Tuesday, December 14, 2010

Formula of a Hydrate Lab (by Bev)

What is a hydrate?

A hydrate is a molecule that has a definite number of water molecules in its crystal structure.  The crystals seem like a dry solid, but they can be decomposed into an anhydrous salt (without water) and water when heated.  A colour change often happens when the hydrate is decomposed

     ex. calcium chloride hexahydrate: CaCl2*6H2O
                6mol H2O are chemically combined to 1mol CaCl2

What are hydrates used for?

Grains of anhydrous CaCl2 can be used to in damp rooms to absorb the moisture out of he air, thereby forming a hydrate.

Lab 4C: Formula of a hydrate

The mass of a dry crucible was found on a centigram balance & the hydrate used in this experiment, MgSO4*7H2O, was heated in a dry crucible over a Bunsen burner and weighed as well.  Using this information, and mass of the anhydrous salt & water was calculated.

%H2O in hydrate: mass of H2O/mass of hydrate = 2.23g/5.02g x 100% = 44.4%

# of mol of MgSO4 in hydrate:
mass after heating-mass of empty crucible = mass of anhydrous salt
28.90g-26.11g = 2.79g

1mol anhydrous salt = mass anhydrous salt x mol/120.4g = 2.79g x mol/120.4g = 0.02mol

mass of hydrate = mass of crucible & hydrate - mass of crucible = 31.13g-26.11g = 5.02g

mass of H2O = mass hydrate - mass of salt = 5.02g - 2.79g = 2.23g

1mol H2O =  mass of H2O x mol/18g = 2.23g x mol/18g = 0.1239mol

mass of salt/1mol salt = 2.79g/120.4g/mol = 0.023mol

mol H2O/mol salt = ratio of H2O to anhydrous salt = 0.1239mol/0.23mol = .539

Thursday, December 2, 2010

Empirical and Molecular Formula - Isabelle Cheng

Isabelle Cheng
Block 2-2 Chemistry 11
Ms. Chen
Empirical and Molecular Formula
Empirical - smallest whole number ratio of atoms which represent the molecular composition of a species. In other words “species” means elements. It is where a number the numbers can be divided by a certain number. For example, C8H16 it would be reduced down to C2H4. That is the lowest terms in this case.
Here is one example of doing an Empirical Formula question.
Say that you have 54.09% of Ca, 43.1% of O and 2.73% H. What is the empirical formula?
The first step:
Separate them into elements. 
The second step:
Now you want to convert grams into moles. This means you have to write the mass on the side now. 
Ca =  54.09
O = 43.1
H = 2.73
The third step:
Now you have to multiply it by one mole and divide it by the mass of the element (species).
Ca = 54.09/40.1 = 1.349
O = 43.16/ 16= 2.699
H = 2.73/1 = 2.73
The fourth step:
Now divide both by the smallest molar amount.
Ca = 1.349/1.349 = 1
O = 2.699/1.352 = 2
H = 2.73/1.352 = 2
The fifth step:
Now put them together:
Ca(OH)2 which is Calcium Hydroxide!
Now moving onto molecular formula!
Molecular Formula - it is a multiple of many empirical formula and has the real number of atoms that combine to form a molecule!
The formula for this is:
multiple N =       Molar mass
Empirical mass
One example of the Molecular Formula is this:
A hydrocarbon is 84.25% carbon and 15.75% hydrogen and has a molecular weight of 114. What is the molecular formula?
C = 84.25/12 = 7.021
H = 15.75/1 = 15.75
Then you go:
7.021/7.021 = 1
15.78/7.021 = 2.25
Next step:

C = 1 X 4 = 4
H = 2.25 X 4 = 9
Now you know that answer!
It is.....: 
4 C = 48
9 H = 9

Wednesday, December 1, 2010

Percent Composition (by Bev)

Percent composition is the percentage by mass of a species in a chemical formula

ex. 1: percent composition of AgOH (assume there is 1mol)

total MM (molar mass): 107.9 + 19.0 + 1.0 = 127.9
MM of Ag: 107.9 (same numerical value as the atomic mass)
MM of O: 16.0
MM of H: 1.0

