NOTE:
***************
In this post, we have included all the missing posts from the past month. The dates are not here, but all the information and class notes should be in this post. If we are missing anything, let us know and we will include it in the next post***********
density, molar volume & molar volume procedure lab, percent composition, empirical formulas, concentration, dilution
- Chem11403 Team (Joanne H, Roxanne S, Christian B)
______________________________________________

DENSITY

Density= Mass/Volume Mass= Volume*Density Volume= Mass/Density



Finding the density of gases at STP is easier then finding it at solids and liquids because we know that the volume of the gas will be 22.4L. We can also find the mass by checking our periodic tables. After you just plug in the numbers and solve for the density of the gas.

Molar Mass g / 22.4L/mol = ?g/L

density of oxygen =
Density = Mass / Volume

32g / 22.4L/mol = 1.43g/L at STP

Finding the density of solids and liquids are much harder you do not have all the information you need such as the volume.

use the formula g / molar mass of element. 6.02 x 10 to the power of 23 subscripts
Density ---> Mass ----> Moles ---> Molecules ----> Atoms

Example: the density of Boron (solid) is 2.34g/mL how many molecules are in a 60ml piece?
2.34 g/ml * 60.0 mL = 140.4g
140.4 g * 1mol / 10.8 g = 13mol
13mol * 6.02 * 10 to the power of 23 / 1mol = 7.84 * 10 to the power of 24

_________________________________________________
MOLAR VOLUME

D=M*V 22.4L/1mol subscripts
Density Volume (STP) Atom
-Molar Mass = g/mol
- Molar Volume = L/mol

Example
-A sample of unknown gas contains 0.635 mol and occupies a volume of 482mL. Determine the molar volume.
482mL x 1L = 0.482L = 0.759 L/mol
1000mL 0.635mol


MOLAR VOLUME PROCEDURE LAB

1. Fill the sink with 3\4 of water and place the lighter in so no air remains.

2. Dry the lighter and weigh it

3. Place the gradulated cylinder into the sink and once again make sure there is no air.

4. Place the lighter in the water under the gradulated cylinder and press the button. Add 10 milileters of gas.

5. Record the volume of the gas and the mass of the lighter.





PERCENT COMPOSITION

Perecent Composition
-means the % mass of each element in a compound.

Examples
Find the % of composition of K2Cr2O7

2K-78.2
2Cr-104
2O-11.2
Total mass: 294.2

Formula: Mass of Element/Total Mass

2K-78.2/294.2=26.6%
1C-104/294.2=35.4%
3O-11.2/294.2=38%
100%

EMPIRICAL FORMULA
Find the total mass of carbon in a 3.0kg sample of Ethanol (C2H6O)

2C --> 24 52.1%
6H --> 6 13%
1O --> 16 34.7%
46g/mol
(0.521)(3.0kg) = 1.57kg

Empirical Formulas = gives the whole number ratios of elements in a given compound. Molecular Formula gives the actual numbers.

Emprical Formulas:

Molecular Emprical
P4O10 P205
C10H22 C5H11
C6H18O3 C2H6O
C5H120 C5H12O
N2O4 NO2

A sample of an unknown compound is analyzed and found to contain 8.4g of C, 2.1 of H, and 5.6g of O

elements mass(g) atomic mass moles Moles/smallest mole
C 8.4 12 0.7mol 0.7/0.35=2
H 2.1 1 2.1mol 2.1/0.35=6
O 5.6 16 0.35mol 0.35/0.35=1

