VSEPR Theory

What does VSEPR Theory stand for?

Valence-Shell Electron-Pair Repulsion Theory

What does it do?

It helps us to predict the geometries of most substances with 3D models.  In other words, it indicates the position of electrons surrounding the nucleus in a molecule.  To do so, we need to first assume that LIKE CHARGES REPEL each other.

How many scenarios are there?

There are three scenarios:

1) all of the electrons pairs in a molecule are shared

2) bonded electron pairs and lone electron pairs are present in the molecule

3) multiple bonds are present in the molecule

Scenario #1: When all of the electrons pairs in a molecule are shared, then...

It can be LINEAR with 180 degree angle between the two identical non-metals.  Example: BeF₂
It can be TRIGONAL PLANAR with 120 degree angles.  Example: SO3
It can be TETRAHEDRAL with 109.5 degree angles.  Example: CH<sub>4


Scenario #2: When bonded electron pairs and lone electron pairs are present in the molecule, then...


A decrease in bond angles will be due to the result of the increased electron-electron repulsion


It can be TRIGONAL PYRAMIDAL with 107.5 degree angles and 1 lone electron pair.  Example:</sub> NH_{3
It can be ANGULAR (BENT) with 118 degree angles and 1 lone electron pair.  Example:} SO2
_{It can be ANGULAR with 104.5 degree angles and 2 lone electron pair.  Example:} H₂O


_{Scenario #3: When multiple bonds are present in the molecule, then...}


There will be double bonds and triple bonds present.  Bond angles will decrease due to the multiple bonds as the electron cloud of a multiple bond takes up more space than single bond electron cloud.


It can be FORMALDEHYDE with 115.8 degree angle and 122.1 degree angle. (well, this scenario do not need to be worried as it is kind of like an enrich)




_{Finally, here is the golden summary of all the VSEPR models!!!}

Note: for our BC chem 11 curriculum, we only need to know: linear, trigonal planar, bent, angular, tetrahedral, and trigonal pyramidal.

Helpful resources and practice question websites:

http://misterguch.brinkster.net/VSEPR.html

http://chemprof.tripod.com/vsepr.htm

http://library.thinkquest.org/C006669/data/Chem/bonding/shapes.html

Here's a video of all the VSEPR model, watch it for... relaxation during study break... :)

Good luck studying! :D

Functional Groups

They are organic compounds made up of elements other than carbon and oxygen

Generally make up the most reactive part of the molecule
Can be either single bonds or group bonds
Single added atoms are generally called HALOGENATED COMPOUNDS, while grouped atoms are NITRO COMPOUNDS.


HALOGENATED:
F: fluoro
Cl: chloro
Br: bromo
I: iodo
General trend:

-Do not dissolve in water
-Hydrocarbons with fluorine are I reactive
-Hydrocarbons with chlorine and bromine are reactive, but only in drastic conditions
-Hydrocarbons with iodine are very reactive

1,3-dibromo-1,1-difluoropropane

NITRO:
NO2: nitro
General trend:
-Usually insoluble in water
-Generally do not react
-Explosive
-Usually have a pleasant odor
**naming**
Follow the same rules as simple hydrocarbons. Can be combined with alkanes, alkenes and alkynes. If there are multiple groups, precede with the prefixes mono-, di-, tri-, tetra-...

2,2-dimethyl-3-nitrobutane

ALCOHOL:
Organic compounds containing hydroxide (OH)
General trend:
-Poisonous to some degree
-While OH is soluble in water, the hydrocarbon chain is not. The longer the chain, the less soluble in water it is.
**naming**
Replace the ending "e" with the suffix "-ol"
If there are multiple, add the previous mentioned prefixes before the "ol" (1 hydroxide bond: propanol. 2 hydroxide bonds: 2,3-propanediol • the "e" remains if there are multiple • identify the placement of the OH with preceding numbers.