% Ag: 107.9g/mol / 127.9g/mol x 100% = 84%
% O: 19.0g/mol / 127.9g/mol x 100% = 15%
% H: 1.0g/mol / 127.9g/mol x 100% = 1%

Check: 84% + 15% + 1% = 100%*
*the percentages should add up to or very close to 100%

ex. 2: percent composition of CO2 (assume there is 1mol)

total MM: 12.0 + 16.0(2) = 44.0g/mol
MM of C: 12.0g/mol
MM of O: 32.0g/mol

% C: 12.0g/mol / 44.0g/mol x 100% = 27.3%
% O: 32.0g/mol / 44.0g/mol x 100% = 72.2%

Check: 27.3% + 72.2% = 100%

ex. 5: percent composition of  O in Cu2CO3

total MM: 123.5g/mol
MM of O: 48.0g/mol

% O: 48.0g/mol / 123.5g/mol x 100% = 39%

ex. 4: a compound contains 5.1 g of Cl, 22.0g of c, an unknown mass of O, & the total mass is 44.1g.  Calculate % composition.

Mass of O: 44.1g-5.1g-22.0g = 17.0g

% Cl: 5.1g / 44.1 g x 100% = 11.6%
% O: 17.0g / 44.1g x 100% = 38.6%
% C: 22.0g / 44.1g x 100% = 49.9%

Check: 11.6% + 38.6 % + 49.9% = 100.1 %

Links to get you started:

http://www.youtube.com/watch?v=qzUMKWhWKm8 (THE BEST LINK!)

Saturday, November 20, 2010

Converting THE MOLE

Mole conversions. They may seem hard, but they're just like unit conversions.

 Things to remember:
-Avogadro's number (SUPER IMPORTANT!!)
-how to calculate molar mass
-the "magic 1" rule

OK. Let's start converting!!

*For these conversions, we'll need to use Avogadro's number: 6.022x10^23*
1. particles-->moles
4.24x10^24 C particle-->moles

4.24x10^24particles               1 mole         = 7.04 moles of C
                                   6.022x10^23 particles

2. moles-->particles
1.34 moles moles CO2-->molecules

1.34moles x 6.022x10^23 molecules = 8.07x10^23 molecules of CO2
                               1 mole

These conversions do not require Avogadro's number. They instead, require a periodic table. Remember how to calculate molar mass!
3. moles-->grams
6.5 moles of C-->grams                                        *Use molar mass of C=12.0g/mol

6.5 moles x 12.0g/mol = 78 grams of C
                      1 mole

4. grams-->moles
8.2 grams of MgCl2-->moles                                *Use molar mass of MgCl2=95.3g/mol

8.2 grams x 1 mole = 0.086 moles of MgCl2
                   95.3 g

YAY!! So now, you should know how to convert from moles<->particles and moles<-> grams.
Just for some extra practice, here are a few questions:

1. How many moles of O2 are in 5.53x10^41 molecules of Ag?

2. How many atoms are in 2 moles of C?

3. How many moles are in 94.0g of Pb?

4. Calculate the mass (g) of 7.42x10^12 atoms of C.

And, lastly, here's a video to summarize the converting stuffs. Don't get distracted by the music!


Written by Jialynn.

Thursday, November 18, 2010

Chapter 4 - Atomic Mass and Avogadro's number - Isabelle Cheng - November 17, 2010

Isabelle Cheng
November 17, 2010
Block 2-2 Chemistry 11
Ms. Chen
Chapter 4 - Avogadro’s Theory and Atomic Masses
The mole is the amount of carbon atoms in 12 grams of carbon. On the other hand, a molar mass is the mass of one mole. For example, a dozen means 12 and so, a mole means 12 grams. Using the periodic table is very necessary for solving these problems. Some examples from using the periodic table is that Iron has an atomic mass of 55.8. Then right away you know that the molar mass of the element is also 55.8g. 
Another example:
AgNO3 - 1 Ag = 1 x 107.9 = 107.9, 1 N = 1 x 14.0, and 3 O = 3 X 16.0 = 48.0 and all of it together is 169.9 grams which means that the molar mass is 169.9 grams.
We also need to include unit conversions in these problems. 
For example, we need to use this equation:
1mol___     molar mass of X
molar mass of X           1mol

