Naming Organic Compounds

Alkane

• An alkane is a hydrocarbon in which all the carbon atoms are connected by single bond.

Rules

• Number the carbon atoms in the parent chain.
• Using the correct substituent name (methyl, ethyl, chloro,etc.), and the carbon atom number on the parent chain.
• If a particular substituent occurs more than once, use a prefix (di-, tri-, tetra-) to indicate the number of those substituents.
• List the alkyl substituents in alphabetical order.

Formula	Name
CH4	Methane
C2H6	Ethane
C3H8	Propane
C4H10	Butane
C5H12	Pentane
C6H14	Hexane
C7H16	Heptane
C8H18	Octane
C9H20	Nonane
C10H22	Decan

 

Cycloalkanes

• Hydrocarbon chains which connect in a head-to-tail circle are called cyclic hydrocarbons or cycloalkanes.

Rules

• A single substituent does not use a number to indicate the position of attachment
• If there are more than on substituent, the first substituent is assumed to be at the "1" position and the remaining substituents are numbered either clockwise or counterclockwise so as to have the lowest set of overall values.

Alkyl Halides

Rules

• Attached F, Cl, Br and I atoms are called "fluoro", "chloro", "bromo" and "Iodo" groups
• If more than one of the same kind of galogen is present, use the prefix di, tri, etc.
• If a compound contains both alkyl and halo groups, list the attached groups in alphabetical order.

Mutiple Bonds

• An alkene is an organic compound containing a carbon-carbon double bond.
• An alkyne is an organic compound containing a carbon-carbon triple bond.

Rules

• If a double bond is present, change the "ane" ending of the parent hydrocarbon to "ene"
• If a triple bond is present, change the "ane" ending of the parent hydrocarbon to "yne"
• Use a number to indicate the lower numbered carbon involved in the bond. The number goes immdiately in front of the name of the parent hydrocarbon.
• Number the parent hydrocarbon to give the double/triple bond the lowest possible mumber. 

Functional groups:

Alcohols

• An alcohol is an organic compound containing an OH group.

Rules

• Number the hydrocarbon chain to give the lowest possible number to the OH group.
• Place the number immediately before the name f the parent hydrocarbon, separated by a dash. Alkyl groups are placed in front of the number for the OH.
• Indicate the presence of an OH group by changing the "e" ending of the hydrocarbon chain to "ol".

Aldehydes

• An Aldehydes is an organic compound containing a C=O group at the end of a hydrocarbon chain.

Rules:

• Number the-longest carbon chain starting with the -CHO group.
• Name the parent compound by using the alkane name and replacing the -e ending with an -al ending.

Ketones

• A ketons is an organic compund containing a C=O group at a psoition other than at the and of a hydrocarbon.

Rules:

• Number the-longest carbon chain starting so that the –C=O group is attached to the carbon atom with the lowest number.
• Name the parent compound by using the alkane name and replacing the -e ending with an -one ending.

Ethers

• An ethers is a organice compound in which an oxygen joins two hydrocarbon groups.

Rules:

• Start with the side group, add "oxy" to the side group name.

Amines

• An amine is an organic compound containing an NH2 group.
• Amines are organice base and react with acids.

Rules:

• Identify the names of the alkyl groups bonded to the nitrogen atom
• Simply replace the alkane -e ending with -amine.
• The format is as follows: (alkyl name)(-amine)

Amides

• An amides is an organic compound containing a CONH2 group.

Rules:

• Amides are commonly named similar to a carboxylic acid, replacing the –oic acid suffix with amide.

Carboxylic Acid

• An carboxylic acid is an organic compound whcih contains COOH group.

Rules:

• Number the-longest carbon chain starting with the -COOH group. Name the parent compound by using the alkane name and replacing the -e ending with an –oic acid ending.

  

VSEPER

  Exclusion of the three electrons in the molecule as follows: solitary the repulsion between the electrons(lone-lone exclusion); solitary electronic and bonding electron repulsion between the (solitary - into exclusion); bonding electron pairs ofrepulsion between the (into the  - into the exclusion).  Molecules will try to avoid such exclusion to maintain stability.  Exclusion can not be avoided, the molecules tend to form a rejection of the weakest structures.  Lone pair repulsion between the electrons is electronic and bonding, which in turn is greater than the repulsion between the bonding electron pairs.  Therefore, the molecules tend to be the weakest into- into exclusion.

Electronegativity and Polarity

Electronegativity

-Electronegativity Table

 

- Electronegativity is a measure of the attraction of an atom for electrons in a covalent bond.

- Fluorine, the most reactive non-metal, is assigned the highest value since it has the greatest attraction for the electron being shared by the other element. Oxygen is also highly electronegative and has a strong attraction for electrons.
- Metals have low electronegativities since they have weak attraction for any shared electrons.
- When two unlike atoms are convalently bonded, the shared electrons will be more strongly attracted to the atom of greater electronegativity. Such a bond is said to be polar. A polar bond results in the unequal sharing of the electrons in the bond.