Hexanol

ALDEHYDES:
Contains carbonyl (a double group)
Has a double bonded oxygen at the end of the chain
General trend:
-Partially soluble in water
-Very reactive
**naming**
Replace "e" at end with -al suffix
Simplest form: methanal

Ethanal 

KETONES:
Double bonded carbon not at end of the chain
General trend:
-Partially soluble in water
-Unreactive
**naming**
Replace e ending with -one suffix

Propanone

CARBOXYLIC ACIDS:
Has a double bonded oxygen and hydroxide at end of chain.
General trend:
-Can be neutralized when placed with an acid.
**naming**
Drop the "e" and replace with -oic acid
Simplest form: methanoic acid

4-bromo-2fluorobutanoic acid

ESTERS:
oxygen bond dividing an acyl from an alkyl. Has a double bonded oxygen attached to closest carbon to oxygen
General trend:
-Pleasant smell
-Can be converted back into alcohol and carboxylic acid
**naming**
Alkyl: name using same rules
Acyl: (the remainder) replace -oic acid suffix of the corresponding carboxylic acid with -oate

Propyl Heptanoate

ETHERS:
Oxygen group connecting two alkyl groups.
General trend:
-Highly flammable
-Insoluble in water
-Makes a good solvent for organize compounds
-Certain compounds (ethoxy ethane) have anesthetic properties
**naming**
O is where side group counting begins
Follow all basic rules, but replace -yl with -oxy to the side group names

1-bromo-2-ethoxy hexane

AMINES:
Contains nitrogen compounds bonded to either hydrogen or carbon
General trend:

-Organic bases ➡ form salts easily when reacted with water
-Closely related to NH3
-Fishy odor
**naming**
Follow standard rules, but the longest chain name is preceded by amino.
As always, prefix numbers are used to identify the location of the amine.

ALCYLICS

Can form cycloalkanes, cycloalkenes and cycloalkynes. Forms a ring with carbons, either a branch or a main chain. Will also support other branches

General trend:

-Less stable (more reactive)

-can be branched

**naming**

numbering can begin anywhere, and go any direction

lowest chains possible

1. count total carbons, and name accordingly. Add cyclo prefix

2. if there is only one side group, it is assumed to be one, and no number is needed

3. If there are multiple, name the first group with number 1

4. side groups are named following basic rules

5. for alkenes, alkynes the bond is assumed to be the first, and no number is needed unless there are multiple

6. if the side groups are tied, give the lowest number to the first alphabetically ordered group

1-ethyl-2-methylcyclohexene

AROMATICS

contains a benzene ring

General Trend:

-Electrons are "delocalized" meaning that they can change position through the ring

-All of the carbon bonds therefore have the same reactivity

-Less reactive that cycloalkenes/ynes

-Because of the delocalization, benzene is very stable

Has a special diagram because of the free movement of electrons

Benzene

**naming**

When it is the main chain, just name the side groups attached to it, followed by benzene

If it makes up a side group, then it is called phenyl.

4-bromo-1-chloro2-ethyl benzene

Still unclear? Watch this YouTube video

Alkenes and Alkynes

Alkenes = double bonds
Alkynes = triple bonds

Alkenes
-simply hydrocarbons with one or more double bonds located between carbon atoms leading to an unsaturated hydrocarbon

ending: change '-ane' to '-ene'

general formula: CnH2n

Geometric Isomers

-same chemitrical formula different geometry

-either cis or trans

larger group are above or below = cis

larger group are across the plane of the bond = trans

if no geometric isomers no need for cis or trans

Alkynes

-triple bonds

ending: change '-ane' to '-yne'

Eg.

4-methyl-2-pentyne

Organic Chemistry

The chemistry of CARBON compounds
PROPERTIES
-Low melting points
-weak or non-electrolytes
-can form chains of carbon atoms that are linked in a:

• straight line
• circular pattern
• branched pattern

-can link with other atoms in:

• single bonds
• double bonds
• triple bonds
• versatility of organic compounds makes it such an important branch of chemistry

Alkanes

-straight/unbranched chain

hydrocarbon: a compound that contains only hydrogen and carbon. 