Avogadro’s number:
- 6.022 X 10 ^ 23
- large number 

Sunday, November 14, 2010

Experiment 2E: Determining Aluminum Foil Thickness by: Mandy

Lab 2E : Determining Aluminum Foil Thickness
Today, we did experiment 2E. This lab is an indirect measurement of the thickness of a piece of aluminum foil.
There were two formulas involved in this experiment:
Volume of a rectangular solid = length X width X height
Density of a substance = mass / volume
We first took three aluminum foils (15 X15)cm. since the measurement of those square aluminum foils might be off by a little bit, we need to measure the length of two widths and take their average; same as the lengths. Then, we took each aluminum foil to the centigram balance for the measurement of mass.  Now that we knew the lengths, widths, mass, and density (2.70 g/cm^3), the height can be calculated.
In this experiment, accuracy and precision are very important.
An ACCURATE measurement is a measurement that is close to the accepted values.
A PRECISE measurement id a reproducible measurement, thus, more precise the measurement is, the more significant digits it has.
Take the average thickness of those three aluminum foils, and calculate experimental error.
Experimental Error is calculated using this formula:
Experimental Error = abs( your measurement – accepted value) /( accepted value) X 100%
In this case, the accepted value is: 1.55*10^-6 cm

household aluminum foil

a little extra information about aluminum foil :) 
        The shiny side is slightly better reflector of heat. To keep things cold, put the shiny side on the outside [that will reflect incoming heat]. To keep things warm, face the shiny side inward toward the hot food [to reflect the heat that is trying to escape back into the food].

THANK YOU for checking out our blog :D

Tuesday, November 2, 2010

So Like, How Much More "Dense" Can You Get?

Things to remember:
1cm^3 of water=1mL
Density of water=1.0g/mL OR 1000g/L

The formula for density is Density=mass/volume

Density Units

for solids--> g/cm^3
for liquids--> g/mL

Some concepts to think about:

If the density of an object is greater than the density of a liquid, the object will sink
If the density of an object is less than the density of a liquid, the object will float.

Try applying the formula to these practice problems. Remember, D=m/v.

1. An iron bar has a mass of 1225g and a volume of 1.2L. What is the iron bar's density?

2. In a balloon, helium occupies 3.8L with a mass of 4.0g. What is the density of helium?

3. An iron bar has a mass of 250g. If iron's density is 7.86x10^3 g/L, what volume does the bar occupy?

4. A block of beeswax has a volume of 210.0mL and a density of 961g/L. What is the mass of the block?

Fun Fact!!
The density of one brain tissue is 1.05g/mL, which is almost equivalent to the density of water!
How "dense" is that??!?! xD

Written by Jialynn

Thursday, October 28, 2010

Accuracy and Precision

Accuracy-->how reproducible a measurement is compared to other similar measurements
     -->can be described as the correctness of a measurement

Precision-->how close the measurement (or average measurement) comes to the accepted or real value
     -->can be described as the exactness of a measurement

*no measurement is exact
-every measurement is only one's best estimate, as it has some degree of uncertainty
-measurements are only exact when they are counted
  eg. a family of 4

  Absolute Uncertainty

-Uncertainty is expressed in the unit of measurement, not as a ratio or fraction
-there are 2 methods:
Method 1:
1. Make at least 3 measurements
2. Calculate the average between the measurements
3.Calculate the different between the average and the lowest or highest reasonal measurement. The largest difference will be the absolute uncertainty.

Method 2:
1. Make the best precise measurement, estimate to a fraction 0.1 if the smallest segment on the instrument scale.

Relative Uncertainty and Sig. Figs.

Relative Uncertainty=Absolute uncertainty/estimated measurement
-relative uncertainty can be expressed as a percent, or with sig. figs.

Practice reading scales, and identifying the uncertainty of each measurement using method 2.

