Alkenes and Alkynes

Carbons can form double or triple bonds with carbon atoms. This, in turn, would mean fewer hydrogens become attached. The naming rules are similar to alkanes.

Try to get the lowest number for indicating the double or triple bond position. An alkene has a double bond.



For example,
             
CH2 = CH - CH2 - CH3

would be butene. Notice that the name will end in -ene.

CH2 = CH2  would be ethene

The general formula is C(N)H(2N)

Some general rules for naming:
1) Find the longest chain. This will be the part that makes up the end of the name.
2) Number the carbon atoms to get the lowest number for the start of the double bond and place this number before the parent name.
3) Write the names and numbers for the side groups, assembling them alphabetically.

Some alkenes have the same structure, but a different geometry (or a different shape). These are called geometric isomers.

We use trans- or cis- to distinguish between these isomers.
If two adjacent carbon atoms are bonded by a double bond and have side chains on them, two chains are possible.

If the larger groups are both on the top or on the bottom, then it is cis.
If the larger groups are diagonal to each other, or across from each other, then it is trans.
If the larger groups are one on top of the other, then there are no isomers, and therefore no cis or trans.



Now, let's do alkynes. These are the same as alkanes or alkenes in terms of naming, but these just end in -yne.

Alkynes have one or more triple bond between carbons which lead to unsaturated hydrocarbons.
For example:
CH ≡ CH would be ethyne.

The general formula is C(N)H(2N-2)
There are no cis or trans forms since they are mostly the same.

Now, let's try a few examples.



The longest chain is 6, and the position of the triple bond is 2, counting from right to left. This is because we get the lowest number possible.

First, write down 2-hexyne. There is a methyl group on the 5th position (since we are already counting right to left, we must stay consistent)

Finally, we end up with 5-methyl-2-hexyne.

Here's a video on alkene.



Alkynes:



http://herh.ccrsb.ca/staff/FarrellL/chem11downloads/org/org2.pdf

Organic Chemistry

Now, we will start a new section on organic chemistry. Organic chemistry is the chemistry of carbon compounds. This is responsible for many of the everyday products that are used around the world, such as polyester, alcohol, cocaine.



Properties of organic compounds are low melting points, with low electricity. The carbon atoms can be linked in straight lines (which are really zigzag), circular structures, or in branches. As well, they can be linked in single, double or triple bonds.

First, we will start with alkanes. These are unbranched, straight chain compounds that only contain H and C. These are all hydrocarbons that are bonded together in single bonds. The names will end in "-ane" because they belong to the "alkane" group.



For the molecular formula C2H6, the full structure is:
      H   H
       |     |
H - C - C - H
       |     |
      H   H

As you see, the two carbons are bonded together, and the remaining bonds are with hydrogens.

The formulas for regular alkanes are:

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

For alkanes, the general formula is C(N)H(2N+2).

When hydrocarbons have side branches, they are called substituted hydrocarbons or branched hydrocarbons. (If you refer to the diagram above, this means that one of the Hydrogens surrounding the C is replaced with, say, a methane)

Carbon has 4 bonds. When a carbon is attached to another carbon, there are 3 hydrogens surrounding the   original carbon. When carbon is attached to two other carbons, there are 2 hydrogens. When attached to 3 carbons, there is 1 hydrogen left. When carbon is attached to 4 carbons, there are no hydrogens left.

For example, if a carbon is attached to another carbon, we would get:
                                              H
                                               |
  CH3                              H - C - H
       |                                       |
H - C - H               or      H - C - H      (the CH3 is the same as an added carbon, with 3 hydrogens
       |                                       |             branching out)
      H                                     H

The CH3 is supposed to be methane, but one H got bonded with the original H that was there, and so CH3 is left.

To name this, we would say this is ethane. We will go through this more later.

Suppose there is a name: 2-methylpentane



The methyl part is an alkyl group, which is an alkane which has lost one hydrogen atom.
The parent hydrocarbon is pentane (which is the longest straight chain, as there are 5 carbons).

The names of the alkyl groups end in "yl" because they are alkyl. If there are 2 or more of the same kind of alkyl group, use prefixes like di, tri, tetra, penta.



If more than one alkyl group (as in different kinds of alkyl groups) are present, list them alphabetically, and put the position number in front, with a dash between each group and number.

Here's a video:



http://www.arps.org/users/hs/thompsom/chemcom/unit_3/Naming_Alkanes_Worksheet_1.pdf
http://www.grossmont.edu/martinlarter/chemistry141/reference/AlkaneWorksheet.pdf

Chemical Bonding

As you know, chemical bonding only involves valence electrons, and compounds continue to gain/lose or share electrons until they have a full closed shell.

