On October the Third: We learned about Ionic and Covalent bonding and the Ionic bonding Lewis dot diagram format.
By Nicole


A Step of the Lewis dot diagram of an ionic compound.

Ionic Compounds!

  • Are made of ions
That is why it is called an ionic compound!
Ions are when atoms have lost or gained electrons.
  • Are made of a non-metal and a metal
Non-metals have a strong enough electronegativity to take a metal’s valence electrons.
This means: Metals lose electrons
Non-Metals gain electrons.
  • When atoms lose or gain electrons they obtain a charge in the process.
This makes the atoms either positive (a cation) or negative (an anion)
Cations and Anions act similarly to magnets and attract one another

Covalent Compounds!

  • Are made of atoms
More specifically non-metals and non-metals
  • Valence electrons are shared
Non-metals have similar electronegativities, meaning they are not strong enough to pull another non-metal’s valence electrons away, thus they are shared.

The number of valence electrons an atom has determines the number of bonds it can make.
Atoms that have 1-4 valence electrons can bond with up to as many bonds as there are electrons.
Atoms that have 5-8 valence electrons can bond to as many atoms as the difference of 8 - the number of valence electrons it has.
(i.e. [ 8 - No. of valence electrons = # of possible bonds])


This is the format for a Lweis dot diagram for ionic compounds.
First you draw a lewis dot diagram of each atom, indicating where electrons are going.
Then you show the ions in square brackets, with their ionic valence shell, if the atom has lost all of its electrons you do not do the electrons in the "new" valence shell.
Make sure the charges are written outside of the brackets.


Homework: page 60 questions numbered 1-8
front of the handout provided

wUnit 1 Daily Notes Continued

By: Pratheep Rajan

[[#|Date]]: Thursday, October 4, 2012

Covalent Bonding-Sharing Electrons

  • What are covalent bonds?
  • Properties of covalent substances
  • Multiple bonds
  • Need & Have Method

Covalent Bonds

Certain elements do not like loosing their [[#|electrons]] so they interact with each other in a different way to [[#|form]] a bond called Covalent Bonding.
In theses bonds instead of [[#|electrons]] being transferred between the elements to form [[#|ions]], the electrons are shared by the elements.
These bonds occur when a non-metal [[#|combines]] with another non-metal and no ions are created.
These elements share electrons so that they have eight outer electrons which include their own and the electrons that are being shared.

Properties of Covalent Substances

  • Tend to form groups of atoms known as [[#|molecules]]
  • Can exist as solids, gases or liquids at room [[#|temperature]]
  • The solids created are usually soft and waxy
  • The solids and liquids are generally volatile( evaporate rapidly or pass of in the form of vapour)
  • They have low melting and boiling points because the bonds are weaker than ionic bonds
  • These substances have low solubility in water and polar solvents( solvents with polar bonds in them), but are soluble in non-polar solvents
  • The liquids and solids do not conduct electricity

Multiple Bonds

It is likely for atoms to share more than one pair of electrons in a bond.
Up to fours pairs of electrons can be shared so that each atom has a full valence orbital.
  1. A single pair of electrons shared [[#|creates]] a single covalent bond.
  2. A double pair of electrons shared creates a double covalent bond.
  3. A triple pair of electrons shared creates a triple covalent bond.
  4. A quadruple pair of electrons shared between creates and quadruple covalent bond.

Covalent bonds can be represented in Lewis Dot Diagrams like these which show each bond:

NH3-Ammonia | HF-Hydrogen Flouride | O2-Oxygen gas
external image edotnh3a.gif external image edothf.gif external image edoto2.gif

Need & Have Method

This method will [[#|help]] you figure out how many bonds are in a covalent bond so that you can draw a Lewis Dot [[#|Diagram]] for it.
It involves these four steps:
Example F2

Step 1. Find the amount of valence electrons that are existent
Have: F + F = 7 electrons + 7 electrons= 14

Step 2. Figure out how many electrons are needed to make a neutral bond
Need: Every atom needs 8 electrons in the valence orbital(except Hyrdogen)
2 Flourine atoms = 2 * 8 = 16 electrons needed

Step 3. To establish the amount of bonds take the difference of steps one and two and divide it by two
Number of Bonds= (Need-Have)/2

Step 4. Draw the Lewis Dot [[#|Diagram]] like this:

[[#|Home Work]]:

Chemical Bonding Worksheet : Covalent Bonding. Do Lewis Dot [[#|Diagrams]]

Important Pages and Sheets:
Read pages 64 - 67 for examples of Covalent Structures
Work sheet; http://sch3uking.wikispaces.com/file/view/Covalent+BondingHandout.pdf

By: Pratheep Rajan

[[#|Date]]: Friday, October 5,2012

Covalent Bonds Continued

  • Coordinate Covalent Bonds
  • Resonance Structures
  • Polyatomic Ions
  • Non-Octet rule for molecular compounds

[[#|Coordinate]] Covalent Bonds:

In most cases of covalent bonding each atom provides a electron to share with another atom so that they can have a full octet.
This is not always the case though as a special type of covalent bond called the coordinate covalent bond exist.
This bond is special because one of the atom provides one whole pair of electrons to be shared( 2 electrons) while the other atom provides no electrons at all.