RATIO:
0.5 2
0.33 or 0.66 3
0.25 or 0.75 4
0.2, 0.4, 0.6, 0.8 5



MOLAR CONCENTRATION
Concentration -- solution-- a homogeneous mixture

Concentration: Amount of solute
Amount of Solvent

Some units for Concentration: g/ml ; g/l ; mg/l ; mg/ml ; ug/l {{Not very useful to us}}
**The most common(and useful) units are mol/l = molarity <-- molar concentration M= mol/L mol= M times L L= Mol/M all of these formula are only for aqueous solutions, not gases Example: Stephanie dissolves 40.0g of NaOH in enough water to make 200ml of the solution. What is the concentration? Concentration= 40.0g = 0.200g/ml 200ml 40g times 1 mol =1.0mol 40.0g M= 1.0mol = 5.0mol/L 0.200L Example: Kira wants to evaporate some 3.0M NaCl to obtain 26.325g of NaCl, What volume should she evaporate? 26.325g times 1mol times 1L = 0.15L <--- = 150ml 58.5g 3mol

DILUTION
-when you add water con’c decreases
-if the volume is doubled con’c is halved
- Volume | Con’c | moles
6.0 | 2.0 | 6L x 2 = 12 mol
12.0 | 1.0 | 12 x 1 = 12 mol
48.0 | .25 | 48 x .25 = 12mol
-N1 = N2; C1V1=C2V2
-Karol adds 150.0mL of water to 50.0mL of .60M HCl. Find [HCl].
V1= 50mL C1= .60M V2=200.0mL C2=?
C1V1=C2V2
C2=.15M
-Jesse adds water to 100.0mL of .35M to a final volume of 400.0mL. Find the [HF].
C1=.25M V1=100mL C2=? V2=400mL
C1V1=C2V2
C2=.0625M
-Cheyenne dilutes 60.0mL of 0.40M HNO3 to 0.15M. What is the volume? How much water did she add?
C1=.40M V1=60mL C2=.15M V2=?
C1V1=C2V2
V2=160mL
160-60=100mL

Here is an interesting video I found on youtube:


DILLUTION PT. 2
(Making directions for experiment procedures)
**First step is to find the amount of mass you will need.
Apply proper unit conversions to achieve this.

Example:
Jeremy is asked to make a 0.55M solution of K2SO4.If he needs 250mL what procedure should he use?
2K-78.2
1S-32.1
4O-64.0
174.3g/mol
250mL x 1L = .25L
1000mL
.25L x .55mol = .1375 mol x 174.3g =24.0g

STEPS:
1. Measure 250mL of water
2. Weigh 424.0g of K2So4
3. Add K2SO4 to water
-Give directions to make 2.00L of 6.0M NaOH
1Na-23.0
1O- 16.0
1H- 1.0
40.0g/mol
2L x 6mol = 12mol x 40g = 480g
ANSWER:
1. Measure 2.00L of water
2. Weigh 480g of NaOH
3. Add NaOH to water and stir until dissolved

OCTOBER 26, 2009 [JOANNE]

GASES AND MOLES:

The volume of a balloon occupied by a certain gas depends on the temperature & pressure.

Standard Ambient Temperature & Pressure (SATP)

• 24.8L / 1mol
• 25°C and 100kPa

*We will focusing much on SATP yet

STP:

The molar volume of any gas at STP is 22.4L

• 22.4L / 1 mol
• 1 mol / 22.4L

EXAMPLES:

Ø Find the volume occupied by 0.060 mol of c02 at STP
0.060mole x 22.4L/1mole = 1.3L

Ø Find the mass of a 200.0 mL sample of NO2 at SATP
STEPS:

o Note that the 200.0mL given is not in L

o Before we can convert it by SATP, we must change the mL to L.

o We divide 200.0mL by 1000 for our conversion factor from mL to L.

o Then we use SATP (22.4L / 1 mol )

o To finish the question in grams, we multiply by 46.0g because N = 14.0g + !6.0(2) O = 46.0g

o The Equation:

§ 200.0mL x 1L/1000mL x 1mol/22.4L x 46 g/ 1 mole = 9200/22400 =.41g

Ø Find the volume occupied by 22.0g of CO2(g)

o 22.0g x 1mol / 44.0g x 22.4L / 1mol = 11.2 L

Oct. 21st /2009 [THE Christian Bondoc]

Herro.