Watch this video! It'll help. (:
How to Read a Ruler

Written by Jialynn

Tuesday, October 26, 2010

Significant Figures - Isabelle Cheng

Isabelle Cheng
October 26,2010
Block 2-2 Chemistry 11
Ms. Chen
Significant Figures
Significant figures are the number of significant digits in an answer to a calculation. They are approximate answers and the more digits it has the more accurate it may be. 
More facts about significant figures: 
non- zero numbers are significant - for example: 355 has three significant digits, and 24.35 has four digits
however with zeros there are different rules:
zeros put between other digits are always significant
zeros put after another digit but behind a decimal are significant
zeros put before other digits are not significant
zeros at the end of a number are significant only when they are after the decimal point
leading zeros aren’t counted though
Exact numbers - Ex. there are exactly 14 dogs. You can’t have 14.34 dogs. 
Inexact numbers - Ex. the length of the table is 124.34 mm then it is an inexact number. 
Rounding numbers - round answers to the approximate - Example: if the number is over 5 then round it to a high number and if it is under five keep it that number. If the number equals to five the put a five
Adding and Subtracting Significant Digits:
round up to the fewest numbers of the decimal places
ex. 14.982+4.2 would be looking like 14.982 = 19.182
+   4. 2 
same thing for subtracting!
round to the thousands place for the first uncertain digit
ex. 24500+7000 = 31,900 - you have to change it to 34,500 = 34,000
Multiplying and Dividing Significant Digits:
round to the nearest number significant digits 
ex. 32.96 x 2.4 = 79.104 - multiplying
ex. 14.59 ÷ 5 = 2.918 - dividing
Some exercises of significant figures:
Write how many significant figures there are!
1.00 x 1000000

Wednesday, October 20, 2010

Lab 3B; Oct 19th, 2010; by: Mandy Xiao

Separation of a Mixture by Paper Chromatpgraphy

Some vocabularies you should know:
Rf value: the ratio of the distance traveled by the solute to the distance traveled by the solvent
Formula: Rf=d1/d2
D1 = distance traveled by solute
D2 = distance traveled by solvent
*Rf values vary from 0 to 1
capillary action: in this case: when water is moving up the paper

solute front: (in this experiment) food colouring
solvent front: (in this experiment) water

Lab procedure:
Part 1: Setting Up
3 large testubes & 3 Erlenmeyer flasks. Label A, B, C
22cm chromatography paper
Use pencil to draw a line across strips 4cm from one end
Trim the end of the strip
Place 2cm deep water into each test tube

Part 2: Rf values of individual food colourings
Use a glass stirring rod to spot the strip with the colour
Write the colour at the top of the strip
Insert strip in test tube A
Observe the sample spot as the water goes up
Observe solute front and solvent front
Remove strip from the test tube. Immediately draw a pencil line across the top edge of the solvent front
Measure d1 and d2 calculate Rf for sample and record

Part 3: Separation of Mixtures onto Their Components
2nd strip with green colour, 3rd strip with unknown LABEL!
Insert strip in test tube B&C (see procedure in Part 2)
Record data on table 3!

Further Information: 
  The substances (solutes) to be analysed must dissolve in the solvent, which is called the mobile phase because it moves. The paper or thin layer of material on which the separation takes place is called the stationary or immobile phase because it doesn't move.

INTERESTING! Paper chromatography Art!

Sunday, October 17, 2010

How to Separate Mixtures (by Bev)

-basis: different components & properties
-come up with a process that differentiates between components with different properties
     ex. high density/low density

-a substance comprised of more than one substance that is not chemically bonded
-the mixture’s components keep their identities
-the more different the properties are, the easier it is to separate them