Resonance Structures

Sometimes the positions of a double bond can change when the central atom is bonded to more than one of the same element.
This allows the double bond to "wander" and be in any of the bonds between the same elements and the electrons are "de-localized".
This then allows each bond to have a bond strength of 1 1/3 which is determined by dividing the number of bond with the available spots for the bonds.( Bond strength=# of bonds / spots for bonds)

They are represented like this showing each bond(dots are not needed):
Each bond is represented by lines ( 2 electrons being shared)
O3-Ozone | C6H6-Benzene

external image images?q=tbn:ANd9GcQOxwUXGpRBXY09GY3tgCMDx1BBgKyXsHmel51o8ncYtaF5Dq-xuw external image images?q=tbn:ANd9GcToLyNUM3TM_cvae1lLkA9NzkzK1qyIFnbpR9j67Ht-UfxRfSSBSA

Polyatomic Ions:

To represent these bonds use these steps.

Step 1. Use the need have method to determine the amount of bonds, but this time for the Have also factor in the charge on the polyatomic ions as electrons(if the charge is 2- add 2 electrons to the Have, if it is 1+ subtract a electron)

Step 2. Draw the Lewis Dot Diagrams following the full octet rule.

Step 3. Put square brackets around the whole diagram and show the charge.

Some example of these diagrams are shown below:

NH4- Ammonium | SO4- Sulphate(ignore the Xs and think of them as dots, they represent the electrons from the charge)

.external image covale9.gifexternal image lewis_dot_16.jpg

Non-Octet Rule for Molecular Compounds:

There are many molecular compounds that do not follow the full-octet rule.
The non-octet rule is used when the central atom( always the element with the lowest electronegativity(EN)) can have either more or less than 8 electrons in the valence orbital.
It is always the central atom that violates the regular full octet rule, while the surrounding atoms still follow the regular octet rule.
In these bonds there will be no coordinate covalent bonds as well as multiple bonds.
The central atom will bond with each and every outer atom once and if there are any left over electrons from after the Need & Method you will add them to the central atom.

You can determine the bonds that do not follow the full octet rule by looking for these signs:
  1. Boron (B)
  2. Noble Gases Ex. XeF4
  3. Two Halogens that do not have a equal ration Ex. FCl7
  4. When there are four or more of an atom Ex. PBr5

To draw this follow these steps:

Step 1. Use the need have method to determine the amount of bonds.If there is a charge add that in aswell

Step 2. Draw the Lewis Dot Diagrams following the full octet rule for the outer electrons.

Step 3. Add up all of your electrons(in your drawing) if there are less electrons in your diagram compared to the "Have", add electrons to the central atom until both figures are equal. If the number of electrons are the same in your diagram and your "Have" leave it as is.

Some example of compounds that follow the non-octet rule are(use dots to represent the electrons instead of the lines):
BCL3- Boron TriChloride


PF5- Phosphorous PentaFlouride

Chemical Bonding Worksheet : Non-Octet Rule Covalent Compunds. Do Lewis Dot [[#|Diagrams]]

Important Pages and Sheets:
Work sheet; http://sch3uking.wikispaces.com/file/view/Covalent+BondingHandout.pdf

Tuesday, October 9th

By: Jennifer Lindeman

Ionic compounds form repeating units.
Covalent compounds form distinct molecules.

Polar Covalent Bonds

Sometimes when atoms of two different elements form a bond by sharing an electron pair there is unequal sharing of electrons. when this occurs the bond is called a POLAR BOND.
The unequal sharing results from the difference in electronegativity of the two atoms. The one with the greater electronegativity exerts a greater attraction for the electrons (making it slightly more positive).
This is shown with a symbol on each element, stating whether it is slightly positive or negative.


Recall that electronegativity is "a number that describes the relative ability of an atom, when bonded, to attract electrons." The periodic table has electronegativity values.
To find the value on the periodic table look on the key, and look at the top right number in black.
We can determine the nature of a bond based on /\EN (electronegativity difference). /\EN=higher EN-lower EN
Example: NBr3
N=3.04 Br=2.96
3.04-2.96=0.08 /\EN *Nitrogen has a higher EN, so it is slightly more negative then Bromine with a lower EN*

/\EN below 0.4= Pure covalent (equal sharing)
/\EN between o.4-1.7= Polar covalent (1.5 is more polar covalent than 0.5, because it is higher)
/\EN above 1.7= Ionic (ALL metal+non-metal is ionic, no matter the /\EN)

Homework: page 73 # 4-7 + the sheet assigned in class (bottom of polar sheet)

Wednesday, October 10th

By: Jennifer Lindeman

Inter and Intramolecular Forces

Intramolecular Forces

Forces of electrostatic attraction within a molecule. Occur between the nuclei of the atoms and their electrons making up the molecule (i.e. covalent bonds). Must be broken by chemical means. Form new substances when broken.