Atomic Mass

• The mass of 1 mole of atoms in an element
• The mass of 1.0 mol of 'C' atoms is 12.0g
• The mass of 1.0mol of 'Ca' atoms 40.1g.

Molecular Mass

• The mass of 1.0 mole of molecules of an element or compound

N2, O2, H2, Br2, Cl2, F2, I2, P4, S8

• Assume the all the rest are monoatomic

(I apologize for the size, when i made it large, part of it got cut off)

Finding the molar mass of compounds

H2O

2 H = 2(1.0) = 2.0

1 O = 1(16.0) = 16.0

18.0g/mol

Find the molar of ammonium phosphate

NH4+

PO4 3-

(NH4)3PO4

3 N = 3(14.0) = 42

12H = 12(1.0) = 12

1P = 1(31.0) = 31

4O = 4(16.0) = 64

149g/mol

Converting mass <----> Moles

Conversion factor g/mol or mol/g

Find the mass of 2.5mols of water

H20 -> 18.0g/mol

(1mol/18.0g) x (1/2.5mol)

(1 mol/18.0g) x (1/2.5mol)

= 1/45g

= 45g

Find the number of moles in a 391g samples of nitrogen dioxide

NO2

1N = 14(1.0) = 14

2O = 2(16.0) = 32

45g/mol

(391g / mol ) (1 / g) =

(391g / mol) (1/ g) =

(391/mol) (1/46) = 8.5mol

1) Find the mass of 2.5 moles of water.

2H: (2)(1)

1O: (1)(16)

= 18g/mol

2.5 mol • 18.0g /mol= 45g

2) Find the number of moles in a 391g sample of nitrogen dioxide.

1N: (1)(14)

3O: (3)(16)

=45g/mol

391g • 1mol/45g = 8.5 mol

3) 3.6kg of sulphur trioxide = ? mol

11S: (1)(32.1)

3O: (3)(16.0)

= 80.1 g/mol

3.6kg • 1000kg/1g = 3600 g

3600g • 1mol/80.1g = 45mol

October 19, 2009 [JOANNE]

THE MOLE:
1 mole = 6.02 x10²³ -- Avogardo's Number

HYDROGEN BOMB EQUATION:
2H₂ + O₂ = 2H₂O

(2H₂ molecules 1 O₂molecule 2 molecules of H₂O)

12,02 x 10²³ of H₂ molecules + 6.02 x 10²³ molecules of O₂ = 12.04 x 10²³ molecules of water
-> 2 mole of H₂, 1 mol of O₂, 2 mol of water

HOW GASES COMBINE
- John Dalton
- Look at masses of gases

(Ex.) 11.1g of H₂ reacts with 88.9g of O₂
(Ex.) 46.7g of N₂ reacts with 53.3g of O₂
(Ex.) 42.9g of C reacts with 57.1g of O₂

JOSEPH GAY-LUSSAC:
- combine gases based on volume
Ex.) 1L of H₂ reacts with 1L of Cl₂ = 2L of HCl
Ex.) 1L of N₂ reacts with 3L of H₂ = 2L NH₃
Ex.) 2L of CO reacts with 1L of O₂ = 2L CO₂
**gases combine in simple whole # ratios


AVAGARDO'S NUMBER
- Equal volumes of any gas at a constant temperature and pressure contain equal number of molecules

OCTOBER 15, 2009 [JOANNE]

We had our chapter 2 nomenclature test.

//.
Here is a random picture of a periodic table tattooed onto an arm!! :)

OCTOBER 13, 2009 [JOANNE]

Today we did an experiment in class.
We started by measuring the weight of a solid hydrate, which was a red colour. We then placed it in a test tube and heated it with a bunsun burner until the hydrate inside melted into a light blue liquid colour.
Our percentage error was 11%

NEXT CLASS:
Test!

-Joanne H.

