Alicyclics and Aromatics

ACYLICS:

• Carbon chains can form two types of "closed" loops
• Alicylic loops usually made with single bonds
• If the parent chain is a loop standard naming rules apply with one addition: "cylo" is added in front of the parent chain. 
• 3 different ways to draw organic compounds: 

                1) Complete structural diagram

                2) Condensed structural diagram 

                3) Line diagram

• Numbering can start anywhere and go c.w. or c.c.w. on the loop but side chains numbers must be the lowest possible 
• Loops can also be a side chain
• Same rules apply but the side chain is given a cyclo -prefix

AROMATICS:

 

• Benzene (C6H6) is a cylic hydrocarbon with unique bonds between the carbon atoms
• Structurally it can be drawn with alternating double bonds 

AROMATIC NOMENCLATURE:

• A Benzene molecule is given a special diagram to show its unique bond structure
• Benzene can be a parent chain or a side chain
• As a side chain is given the name phenyl

-Ben Suratos

Alkenes and Alkynes

Example of an Alkene

Example of an Alkyne


The naming rules for double and triple bonds are almost the same, except there’s one thing that we need to add.

Double bonds take priority.
There are two possible ways to number the carbon atoms.
When we’re numbering we must choose the lowest number

Double bonds (alkenes) end in –ene

Triple bonds (alkynes) end in –yne

Multiple double bonds:
As we all know more than one double bond can exist in a molecule.
The only difference here is that we use the same multipliers inside the parent chain.
EXAMPLE: (1,3 butadiene)

Here's a great video that teaches how to name Alkenes and Alkynes.

-Ben Suratos

Classes of Organic Compounds: Part Two

Amines

• Functional Group: 
• -NH₂ 
• General Formula:  
• R-NH₂ 
• Naming:  
• Name the parent chain (methane, ethane, propane...)  
• Substitute the "-e" ending with the suffix "-amine" (methanamine, ethanamine, propanamine...) 
• Examples:  
• C₆H₁₅N --> hexylamine 
• _CH3CH2NH2 --> ethylamine 
• C₄H₁₁N --> butylamine 
• CH₃CH₂CH₂CH₂CH₂CH₂CH₂CH₂NH₂  --> octylamine

• Above diagram --> nonylamine

Amides 

• Naming:  
• Name the parent chain (methane, ethane, propane...)  
• Substitute the "-e" ending with the suffix "-amide" (methanamide, ethanamide, propanamide...) 
• Examples:  
• CH₃NO --> methanamide 
• _CH3CH2CONH2 --> propanamide
• C₁₀H₂₁NO --> decanamide 
• CH₃CH₂CH₂CH₂CH₂CH₂CH₂CONH₂ --> octanamide

• Above diagram --> heptanamide  

Nitro

• Functional Group: 
• C≡N 
• General Formula:  
• R-C≡N
• Naming:  
• Name the parent chain (methane, ethane, propane...)  
• Add to the "-e" ending,  the suffix "-nitrile" (methanenitrile, ethanenitrile, propanenitrile...) 

• Examples:  
• CH₃CH₂CH₂CN --> butanenitrile 
• CH₃CH₂CH₂CH₂CH₂CN --> hexanenitrile 
• CH₃CH₂CH₂CH₂CH₂CH₂CN  --> heptanenitrile

• Above diagram --> octanenitrile  

Esters

• Naming:  
• Name the parent chain (methane, ethane, propane...)  
• Substitute the "-e" ending with the suffix "-oate" (methanoate, ethanoate, propanoate...) 
• Name the side chain that appears past the oxygen bond (methyl, ethyl, propyl...
• Place the side chain name in front of the modified parent chain name (methyl methanoate)
• Examples:  
• C₂H₄O₂ --> methyl methanoate 
• _C5H10O2 --> methyl butanoate 
• C₉H₁₈O₂ --> propyl hexanoate 
• C₆H₁₂O₂--> ethyl butanoate

• Above diagram --> methyl hexanoate  

• Above diagram --> propyl propanoate 

Below is a video lesson further expanding what we have covered in the past several lessons of the many classes of organic compounds: 

-Simon Sierra

Classes of Organic Compounds: Part One

What are Functional Groups?

