Calculating Empirical Formula of an Organic Compound

Did you know that you can use the empirical formula on organic compounds?
Now what is an organic compound?
An organic compound is any member of a large class of gaseous, liquid, or solid chemical compound whose molecules contain carbon.

Ex. When a 3.79 grams of an organic compound is burned, 6.61 grams of CO2 and 3.59 grams of H2O are produced.  What is the empirical formula of this compound?

Molar mass of CO2: 12 + 16 + 16 = 44 g/mol
Mole C = 6.61g x 1 mole CO2     1 mole of C   =  0.15023 mole of C
                              44.0 g CO2      1 mole of CO2

Molar mass of H20: 1.0 + 1.0 + 16.0 = 18.0 g/mol
Mole H = 3.59 g x 1 mole of H2O  x     2 moles of H  =  0.39889 mole of H
                              18.0 g H2O          1 mole of H20

Check the mass of C and H
Mass of C = 0.15023 x   12.0 g C = 1.80 g C
                                      1 mole C

Mass of H = 0.39889 x   1.0 g H   =  0.39889 g H
                                     1 mole H
Since the masses of C and H do not add up to 3.79, the rest of the mass must be from O.

Mass of O = 3.79 - 1.80 - 0.39889 = 1.59111

Mole O = 1.59 x 1 mole O  =  0.0994 g O
                           16.0 g O

Now we divide by the smallest molar amount. That is 0.0994

Carbon: 0.15023/0.0994 = 1.5 x 2 = 3
Hydrogen: 0.39889/0.0994 = 4.0 x 2 = 8
Oxygen: 0.0994/0.0994 = 1 x 2 = 2

Empirical formula: C2H8O2

This video will explain determining the empirical formula of an organic compound:

Empirical and Molecular Fomula

What is the empirical formula?
It gives the lowest term ratio of atoms (or moles) in a formula.

Ex. C4H10 (butane) --> molecular fomula
      Reduce the subscripts to lowest terms to get the empirical formula
      C2H5 --> empirical formula

Ex. Consider that we have 10.87 grams of Fe and 4.66 g of O.  What is the empirical formula?
First, we must convert the grams to moles.
10.87g x 1mole  =  0.195 moles of Fe

4.66g x 1 mole  =  0.291 moles of O

Now, we divide the mole amount by the smallest mole which is 0.195

Fe: 0.195  =  1

O: 0.291  =  1.49 = 1.5
Next, we must get those two numbers to the smallest whole number
Fe 1 x 2 = 2
O 1.5 x 2 = 3
The empirical formula is: Fe2O3

Ex. A compound contains 31.9% of K, 28.9% of Cl, and 39.2% of O.  What is the empirical formula?
*assume we have 100g*

31.9 x 1 mol  =  0.816 moles of K

28.9 x 1 mol  =  0.814 moles of Cl

39.2 x 1 mol  =  2.45 moles of O

lowest mole: 0.814
0.816  =  1.00

0.814  =  1

  2.45  =  3

Empirical formula: KClO3

What is the molecular formula?
It is a multiple of the empirical formula and shows the actual numbers of atoms that combine to form a molecule.

To calculate the multiple, we use this formula:
n =      molar mass of the compound    
      molar mass of the empirical formula

Ex. A molecule has an empirical formula of C2H5 and the molar mass is 58 g/mol. What is the molecular formula?
Molar mass is C2H5 = 58 g/mol
Molar mass of the empirical formula (total molar mass of all elements in the compound):
C2 = 12 x 2 = 24
H5 = 1 x 5 = 5
C2H5 = 29 g/mol

58 g/mol  =  2
29 g/mol

2 x C2H5 = C4H10

Video about the empirical formula:

Video about the molecular formula:

Percent Composition

Percent composition is the percentage by mass of a "species" in a chemical formula.  You are trying to find the percentage of how much one element is in a compound.  You can find this by using the molar mass of each compound.
**(sometimes the percentage may not add up to 100% because of the rounding.  If the percent is somewhere near 99.9% its ok.)**