-filtration: chooses components by particle size
-flotation: chooses components by density
-crystallization & extraction: chooses components by solubility
-distillation: chooses components by boiling point
-chromatography: chooses & absorbs components at different rates in a fluid mixture
Hand separation
-for solids
-mechanical/heterogeneous mixture can be separated with magnet/sieve
-evaporation: a solid dissolved in a liquid solution
                -liquid evaporates (from boiling) & the solid is left
-for solids that are not dissolved in liquids
-using a permeable filter, pass a mixture with solid particles through
-the solid particles stay on top of the filter because they are bigger than the pores
Filtration Apparatus Diagram
-the filtrate permeates but the residue remains in the filter
-a solid in a liquid
-precipitation: from physical/chemical change, a solute (dissolved substance) is converted into a solid
-flotation/filtration separate the solids
-the desired solid becomes a saturated solution, containing the maximum amount of solute (which the liquid can no longer dissolve
-evaporate/cool: solid becomes pure crystals, which are filtered
Gravity separation
-for solids based on density
-centrifuge rapidly spins a test tube & separates substances of different densities, forcing the denser materials to the bottom
-works best with small quantities
Solvent extraction
-a component moves to a solvent shaken with a mixture
-works best with solvents that dissolve only one component
-mechanical mixture (2 solids): only one solid dissolves in the liquid & the desired solid is left behind
-solution: the solvent is insoluble because it is already present.  It dissolves at least 2 substances & the unwanted substances remain
                -if shaken in the separatory funnel, the liquids from layers
-for a solution of 2 liquids
-heating the mixture triggers the low-boiling components to volatize (vapourize)
-evaporated components collect & condense
-the liquid with the lowest boiling point boils first, the vapour ascends to distillation flask & enters condenser, gas cools to a liquid, & distillate (condensed liquid formed from boiling) is dropped as a purified liquid
-a mixture is passed over a material that absorb some components more than others
-different components pass over the material at different speeds
-mobile phase: sweeps the sample over the stationary phase (ex. Wind sweeping swarm of bees over flower bed)
-can separate extremely complex mixtures
                Ex. Drugs, plastics, flavourings, foods, pesticides
-using very small sample sizes, the analysis can be highly accurate & precies
-the separated components can be collected individually
Sheet/paper chromatography
-stationary phase is a liquid soaked into sheet of paper & mobile phase is a liquid solvent
-some components spend more time in the stationary phase than others
-components appear as separate spots spread out on the paper after drying/”developing”
Thin layer chromatography
-the stationary phase is a thin layer of absorbent (often SiO2 or Al2O3) coating a sheet of plastic/glass
-some components bond to the absorbent strongly; others, more weakly
-components appear as spots on sheet

Practice Problems: Separate!
1)      Coins: by hand
2)      Sand & copper sulphate: solvent extraction with filtration
3)      Salt in water: evaporation/distillation
4)      Sulphur 8 iron fillings: magnet (by hand)
5)      Ink: chromatography

Friday, October 15, 2010

How to name acids (by Bev)

-ACID: a covalent (non-metal + non-metal) bond formed from a negatively charged ion & a hydrogen ion dissolved in water
     -when dissolved in water, ions separate
     -chemical formula for an acid starts with a H (hydrogen)

1) prefix: "hydro"
2) the last syllabe of the non-metal is replaced with the suffix "ic"
3) add the word "acid" at the end

ex. HF = hydrofluoric acid
HCl = hydrochloric acid
HBr = hydrobromic acid

1) "-ate" of the polyatomic anion is replaced with the suffix "-ic"
2) "-ite" of the polyatomic anion is replaced with the suffix "-ous"

3) add the word "acid" at the end

*REMEMBER (this actually works!): "We ate ic-y sushi & got appendic -ite-ous."

ex. HCH3COO = acetic acid
HClO3 = chloric acid
HNO2 = nitrous acid
HClO4 = perchoric acid

beware the exception!
HCN = hydrocyanic acid (it is a simple acid!)


Thursday, October 7, 2010

Lab 2C: Melting & Freezing Points of Pure Substances (by Bev)

In the experiment, liquid dodecanoic acid (C12H34O2) was cooled in cold tap water until it reached a temperature of 25 degrees Celsius, and observations and temperature readings were taken every 30 seconds.  The same process was used for the heating process, except the solid dodecanoic acid was placed in 55 degrees Celsius water.  The purpose of the lab was to find the melting and freezing point of this substance (40 degrees Celsius) by graphing the results and finding the intersection between the heating and cooling processes.  It was also discovered that melting point and freezing point are the same and interchangeable terms.