Halides:

• Functional Group: 
• -F, -Cl, -Br, -I 
• General Formula:  
• R-X
• Naming:  
• Name the parent chain (methane, ethane, propane...) 
• Name halogen groups as side chains (fluoro-, chloro-, bromo-, iodo-)
• Name other side chains
• Examples:  
• CH₃Br --> bromomethane 
• CCl₄ --> tetrachloromethane 
• H₂FCCHFCH₂CH₃ --> 1,2-difluorobutane 
• CH₃CHFCHFCH₃  --> 2,3-difluorobutane

• Above diagram --> trichloromethane  

• Above diagram  --> 1,2-dichlorobenzene

Alcohol:

• Functional Group: 
• -OH 
• General Formula:  
• R-OH
• Naming:  
• Name the parent chain (methane, ethane, propane...) 
• Substitute the "-e" ending with "-ol" (propanol, butanol, pentanol...) 
  •  
• Name other side chains
• Number the locations of the "-OH" on the structure
• Examples:  
• C₆H₅CH₂OH --> benzyl alcohol 
• CH₃ OH --> methanol (methyl alcohol) 

• Above diagram --> cyclohexanol

• Above diagram  --> ethanol (ethyl alcohol)

Ether: 

• Functional Group: 
• -O- 
• General Formula:  
• R-O-R'
• Naming:  
• Visualize the oxygen interrupting the structure, so there are going to be three pieces (in the most basic forms)  
• The first part of the structure (for this, we use the side chain prefix: meth-, prop-...) 
• The interrupting oxygen (list it as "oxy") 
• The last part of the structure (for this, we name it as the parent chain: methane, ethane)
• Combine the first part and the oxygen into one term (methoxy, propoxy) 
• And list the final piece as a parent chain (methoxy-methane, propoxy-ethane)
• Examples:  
• C₃H₈O --> methoxy-ethane 
• C₄H₁₀O--> methoxy-propane
• C₅H₁₂O --> methoxy-butane 

• Above diagram --> propoxy-pentane  

• Above diagram  --> methoxy-pentane

Aldehyde:

• Naming:  
• Name the parent chain (methane, ethane, propane...) 
• Substitute the "-e" ending with "-al" (propanal, butanal, pentanal...) 
   
  •  
• Name other side chains
• Number the locations of the "-OH" on the structure
• Examples:  
• CH₃ CHO --> ethanal 
• HCHO --> methanal
• _CH3CH2CHO --> propanal 

• Above diagram --> nonanal  

• Above diagram  --> pentanal

Ketone:

• Naming:  
• Take the prefix of the parent chain (metha-, etha-, propa-...) 
• End with the suffix, "-none" (methanone, ethanone, propanone...) 
   
  •  
• Name other side chains
• Number the locations of the double-bonded oxygen on the structure
• Examples:  
• C₇H₁₄O --> 3-heptanone 
• _C10H20O --> 2-heptanone

• Above diagram --> 2-heptanone  

• Above diagram  --> 3-nonanone

Carboxylic Acid: 

• Naming:  
• Name the parent chain (methane, ethane, propane...) 
• Substitute the "-e" ending with"-oic acid" (methanoic acid, ethanoic acid, propanoic acid...)  
•  
• Name other side chains
• Number the locations of the "-OH" on the structure
• Examples:  
• CH₃CO2H--> acetic acid (ethanoic acid) 
• _C8H16O2 --> octanoic acid 
• _C10H20O2 --> decanoic acid 
• _C4H8O2  --> butanoic acid

• Above diagram --> nonanoic acid  

• Above diagram  --> propanoic acid 

-Simon Sierra

Introduction to Organic Chemistry

Organic chemistry is the study of carbon compounds. Carbons also form multiple covalent bonds. Carbon compounds can form chains, rings or branches. There are less than 100,000 non-organic compounds. Organic compounds number more than 17 million. The simplest organic compounds are made of carbon and hydrogen.

