Functioning Groups

Functioning groups are carbon compounds that contain other atoms besides just carbon and hydrogen.
They are very specific and each one is different with its own properties and reactivity.

The main Functioning Groups being focused on are
1) Halides and Nitro Compounds
2)Aldehydes and Ketones
3) Alcohols

 Halides and Nitro Compounds

How to name Halides:
Remember these prefixes.
Br-bromo
I-iodo
Cl-chloro
F-fluoro

For nitro compounds-

NO2-nitro

The same rules as naming alkanes apply.
In alphabetical order, state where its attached to the main parent. (e.g. 4) Then use the prefix.
Do this with as many as needed. Then finish with the length of the parent hydrocarbon.

E.g.

Naming: This has a parent hydrocarbon of three. Which makes it prop-. Because there are only single bonds, the ending is -ane(PROPANE) NO2 is present so we name it nitro and then place the position of the nitro in front of the word nitro.( in this case, 2)
So this is called 2-nitropropane.



Properties of halides:
1. Not soluble in water
2. I- compounds are more reactive
3. F-compounds unreactive
4. Cl and Br compounds vulnerable to chemical attack

Nitro compounds are nice to smell, not vulnerable to attack(unless very drastic conditions), water soluble and explosive!

Aldehydes and Ketones

Naming: Remember than an aldehyde always ends with -al and a ketone always ends with -one.
In an aldehyde the double bond with oxygen will always be on the last carbon in the chain.
In ketones it isn't.
It is named the same way all the others are named except with the new endings!!!!!

Alcohols
Alcohols are a functional group containing (OH) hydroxyl group.
NAMING:
First name the parent hydrocarbon group with ending-ol.
Then place the position of the OH group(the shortest possible) in front. Then do the rest how it is normally named with the alkyl groups.

First, name the parent hydrocarbon. So.. pentanol. Then check the position of the OH. In this case it is 3.

This is called 3-pentanol.

Note: OH makes alcohols soluble in water. The alcohols are all poisonous! Methanol, propanol and ethanol are highly soluble.

http://www.youtube.com/watch?v=tyMD_1gBLmg

http://www.youtube.com/watch?v=XHn0lLH9SSY&feature=related

We found these helpful.!!!! ENJOY

Alkenes and Alkanes

Today we learned about alkenes and alkynes.

They are different from alkanes because they can form double(alkene) and triple(alkyne) bonds.

Alkenes are formed with double bonds between carbon atoms and the ending is -ene.

Example.

Let us name this.
If this was an alkane, it would end in -ane and its name would be 2-methylpentane.
However it is an alkene as evidenced by the double bond.

Therefore the name is 4-methyl
Then you write the number of the carbon# the double bond comes after so 2-pent-ene

so the name of this compound is 4-methyl-2-pentene.

Alkynes are exactly the same as alkenes except when you name it, change the ending to -yne. They use triple bonds.

Name this.

First count the parent hydrocarbon.
=4 so Butane.
Then you take into account the triple bond so the name changes to butyne.

Since there are no branched hydrocarbons, nothing needs to be added except the position of the triple bond.
This is added in front of  the word butyne

It is 2-butyne.

http://www.youtube.com/watch?v=7XbYhjyUI-M

Organic Chemistry

Organic Chemistry is the study of compounds containing carbon. These compounds can be formed in several ways, and may appear in a straight line, circular or branched pattern. They also have low melting points. To combine with other atoms they can form single,double or triple bonds.

 A hydrocarbon is a compound containing only carbon and hydrogen.
Alkane: A saturated hydrocarbon involving only single bonds.

Here are some examples of alkanes:

The chemical formula for alkanes is CNH2N+2

Methane: CH4
Ethane: C2H6
Propane : C3H8
Butane: C4H10

* An alkyl group is an alkane that loses an atom of hydrogen.

A branched hydrocarbon is a hydrocarbon attached to the original carbon atom.

Chemical Bonding

Chemical bonding is molecules and atoms being attracted to each other and this allows them to combine to form chemical compounds.