-non-polar molecules ==> immiscible with water

-Geometry: Tetrahedron

-saturated hydrocarbons bonded by single bonds

• Saturated: not possible for another atom to bond to the structure

-naming: "-ane" endings because they are Alkanes

homologous series: a series of organic compounds with a similar general formula, possessing similar chemical properties

meth = 1

Eth = 2

Prop = 3

But = 4

Branched Hydrocarbons

-"side branches" = hydro chains

• called substituted hydrocarbons or branched hydrocarbons

ending: "-yl"

-Alkyl group: an alkane which LOST one hydrogen atom

Ex 2-methylpentane

Rules for naming:

1) find and name the longest continuous carbon chain and place at the end of the name

2) Identify and name groups attached to this chain

3) number the chain consecutively, starting at the end nearest a side group

4) designate the location of each side group by an appropriate #

5) Assemble the name in alphabetical order

Electronegativity and Polarity

Hello all, and welcome to one of the less interesting blogs that we have to offer you.  Unfortunately it is important information so we'll try to make it as short and sweet as possible.  This is an electronegativity chart showing the charges for how easy/hard it is for each of these elements to loose an electron(s).  For instance, elements like Cesium is much more likely to give away its electron to another element (such as fluoride).  Make sense so far?  I hope so.














These charges can be used to determine what kind of bond you are looking at.  So, for example, you were looking at what kind of bond sodium and chlorine would make.  You would find their electronegativity difference.  


Na= 0.9                                              3.0-0.9= 2.1
Cl= 3.0
(if you get a negative answer, then think of what you have as having absolute value bars around it, just make it positive)
Based on this chart...and on prior knowledge, we know that this is an ionic bond.  However, when you do the same process with say, Nickel and oxygen, you will see that it instead comes up as a covalent bond.  That is something that will be explained more next year.

Here is the promised video to help keep the difference between the three rememberable.  If you can't deal with bad corny science songs, then don't click play.  Otherwise, I hope it helps!
http://www.youtube.com/watch?v=oNBzyM6TcK8

Part 2: knowing how to draw these.

Luckily, for the most parts, it is drawn as a regular lewis dot structure; which you already know how to do right? Right.  The only difference, is you need to indicate which atoms would become positive (loose electrons)  and which would be come negative (gain electrons)  Then draw an arrow pointing toward the negative one.  Have no idea what is being said, well, here is a nice long video that takes you back to periodic trends and explains it a bit more.

Best of luck :)

Chemical Bonding

Here starts the tedious introduction blog of the last unit we will do this year.  Many things that are mentioned here will come up later so if you don't understand this one....take the time to figure it out.

Electrostatic Force

Is a force existing as a result of the attraction or repulsion (two positive or two negative particles repel each other but one of each attract each other) between two charged particles.  It is a important to remember all bonds are based on the electrostatic relationships formed between particles.  When these particles are farther from each other, the attraction is greater.  If there is more charge, then there is more force.  This lovely force that occurs acts in all directions which is why that in both theory and practice, bonds can be made on any electron in the valance shell.  

Ionic Bonds

Ionic bonds are formed between positive and negative ions (most often medal and non-medal).  In an ionic bond, the electrons are transferred from what atom to another.  If you need a good way to remember that, just look at the chemistry cat on the side of our blog.  These bonds are always made to have a closed shell.  Ionic bonds are:

• very strong
• have high boiling points
• high electronegativity
• ionization energy

After the bond, all atoms become ions.  The positive ions are called cations and the negative ions are called anions.  

Then there is the crystal lattice...

In a crystal lattice, the ions in the bonds line up so the positive side of one molecule line up with the negative side of the next.  This is what allows them to connect and make things such as table salt. 

[Non-Polar] Covalent Bonding

These bonds, unlike ionic bonds, share their electrons amongst that atoms in the molecule.  As a result of this, the electrons are stuck in the middle between the two atoms.  These bonds tend to be very strong with lower melting points.  