If electrons are shared equally, we say the covalent bond is non-polar. If they are shared unequally, a polar covalent bond is formed.

But when electrons are transferred, it is totally different, and it is an ionic bond.



Recall from the periodic table trends the term electronegativity. It is the tendency to attract electrons, with non-metals having a high number of electronegativity.

In non-polar bonding, equal sharing is observed to get a full shell. Electrons are attracted to the nuclei, and there are some in between the two atoms. The bonds are very strong, and require a high amount of energy to break them.

Intramolecular forces hold atoms of the molecule together. Intermolecular forces bond the molecules together. When melting occurs, the bond between the atoms is not broken, and only the intermolecular, weak bonds are affected.

In polar bonding, atoms with greater electronegativity pull electrons toward itself, so the shared electrons will actually be closer to this atom.



Just remember, atoms with the higher electronegativity form a PARTIAL NEGATIVE CHARGE, while the lower electronegativity atom will form a PARTIAL POSITIVE CHARGE.
Then, add an arrow to indicate which way the electrons will tend to move.  (in other words, draw the arrow pointing to the partial negative atom)

E.g. What is the bond between C and O?

To do this, we will use a Table of Electronegativities. If the difference is < 0.5, it is covalent. If the difference in EN is >0.5 and <1.8, it is polar covalent. If the difference is >1.8, it is ionic.



So, the EN of C is 2.55, and the EN of O is 3.44, giving a difference of 0.89, which is polar covalent.

This means our diagram would look like:

C δ+ ---------> O δ-


Here's a video:



On related videos, there are continuations of this video.

Here is some practice:
http://chemistry.sswiki.com/file/view/8.4+Review.pdf
http://www.dorjegurung.com/chemistry/IB_year1/worksheets/wkst_hybridization_shape_polarity.pdf

Test next day!

Today, we just had a review day for our test on Monday on Atomic Theory and Periodic Table and trends.

Good luck to all!

Electron Dot Diagrams

How to draw them

Drawing electon dot diagrams (or Lewis diagrams) are quite easy if you follow these steps/rules.

1. The nucleus is represented by the atomic symbol
2. Each electron is represented by a dot which are around the atomic symbol
3. There are four orbitals on each side of the nucleus. Each orbit can hold a maximum of 2 electrons
4. 8 electrons represent a closed shell, except for H, which needs 2 electrons to become stable
5. Determine the # of valence electrons

For example sodium is in group 1 therefore it has 1 valence electron

Here are some examples that include double/triple bonding

There are many types of electron diagrams, such as Bohr, Lewis or structural.

In a structural diagram, each bonded pair is represented as a line.
Here's an example:

Now, if we want to represent ionic bonds, we would look at the charge of the ion. For example, for NaCl:

You would put the negative ion in square brackets, and write the positive ion right beside it. So, you would write Na+ [Cl]- (but you would include dots around the four orbits to symbolize that the compound has satisfied the octet rule).

In Lewis structures, the number of bonds is shown, along with the lone paired electrons, which are electrons that do not participate in sharing.

But how do we determine which element is the central atom?
Here are a few rules to follow:
1. H and F are never in the middle
2. If there is a metal, the metal is always the central one
3. If a molecule contains only one atom of one element, and many of another, then the single atom is in the centre.
4. Atoms that need the most electrons to complete their valence shell is in the centre.

Let's try an example:
C2H2:

First, how many valence electrons do we have?

4x2 + 1x2 = 10

Now, we have two C's in the middle, with a bond in between them.

C - C

Then, we have two H's branching out from the C's.

H - C - C - H

That makes 6 electrons, so we have to distribute 4 more to the C's (since H's are already full)

If we distributed them evenly, we would have two lone electrons on each of the C's. (preferably one on each of the orbits)

Then, if we took one of the electrons and bonded it with an electron from the other C, we would have a double bond between the C's and one lone electron remaining on each C.

Now, if we take those ones and bond them together, we will successfully get our full shell, with a triple bond between the Carbons.




Here's some practice:
http://misterguch.brinkster.net/PRA017.pdf
http://www.arps.org/users/hs/thompsom/honors/chap_07/Lewis_Structure_Worksheet_1.pdf

Bohr Model Diagram

The Bohr model was approximately correct to what we know today, it was close in the theory of quantum mechanics. Bohr proposed that electrons occupied shells, otherwise known as orbitals. He believed that electrons had certain "energy levels", the lowest energy state being the |ground state. Bohr thought that the electron can "jump" to higher levels when they are "excited"and vice versa to when the electrons fall to the lower lessons. He thought that each "jump" or "fall" would give off a quantum of light energy, emitting a spectrum of light.