Ex.1 What is the percentage composition of NaCl?
First we need the following:
Total Molar Mass of NaCl: 58.5g/mol
Molar Mass of Na: 23.0 g/mol
Molar Mass of Cl: 35.5 g/mol
Then we do the calculations like so:
% of Na = 23    g/mol   x  100% = 39.3%
                 58.5 g/mol

% of Cl = 35.5 g/mol   x   100% = 60.7%
                58.5 g/mol
Total percentage: 100%

Ex.2 What is the percentage composition of ZnSO3?
Total Molar Mass of ZnSO3: 145.5 g/mol
Molar Mass of Zn: 65.4 g/mol
Molar Mass of S: 32.1 g/mol
Molar Mass of O: 16.0 x 3 = 48.0 g/mol

% of Zn = 65.4  g/mol  x  100% = 44.9 %
                145.5 g/mol

% of S = 32.1   g/mol  x  100% = 22.1%
               145.5 g/mol

% of O = 48.0    g/mol  x  100% = 33.0%
                145.5 g/mol
Total percentage: 100.0%

Ex. 3 What is the percentage composition of Zn3(PO4)2?
Total Molar Mass of Zn3(PO4)2: 386.2 g/mol
Molar Mass of Zn3: 65.4 x 3 = 196.2 g/mol
Molar Mass of P: 31.0 x 2 = 62.0 g/mol
Molar Mass of O: 16.0 x 8 = 128.0 g/mol

% of Zn = 196.2 g/mol  x  100% = 50.8%
                 386.2 g/mol

% of P = 62.0    g/mol    x  100% = 16.1%
               386.2 g/mol         

% of O = 128.0 g/mol  x  100% = 33.1%
                386.2 g/mol
Total percentage: 100%

Ex. 4 If a compound contains 54 grams of Al, 36 grams of C, 56 grams of N, some amount of O, and has a total molar mass of 226 grams, what is the percentage composition? 
First, we need to find the molar mass of the amount of O.  To do this we do:
226 - 54 - 36 - 56 = 80  <--- this is the molar mass of O.

% of Al = 54.0   g/mol  x  100% = 23.9%
                226.0 g/mol

% of C = 36.0   g/mol  x  100% = 15.9%
               226.0 g/mol

% of N = 56.0   g/mol  x  100% = 24.8%
               226.0 g/mol

% of O = 80.0    g/mol  x 100% = 35.4%
                226.0 g/mol 
Total percentage: 100%

Ex. 5 If a compound contains 21.6 grams of B, 57.0 grams of F, 80.8 grams of Ne, some amount of He, and has a total molar mass of 179.4, what is the percentage composition? 
What is the molar mass of He?
179.4 - 21.6 - 57.0 - 80.8 = 20 <--- this is the molar mass of He. 

% of B = 21.6   g/mol  x  100% = 12.0%
               179.4 g/mol

% of F = 57.0   g/mol  x  100% = 31.8%
               179.4 g/mol

% of Ne = 80.8   g/mol  x  100% = 45.0%
                 179.4 g/mol

% of He = 20.0   g/mol  x  100% = 11.1%
                 179.4 g/mol
Total percentage: 99.9%
(In this case, the total percentage was not 100% but was 99.9%.  It is just a 0.1% difference and not that big of a deal.  This is the result due to rounding.)

Video similar to Ex 2 & 3:

Video almost similar to Ex. 4 & 5

More Mole Conversions

During class, we learned that there are more mole conversions.  You can convert from grams to particles!! Or moles to # of atoms in particles!! Seems like a lot...
The units in these conversions are:
-Particles/Atoms/Molecules/Formula Units
-# of Atoms in Particles

From Grams to Moles, you multiply using 1 Mole

From Moles to Grams, you multiply using   MMg 
                                                              1 Mole
MMg = Molar Mass grams

From Moles to Particles/Atoms/Molecules/Formula Units, you multiply using
6.022 x 10^23 Particles
          1 Mole

From Particles/Atoms/Molecules/Formula Units to Moles, you multiply using
            1 Mole             
6.022 x 10^23 Particles