Individual Molecules

Individual molecules are actually quite strong.  During melting, it is not the molecules themselves breaking apart, but the different molecules breaking apart from each other.  They are held together only by the intermolecular force.  This is especially true for covalent bonds.

The next blog will have a video that will help to remember the differences between the three.  (Yes, there is one more; covalent bonds aren't always non-polar...)

A Brief History of the Atom

Democritus: Everything is made up of infinitely small particles. Titled it atomos, meaning indivisible.

Dalton: Created modern atomic theory. All atoms are hooked together, and form specific ratios. Did not realize that atoms are made up of different particles

Thomson: Discovered electron with cathode ray tube. Applied a magnetic force, and noticed that the ray deviated, due to opposing magnetism. Also created a new idea of appearance, titled plum pudding formation.

Rutherford: Fired alpha particles through gold, some rays deviated from center, while others bounced right back. This was because of the positively charged nucleus. He also realized that a lot of the volume of the atom is empty space.

Bohr: Worked in Rutherford's lab, and realized if the charge on the inside was positive, that the negative and positive charges would attract. Created the Bohr model on the assumption that electrons followed specific paths.

Schroedinger: claimed that electrons weren't particles but were in fact waves.

Chadwick: Discovered the neutron by disintegrating atoms through bombardment of  alpha particles. First to create artificial nuclear reaction
Heisenberg: Confirmed both of Bohr and Schroedinger's hypothesis. The electron is indeed a particle, that travels in a wave like motion, which is described by quantum theory.

History of the Periodic Table

Since the 5th century, a few elements had been discovered, including gold, tin, silver, lead and mercury, as these were easy to find. The first discovery was by Henning Brand, a German man, on a wild chase for the philosophers stone. In 1669, he discovered phosphorus, but kept it a secret, and the rock was later publicized in 1680 by Robert Boyle.
A hundred years later, a chemist named Antoine Lavoisier lived. He created a text book of elements that could not be broken down, which also included light, and caloric, which were then believed to be material substances.
Lavoisier's elements

Gases New names (French) Old names (English translation) Lumière Light Calorique Heat Principle of heat Igneous fluid Fire Matter of fire and of heat Oxygène Dephlogisticated air Empyreal air Vital air Base of vital air Azote Phlogisticated gas Mephitis Base of mephitis Hydrogène Inflammable air or gas Base of inflammable air

Metals New names (French) Old names (English translation) Antimoine Antimony Argent Silver Arsenic Arsenic Bismuth Bismuth Cobolt Cobalt Cuivre Copper Étain Tin Fer Iron Manganèse Manganese Mercure Mercury Molybdène Molybdena Nickel Nickel Or Gold Platine Platina Plomb Lead Tungstène Tungsten Zinc Zinc

Nonmetals New names (French) Old names (English translation) Soufre Sulphur Phosphore Phosphorus Carbone Pure charcoal Radical muriatique Unknown Radical fluorique Unknown Radical boracique Unknown

Earths New names (French) Old names (English translation) Chaux Chalk, calcareous earth Magnésie Magnesia, base of Epsom salt Baryte Barote, or heavy earth Alumine Clay, earth of alum, base of alum Silice Siliceous earth, vitrifiable earth

 By 1809, 47 elements had been discovered. Chemists noted pattern in reactions, and attempted to classify them. 

In 1862, the first form of element organization is published by a French chemist, Alexandre-Émile Béguyer de Chancourtois, that appeared to be a spiral, with increasing mass downward.

1864 John Newland proposes the Law of Octaves, which he found by assigning a mass of 1 to helium, and then ordering the rest by mass. This lead to the discovery that every 8th element had similar properties, 

in 1869, Mendeleev published the first official periodic table. He ignored certain previous ideas, and instead grouped certain elements together according to properties, vs the previous method of weight organization. Mendeleev predicted atomic weights of yet to be discovered elements, and left empty sections in his table to accommodate them

In 1911, Ernest Rutherford publishes a paper that explained nuclear charges. Later that year, Antonius Van Den Broek publishes a paper relating the atomic weight of an element to the charge of the atom. This became the basis of organization of the table, the atomic number.