There are 3 types of bonding:

Ionic Bonding: Is the transfer of atoms to form positive or negative ions.
Non-Polar Covalent Bonding- Is the unequal sharing of electrons
Polar Covalent Bonding : sharing of electrons equally.

IONIC:

Ionic bonding involves metal and non-metals. The force that powers ionic bonding is called the electrostatic force, which is the attraction or repulsion of charged particles. Electronegativity is also an important concept behind this bonding. Metal electronegativity is low and non metal electronegativity is high. As result, a high ionization energy is created. To determine how electrons are shared, electronegativity difference must be found.

ENeg Diff= ENeg 1 - ENeg 2
ENeg Diff  less than 0.5 it is a non polar
ENeg Diff greater than 0.5 and less than 1.8 is is polar covalent bond
 ENeg Diff greater than 1.8 is it ionic

Polarity: is the electrical balance or imbalance of the molecule. If the molecule is electrically unbalanced, then it is polar. This is caused by electronegativity. If the EN is high, the shared electrons will be pulled more to the atom. A partial negative results from high electronegativity and a partial positive results from a  low one.

In a Non- polar bond, it is one of the simplest, only involving the sharing of electrons to create full valence shells. Intramolecular forces in the molecule hold the atoms of the molecule together.
However, intermolecular forces act between molecules and are responsible for bonding.

A London force is a weak intermolecular force. They occur because of temporary dipolar attractions. A dipole is when one side is slightly positive and the other side is slightly negative ,creating partial separation.

Drawing Electron Dot Diagrams and Lewis Dot Diagrams

Let us examine an example of an electron dot or Lewis diagram.

1. The centre of the diagram, the atomic symbol represents the nucleus.
2. Each dot represents an electron in an electron dot diagram. These are only the valence electrons, or the electrons in the outermost shell.
3. In a Lewis diagram, lines represent a  bond between two electrons.

Here is an example. Because Nitrogen has 5 valence electrons and hydrogen has 1. Hydrogen needs 1 electron to become full and nitrogen needs 3.
Therefore, in NH3, N and H share electrons so that H is full with 2. and N, being surrounded on 3 sides gets 3 electrons.


In a structural diagram, electron bonds can be represented with a ---(line)


In Ionic compounds:

As you can see here, sodium has 1 valence electron and chlorine has 7.
Sodium and chlorine both want to become happy, and they are attracted to each other. The sodium loses its valence electron and becomes (+)1 charge. Chlorine takes this electron and fills it valence shell, giving it a (-1) charge.  Therefore, this is how the Lewis diagram is shown.



Extra help!

http://www.youtube.com/watch?v=y6QZRBIO0-o
http://www.youtube.com/watch?v=QKoA3fZ29B0

Periodic Trends

As we examine the periodic table each elements have certain characteristics and looking at the periodic table as a whole, some trends are noticed.

Recall that Dmitri Mendeleev created a periodic table of the 63 known elements which he arranged by atomic mass.

First we must know some important vocab.

Density: of an object is its mass per unit volume. It increases as the atomic numbers increase.
Ionization energy: is the energy needed to remove electrons from an atom. It increases as you go across a period(left----> right) and decreases as you go down in families.

Melting and Boiling Points: This is the temperature at which an object liquefies and  becomes a gas respectively.
These properties increase as you go from bottom to top in a group.

Electronegativity: This is the tendency of an atom or functional group to attract electrons towards itself(form negative ions) In increases left------------->right

Atomic Radius: Is the measure of the size of an atom. It decreases left to right and increases going down.

Valence Electrons

Valence electrons are the electrons that exist on the outermost shell of an atom.
Are the ones involved in chemical reactions

An open shell is a shell that has less than the maximum amount of electrons in it.
A closed shell is a shell that is completely full with the max. amount of electrons that can occupy it.

How to determine the number of valence electrons from electronic configurations:

Sodium: has 11 electrons.

1s^2 2s^2 2p^6 3s^1

Now we must put it into core notation:
[Ne] 3s^1

When counting valence electrons, only the s and p subshells are considered, unless the d and f subshells are full.