From Particles/Atoms/Molecules/Formula Units to # of Atoms in Particles, you multiply using
# of Atoms 
1 molecule

From # of Atoms in Particles to Particles/Atoms/Molecules/Formula Units, you multiply using
1 molecule 
# of Atoms

Conversions from Particles to Mass
Ex. What is the mass of 2.78 x 10^22 Fe Atoms?
2.78 x 10^22 Fe Atoms x                 1 Mole                x      55.8 g (Molar Mass)
                                         6.022 x 10^23 Fe Atoms                 1 Mole
= 2.58 g

Ex. What is the mass of 8.4 x 10^18 SO3 molecules?
8.4 x 10^18 molecules x                 1 Mole                 x 8.01 g (Molar Mass)
                                        6.022 x 10^23 Molecules               1 Mole
= 1.1 x 10^-3 g OR 0.0011 g

Conversions from Grams to Particles
Ex. How many atoms of Iron in 20.0 g of Iron?
20.0 g x           1 Mole               x  6.022 x 10^23 Atoms
              55.8 g (Molar Mass)               1 Mole
= 2.16 x 10^23 Atoms of Fe

Here is a video that demonstrates from Grams to Atoms (Skip to 4:10 in the video)

Ex. How many formula units of KMnO4 in 0.240 g of KMnO4?
0.240 g x 1 Mole    x  6.022 x 10 ^23
                158.0 g             1 Mole
= 9.15 x 10^20 Formula Units


Mole Conversions

1. Conversions from particles <--> moles
(recall 6.022 x 10^23)

From particles --> moles
3.01 x 10^24 particles x (1 mole/ 6.022 x 10^23) = 5.00 moles

From moles --> particles/molecules/formula units
Ex. 0.75 moles of CO2 --> molecules
0.75 moles x (6.022x10^23/ 1 mole) = 4.5 x 10^23 molecules of CO2

Ex. 0.75 moles of CO2 --> atoms of Oxygen
4.5 x 10^23 molecules x 2 atoms of O = 9.0 x 10^23 atoms of oxygen.

2. Conversions between grams <-->moles

From moles --> grams 
Ex. 2.04 moles of carbon --> grams
molar mass of Carbon = 12.0 g / mol
2.04 mole x 12.0 g/ 1 mole = 24.5 grams of Carbon

Ex. 0.341 moles of NO2 --> grams
Molar mass of NO2 = 46.0 g / mol
0.341 moles x (46.0 g / 1 mole) =  15.7 grams of NO2

Here is a video explaining how to convert moles to grams:

From grams --> moles
Ex. 3.45 grams of Carbon --> moles
Atomic mass of Carbon = 12 grams
3.45 grams x (1 mole / 12 grams) = 0.288 moles

Ex. 6.2 grams of MgCL2 --> moles
Molecular mass of MgCL2 = 95.3 grams
6.2 grams x (1 mole / 95.3 grams) = 0.065 moles

Here is a video explaining how to convert grams to moles:

The Mole

 Equal volumes of different gases have a constant ratio:
Oxygen             :  Hydrogen   16:1
Carbon Dioxide :  Hydrogen   21:1

Relative Mass:
-Expressed by comparing it mathematically to mass to the mass of another object.

Avogadro's Hypothesis:
-Equal volumes of different gases at the same temperature and pressure have the same number of particles.
-If they have the same number of particles, the mass ratio is due to the mass of the particles.
-We now use this principle for the relative masses of all atoms on the periodic table.

Atomic Mass:
-The mass of a specific isotope of a given atom. Expressed as amu (atomic mass unit)

Formula Mass:
All the atoms of a formula of an ionic compound.
Example:  Potassium Fluoride
                K            +         F
                39.1        +    19.0
                KF = 58.1 amu

Molecular Mass:
-The mass of one molecule of a substance.  Expressed as amu.