OCTOBER 8, 2009 [JOANNE]

Today, Mr Doktor showed the class how to use a bunsun burner. He went over some basic safety rules and told us to always inspect the hose before using it and to never leave the gas on.
When he turned off the classroom lights, the bunsun burner emited a blue flame.

Mr. Doktor also did an experiment in class where he mixed sulfuric acid with sugar.

Classmates were complaining about the concentrated smell of burning sugar!

ACIDS
-Solid, liquid, or gas at SATP (standard ambient temp & pressure)
-form conducting aqueous solutions
-dissolve in water to produce H+
-taste sour
BASES
-turn red litmus blue
-slippery
-nonconductive
-dissolve in water to produce OH-
NAMING ACIDS
~acids are aqueous (dissolved in water)
~hydrogen compounds are acids
-HCI(aq)-->Hydrochloric acid
-H(2)SO(4)(aq)-->sulfuric acid
~Hydrogen appears first in the formula unless it is part of a polyatomic group
~CH(3)COOH(aq)-->acetic acid

HI(aq)--> Hydro Iodic Acid
-classical rules use the suffix IC and/or the prefix HYDRO-

sulfuric acid
hydrochloric acid
~IUPAC system uses the aqueous hydrogen compound

HCI(aq)Aqueous hydrogen chloride

NAMING BASES
~for now, all bases will be aqueous solutions of ionic hydroxides
-NaOH
-Ba(OH)2
~use the cation name followed by hydroxide
-sodium hydroxide
-barium hydroxide

these are some examples we went over in class
-H(3)PO(4)(aq) Phosphoric acid
-HNO(3)(aq) Nitric acid
-HNO(aq) Nitrous acid
-Mg(aq) Magnesium hydroxide
-HBr(aq) Hydrobromic acid
-HOOCCOOH(aq) Oxalic acid

HOMEWORK:
Read over lab

OCTOBER 6, 2009 [ROXANNE]

Today in class, we learnt about naming hydrates and their prefixes.

• copper sulfate & sodium sulfate --> without water the compound is often preceded as 'anhydrous'
  -these crystals contain water inside them which can be released by heating

TO NAME HYDRATES:

• write the name of the chemical formula
• add a prefix indicating the number of water molecules
  1-mono
  2-di
  3-tri
  4-tetra
  5-penta
  6-hexa
  7-septa
  8-octa
  9-nona
  10-deca
  
  

Examples:

Name the following compounds:

1. Cu(SO4) - 5H2O ---> copper (2) PENTAhydrate
2. Li(ClO4) - 3H2O ---> Lithium perchlorate - TRIhydrate
3. what is the chemical formula of Nickel(2)sulfate hexahydrate?
   Ni(SO4) - 6H2O

MOLECULAR COMPOUNDS

• composted of 2 or more non-metals
• low melting point and boiling point
• share (not exchange) electrons
• usually end in 'gen' (hydroGEN, oxyGEN, nitroGEN)
• 7 molecules are DIATOMIC - 2 of the same elements
  -H2, N2, O2, F, Cl2, Br2,I2
• P4, S8 --> Polyatomic
• water - H2O
• hydrogen peroxide - H2O2
• ammonia - NH3
• glucose - C6H12O6
• sucrose - C12H22O11
  *KNOW THE FIRST FIVE!!!!!*
• methane - CH4
• propane - C3H8
• octane - C8H18
• methanol - CH3OH
• ethanol - C2H5OH
• ethane - C2H6

OCTOBER 2, 2009 [CHRISTIAN]

SEPARATING MIXTURES
Heterogeneous Mixtures: (done by hand)

• filtration


• distillation:


• crystallization


• chromatography
  *All are physical changes

REVIEW OF ATOMS

1. matter is made up of atoms


2. molecules are groups of atoms held together by electrical bonds


3. ions are atoms or molecules that have an electric charge
   -positive ions are cations
   -negative ions are anions


4. atoms are made up of 3-subatomic particles


5. Protons:
   -positive charges
   -inside nucleus
   -each element has a different number of protons
   -protons = atomic number


