluble it should be written in ionized form as it shown in the example above and if not soluble it should be written as a whole molecule.

Now to determine the net ionic equation, spectator ions are being canceled out or eliminated.  Spectator ions are ions which are not included in the overall reaction  These are ions that are present both in reactants side and in the products side.  After the spectator ions are eliminated, what will be left is the net ionic equation.

After all spectator ions are eliminated the net ionic equation for the above reaction will be:

Net ionic equation shows only the species that take part in the reaction.

In summary, there are several rules to follow in order to come up with a net ionic equation:

1.  A reaction must be written in molecular form specifying the reactants and products if soluble or not in water.  If soluble in water aq is used which means that the substance is dissociated in water, if not soluble, s is written as subscript which means that precipitate is formed.   And the equation must be balanced.

2.  Write the ionic equation.  Substances that are soluble in water must be written in ionized form while the substance which is not soluble in water should be written as a whole molecule.  

3.  Cancel all the spectator ions both in the reactants and products side.  Then write the net ionic equation.

4.  Finalize by checking the charges and number of atoms in the net ionic equation.

Additional Example:

1.  Predict what happens when a potassium phosphate (K3PO4) solution is mixed with calcium nitrate

[Ca(NO3)2] solution. Write the net ionic equation for this reaction.

Step 1.  Write the balanced equation of the reaction.

Step 2.  Write the ionic equation by writing the soluble substances in ionized form.

Step 3.  Cancel all the spectator ions and write the net ionic equation:

Step 4.  Check if the net ionic equation has balance charges and number of atoms

                            Reactants,      +6 and -6 = 0    ,   Products  = 0

TRY THIS:

Write the balanced molecular equation, ionic equation and net ionic equation of the following:

1.  Na2S(aq)   +     ZnCl2(aq)  >>>

2.  K3PO4(aq)   +   Sr(NO3)2(aq)  >>>

3.  Mg(NO3)2 (aq)  +   NaOH(aq)  >>>>

General Properties of Aqueous Solution

Solution is a homogeneous mixture of solute and solvent.  Solute is a substance or component present in smaller quantity or proportion, while solvent is a substance or component present in larger amount or proportion.  There are different types of solution, these are the solid solution, liquid solution and the gaseous solution.  Aqueous solution is a type of liquid solution where the solvent is water. Solutes can either be solid, liquid, or gas.  Example of aqueous solution is sugar dissolved in water, salt dissolved in water, alcohol in water and more.

Aqueous solution can be electrolyte and nonelectrolyte.  Electrolyte is a substance that when dissolve in water it conducts electricity while nonelectrolyte is a substance that does not conduct electricity when dissolve in water.  Electrolytes conduct electricity since the substances is completely breaks up into ions in water while nonelectrolytes, the substances do not dissociate in water.

Electrolytes can be classified into strong electrolytes and weak electrolytes:  Strong electrolytes are those substances that are completely or 100% dissociated in water, while weak electrolytes are substances not completely dissociated in water.  Below are examples of weak and strong electrolytes:

How can we determine if an electrolyte is a strong, weak and nonelectrolyte?

An electrical conductivity apparatus is used to determine if an electrolyte is weak, strong and nonelectrolye.  The two electrodes of electrical conductivity apparatus are placed in the solution, once the bulb light very bright it means that the solution is a strong electrolyte and if light dimly, solution is said to be weak electrolyte.  But if no light produced in the electrical conductivity apparatus then the solution is a nonelectrolyte.

Electrical Conductivity Setup

Difference between strong, weak and nonelectrolyte by means of light produced in an electrical conductivity apparatus.

The  first bulb don't show light which means that the first solution is a nonelectrolyte.  Second bulb lights dimly which indicates that the solution is weak electrolyte and the third bulb lights brightly which indicates that the solution is a strong electrolyte.

What happens when substances dissociate in water?

Water is a polar substance, it can dissolve both ionic and polar substances. Being polar, water has two ends, the positive and negative.  The oxygen is the negative end  and the hydrogen is the positive end.  When an ionic substance such NaCl is dissolved in water, the three-dimensional structure of solid in NaCl is destroyed, and the Na+ and Cl- ions are separated from each other. In solution, each Na+ ion is surrounded by a number of several water molecules orienting their negative end towards the cation.  In the same manner, each Cl- ion is surrounded with several water molecules with their positive ends oriented towards the anion.     The process is called hydration.  A process where ions are surrounded by water molecules arranged in a specific manner.  It helps in the stabilization of the solution which prevents the cations to combine with anions.

Acids and bases are also electrolytes.  Some are 100% ionized in water (strong electrolytes) like HCl (hydrochloric acid)  and HNO3 (nitric acid) while other are just slightly ionized in water (weak electrolytes) like CH3COOH (acetic acid) and HF (hydrofluoric acid).  In water, HCl gas is ionized into H+ ions and Cl- ions completely as shown in the equation below:

Acetic acid (CH3COOH) in water is slightly ionized  in water as shown in the equation below:

The double arrow in the equation above shows that the reaction is reversible, meaning the reaction can be in the forward reaction and reverse reaction.  Initially, acetic acid is ionized to formed into CH3COO- ions and H+ ions,  As time goes on some of the CH3COO- ions recombine with H+ ions forming again CH3COOH molecules.  There comes a time when the rate of forward reaction equalizes the rate of backward reaction, in this point a state of chemical equilibrium is reached.

Note:

Dissociation is used for ionic substances dissolved in water.

Ionized is used for acids and bases dissolved in water.