Molar Mass
-Atomic/Molecular/Formula mass of any pure substance (in grams per mole)
Example: 1 mole of oxygen = 16.0 g/mol
               1 mole of carbon = 12.0 g/ml
-The molar atomic mass of an element is the mass of 1 mole of that element.
This video explains how to find molar mass:

Avogadro's Number
-The number of particles in 1 mole of any amount of substance is 6.022 x 10^23 particles/mol

-The mole is important because it allows chemist to count atoms and molecules 

A chemistry song

Density of Cold and Hot Water

The class went to the crick lab and were presented with a open office document that displayed numbers.  The next step was input the formula for density.  What would happen was that the formula used to calculate the density and it was then displayed on the graph.  Then we would change the colors of the graph to our personal preferences, add a title, and label the x and y axis.  Then we saved the file and handed it into  Mrs. Chen's handin folder.

Accuracy and Precision

Precision is how reproducible a measurement is compared to other similar measurements.
Accuracy is how close the measurement (or average measurement) comes to the accepted or real measurement.
Uncertainty is the margin of error and there are two types of uncertainty, absolute and relative.

There are 2 methods of finding absolute uncertainty

Method 1: Discard any unreasonable information.  Then find the average of the 3 numbers and find the largest difference between the average and either the lowest or highest reasonable measurement.

Trial 1 - 29.86 g (this would be discarded because the other numbers include numbers from 0.90-0.96)
Trial 2 - 29.97 g
Trial 3 - 29.93 g
Trial 4 - 29.94 g
Trial 5 - 29.96 g
Average of the numbers: 29.95 g

Difference between average and lowest number =  29.95g - 29.93g = 0.02g
Difference between average and highest number = 29.97g - 29.95g = 0.02g
Absolute certainty based on average: 29.95 ± 0.2 g

Method 2: Measure to the most precise measurement then estimate to the 0.1 of the smallest measurement on a measuring instrument. 

For example, if the measuring instrument is a ruler, the smallest measurement is 0.1mm.  The data recorded for a measurement of a ruler would end up looking  like  ________(best precise measurement) ±0.1mm
For relative uncertainty, it is the ratio of absolute uncertainty to the estimated measurement. For example, if the absolute uncertainty is 29.95 ± 0.2 g, the relative uncertainty would be 0.2 / 29.95.  It can be expressed as a percentage or in significant figures. 
Significant Digits
The last digit in a measurement is always the uncertain digit(it could be one digit higher or lower, we don't know). The significant figure in the measurement includes all of the certain digits and the uncertain digits.

Significant Figure Rules
1. Leading zeros are not counted
Example: 0.01 = 1 SF
2. Trailing zeros after the decimal point are counted
Example 10.050 = 5 SF
3. Trailing zeros in front of the decimal point are not counted
Example  1200000000000 = 2 SF
4. Zeros between significant figures are counted
Example  123000.000123 = 12 SF

Rounding Rules
To start off, rounding just almost like we do it in math
1.Look at the digit to the right the the digit you are rounding (The uncertain digit)
2. If that digit > 5 then you round up
3. If that digit < 5 then you keep it the same
4. If that digit = 5 and it has more non-zero digits following behind it, round up
5. If that digit = 5 and ends at that 5, round to make digit your rounding even (2,4,6,8)

Math "+" & "-" Rules
When you add or subtract, round to the number with the fewest DECIMAL places
For Example
                                               +    1.1          
Which becomes 223.2 because the 1.1 had the lowest decimal point

Math "X" & "÷" Rules
When you multiply or divide, round to the number with the fewest number of significant figures
For Example
                      (123)(1)  =  123
Which becomes 100 because the number with the lowest significant figures (that 1) only had 1 SF.

Separation of a Mixture by Paper Chromatography

Before this lab, 3 tables from the text book were to be copied down onto paper.  This lab had 6 objectives:
To assemble and operate a paper chromatography apparatus
To study the meaning and significance of Rf values
To test various food colorings and to calculate their Rf values
To compare measured Rf values with standard Rf values
To separate mixture of food colorings into their components
To identify the components of mixtures by means of their Rf values

The supplies were:
5 stirring rods
3 large test tubes
3 Erlenmeyer flasks
Chromatography paper strips
Food colorings (assigned either yellow, green, blue, or red)
Unknown mixture of food colorings

Part I: Setting Up
Obtain 3 Erlenmeyer flasks and 3 pieces of chromatography paper.  Draw a line at 4 cm into the paper and then cut the edges to make it into a arrow shape. Then place 2 cm of solvent into the flasks.