Therefore, Na has one valence electron.

http://www.youtube.com/watch?v=1TZA171yxY4&feature=fvwrel

Cheers

Electrons in an Atom

Niels Bohr proposed that atoms exist in energy levels surrounding the nucleus.

Electrons exist in two states:
1. Ground state- when electrons are in their lowest possible energy level
2-Excited State-when one or more electrons of an atom are in energy levels above their lowest possible level.

An energy level is the amount of energy an electron can possess.
The quantum of energy is the difference in energy between two levels.

Recall that electrons exist in shells,which are sets of orbitals having the same 'n' value
and an orbital is  the space an electron takes up in a particular energy level.
A subshell is a set of orbitals of the same type.

There are 4 different types of orbitals
S, D, P and F.
Each subshell has

1 S Orbital
3 P orbitals
5 D orbitals
7 F orbitals

This means that the most electrons you can put in

S subshell- 2
P subshell- 6
D subshell-10
F subshell-14

This is governed by Pauli's Exclusion Principle which states that each orbital contains 2 electrons max.

This is a diagram which will help when filling orbitals.

To write the electronic configuration of neutral atoms:

1) The Aufbau Principle states that you must start with the lowest energy level.
2) Figure out how many electrons you have.
3)According to the diagram, that is the order you follow, so 1s,2s,2p etc..
4. Remember how many electrons in each subshell
5) Keep filling until you run out of electrons.

Example 1:

Oxygen- has 8 electrons.

So 1s^2 2s^2 2p^4

You write the name of the subshell and then the number of electrons in it as an exponent.

Example 2:

Krypton has 36 electrons.

1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6

Note that 4p^6 is filled up. This is because krypton is a noble gas.

This is the visual way to display electron configuration. Note how each shell is filled with up and down arrows because they have opposite spins. Remember to fill each subshell with all up arrows and then go back and fill the down arrows.

Another way to display electron configuration is to use core notation.

The core-the set of electrons with the configuration of the nearest noble gas

The outer part of an atom- the electrons other than the core.


How:

1) Locate the closest noble gas to that element on the periodic table.
2) Cite the noble gas in [ ] and then write the rest of the configuration.

Example 3:

Nitrogen:
Has 7 electrons= 1s^2 2s^2 2p^3

What you do is backtrack and find the nearest noble gas. Go backward across the periodic table(left) until you run off the table and come out one level up on the right side.

In this case the nearest NG is helium.

Its atomic number is 2 so, [He].
Then eliminate the first 2 electrons from the configuration and add the rest.

2s^2 2p^3

[He] 2s^2 2p^3.


Exceptions
Instead of: Cr --> [Ar] 4s^23d^4
Cu--> [Ar] 4s^23d^9
In actuality:
Cr--> [Ar] 4s^13d^5
Cu--> [Ar] 4s^13d^10 

http://www.youtube.com/watch?v=YURReI6OJsg&feature=relmfu 

That should help!

Atomic Number, Mass and Mass Number- Subatomic Particles

Today we learned about the different characteristics of modern atoms.

Recall that an atom is made up of 3 types  of subatomic particles.

The Proton
Has a charge of +1.
Has a mass of 1.
Location: Nucleus

The Electron
Has a charge of -1
Has a mass of (not equal to) zero
Location: Energy shells surrounding nucleus

The Neutron
Has a charge of 0
Mass of >1.
Location: Nucleus

The atomic number(Z) of an element is the number of protons in its nucleus.

Atoms can form ions, which are charged atoms.
To form a negative ion(anion) electrons must be added to a neutral atom.
To form a positive ion(cation) electrons must be taken away from a neutral atom.

To find the charge of an atom: # Protons- #Electrons.

The mass number of an element is the sum of its protons and neutrons
The atomic mass is the average of masses of existing isotopes of an element.

An isotope of an element always has the same number of protons but a different number of neutrons.

How to calculate atomic mass.

Three naturally occurring isotopes with their percent abundance.

12(60%), 13(36%), 14(4%)

Calculate atomic mass.

