Types of Bonds

A bond, as its names suggests, holds atoms together to form molecules.

In this article we study a few elementary types of bonds.


Ionic Bond

An ionic bond is formed as consequence of the electrostatic forces of attraction between oppositely charged ions. 

If the difference in electronegativity of the atoms is over 1.7 the bond is likely to be ionic.

Covalent Bond

A covalent bond is formed due to the sharing of valence electrons or in terms of the orbital concept, a covalent bond is formed due to the overlapping of atomic orbitals. 

If the difference in electronegativity of the atoms is less than 1.7 the bond is likely to be covalent. 

In practice most bonds are not 100% covalent or ionic.

If the difference in electronegativity between the 2 bonded atoms is zero, then the bond is purely covalent. If the difference in electronegativity between the 2 bonded atoms ranges from zero to 1.7 the bond is a polar covalent bond and exhibits some ionic character. If the difference in electronegativity between the 2 bonded atoms is 1.7 then the bond is 50% covalent and 50% ionic and if the difference in electronegativity is over 1.7 then the bond is said to be ionic.

Ready-Reference Chart (ΔE is the difference in electronegativity)

ΔE Value             Bond Character 

ΔE = 0;                 100% Covalent

ΔE =  0-1.7;          polar covalent

ΔE = 1.7;              50% covalent 50% ionic
ΔE > 1.7;              ionic

Dipole Moment: Accounts for the ionic character of covalent bonds.

Homo-nuclear, diatomic molecules (eg. H_{2  ,}  Cl₂ etc.) are non-polar.

In case of hetero-nuclear molecules, for a molecule to be non polar, the central atom mustn't have a lone pair of electrons and the central atom must be surrounded by the same atoms i.e the molecule must be symmetrical (explained below). (both conditions must be met).

However, if these conditions aren't met, then the more electronegative element will pull the electrons towards itself and thus acquire a slightly negative charge (given as  δ- ) and the other atom shall acquire a slight positive charge (given as  δ+ ).

Quick Tip: Think of the bond as a match of tug of war, the stronger (in this case more electronegative element) player will pull the rope (in this case electrons) towards himself.

Such a bond is termed a polar covalent bond and its polarity is given in terms of dipole moment. 

Dipole moment (m) = electric charge x distance of charge seperation (bond length)

_{m = q x d}

_{Dipole moment is measured in ‘Debye’ unit (D)}

Dipole moment is a vector quantity, therefore if a molecule with polar bonds is symmetrical then the opposite pulls will cancel each other resulting in zero net dipole. eg. Carbon dioxide, pictured below.

 

Fajans' Rules: Account for the covalent nature of ionic bonds. According to these small cations (or high positive charge) have greater polarising power and large anions have greater polarisibility, thus bonds formed between such species involves overlap of their electron clouds resulting in covalent character.

Quick Tip: Picture a large bunch of teenage girls (electrons) on a field trip with their teacher (nucleus) and let's call this group our anion. Now a small cation, hmm Justin Beiber for instance, walks in, and some of the girls move away from the group towards him. Thus the small cation, Justin Bieber in this case has pulled the electron cloud of the large anion towards himself and this mingling of electron clouds accounts for the covalent character of the bond.

If the group were smaller (i.e the anion was smaller) the teacher would've been able to check the movement of the girls. And  if the celebrity was somebody more important (i.e a larger cation), like Kazimierz Fajans(yes, he is the guy behind fajans rule) for instance, again fewer girls (electrons) would've strayed away from the group.


Ready-Reference Chart

Ionic                               Covalent

Low +ve charge              High +ve charge
small anion                      large anion
large cation                     small cation 



Metallic Bonding

Since Ionic and Covalent bonds cannot account for many physical properties of metals, such as strength, malleability, ductility, thermal and electrical conductivity, opacity, and luster, the concept of metallic bonding has been devised.

The positive charge in a metal aggregates to form a "kernel" whereas the electron separate out forming an electron cloud.

Metallic bonding constitutes the electrostatic attractive forces between these delocalized electrons, gathered in an electron cloud, and the kernel of positively charged metal ions.

Shitty illustration depicting kernel of positive charge and cloud of electrons.

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That's all folks!

The Passive Observer

What is an Acid?

The Arrhenius Definition

Svant Arrhenius

According to Arrhenius (pictured left) an acid disassociates in an aqueous solution to give H₃O⁺ ions (Hydronium ion, which is a protonated water molecule as a bare proton H+ cannot exist in aqueous solution as a free species).

An Arrhenius base on the other hand disassociates in aqueous solution to give OH- ions

According to this definition:

Acids                                    Bases

H₂SO₄ ,                               NaOH

HCl ,

HNO_{3 etc}                                      Ca(OH)_{2 etc}

Credits: Z(J)ubin

What The Fact... Most students, thus tend to believe that bases contain an OH group, but they fail to realize that some acids too have an OH group attached to them, but still disassociate to give H+ (or H₃O⁺ ions).  eg H₂SO₄

Well this anomaly of sorts can be explained on the basis of difference in electronegativity.