Part II: Rf Values of Individual Food Colorings
One food coloring was assigned either red yellow or blue.  On the center of the line that was drawn on the chromatography paper, the color that was assigned was to be spotted onto the center of the line. Next the strip was to be inserted into the test tube.  Once that was placed, observe the sample spot and make any observations in the next 20 minutes.  While that was happening, move onto the next part of the lab, the green and unknown food colorings.  Observe any changes with those two colorings as well.  After the solvent and food coloring sample have moved up and no longer moved, take them out and using the ruler measure how much cm the solvent and food coloring sample moved.  Record these on the tables that were drawn before the start of the lab.  Then clean up and continue onto the lab report. 
Naming Acids
Acids are formed when a compund of Hydrogen Ions and a negitively charged ion are dissolved in water(aqueous). Ions separate when they are dissolved in water.

For Example:
H + Cl  ->  HCl(g)
HCl(g) + H2O(l) ->  H3O(aq) + Cl(aq)

Naming Acid Guidlines
1. Use "hydro" as the beginning
2. Last syllable of the non-metal is dropped and replaced with "-ic"
3. Add "acid" at the end
*                ide ->  hydro                ic  acid*
For Example:
                      HF        becomes      hydro flouric acid

Naming Complex Acids
1. -ate is replaced with -ic
    -ite is replaced with -ous
2. "acid" at the end of the name
For Example:
                      HCH3COO    becomes Acetic Acid

 Try This Video, It should help

Heating and Cooling of a Pure Substance experiment

Last class was when we did the heating and cooling of a pure substance.  The materials used were a pair of safety goggles, a test tube with the pure substance, a beaker with warm water, 2 thermometers, a mini stove like device, and a device that held the test tube that could also adjust to move the test tube up or down. 

The Cooling Process

The test tube was adjusted on the device and lowed into a beaker of room temperature water.  In every 30 seconds, the thermometer would either drop quickly or slowly.  The original temperature of the pure substance was 45 degrees.  The purpose was to see how long it would take for the substance to have a temperature of 25 degrees.  The whole process was to be recorded on the pre drawn table.  Next was the heating process.

The Heating Process

By now the pure substance should be a solid from the cooling process.  The next step was to bring out the mini stove like device and place a beaker of water on top of it.  The mini stove like device has a nob which was turned to 200 for the amount of heat used.  By now the substance's test tube is 25 degrees.  Place it in the beaker that is sitting on top of the mini stove like device.  The result should be a increase of temperature.  The temperature may move rapidly or slowly depending on how cold the water of the beaker is.  The temperature was to be recorded every 30 seconds until the temperature of the pure substance reached 50 degrees.  When the temperature reaches close to 50 degrees, the pure substance should start to change state.  By 50 degrees, the pure substance would be completely dissolved.  By the end of the heating process, your table should have the amount of time it took to reach the required temperature of both the heating and cooling processes.  When you have both of them, clean up the station and move onto the lab report.
Law of Definite Composition
Compunds will have a definite composition
Ex. H2O will be H2O anywhere (It will always have 2 H's and 1 O)

Law of Multiple Preportions
When 2 or more compunds with different proportions of the same element can be made
Ex.  CO2   ->   C2O4  -> C4O6

The Heating/ Cooling Curve of a Pure Substance
Picture made with Firworks

A: It is now in a solid state ( any temperature below melting point ). The particles are packed very close and neatly together and the forces between the particles are very strong. The particles can only vibrate at a fixed position.