12 x 60%= 7.2
13 x 36% = 4.68
14 x 4% = 0.56

Now add: 7.2 + 4.68 + 0.56 = 12.4.

Atomic Theory

The history of atomic theory has evolved over thousands of years. Several new ideas, concepts and models have been put forth and have contributed to our understanding of the structure of an atom.

Democritus- believed that matter was made up of atoms. Had a theory that atoms remain unchanged but move about in space to form macroscopic objects.

Aristotle- His theory was that atoms were made up of 4 elements; earth, wind, fire and water. When these elements combine they created dryness, wetness, coldness and hotness.

Antoine Lavoisier- accidentally discovered oxygen and created an incomplete table of the elements. Created early version of Law of 1) Definite Proportions 2) Conservation of Mass

Joseph Proust- created the Law of Definite Proportions-a given compound always has the same elements in the same proportions by mass.

John Dalton- he proposed Dalton's Atomic Theory
1. elements are made up of atoms
2.atoms of same elements are identical
3.atoms of different elements can be told apart by atomic weight
4.atoms of different elements can combine in a reaction to form chemical compounds in fixed ratios.

JJ Thomson-proved the existence of the electron, proposed the raisin bun model of the atom.

 Ernest Rutherford- discovered the positively charged, central, dense nucleus.
Did this using his gold foil experiment in which he fired radioactive particles through thin foil; most passed through but some particles were reflected.

Niels Bohr- proposed the planetary model of the atom
-Electrons exist in energy levels surrounding the nucleus.
-studied gaseous samples of atoms

The Atom today:
It is the smallest particle of an element and cannot be broken down.
It contains 3 subatomic particles- the proton(+), the electron(-) and the neutron (0).
The protons and the neutrons occupy the nucleus and electrons exist in levels around the nucleus.

Percent Yield and Percent Purity

Today we learned the last things for the test.

1) Percent Yield: it is the amount of actual product obtained as a percentage of the theoretical amount of product.

It is found using the formula:

Actual mass produced(g) / theoretical mass produced(g)

Lets try some problems!!!!

Example :

The reaction SiO2(s) + 4 HF(g) -----> SiF4(g) + 2H2O(l) produces 3.0 grams of H2O when 13.30 g of SiO2 is treated with small excess of HF.

What mass of SiF4 is formed?

Looking back at previous lessons this is fairly straightforward.


Mass of SiF4 formed:

3.0 g H20 x 1mole/ 18.0 x 1SiF4/2H20 x 104.1/1 mole=8.7 g of SiF4


What mass of SiO2 is left unreacted?

Mass of SiO2 used - 3.0 g H2O x 1mole/18.0 x 1 SiO2/2H2O x 60.1/1mole=5.0 g SiO2

13.30-5.0 g = 8.3 g SiO2 in excess

What is percent yield of reaction?

amount of SiF4 expected:  13.30 g SiO2 x 1 mole/ 60.1 x 1 SiF4/1SiO2 x 104.1/1mole = 23.0 g SiF4

Now remember Percent Yield is amount obtained/amount expected x 100

8.7/23 x 100 = 37 .8% yield.


Percent Purity:

(Mass of pure substance/mass of impure substance) x100

If the mass of silver ore is 66.4 grams and contains 12.0 g of pure silver what is percent purity?

12.0/66.4

x 100 = 18.1 % purity





http://www.youtube.com/watch?v=LicEaaXhlEY

Lab 6D: Determining the Limiting Reactant and Percent Yield in a Precipitation Reaction

CaCl2(aq) + Na2CO3(aq) -----> CaCO3(s) + NaCl(aq)

In this lab we observed the double replacement reaction between sodium carbonate and calcium chloride. We determined the excess and limiting reagents of the reaction, a theoretical and actual mass. After doing this we determined percent yield of the reaction.