6. Neutrons:
   -neutral
   -inside nucleus
   -nearly same mass as protons
   -adding or removing neutrons does not change the element


7. Electrons:
   -negatively charged
   -located outside the nucleus
   -1800 times smaller than protons
   -chemical reactions occur between electrons in different atoms/compounds

PROPERTIES OF THE PERIODIC TABLE

• FAMILIES(or groups) form vertical columns
  -all elements of a family have similar traits and characteristics


• PERIODS are horizontal rows. Elements gradually chane from metals to non-metals as you move from left to right

ELEMENTAL INFORMATION

• exceptions are:
  copper - cuprum
  gold - aurum
  iron - ferrum
  lead - plumbum
  silver - argentum


• first letter is always CAPITAL


• second letter is always LOWER CASE

CHEMICAL NOMENCLATURE

• naming chemical compounds has been a very difficult task and different systems have been used through the centuries


• today the most common system in IUPAC for most chemicals
  -ions
  -binary ionic
  -polyatomic ions
  -molecular compounds
  -acids

CHEMICAL FORMULAS

• be aware of the differences between ion and compound formulas
  - Zn^2+ (2+ is the ion charge)
  -BaCl2 (the 2 is lower)(number of ions)

NAMING IONS

• for metals use the name of the element and add ion
  -Al^3+ = aluminum ion


• for non-metals, remove the original ending and add "-ide"
  -F- = Fluorine becomes Fluoride


• polyatomic ions have special names

BINARY IONIC COMPOUNDS

• ionic compounds contain two elements - one metal and one non-metal


• metallic and non-metallic ions bond together


• election is transferred from the metal to the non-metal


• net charge must be zero - total positive charge - total negative charge

NAMING BINARY IONIC COMPOUNDS

• metal name + first part of non-metal "-ide"

STEPS

• write formula


• criss-cross charges


• reduce ion numbers to lowest common multiples

SEPT 28 + 30, 2009 [JOANNE]

SEPTEMBER 28, 2009

Today we watched a video.
The class was to identify any chemical / physical changes shown in the video.
________________________
PAGE 73: SECTION REVIEW 2-3
_________________________

LAW OF CONSERVATION OF MATTER:
- Matter can neither be created nor destroyed

A. CLASSIFICATION OF MATTER

• We can divide matter into two types:
• Homogeneous Substances
  
• Heterogeneous Substances

• HOMOGENEOUS: consists of only one visible component
  -distilled water, oxygen, graphite


• HETEROGENEOUS: contain more than one visible component
  -chocolate chip cookie, granite

B. PURE SUBSTANCES

there are two types of pure substances:
-Elements: substances cannot be broken down into simpler substances by chemical reactions (oxygen, iron, magnesium)
-Compounds: substances that are made up of two or more elements and can be changed into elements (or other compounds) by chemical reactions (water, chocolate chips, sugar)

FOUR STATES OF MATTER:

STATE:	EXAMPLE:	CHARACTERISTICS:
1.) SOLID	Gold	- High density - Density little affected by pressure - Holds its own shape in a container
2.) LIQUID	Water	- High density + little affected by pressure - Adopts shape of its container
3.) GAS	Nitogen	- Low density - Density depends on pressure - Expands to fill its container
4.) PLASMA	Interior of the sun	- Density depends on pressure - Expands to fill its container - Exists only at high temperature

CHEMICAL PROPERTIES
-Substance cannot be observed without altering the substance
- Ex. Flammability

CHEMICAL CHANGES:
Ex.// wood burning, food cooking, rust

PHYSICAL PROPERTIES
-Substance can be observed without altering the identity of the substance
Ex.// Density, colour, melting point

PHYSICAL CHANGES:
Ex.// crushing, tearing

_________________________________

SEPTEMBER 30, 2009

ELEMENT:
-Substance can not be separated into simpler substances

COMPOUND:
- Two or more elements combined together in a chemical reaction

ELECTROLYSIS: - "to tear apart with electricity"
-An electric current passes through a substance. If the substance is a compound, it will split into two or more elements
Ex.// Waster ---> Hydogen + Water

Elements and compounds are pure substances because it has a unique set of chemical and physical properties.