 

In the first shell, only 2 electrons are able to occupy it. The second shell there can be 8 electrons, and the third shell, there can also be 8 electrons, and this is called the octet. Bohr also suggested that electrons were not able to move freely in the atom and that when electrons are heated, they will give off a certain wavelength that is unique for each element.

Bohr also wrote something in the middle of the diagram. He wrote the number of protons and neutrons in the atom. For example, the diagram above is a representation of a Boron atom. The Atomic number is 5, meaning the proton is also 5... So in the middle it says P:5. The atomic number of this element is 10.8, and to get the neutrons you do 10.8-5  which equals 5.8, which is 6 when rounded. There fore you write N:6 in the middle. After drawing the thing in the middle, don't forget to draw the electrons that are in the orbitals. In this case, there are 5 electrons, so you put 2 electrons in the first shell, and 3 in the next.

and... BAYUM u just got yourself a Bohr Model Diagram! how easy was that! YAY! success ^^ WHOOP WHOOP!! :)

For more practice or more knowledge on Bohr Diagrams.. refer to these sites/videos:

http://needham.wikispaces.com/file/view/03+BohrModelPractice.pdf
http://pw.vsb.bc.ca/wyper/sci9/2-3_bohr_diagram_worksheet.pdf

Periodic Table Trends

In this blog, we will cover some patterns and trends in the periodic table.

Some of the patterns that we will discuss include
1. Metallic properties
2. Atomic Radius
3. Ionization energy
4. Electronegativity
5. Reactivity
6. Ion charge
7. Melting/boiling point
8. Density

First, let's start with metallic properties.
Although this one seems obvious, it is still good to point it out. When moving from left to right in the periodic table, the properties of the elements change from metallic to non-metallic.



Also, when going down a family (column), elements become more metallic, or better metals.

Secondly, let's study about the atomic radii.
When moving from left to right, the atomic radii decreases. This is because the atomic number and the positive charge increase. The increase in atomic number means an increase in both electrons and protons, making the force of attraction much stronger, and decreasing the distance between each other.



When going from up to down in a column, the radii will increase. This is because there will be more orbits taken up by the electrons. The inner electrons also repel the outer electrons, increasing the distance between the outermost electrons and the nucleus.

To recap atomic radii: left to right, decrease; up to down, increase.

Now, ionization energy. First, let's define the term. Ionization energy is the energy required to remove an electron from the neutral atom. Measured in kJ/mol, IE is basically the opposite of the atomic radius. Helium has the highest IE, and Francium has the lowest IE.



Usually, the outermost electron is removed, and so it should always be a valence electron being removed, unless it is a closed shell.

When moving left to right across a period, ionization energy increases. This is because the radii has been decreased, meaning a very strong attraction between the electrons and the nucleons (protons and neutrons). This means that it will be harder to remove an electron.

When moving up to down, the ionization energy decreases due to a larger radius. Attraction between the nucleus and the outer electrons is also decreased because there are more orbits in between blocking the way.

We can refer to removing the first electron as the first ionization energy, and removing the second as second ionization energy, etc.

Number 4, electronegativity. Definition: how much atoms want to gain electrons, or tendency to attract electrons from a neighbouring atom.



This means that if an atom has high electronegativity, it strongly attracts electrons from a neighbouring atom, and could almost "steal" an electron from its neighbour. As well, this means that the atom has a strong attraction with its valence electrons, so the electrons are harder to remove and thus has a higher ionization energy.

In the same way, lower electronegativity = lower ionization energy, because the atom does not have a strong attraction with its own valence electrons, then electrons are easier to remove.

The top right (Fluorine), except for noble gases, has the highest electronegativity.

To recap: left to right, increase in electronegativity; up to down, decrease in electronegativity.

5. Reactivity
In metals:
left to right, decreases
up to down, increases, because it is easier for electrons to be given away, meaning higher reactivity.

In non-metals, left to right, increase
up to down, decreases, because non-metals have higher electronegativity.

6. Ion Charge
The charges depend on its group.
For example,
group 1, +
group 2, 2+
group 13, 3+
group 14, 4+
group 15, 3-
group 16, 2-
group 17, 1-
group 18, 0
In the transition metals, the charges are variable.

7. Melting/boiling point
The noble gases have the lowest melting points, and the elements in the center have the highest. Melting point increases from left to right, except for in the middle.
In metals, going down the group decreases the melting and boiling points.
In non-metals, going down the group increases the melting and boiling points.



8. Density
As you go down a group in the periodic table, density increases. This is because, as you go down, atomic radius increases, and volume increases, so it will be more dense.

In summary,


http://mysite.oswego308.org/schools/uploads/files/3231/ws_periodic_table_and_trends.pdf
http://butane.chem.uiuc.edu/cyerkes/Chem102AEFa07/worksheets/Worksheet%2012.pdf