Polarity

- Polarity is a physical property of compounds which relates other physical properties such as melting and boiling points, solubility, and inter molecular interactions between molecules.

- For the most part, there is a direct correlation between the polarity of a molecule and number and types of polar or non-polar covalent bonds which are present.

- In a few cases, a molecule may have polar bonds, but in a symmetrical arrangement which then gives rise to a non-polar molecule such as carbon dioxide.

              

Chemical Bonding

Electrostatic Force

-Force existing as a result of the attraction or repulsion between two charged particles.

All bonding is based on the electrostatic relationship.

1) Opposite charge attract each other.

2) Like charges repel each other.

3) The greater the distance between two charged particles, the smaller the attractive force             existing between them.

4) The greater the charge on the particles, the greater the force of attraction between them.

   - This force operates equally in all directions meaning that positively - charged particles attract negatively - charged particles every way.

Ionic Bond

- An ionic bond is formed by the attraction of a positive ion to a negative ion and is formed when an electron from one atom is transferred to another atom, so as to create on positive and one negative ion.

Polar and Non Polar Covalent Bond

polar Covalent Bond

- A bond between 2 nonmetal atoms that have different electronegativities and therefore have unequal sharing of the bonding electron pair

- Example: In H-Cl, the electronegativity of the Cl atom is 3.0, while that of the H atom is 2.1

- The result is a bond where the electron pair is displaced toward the more electronegative atom. This atom then obtains a partial-negative charge while the less electronegative atom has a partial-positive charge.This separation of charge or bond dipole can be illustrated using an arrow with the arrowhead directed toward the more electronegative atom.

The Greek letter delta indicates "partially".

Non-polar Covalent Bond

- A bond between 2 nonmetal atoms that have the same electronegativity and therefore have equal sharing of the bonding electron pair

- Example: In H-H each H atom has an electronegativity value of 2.1, therefore the covalent bond between them is considered nonpolar

 

Bohr, Lewis and Electron Dot Diagrams

Electronic Structure of the Atom

Electronic Structure of the Atom

The electronic configuration of an atom is notation that describe the orbitals in which the electrons occupy and total number of electron in each orbitals.

Niels Bohr proposed that electrons only exist in specific energy states when an electron absorb or emits a specific amount of energy it instantaneously moves from one orbital to another.

 

Energy Levels of Electrons

As you may remember from chemistry, an atom consists of electrons orbiting around a nucleus. However, the electrons cannot choose any orbit they wish. They are restricted to orbits with only certain energies. Electrons can jump from one energy level to another, but they can never have orbits with energies other than the allowed energy levels.

The Energy difference between two particular energy level is called the quantum of energy.

- Ground State: When all the electron of an atom are in their lowest possible energy levels.

- Excited State: When on or more of an atom's electrons are in energy levels other than the lowest available level.

Electron Configuration

The letters S,P,D,F are different types of orbital.

Example

For example, let's say you have an atom of lithium. Lithium's atomic number is 3. So a neutral atom of lithium has 3 protons and 3 electrons. We would need space for 3 electrons. Write:

There are 2 electrons in the 1s sublevel and 1 electron in the 2s sub-level. 2 + 1 = 3. That's it!

Now let's try fluorine, which has an atomic number of 9. A neutral atom of fluorine has 9 protons and 9 electrons. We need enough space for 9 electrons. The 1s orbital can hold 2 electrons and the 2s orbital can hold 2 more electrons. The five remaining electrons must go into the next orbital, the 2p orbital. The 2p orbital can hold up to 6, but we only have 5. So the following would be the correct electron configuration for a neutral atom of fluorine. 2 + 2 + 5 = 9.

Writing Electron Configuration for Ion

For a negative ion: Add electrons (equal to the charge) to the last unfilled sub-shell starting where the neutral atom left off.

For a positive ion: 1) Start with the neutral configuration, remove electrons from the outermost shell first. 2) If there 's are electrons on both the S and P orbitals of the outermost shell, the electrons in p orbitals are removed first.

Valance Electrons

The number of valence electrons is just how many electrons an atom has in its outer shell. It's easy to figure out if you've got a periodic table.

All the elements in each column have the same number of electrons in their outer shells. All the elements in the first column all have a single valence electron (H, Li, Na, K, etc.).

The second column elements all have 2 valence electrons (Be, Mg, Ca, Sr, etc.).
Skipping over the gap, go to the Group 3 elements, which all have 3 valence electrons (B, Al, Ga, etc.).

The elements in the next column (C, Si, Ge, etc.) all have 4 valence electrons.
The elements in the next column (N, P, As, etc.) all have, yes, you guessed it, 5 valence electrons.
O, S, Se, and the others in this column have 6 valence electrons.
The halogens in the next-to-last column (F, Cl, Br, etc.) have 7 valence electrons.