In case of the H₂SO₄   molecule, the difference in electronegativity between Oxygen and Hydrogen is greater than the difference in electronegativtiy between Oxygen and Sulfur. Therefore the bond between Oxygen and Hydrogen polarizes giving HSO₄^-  and H⁺   Refer to structure pictured below.

In case of a base (NaOH) the difference in electronegativity between Sodium and Oxygen is greater than the difference in electronegativity between Hydrogen and Oxygen, thus the bond between Oxygen and Sodium polarizes giving Na⁺   and OH^-  Refer to structure pictured below.

<Tip: Think of a bond as a match of tugowar in which the more electronegative element will pull the electrons towards itself>

Drawbacks: This theory fails to account for the acidic behavior of compounds like BCl3 and basic nature of compounds like NH3.

Brønsted–Lowry Definition

Brønsted and Lowry defined an acid as a proton donor and a base as a proton acceptor.

The removal of a proton (hydrogen ion) from an acid gives its conjugate base, and the addition of a proton to a base produces its conjugate acid. 

For instance:

• HCl + H₂O   Cl^- + H₃O⁺            

Conjugate acid:  H₃O⁺

Conjugate base: Cl^- 

• NH₃ + H₂O    OH^- + NH₄⁺

Conjugate acid:  OH^-

Conjugate base: NH₄⁺

Note: Weak acids give strong conjugate bases and weak bases give strong conjugate acids. Whereas strong acids give weak conjugate bases and strong bases give weak conjugate acids.

Water being amphoteric acts as both acid and a base. For instance in the reaction with Hydrochloric acid, water acts as a base but in the reaction with Ammonia water acts as an acid.

H₂O + H₂O  H₃O⁺ + OH⁻

Lewis Definition

Gilbert N Lewis removed the hydrogen (or proton requirement) of the Arrhenius concept and Bronsted-Lowry Definition and instead based his definition on electron pairs.

Lewis acids thus are compounds that are electron deficient and thus can accept a lone pair of electrons 

eg: BCl₃ 

In BCl₃  Boron is sp2 hybridised and it is made apparent from the diagram pictured below it has a vacant p-orbital and is thus electron deficient.

Relation between alkalis and bases

And Lewis base is a compound that can donate a lone pair of electron 

eg. NH₃

Note: On an unrelated note, a lot of people tend to use the terms alkali and base interchangeably, this however is incorrect. All alkalis are bases, but only water soluble bases are alkalis. 

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That's all folks.

The Passive Observer Out!

Mole Concept

What Is A Mole?

This Is a Mole!

And so is this..A mole of Water

• A mole is a unit of measurement often used in chemistry to measure the amount of substance and is defined as as the amount of any substance that contains as many elementary entities (e.g., atoms, molecules, ions, electrons) as there are atoms in 12 grams of pure carbon-12 (which works out to be 6.022 x 10^23 particles)

Well think of a banana, or better yet, think of a dozen bananas...how many bananas do you have? 12..right.

So now think of 1 mole of bananas, how many do you have now? well you have 6.022 x 10^23 bananas.

It would be incredibly inconvenient (and stupid) to measure bananas (or other fruit) in moles and it would even more inconvenient (and stupid) to measure atoms, molecules or other such species in terms of dozens, hence the unit mole has found great favor with (and is often restricted to) chemistry.

• Therefore, 1 mole of Hydrogen will contain 6.022 x 10^23 molecules of Hydrogen. 

Note: 1 mole of Hydrogen will contain 2 x 6.022 x 10^23 atoms of Hydrogen, since Hydrogen is a diatomic molecule, i.e one hydrogen molecule( H₂) contains 2 hydrogen atoms.

What does a Mole Weigh?

Take a dozen apples and a dozen bananas and you have 12 apples and 12 bananas (i.e the two quantities are numerically equal), but, A dozen apples doesn't weigh the same as a dozen bananas.

Similarly a mole of oxygen and a mole of hydrogen will contain the same no. of molecules but will not be equal in terms of weight.

• The molar mass of a substance or mass per mole is numerically equal to its atomic mass or molecular mass.

Therefore 1 mole of Hydrogen will weigh 2 gm and and 1 mole of Oxygen will weigh 32 gms.

1 mole of water will weigh 18 gms and since the density of water is 1g/cc, it will amount to 18mL (refer to the pic)

A mole of an Ideal Gas

1 mole of an ideal gas will occupy 22.4L at STP (0 C and 1 atm.)

Final Definition: 1 mole of any substance can now be defined as follows:

• 1 mole= 6.022 x 10^23 particles = gram molar mass of substance= 22.4L of gas at STP

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That's all folks!

