Intermolecular Bonds

Intramolecular Bonds exist inside molecules, for example Ionic and Covalent bonds.

Intermolecular Bonds exist between molecules. The stronger the intermolecular bonds the higher the boiling point. Two types of intermolecular bonds are Van der Waals bonds and Hydrogen bonds 
 

A. Van der Waals bonds are based on electron distribution. Van de Waals bonds can be categorized in two categories: a weak bond created by the London Dispersion Force (LDF) or dipole-dipole bonds.

1. London Dispersion Forces (LDF)

These are the weakest intermolecular bonds. London Dispersion Forces are present in every molecule and are caused by the random movements of electrons inside atoms. Sometimes a large of electrons congregates on one side of an atom, causing a temporary dipole. The more electrons in the molecule the stronger the LDF.




2. Dipole-Dipole Bonds

These exist only in polar molecules, where the negative and positive ends of molecules are attracted to the negative and positive ends of other molecules. These are stronger than London Dispersion Forces but weaker than Hydrogen Bonds.




A Polar Molecule

B. Hydrogen Bonds

When Hydrogen bonds with certain elements( Oxygen, Fluorine, Nitrogen, and in some cases Chlorine). Hydrogen Bonds are very strong and highly polar.



A Hydrogen Bond

-Ben Suratos

Polar Molecules

In the past lesson we learned how to determine whether a bond was polar or non-polar and whether it was covalent or ionic. The "polar" sub-types were found only in covalent bonds and the difference lay in how they shared their bonds.

• In a polar bond, the electrons are shared unequally between two atoms. 
• The electrons are pulled closer to the more electronegative atom, giving that atom a slight negative charge and the other atom a slight positive charge. 
• In a non-polar bond, the electrons are shared equally between two atoms. 
• The electrons are not charged meaning the bond has no positive or negative end.  

In the same way that we can differentiate between polar and non-polar bonds, we can classify molecules either as polar or non-polar.   

• A polar molecule has one end with a positive charge and another end with a negative charge 
• This means polar molecules have an overall charge separation
• Polar molecules are also called dipoles (the prefix di- means two) because of its two charged ends 
• A non-polar molecule has neither positive or negatives charges on its ends 
• This means it is not a dipole  

Determining Polarity 

Being polar on non-polar gives a molecule a variety of different properties. If a molecule contains only non-polar bonds, it will be a non-polar molecule. However, a molecule that contains polar bonds is not necessarily a polar molecule. 

To determine whether a molecule is polar, you need to look at more than just the polarity of its bonds. You need to look at the shape of the molecule. 

*The shape of the molecule and the polarity of its bonds together determine whether the molecule is polar or non-polar* 

  
But the shapes of molecules can get quite convoluted and require further learning on our part so to compensate for that we can look at the molecules symmetry. When observing the symmetry of the molecule to determine its polarity a good rule to keep in mind is this:

• Polar molecules are unsymmetrical  
• And molecules can be unsymmetrical in two ways: 
• Different atoms 
• Different numbers of atoms
• Non-polar molecules are symmetrical (usually) 

The symmetry of a molecule is found by drawing the molecule's Lewis dot diagram or bond diagram and inspecting both the vertical and horizontal symmetries. Remember, symmetry is the quality of being made up of exactly similar parts facing each other or around an axis.  

 

Examples:  

Given the following compounds determine if it is a polar or non-polar molecule: 

• NH₃
• Lewis Diagram:  
• 
• # of Lone Pairs Around Central Atom  
• 1
• # of Bonding Electron Groups Around Central Atom  
• 3
• Name of Shape  
• Pyramidal
• Shape Diagram and Bond Dipoles   
• 
• Symmetric? 
• Asymmetric
• Polar? 
• Polar Molecule

• C₂H₂ 
• Lewis Diagram: 
• H : C ::: C : H
• # of Lone Pairs Around Central Atom  
• 0
• # of Bonding Electron Groups Around Central Atom  
• 2 
• Name of Shape 
• Linear 
• Shape Diagram and Bond Dipoles   
• 
• Symmetric? 
• Symmetric
• Polar? 
• Non-Polar Molecule 

Below is a video that provides a more comprehensive explanation on this subject:  

 

Here's a great acronym for this lesson: 

Symmetric 

Non-Polar

Asymmetric

Polar 

*SNAP*

- Simon Sierra

Bonds and Electronegativity

Bonding:
Previous studies in chemistry have shown us three main types of bonds.

• Ionic bonds which exist between a metal and a non-metal. In this bond the electrons are transferred. 
• Covalent bonds which exist between a non-metal and a non-metal. In this bond the electrons are shared.  
• Metallic bonds which exist between metals and metals. In this bond pure metals are held together by electrostatic attraction.

We have worked largely with covalent and ionic bonds and know them to be illustrated as such:

 

Electronegativity: 

An atom's electronegativity reflects its ability to attract electrons in a chemical bond. As we learned before, there is a distinct trend in the periodic table when reading for the electronegativity of an element: 

 

And the values are as follows:  

From these two tables we can conclude that fluorine is the most electronegative element with an electronegativity of 4.0 whereas caesium and francium are the least electronegative elements with an a shared electronegativity of 0.7. And these values are useful as they help us distinguish between ionic and covalent bonds as well as the sub-bonds of a covalent bond. These two sub-bonds are: 

• Polar covalent bonds which form an uneven sharing of electrons 
• Non-polar covalent bonds which form an equal sharing of electrons. 

By calculating the difference in electronegativity between the elements involved in a bond, we can predict the type of bond. The results are held within these ranges: 

• en > 1.7 = ionic bond; the electrons are transferred 
• en < 1.7 = polar covalent bond; the electrons are shared, but not equally
• en = 1.7 = non-polar covalent bond; the electrons are equally shared 

Examples: 

Predict the type of bonds formed by calculating the electronegative difference in the following compounds: 

• H-O
• Electronegative Difference: 3.44 - 2.20 = 1.24
• Type of Bond: Polar Covalent Bond
• C-H  
• Electronegative Difference: 2.55 - 2.20 = 0.35
• Type of Bond: Polar Covalent Bond
• K-F  
• Electronegative Difference: 3.98 - 0.82 = 3.16  
• Type of Bond: Ionic Bond
• N-H  
• Electronegative Difference: 3.04 - 2.20 = 0.84 
• Type of Bond: Polar Covalent Bond
• Na-F  
• Electronegative Difference: 3.98 - 0.93 = 3.05  
• Type of Bond: Ionic Bond
• O-Cl 
• Electronegative Difference: 3.44 - 3.16 = 0.28  
• Type of Bond: Polar Covalent Bond 
• O-O 
• Electronegative Difference: 3.44 - 3.44 = 0.00 
• Type of Bond: Non-Polar Covalent Bond  

Further explanation on this topic can be found with the following videos, one focused on the electronegativity trend and other deals with distinguishing the bonds:

 

- George Spencer, Simon Sierra, and Benedict Suratos