Intermolecular Forces

Forces of attraction between two molecules (i.e. London dispersion forces, dipole-dipole interactions or hydrogen bonds). These forces are much weaker then Intramolecular forces or bands and are much easier to break. Physical changes (changes of state) break or weaken these forces. Do not form new substances when broken. These forces affect the melting and boiling points of substances, the capillary action and surface tension, s well as the volatility and solubility of substances.
Capillary: The ability for a substance to move upward (water up a tree)
Volatility: Ability to change into a gas

Types of intermolecular forces (only for molecules/covalent compounds

London dispersion Forces (In ALL molecules)

- Result form a type of tiny dipole (2 different charged ends).
- exist between all molecules
- masked my stronger forces (dipole-dipole)
- important in all non-polar molecules
- Greater number of electrons, greater the force (Molecule size increases, so does the force. EX: A has a higher melting point than B because A has more electrons/is bigger)
- Van der waal force

Dipole-dipole Interactions

- Unequal distribution causing two different charged ends
- Van der waal as well

Negative and positive ends attract

Hydrogen Bonds

- Type of dipole-dipoe interaction
- Occur between Hydrogen and a highly electronegative atom, mainly [F, O, and N]
- Negatives attach to positives and positives attach to negatives
Solid lines show the intermolecular bonds. Broken lines show the Intramoleculat bonds


Ionic Forces

- Ionic may be both inter and intra since a crystalline lattice is formed
- Ionic is stronger because you have to break that bond, while covalent you just have to weaken it
Melt salt= break ionic bonds
Melt sugar= weaken covalent bonds

HOMEWORK: Read pg 102, 104-105
pg 108 #1, 4, 5 (with pictures showing all bonds/forces)

Thursday, October 11th

By: Jennifer Lindeman

Polar and Non-polar Molecules

- Polar molecules have a negative and a positive side (HCl on the diagram 3)
- H20 is polar because it is drawn on an angle, it is bent!

- Non-polar molecules do not have one side either positive or negative
- Either all the positive or negative atoms are on the outside (charges[arrows] go from positive to negative)

Examples of polar and non-polar diagrams with arrows


Drawing 3D diagrams

Looking straight onto a 3D model, draw the center element (carbon) and the elements on the sides (chlorine)
The chlorine behind/further back is drawn with dashes
The chlorine in front/closer is drawn with a triangle/rectangular shape

Homework: Finish the whole sheet from class today

MODEL BUILDING- Friday October 12, 2012

By: Siham Sirage

  • Continuation of worksheet and building 3D models of different bonds and drawing diagrams to go along with the models
  • When activity sheet was completed, "Testing concepts" sheet was to be worked on

Example 3D Structure Drawing and ball model for CH4 (Methane)
external image 3D-Molecular-Structure_of_Methane.jpg
The triangular wedge represents the hydrogen bond in front and the dotted line represents the
hydrogen bond at the back. The two straight lines represent the bonds at top and side of carbon.

  • "Testing concepts" worksheet is to be completed for homework (# 1-11)
  • Chemistry textbook is needed to complete question 11 of the worksheet
  • Do not work on the back side of the worksheet as it is to be completed during mondays computer lab

MOLECULE POLARITY PhET LAB- Monday October 15, 2012

By: Siham Sirage

  • Worked in the computer lab on molecule polarity worksheet given on friday
  • The worksheet given focused on electronegativity, bond polarity, and molecular polarity

Heres a link to the website if you have not finished the worksheet during the class time provided -

external image ElectronegativityTrends.gif
This periodic table shows the trends in electronegativity, the electronegativity increases as you move across from left to right, it also increases as you move up the periodic table.

external image bonding_fig03.gif
In this picture of Hydrogen Bromide, bromide has a greater electronegativity that makes it slightly negative with more electrons to itself when compared to hydrogen with a lower electronegativity and a slightly positive charge. This is an example of a POLAR bond.
  • If you did not finish the worksheet in the time provided in the computer lab, you are to finish that for homework as well as the "testing concepts" questions on the back.

October 16 2012
By: Luxena Sribaskaran

  • Discussed how to determine relative boiling and melting points
  • I Love water I love Oil- Lab


  • Melting points and boiling points are determined by looking at the intramolecular forces. Ionic compounds have HIGH boiling and melting points because the physical ionic bond is being disrupted. Ionic compounds include a metal and a non-metal.
  • For two different ionic compounds, to determine which has the higher boiling point you look at the charges; the one with the higher charges has the higher boiling point.
  • Covalent compounds include two non-metals.
  • TO determine which covalent compound has the higher boiling point, look at the strengths f the intramolecular forces between the bonds ( London dispersion forces, Dipole Dipole, Hydrogen bonds)
  • London dispersion is the weakest
  • Dipole Dipole is in the middle
  • Hydrogen bonds are the strongest
  • ** For two molecules that only have London dispersion forces, look at the total number of electrons; the one with more electrons, has a higher boiling point. This is because they have more electrons which result in stronger London dispersion forces.

This is an example if boiling points increasing in various compounds.
external image HydCmpd_bp.jpg

  • For more information on boiling points refer to pages 110, 111 and 116 of your textbook.

    I Love water I love Oil- Lab
Here is a link to the Lab sheet.