• Saturated compounds have no double or triple bonds.
• Compounds with only: single bonds -> Alkanes, double bonds -> Alkenes, triple bonds -> Alkynes
• ISOMER: Two compounds with the same empirical formula

Example:

• Name the alkane:

                       

Nomenclature:

There are 3 categories of organic compound:

• Straight
• Cyclic chains
• Aromatics

Straight Chains:

1.) Circle the longest continual chain and name this as the base chain. They can bend and twist.
2.) Number the base chains so side chains have the lowest possible numbers.
3.) Name each side chain using the "-yl" ending.
4.) Give each side chain the appropriate number.
      ->if there is more than one identical
          side chain, number/labels are slightly different.
5.) List side chains alphabetically.

A video on straight chains:

- George Spencer

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

Mixing Acids and Bases

Acids and pH: 
A measure of the hydrogen ion concentration [H⁺], pH is calculated using the following formula: 

• pH = -log₁₀[H⁺]   

Hydrogen ion concentration [H⁺] can be calculated using the following formula:  

• [H⁺] = 10^-pH

Example One (pH):

Find the pH of a 0.2 M solution of HCl: 

• Write the balanced equation for the dissociation of the acid:
• HCL ==> H⁺(aq) + Cl^-(aq) 
• Use the equation to find the [H⁺]: 
• 0.2 M x (1/1) = 0.2 M of [H⁺] **HCl is a strong acid which fully dissociates** 
• Calculate pH: 
• pH = -log₁₀[H⁺]   
• pH = -log₁₀[0.2]  
• pH = 0.7 

Example Two (pH):

Find the pH of a 0.2 M solution of H₂SO₄: 

• Write the balanced equation for the dissociation of the acid:
• H₂SO₄ ==> 2H⁺(aq) + SO₄^2- (aq) 
• Use the equation to find the [H⁺]: 
• 0.2 M x (2) = 0.4 M of [H⁺] **H₂SO₄ is another strong acid which fully dissociates** 
• Calculate pH: 
• pH = -log₁₀[H⁺]   
• pH = -log₁₀[0.4]  
• pH = 0.4 

Example Three (pH):

Find the [H⁺]  of a nitric acid solution with a pH of 3.0: 

• Use the formula to find the [H⁺]: 
• pH = 3.0 
• [H⁺]  = 10^-pH 
• [H⁺]  = 10^-3.0 
• [H⁺]  = 0.001 M

• Check it: 
• pH = -log₁₀[H⁺]   
• pH = -log₁₀[0.001]  
• pH = 3.0 

Bases and pOH: 

A measure of the hydroxide ion concentration [OH^-], pOH is calculated using the following formula: 

• pOH = -log₁₀[OH^-]    

Hydroxide ion concentration [OH^-] can be calculated using the following formula:  

• [OH^-] = 10^-pOH 

Example One (pOH):

Find the pOH of a 0.1 M solution of NaOH: 

• Write the balanced equation for the dissociation of the base:
• NaOH ==> Na⁺(aq) + OH^-(aq) 
• Use the equation to find the [OH^-]: 
• 0.1 M x (1) = 0.1 M of [OH^-] **NaOH is a strong base which fully dissociates** 
• Calculate pH: 
• pH = -log₁₀[OH^-]   
• pH = -log₁₀[0.1]  
• pH = 1 

Example Two (pOH):

Find the pOH of a 0.1 M solution of Ba(OH)₂: 

• Write the balanced equation for the dissociation of the base:
• Ba(OH)₂ ==> 2OH^-(aq) + Ba²⁺ (aq) 
• Use the equation to find the [OH^-]: 
• 0.1 M x (2) = 0.2 M of [OH^-] **Ba(OH)₂ is another strong base which fully dissociates** 
• Calculate pH: 
• pH = -log₁₀[OH^-]   
• pH = -log₁₀[0.2]  
• pH = 0.7 

Example Three (pOH):

Find the [OH^-]  of a sodium hydroxide solution with a pOH of 1.0: 

• Use the formula to find the [OH^-]: 
• pH = 3.0 
• [OH^-]  = 10^-pOH 
• [OH^-]  = 10^-1.0 
• [OH^-]  = 0.1 M

• Check it: 
• pH = -log₁₀[OH^-]   
• pH = -log₁₀[0.1]  
• pH = 1.0   

Another formula to know from this unit that greatly simplifies things is:

• **pH and pOH = 14**

- Simon Sierra, George Spencer and Ben Suratos