A-B: As the particles are heated, the heat energy is converted to kinetic energy. The Kinetic energy increases and the molocules vibrate faster around their fixed position. This also increases the temperature

B: By here the solid will change into a liquid or a liquid turned into a solid (Frozen)

B-C: In Both solid and liquid states ( Ice and water ), the temperature remains constant because the heat supplied is used to overcome the forces of atractiong between particles. This constant temperature is called the melting point.

C: All melted, solid has now become liquid

C-D: It is still a liquid, as it is heated the molocules hgain more heat energy and the temperature increases. The particles move faster as the kinetic energy increases.

D: It still exists in a liquid state. The molocules have recieved enough energy to overcome forces of attraction between particles and they start moving freely. The liquid starts becoming gas.

D-E: The liquid is changing into a gas (Is both a liquid and gas at the moment). The temperature remains unchanged which we call the boiling point. The absorbed heat is being used to overcome the forces of atraction instead is raising the temperature.

E: all liquid is now gas
E-F: The gas particles continue to absorb energy and move faster. The temperature increases as the heating continues.

Matter in the Macroscopic World
When you look at anything, you will see that you have used the tools of science to observe and order observations through classification. For example you could say that from what you've seen, people are happy and are preparing to eat. Science uses the same skills as looking except you make much more precise and detailed observations requiring much time and effort to specialize in something. Chemists specialize in matter with questions asking what is it, how does one differ from another, what different things have in common, and how matter can be kept the same?

What we know about Matter
We've learned a lot about matter as we've grown up like water when cooled turns into ice and the water can be freely poured to fill any solid container. But if you took water from a muddy stream and water from a river you could say their different right? Properties can help us identify things in many ways including touch and taste. Part of the study of science is learning of all the ways you can distinguish one piece of matter from another like the temperature at which liquid changes into gas is which called the Boiling Point

Purifying Matter
There are 2 classifications of matter. Mixtures which are 2 or more kinds of matter that have separate identities or Pure substances which cannot be separated. Understanding which samples can be separated and which samples can be not is something that comes with experience. If you have dirty water then put it in the light, it will scatter the light. If you check with clean water, it won't scatter the light. That means you should be able to separate the mud from the water by how? If you go to a water processing plant, clean the water, and then see if it scatters the light which it doesn't, is it pure water? If we check the water there could be salt or sugar added to the water that do not scatter light which is when we call the water a Solution. This is still a mixture though so you should be able to separate it right? Then if we boil the water so that we still will have the water and a white solid of sugar or salt separate, this procedure is what we call Distillation. After you have boiled the water you can conlude that the water and the sugar or salt is pure. Many mixtures are too difficult to separate thus are considered pure for years until a way is found to separate them.

Characteristics of Pure Substances
An easy way to detect an important difference between mixtures and pure substances is that pure substances have a CONSTANT boiling point unlike mixtures that ordinarily do not. The problem is nature likes to defy our simple definition of saying the mixtures never have a constant boiling point. The reason the mixture of water and grain alchohol cannot be separated is that it has a constant boiling of 78.2 degrees celcius. The temperature when a liquid changes to a solid is call the Freezing Point. The temperature at which a solid becomes a liquid is called the Melting Point.

Chemecal and Physical Changes
Density is the property of matter that describes its mass per volume unit. A chemical change when a new kind of matter with different properties is formed. When a new substance is formed, thats a good way to tell if a chemical change occurred. Decomposition is a chemical change where one kind of matter come apart to for 2 or more kinds of matter. Changes that are easy to reverse or reversable are called Physical Changes. Physical changes do not produce new kinds of matter.

Compunds & ElementsElectrolysis is when a passing electric current throught a substance causes it to decompose into new kinds of matter. Usually electrolysis represents a chemecal change in the form of decomposition. Distillation is a physical change that separates two or more things that already exist like separated components that are still in the original mixture as separate substances. Pure substances that can be decomposed into new kinds of matter are called Compounds. The word comes from the fact that they appear to be "compunded" or put together from simpler substances. Pure substances that cannot be decomposed are called Elements. There are 109 known elements, 85 exist in nature and the others have been made in atomic reactors(they do not exist in nature).