Equipment:

-2 25 mL graduate cylinders
-1 250 mL beaker
-filter paper
-ringstand

After filling the cylinders with sodium carbonate and calcium chloride we added the calcium chloride to the sodium carbonate in a beaker. We then obtained a piece of filter paper, and weighed it. We wrote our names on the paper and then set up an apparatus for filtering using a ringstand. We poured the new solution carefully, bit by bit into the funnel. The filter paper filtered out the new precipitate, which was  CaCO3. We then let the filter paper and precipitate dry and then weighed it again.

NEXT CLASS WE WILL LEARN PERCENT YIELD!!

Excess and Limiting Reactants

All the reactions we've done up until now assume that all reactants are completely used up during reaction. This is not the case. Often, reactants are in excess; there is some left over.

Here is an example:

What mass of CS2 is produced when 18.0 g of C are reacted with 41.0 g of SO2?

5 C + 2 SO2-------> CS2 + 4 CO

What mass of excess reactant is left over?

Here try and find mass of CS2 by using the masses of the reactants C and SO2.

Mass of CS2 (based on C) 18.0 g C x 1 mole/ 12.0 x 1 CS2/5 moles C x 76.2 / 1 mole= 22.9 g CS2

Mass of CS2 (based on SO2)  41.0 g SO2 x 1 mole/ 64.1 x 1 CS2/2 SO2 x 76.2/ 1mole = 24.4 g CS2

Since there is too much SO2 and not enough C, SO2 is the excess reactant because in the reaction all the C will be used  up and the SO2 will be in excess.

In this reaction, C is the limiting reactant because it limits the amount of CS2 that can be formed, there is lots of SO2 and not enough C to make that much SO2.

Example #2. What mass of P4 is produced when 38.0 g of Ca3(PO4)2, 25.0 g of SiO2 and 9.70 g of C are reacted?

2 Ca3(PO4)2 + 6 SiO2 + 10 C -----> P4 + 6 CaSiO3 + 10 CO

Use the mass of each reactant and find the mass of P4.

Mass of P4( based on Ca3(PO4)2)38.0 g Ca3(PO4)2 x 1 mole/ 310.3 x 1 P4/2 Ca3(PO4)2 x 124.0/1 mole = 7.59 g P4

Mass of P4 ( based on SiO2) 25.0 g SiO2 x 1 mole/60.1 x 1 P4/ 6 SiO2 x 124.0/1mole = 8.60 g P4

Mass of P4 (based on C) 9.70 g C x 1mole/12.0 x 1 P4/ 10 C x 124/1mole= 10.0 g P4.

Ca3(PO4)2 produces the least amount of P4 = 7.59 g

The excess reactants are  SiO2 and C.

How many grams of excess reactant are left over?

38.0 g Ca3(PO4)2 x 1mole/ 310.3 x 6 SiO2/2 Ca3(PO4)2 x 60.1/1 mole = 22.1 g SiO2.

To calculate excess, subtract this number from the original amount.

25.0 - 22.1 = 2.9 g SiO2 in  excess.

38.0 g Ca3(PO4)2 x 1 mole/310.3 x 10 C/ 2 Ca3(PO4)2 x 12.0 / 1mole =7.35 g C

Subtract from original.

9.70 g- 7.35 = 2.35 g excess C







http://www.youtube.com/watch?v=JOQz2rIFpVM

Xtra help above :)

Stoichiometry with Molarity and STP

Today we learned how to solve stoichiometry problems involving molarity and STP.


Remember Molarity =M


M=moles/Litres or moles= M/L  or moles/M = L

Example:
A tablet of Tums has a mass of 0.750 g. What volume of stomach acid(HCl) = 0.010 M is neutralized by a 0.750g portion of CaCO3?

1CaCO3(s) + 2 HCl(aq) ----> CaCl2(aq) + CO2(g) + H2O(l)

First find moles of HCl using stoichiometry.
0.750 g CaCO3 x 1 mole/100.1 g CaCO3 x 2 mol HCl/1 mol CaCO3 = 0.0150 mol

Now find the Litres using the molarity equation.
L= moles/Molarity
0.0150 mol/0.0010 mol/L = 15 L.

Now remember STP.