Posted by: Joanne

SEPT 17, 2009 - Unit Conversions review / Salt water lab [CHRISTIAN]

Class started out, with a review on unit conversions, giving us a few examples on the board.

Examples:

1. Change 120kg into mg.

(120kg/1)(1000g/1kg)

(120kg)(1000g/1kg)

= 120 000 000mg(in SD: 1.2×10^8)

2. 7.25 L/S into mL/min

(7.25L/S)(60 S/min)(1000mL/L)

(7.25L/S)(60 S/min)(1000mL/L)

(7.25/s)(60 /min)(1000mL)
= 435000mL/min(SD: 4.35×10^5)

3. 174kg/s into Mg /h

(174kg/s) (3600s/h)(MG/1000kg)

(174kg/s)(3600s/h)(MG/1000kg)

(174)(3600)(MG/1000)
= 626.4 Mg/h

Lab:
The experiment consisted of salt, and water, and to figure out what is the maximum amount of table salk you candissolve in 200mL of water(Although wec hanged the volume of water)

The lab was to figure out the amount of salt that can be a volume of water. First trial was with 10mLs, 2nd trail was with 20mLs, 3rd was with 30mLs and the last, and 4th trail was with using 40mL. Then the groups were suppose add salt slowly and stir, while using the electronic scale until salt cannot be mixed anymore and you could see bits of salt floating in the water. At that point, you minus the previous weight with the current weight, and you get the mass of salt. Groups were suppose to do that with 4 trials and draw out a graph, and using that graph make a line of best fit. And using that line of best fit, take an educated guess at the amount of salt 200mL of water can hold.

Posted By: Christian B

SEPT 15, 2009 [ROXANNE]

Significant Figures and 0’s

Two tutorials on significant figures two supplement the “Significant Figures and 0’s” section of these notes.
[youtube=http://www.youtube.com/watch?v=5UjwJ9PIUvE&feature=related]
[youtube=http://www.youtube.com/watch?v=PNH7_nDE6SQ]

For a more in-depth explanation of “Significant Figures and 0’s” http://en.wikipedia.org/wiki/Significant_figures#Identifying_significant_digits is a good resource.

► All digits from 1 – 9 are significant, no matter where they are in a number.
► Zeroes between the digits 1 – 9 are significant.
→e.g. 3009 has 4 sig digs 140012 has 6 sig digs
► “Leading zeroes” (zeroes in front of a number) are not significant. They are “place holders”.
→e.g. 0.00231 has 3 sig digs 0.1003 has 4 sig digs
► If there is NO decimal point in the number, then trailing zeroes (zeroes at the end of a number) are not significant.
→e.g. 100 has only 1 sig dig 45300 has 3 sig digs
► If there IS a decimal point in the number, then trailing zeroes (zeroes at the end of a number) are significant.
→e.g. 103.00 has 5 sigs digs 0.02480 has 4 sig digs 250. has 3 sig digs

Concerning Calculators:
►Do not round with your calculator
→Round your answer only once
→Round only to the correct number of significant digits
►Do not use ^ on your calc, it does not recognize the order of operations in all cases
►Buttons used for scientific notation in your calc can go by these names: EXP, EE, x10, S.N.