The noble gases in the right-most column (Ne, Ar, Kr, etc.) all have 8 electrons in their out except for He, which only has 2 electrons.

If an atom is an ion, you must include the charge also:
For a positive ion, for each charge subtract one electron, *for instance, Na⁺ has 1-1 = 0, BUT it has 8 valence electrons because it has the same electron configuration as Ne. Just as K⁺ has the same configuration as Ar. Therefore, the alkali metal ions with a single positive charge will have 8 valence electrons.

For a negative ion, add one electron for each charge, for instance, O^2- has 6+2 = 8 valence electrons

Periodic Table Trends

Trends on Periodic Table!

Metallic Properties

• decrease across a period with increase in number of valence electrons as well as a decrease in atomic radius
• increase down the group with increase in number of shells and atomic radius

Atomic Radius

•  the distance from the atomic nucleus to the outermost stable electron orbital in an atom
• decrease as one progresses across a period from left to right because the effective nuclear charge increases, thereby attracting the orbiting electrons and lessening the radius
• usually increases while going down a group due to the addition of a new energy level (shell)

Ionization energy

• minimum amount of energy required to remove one electron from each atom in a mole of atoms in the gaseous state
• increase while one progresses across a period because the greater number of protons (higher nuclear charge) attract the orbiting electrons more strongly
• As one progresses down a group on the periodic table, the ionization energy will likely decrease since the valence electrons are farther away from the nucleus and experience a weaker attraction to the nucleus

Electronegativity

• a measure of the ability of an atom or molecule to attract pairs of electrons in the context of a chemical bond
• as one moves from left to right across a period in the periodic table, the electronegativity increases due to the stronger attraction that the atoms obtain as the nuclear charge increases
• Moving down a group, the electronegativity decreases due to the longer distance between the nucleus and the valence electron shell, thereby decreasing the attraction, making the atom have less of an attraction for electrons or protons

Reactivity

• how likely or vigorously an atom is to react with other substances
• Metals
       Period - reactivity decreases as you go from left to right across a period
       Group - reactivity increases as you go down a group
• Non-metals
Period - reactivity increases as you go from the left to the right across a period
Group - reactivity decreases as you go down the group

Ion Charge

Melting Point and Boiling Point

• Metals - the melting point for metals generally decreases as you go down a group.
• Non-metals - the melting point for non-metals generally increases as you go down a group.

  

History of Periodic Table

Aristotle in 330 BC

• Four element theory: earth, air, fire & water

John Newlands in 1864



• The known elements (>60) were arranged in order of atomic weights and observed similarities between the first and ninth elements, the second and tenth elements etc.
• He proposed the 'Law of Octaves'.
• Also he assigned Hydrogen an arbitary mass of 1.

Dmitri Mendeleev (Russian) in 1869

• First produced a table based on atomic weights but arranged 'periodically' with elements with similar properties under each other
• Gaps were left for elements that were unknown at that time and their properties predicted (the elements were gallium, scandium and germanium)
• The order of elements was re-arranged if their properties dictated it, eg, tellerium is heavier than iodine but comes before it in the Periodic Table
• Mendeleev's Periodic Table was important because it enabled the properties of elements to be predicted by means of the 'periodic law': properties of the elements vary periodically with their atomic weights.

Mendeleev's 1871 periodic table

Modern Periodic Table

Our modern day periodic table is expanded beyond Mendeleev's initial 63 elements. Most of the current periodic tables include 108 or 109 elements.

        

Groups

• The modern periodic table of the elements contains 18 groups, or vertical columns.

• Elements in a group have similar chemical and physical properties because they have the same number of outer electrons.
• Elements in a group are like members of a family--each is different, but all are related by common characteristics.
• To avoid confusion, the Roman numerals and letters designating groups will eventually be replaced by the numerals from one to eighteen.

Periods

• Each of the table's horizontal rows is called a period.

• Along a period, a gradual change in chemical properties occurs
  from one element to another.
• Changes in the properties occur because the number of protons and electrons increases from left to right across a period or row.
• The increase in number of electrons is important because the outer electrons determine the element's chemical properties.

Metals

• They are usually shiny, very dense, and only melt at high temperatures. 
• Their shape can be easily changed into thin wires or sheets without breaking. 
• Metals will corrode, gradually wearing away, like rusting iron.
• Heat and electricity travel easily through metals, which is why it is not wise to stand next to a flagpole during a thunderstorm!

Nonmetals

• On the right side of the periodic table, are very different from metals.
• Their surface is dull and they don’t conduct heat and electricity. 
• As compared to metals, they have low density and will melt at low temperatures. 
• The shape of nonmetals cannot be changed easily because they are brittle and will break. 

Metalloids

• Elements that have properties of both metals and nonmetals.
• They can be shiny or dull and their shape is easily changed. 
• Electricity and heat can travel through metalloids but not as easily as they travel through metals.