STP stands for Standard Temperature and Pressure

The conversion factor is 22.4 L / 1 mole or 1 mole/22.4 L.

You do the exact same thing with the other problems but when you have your desired substance in moles, you just multiply by 22.4 L / 1mole to find volume at standard temperature and pressure.

Example:
What volume of CO2(g) at STP is produced if 1.25 L of 0.0055 M HCl reacts with an excess of CaCO3?

First you find moles of the HCl.

moles = Molarity x litres

0.0055 M x 1.25 = 0.006875 mol HCl

Now just find the moles of CO2 using stoichiometry.

0.006875 x 1 mol CO2/ 2mol HCl

Now multiply by 22.4 L/ 1 mole.

Your answer will be 0.077 L CO2.

http://www.youtube.com/watch?v=z4nPxk1deuw

Here is some extra help!

Stoichiometry

Stoichiometry is the measurement and study of the relationships and quantities between products and reactants.

Let's examine a chemical equation:

__C4H10(g) + __ O2(g) ---------> __CO2(g) + __H2O(l)
First we must balance it!

2 C4H10 + 13 O2(g) -------> 8 CO2(g) + 10 H2O(l)
The numbers in front of the products and the reactants are called mole ratios.
We use this in stoichiometry to convert between grams and moles of the products and reactants.

Example:

If 24.0 grams of C4H10 was reacted with O2, how many grams of CO2 would be formed??

Make a mole map!!!!

You cannot convert directly from grams of C4H10 to grams of CO2 so first you must convert to moles.

Here is how you set up your equation:

24.0 g C4H10 x 1 mole/molar mass C4H10- This converts grams of C4H10 into moles.
Now you use stoichiometry to determine moles of the CO2

Remember: What you want/what you have

So: multiply moles of C4H10 by 8 mol CO2/ 2 mol C4H10
Now you have moles of CO2. All thats left is to convert moles of CO2 into grams of CO2

Multiply moles of CO2 by the molar mass of CO2/1 mole.

Your answer will be: 72.8 g CO2

Ex. 2

2 KClO3(s) -------> 2KCl(s) + 3O2(g)

How many molecules of KCl will be formed when 45.0 g KClO3 is decomposed?

45.0 g KClO3 x 1mole/122.6 x 2 KCl/2 KClO3 x 6.022 x 10^23/1 mole=2.21 x 10^23 molecules of KCl

http://www.youtube.com/watch?v=nf8J2g_A8uE
Here is some extra help!

Simple Stoichiometry

Stoichiometry is the measurement of the different quantities and relations between products and reactants.

It is very simple. In a chemical equation, you can obtain moles of a reactant from moles of a product and vice versa.

The stoichiometry ratio is WHAT YOU NEED/WHAT YOU GET
THIS IS THE MOLE RATIO.

For example: __H2 + __O2 -----> __H2O
BALANCE: 2H2 + 1O2 ----> 2H2O
Use the numbers in front to use the mole ration for stoichiometry.
If you have moles of oxygen( lets say 4.1 moles of oxygen) you can convert to moles of H2O.

4.1 mol O2 x 2 mol H2O/4.1 mol O2 = 2 mol H2O

It is very straightforward but harder stuff is to come!!

Mole Calculations with Energy

Today we learned about mole calculations with energy.

This involves enthalpy, delta H. Delta H is expressed as a kJ/mole.

 To convert:
Moles to energy ----> 1kJ/mole

Energy to moles----> 1 mole/ 1 kJ


Here are some examples.

E.g.
CuO (s) + H2(g) + 135.0 kJ ----> Cu(s) + H2) (l)

Calculate energy absorbed for 6.0 moles of  CuO reacting.

130.5 kJ/mole x 6 moles = 780 kJ

Calculate the amount of energy released when 6.00 grams of Carbon are combined with O2 to form CO2 gas.

C(s) + O2(g) -----> CO2(g) + 393.5 kJ
First you convert grams to moles:

6.00g C (1 mole C/ 12.0107 g)

Then you multiply by (393.5 kJ/ 1mole C) and you get 196.57 kJ.

Piece of cake!