The Passive Observer Out!

Redox Reactions

Redox (portmanteau of the words: reduction-oxidation) reactions include all chemical reactions in which atoms have their oxidation state changed.

Simply put, redox reactions are a family of reactions that are concerned with the transfer of electrons between species.

• Oxidation: the loss of electrons and/or increase in oxidation state

• Reduction: the gain of electrons or a decrease in oxidation state 

• Oxidizing agent : Is an “electron acceptor” i.e removes electrons from a substance and is reduced in the process.

•  Reducing agent : Donates electrons and gets oxidized in the process.

Quick Tip: Think 69..go it? Good, now stop thinking and concentrate (thank you Benny sir! and a shout out to Z(J)ubin)

• If an element is in its highest oxidation state then it cannot be oxidized any further therefore it cannot act as a reducing agent and will always behave like an oxidizing agent (i.e will get reduced)

eg: KMnO₄  Mn is in +7 oxidation state

• Similarly, if an element is in its lowest oxidation state cannot be reduced any further and thus will always act as a reducing agent and be oxidized.

eg: H₂S S is in -2 oxidation state

• Element which is present in an intermediate oxidation state will act as a reducing as well as an oxidizing agent.

eg: Peroxide

Balancing Redox Reactions

• In acidic medium

Consider the following reaction:

Fe²⁺ + MnO₄^-  è Mn²⁺ + Fe₂³⁺     

Step 1: Make note of atoms that have undergone change (in oxidation state)

Fe²⁺  è  Fe₂³⁺          

Mn⁺⁷ è Mn²⁺

# Fe has been oxidised ( reducing agent) and Mn has been reduced (oxidizing agent) 

^{Step 2: Balance the atoms}

2 x Fe²⁺  è  Fe₂³⁺          

Mn⁺⁷ è Mn²⁺

Step 3: Make note of the the electron involved and multiply by suitable whole nos. to balance the electrons.

 2 x Fe²⁺  è  Fe₂^{3+     | 2 e-  multiplying by 5}

Mn⁺⁷ è Mn^{2+              |  5 e-  multiplying by 2}

Step 4: Add the two equations and compare with original equation

10Fe²⁺ + 2MnO₄^-  è 2Mn²⁺ + 5Fe₂³⁺

STOP: Before proceeding ensure that all atoms except oxygen and hydrogen are balanced.

Step 5: Balance Oxygen by adding water to the oxygen deficient side. (LHS in this case) (Since, the RHS contains 8 oxygen atoms we add 8 molecules of water to the LHS)

10Fe²⁺ + 2MnO₄^-  è 2Mn²⁺ + 5Fe₂³⁺  + 8H₂O

Step 6: Balance Hydrogen by adding H+ to the hydrogen deficient side.

Final Balanced Equation (acid medium):  16H⁺ + 10Fe²⁺ + 2MnO₄^-  è 2Mn²⁺ + 5Fe₂³⁺  + 8H₂O

Checking Your Work: Ensure that the no. of atoms and the net charge on both the sides are equal

• In basic medium

Consider the following reaction:

Al + NO₃^--  è Al(OH)₄^-  +NH₃

Follow steps 1 to 4: We obtain

8Al + 3NO₃^--  è 8Al(OH)₄^-  + 3NH₃ 

Step 5: Balance Oxygen by adding water to the oxygen deficient side. 

23H₂O +  8Al + 3NO₃^--  è 8Al(OH)₄^-  + 3NH₃

Step 6: Balance Hydrogen also by adding water to the the hydrogen deficient side and then add an equal amount of of OH- to the opposite side. ( In this case we add 5H₂O to the LHS and 5 OH- to the RHS)

5OH^- + 23H₂O +  8Al + 3NO₃^--  è 8Al(OH)₄^-  + 3NH₃  + 5H₂O

Upon simplification,

 Final Balanced Equation (basic medium):  5OH^- + 18H₂O +  8Al + 3NO₃^--  è 8Al(OH)₄^-  + 3NH₃  

Checking Your Work: Ensure that the no. of atoms and the net charge on both the sides are equal

Disproportionation Reactions: Oxidation and reduction in the same element. An element must exist in atleast 3 oxidation states for disproportionation to take place.

Consider the following disproportionation reaction (follow the above described steps to balance this equation)

HNO₂  è HNO₃ + NO

N⁺³ è N⁺⁵ (oxidized to HNO₃)       | 2 e- multiplying by 1

N⁺³  è N⁺² (reduced to NO)           | 1 e- multiplying by 2  

________________________________adding the two eqns.

3 N⁺³ è N⁺⁵ + 2 N⁺²

Comparing with original equation we get,

3HNO₂  è HNO₃ + 2NO

Balancing Oxygen (Final Equation)

3HNO₂  è HNO₃ + 2NO +H₂O

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That's all folks!

The Passive Observer Out!