Writing and Naming Ionic and Covalent Compounds - Isabelle Cheng

Isabelle Cheng
October 7, 2010
Block 1-2 Chemistry 11
Writing and Naming Ionic and Covalent Compounds 
composed of two or more particles (ions) --> oppositely changed
held together by electrostatic forces
electrons are transferred from a metal to a non-metal
Ex. Li+1     O-2
= Li2O
Ex. Ti has 4+ or 3+
Titanium (IV) fluoride ---> TiFl4
Ex. Complex Ions: a group of atoms that behave as one atom.
Na2SO4 ---> Sodium Sulphate
Na3PO4 ---> Sodium Phosphate

share electrons
non-metal with non-metal
- use GREEK prefixes to indicate the number of atoms
*Diatomic molecules: H2,O2,F2,Br2,N2,Cl2,I2
mono - 1
di - 2
tri - 3
tetra - 4 
penta- 5
hexa - 6
hepta - 7
octa - 8
nona - 9
deca - 10
a) CO2 - Carbon dioxide
b) N2O4 - Dinitrogen Tetraoxide

Monday, October 4, 2010

Finding Out About Matter

Science requires increasingly precise observations and detailed inferences. Observation takes too much time that science must be separated into different specialization. Individuals study matter in chemistry. One way to get a better understanding about matter is that you can carefully observe familiar substances, classify and generalize them on regular bases.

We learnt that few of the ways to identify matter is by their colour and taste. The characteristics of matter can also be recognized by the temperature at which matter changes from a liquid to a gas (boiling point) A mixture is twp or more kinds of matter that have separate identities. (impure) The idea of mixture and pure substance are used to describe matter.

Scatter light will appear if one shines a strong light on impure water and solutions Adding alum and lime to the water, then remove suspended particles from water. The water will not appear to have scattered lights Solutions are mixtures that look uniform throughout and do not scatter light Solutions such as sugar and salt in water are to be separated into their component parts using distillation
New techniques are still being developed in order to meet the environmental problems concerning today’s society. Pure substances have constant boiling point

Characteristics of Pure Substance
·         Have a constant boiling point & melting (or freezing) point, unlike most ordinary mixtures
o   Ex. As pure methanol is heated, the temperature gradually rises to 65C & begins to boil.  The temperature remains constant throughout boiling.  However, a mixture of 25% methanol & 75% water boils at around 86C, but the temperature continues to rise as it boils.
Chemical & Physical Changes
·         Chemical changes: produce a new kind of matter with different properties
o   Ex. When sugar is heated in a test tube, it bubbles & forms a black solid while a colourless liquid forms on the walls of the test tube.  The black solid & colourless liquid is both formed from an irreversible process of decomposition that forms new substances with new properties. 
·         Decomposition: type of chemical change when one kind of matter decomposes decomposes (comes apart), forming at least 2 types of matter.
·         Physical changes: no new substances produced
o   Ex. Melting, boiling, & freezing are types of physical changes because the substance has all the same properties after a change of state, and the process is reversible.
Compounds & Elements
·         Element: pure substances that cannot be decomposed
o   Smallest particle: atom
·         Compound: chemically combined elements & can be decomposed
o   Smallest particle: molecule

p.36-39 Macroscopic observations are observations from what an individual can see, feel or smell. Whereas melting point, boiling point, heat of fusion, temperature and mass are all called macroscopic properties. On the other hand, microscopic model is to explain the performance of matter. Matter is made up of atoms. Atoms are usually expressed as spheres because chemists use spheres to help them
understand the microscopic world. Elements only contain one kind of atom. If the particles in an element vibrate, temperature will increase, therefore, an element can exist as a solid, liquid, or gas. Particles made of more than one atom are called molecules. The more outsized the particle is, the higher the boiling point will be. Energy such as heat and electricity are used to break down compound
Compounds can be formed by molecules or ions. Chemists need to check their conductivity in order to determine rather the compound is ionic or molecular.