Unfortunately no videos were found. But...

Exothermic and Endothermic Reactions

Today we learned about endothermic and exothermic reactions.

Exothermic reactions are reactions that give off energy.
Endothermic reactions absorb energy. 

Today we also discovered the joy of energy diagrams. These show the energy of the reactants and the energy of the products.

 The reaction is exothermic because, the energy of products is less than energy of reactants. This means energy was released.

The x axis represents the time needed to undergo the reaction and the y-axis represents potential energy in kJ.

The diagram is very straightforward.

The initial flat line represents the energy of the reactants. As you can see, the line transforms into a hill. The distance between the initial flat line and the top of the hill is called the activation energy.This is the amount of energy needed to carry out the reaction.  The level of the hill (on the y axis) is called the activated complex. It is the highest amount of energy that the reactants can reach in the reaction process. The curve then drops off and then evens out. The distance between the energy of the reactants line and the energy of the products is known as delta H. Delta H is the change in enthalpy,(heat contained in the system).

This is found using the equation
Delta H = energy of products - energy of the reactants.
A negative value of H  signifies an exothermic reaction because you end up with less energy than you started with.

A positive value of H signifies and endothermic reaction because you end up with less energy than you started with.

Once H is found, you use it in chemical equations like this

Mg + Cl2 ------> MgCl2 + 347 kJ

Notice that the 347 is on the side of the reactants. This means that the reaction is exothermic because:

Energy on one side = energy on the other
So the MgCl2 needs the 347 kJ to make up the energy so it equals the reactants.
This means that the reaction released energy.

Quiz

We had a quiz today in Chemistry.

It was on balancing and predicting products of chemical equations
It covered everything we had been learning until now.
Types of reactions
Prediction Products
Net Ionic Equations
Solubility
Balancing equations


It was fun!

Advanced Double Replacement

Remember double replacement?
Now we have double replacement with solubility.

Now we have to put the states of matter signs beside in compound in a balanced equation
E.g. (g), (l), (aq), (s)

Lets take a double replacement reaction.

2 K3PO3 + 3Cu(NO3)2 -------> 1Cu3(PO3)2 + 6 KNO3

http://nobel.scas.bcit.ca/chem0010/unit8/solubilityRules.htm Here are the solubility rules:

Look at the products. Use the solubility sheet to see whether Cu will react with PO3. Do the same for K and NO3
Since Cu and PO3 do not react, they are a solid (s).
K and NO3 do react and will be (aq)
It will look like this:

2K3PO3 + 3 Cu(NO3)2 ------> 1 Cu3(PO3)2 (s) + 6 KNO3(aq)

The equation will look like this

2K3PO3 (aq) + Cu(NO3)2(aq) -----> 1 Cu3(PO3)2 (s)  + 6 KNO3(aq)

Lab 5 B

Today we did a great lab in class:

There were 7 reactions that we observed and each one was amazing!

Here they are:

Reaction 1: We held a piece of copper wire in the hottest part of a bunsen burner flame for several minutes and recorded our observations:

Reaction 2: We took a nail, and cleaned it, then we submerged half the nail in liquid copper (II) sulfate solution. 15 minutes later, the part of the nail that was submerged was coated with copper.

Reaction 3: We heated some copper (II) sulfate(s) as we have already done(we are experts) and dried it out completely. It turned greyish,blue and was grainy and dry

Reaction 4: We added water to the anhydrous salt! Remember what happens:??? You get the hydrate back.

Reaction 5: We took 2 test tubes, one contained CaCl2(l) and the other Na2CO3(l) and we poured the calcium chloride into the sodium carbonate.

Reaction 6: We took mossy, fuzzy zinc and submerged it in hydrochloric acid.

Reaction 7: Manganese (IV) oxide is added to hydrogen peroxide. We tested the gas by placing a burning splint into the test tube. It lit up!

http://www.youtube.com/watch?v=ARRCD_Ry4eI  #1
http://www.youtube.com/watch?v=oQz5YEsx7Fo #6

Sorry we couldn't find them all!