Compunds have a Definite Composition
Not all combinations of elements are compounds. The important difference between mixtures of elements and compunds of elements is that the mixtures can have almost any composition desired but the compounds will have a definite composition which is called The Law of Definite Composition. The fact of 2 or more compunds with different amounts of the same elements is called The Law of Multiple Proportions.

 Matter is made of Atoms
When you have made an observation using feel, touch, or etc... is is called a Macroscopic Observation. Macro means large and scopic means viewing or observing. Melting point, boiling point, heat of fusion, temperture, etc... are properties of many parts of matter which we call Macroscopic Properties.

Atom means the smallest possible piece of something like saying you take a iron nail and you break it again and again. It only seams reasonable that what you get from the broken nail is to obtain the "smallest possible piece" of iron. Atoms are usually represented by spheres of various shapes and sizes. There are as many atoms as there are elements. Elements should be made of only one atom as if was made of 2 different atoms, the thought of we should be able to separate it would occur. Atoms vibrate even though stay in the shape they hold. As the temperature increases, they start vibrating more and they should vibrate with such force that they overcome the forces of attraction. Since atoms vibrate, they are able to move past each other from one side of a container to another. Particles made up of more then one atom is called a Molocule. Elements have different melting points and boiling points from the fact the atoms vary in size and mass which means more energy is required to get large atoms to break the forces of attraction.

CompundsAll compunds are made by combining elements in definite proportions which means all compunds must be made of 2 or more kinds of atoms. Since compunds are made of 2 or more kinds od atoms, if enough energy is supplied, you can break the compund apart. Heat and electricity supply the energy needed to do that. Not all compunds are made of molocules. Ions are particles that have an electrical charge. Compounds that melt to form ions conduct an electric current and compunds that don't melt to for ions do not conduct electricity.

Laboratory Report Format

1.Cover Sheet
  •  Name, date, block, name of partner in top right corner 
  • Title of lab in center of page
  • brief abstract underneath title
2.Report write-up
  • purpose: Write a statement that indicates what you want to find out by doing the lab
  • Materials and equipment:
  • Procedure:
  • Data and observation
  • Sample Calculations
  • Discussion
  • Sources of Error
  • Conclusion: Include a brief formal statement, which lists your results and/ or what you have learned in the lab



What is matter?
Matter is anything that has mass, volume, and take up space.

Matter can be classified in two categories: Pure substances and Mixtures.


Pure substances has one set of properties and are one kind of particle.  They can be either elements or compounds.

A element is made of atoms and is the simplest form of a pure substance and con not be decomposed.  Elements are also separated in two categories of metal and non-metal.

Compounds are made of elements that are chemically combined.  They are the smallest particles of a molecule.  Compounds are classified as ionic and covalent compounds depending on what element that are combining with.


Mixtures have sets of properties and are a mix of substances.  Two categories of mixtures are Homogeneous and Heterogeneous.  Homogeneous are uniform and appear to have only one component.  Heterogeneous are not uniform and appear to have more than one component. 

Physical Vs. Chemical Change

Physical changes are changes that result in no new substances are formed and are reversible.  Chemical changes are changes that new substances are formed and are not reversible.  For example, freezing water into ice is a physical change because no new substance is formed and the ice can be changed back into water.  Burning wood is a chemical change because it makes ashes and the new substance can not be changed back into wood.

Scientific Notation Review / Unitary Rates

Scientific Notation Review
Scientific Notation is used to express very large or very small numbers by using the powers of 10
For Example say that we want to express a in scientific notation.
A) 3210000 
First we look at the number and check the important numbers needed then we bring the decimal to the left
Now remember how many spaces we moved to the left (6 times)  and then remove the zeros
Now we times it by 10 to the power of 6
3.21 X 106                
and now your done =) (Still don't understand it? then check out thie video

Unitary Rates
Example:        If        1m     =    100 cm           
                                1m2     =    10000 cm                (100 X 100)
                                1m3     =    1000000 cm            (100 X 100 X 100)
Remember that when you square or cube the units, you also square the powers.