Two years later, Henry Moseley proposes that the wavelengths of x-ray emissions of elements is directly proportional to that of the atomic number.

The latest changes made were by Glen Seaborg, who discovered plutonium, and elements 94-102. It was his doing that resulted in the lanthanide and actanide series' being placed below the table.

Periodic Table Trends

The periodic table is not organized randomly, it follows certain trends.

There are some trends we need to know.

Here is a youtube video that pretty much covers EVERY trends we need to know.

Below are the trends summarized in point form :D

1) Metallic properties

• From left to right across the table, the elements change from metal to non-metal
• From top to bottom down the table, elements become more metallic, or better metals

2) Atomic radius

• From left to right across a row, the radius of atom decreases because increase in number of protons brings the electrons closer to the nucleus
• From top to bottom down a group, the radius increases

3) Atomic size

• From left to right across a period, the size decreases
• From top to bottom down a group, the size increases

4) Reactivity

• For metal:
• From top to bottom down a group, the elements become more reactive
• in transitional metals, the middle part of the table is the least reactive and to the left or to the right become more reactive
• the most reactive metal is Francium
• For non-metal (excluding noble gas):
• From top to bottom down a group, the elements become less reactive

5) Ion charge

• Metals tend to have positive charges
• Non-metals tend to have negative charges
• the transitional metals have variable charge

6) Melting & boiling point

• elements in the centre of the table have the highest melting and boiling point
• Noble gases have the lowest M & B point
• from left to the right across the table, the M & B point increases until the middle of the table

7) Ionization energy

• this is the energy required to completely remove an electron from an atom
• From bottom to up, left to right cross the table, the energy increases
• All noble gases have high ionization energy
• Helium has the highest ionization energy while Francium has the lowest
• Note: this trend is opposite from atomic radius trend

8) Electronegativity

• Electronegativity is basically how much atoms want to gain electrons
• Note: this trend is the same as the trend for ionization energy
• Noble gases are excluded from the trend as they already have full shells (so they don't want to gain or loss electrons)

Writing Electronic Configurations

Neutral Ions

-always start with the lowest level first (Aufbau Principle)

-Figure out how many electrons you have (Neutral atom = Atomic #)

-start with the lowest energy level (1s) and keep adding until you have no electrons left

-Each electron has an opposite spin designated 

Ex. Sillicon: 14 electrons

2 electrons in 1s, 2 electrons in 2s, 6 electrons in 2p, 2 electrons in 3s, 2 electrons in 3p

A-E shows the steps taken to write the electron configuration

The 2 electrons in the 3p subshell aren't paired because of Hund's rule: when electrons occupy orbitals of the same energy, they can't be paired up until they have to.

written as: 1s²2s²2p⁶3s²3p²

Ions

negative charged ions: add electrons to the original number of electrons according to its charge

ex. P3- : 15+3

15 being its original # of electrons and 3 added because of its charge

total electrons = 18

written as: 1s²2s²2p⁶3s²3p³

positive charged ions: remove electrons from its original number of electrons according to its charge

ex. Ba2+: 56-2

total electrons: 54

written as: 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s^24d105p6

Core Notation

-set of electrons that can be divided into two subsets: core electrons and outer electrons

core: set of electrons with the configuration of the nearest noble gas that comes before it

outer: consists of all electrons outside the core (normally takes part in chemical reactions)

ex. Ca

1)the noble gas before calcium is Argon, so you put argon in square brackets

[Ar]

2)then add the remaining electrons

[Ar]4s²

2 EXCEPTIONS: Copper (Cu) and Chromium (Cr)

Cu:  [Ar]4s3d_­¹⁰

^{Cr: [Ar]4s3d5}