Expressing Error
►Error is a fundamental part of science
►you need to have a way to tell in which cases an error is and is not important
e.g. If you are 80cm off measuring a person’s height it makes a big difference, but not when you are 80cm off when measuring the distance from Vancouver to Kamloops
►There are usually 3 reasons for error
→Physical errors in the measuring device (the device is not accurate)
→Sloppy measuring (you can avoid this one with care)
→Changing ambient conditions
→→e.g. This means the measuring device is altered(such as a metal meter stick because they expand or contract in different temps)
►Error is taken to be half the smallest division on your measuring device

Calculated Errors
►There are two different possibilities:
1.Absolute error
2.Percentage error

Absolute Errror
►This is how off you are from the actual answer
ex.50km off from 400k to Kamloops
►Equation: absolute error= measured-accepted
►positive number means your over the accepted value
►negative means your under the accepted value

►Accepted value can also be your predicted value
►Percentage error is used to measure to determine the importance of the difference

Percent Error
(Song video to remember the equation)
[youtube=http://www.youtube.com/watch?v=DmB5ZuYhFmE&feature=related]
►Most common mistakes are made here
►Equation:
Percent error= absolute error/ accepted value
►Equation an a calculator:
%error= ([measured-accepted]/accepted)x100
%error= ([observed-theoretical]/theoretical)x100
►The “( )” in the equations above represent absolute value(even if answer is negative it can be switched to a positive)

Practice/Example Question:
You measure the weight of an orange to be 15n. the actual weight is 17.5 N, what is the % difference? Round to the nearest tenth.
ANSWER
-14.28%>round to>14.30%
Process

Dimensional Analysis
want to know what 100km/h is in miles/hour? Read these notes!
►Conversion rates help us find this out, they never change unlike currency
► Just like converting between currencies in chemistry it is usually necessary to convert between units
►This process is called dimensional analysis

STEPS
1. Find a unit equality
2. Find the conversion factors
3. Apply conversion factors
4. Cancel units
►You can skip 1,2,3 as long as 4 is correct

Practice/Example Question:
How many miles are there in 120km?
1.[unit equality] 1mi=1.6km
2.[conversion factor] 1=(1mi/1.6km)
3.[apply the conversion factor] (120km)(1mi/1.6km)
4.[cancel units] (120)(1/1.6)
[answer] 75

Sept 11, 2009 [JOANNE]

Todays class we got 2 handouts: a picture of the periodic table (have not yet discussed) and a picture of some whimis symbols for lab safety. We are expected learn and memorize them because it will be on the chapter test.

____________________________________________________________________________________________

Class Notes:

MEASUREMENT AND CHEMISTRY:

• In Scientific Method, accurate measurements are essential.
• The most common system in use today is the SI system.
• The metric system started about 300 years ago in France.

7 FUNDAMENTAL UNITS:

1. Mass (kg) - Kilograms
2. Distance (m) - Meters
3. Time (s) - Seconds
   
4. Temperature (mol) - Billions
   
5. Amount (K) - Kelvin
6. Current (A) - Amps
7. Luminosity (cd) - Candela

PREFIXES & SI UNITS:

Meter ---> Kilometers
1000m -> 1km

1 ---------> m

10^1 = (da) m
10^2 = (h)m
10^3 = (k)m
10^6 = (M)m
10^9 = (G)m

0.1m------------------> 10^-1 dm
0.01- 10^-2 = cm
0.001- 10^-3 = mm
0.0000001- 10^-6 = um - micro
0.00000000001- 10^-9 nm - nano
0.0000000000001- 10^-12 (p)m - pico

SIGNIFICANT DIGITS:

1. Digits 1-9 are always significant
2. Zeros are significant if they are right of a decimal.
   • Example: 2.010 - significant
     0.002 - insignificant

_____

CLASS HOMEWORK:

Practice questions #7-8 Page 30 and Page 33 #9-10

September 9, 2009

Today we learned about lab safety and scientific method.

Scientific method goes by this flowchart
Observation-> Question-> Hypothesis-> Design experiment-> Test it-> Conclusion

If the results are different then simply design a different experiment for the next question, if the answer is correct redo the experiment to see if the answers are consistent.

For lab safety the most important rules are:
- No horse play
-listen to instructions
-wear safety gear
-don’t eat/drink in the lab
- be aware of where the safety equipment is (ex. eye wash)
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Class HW:
PG 14-15 read
Do questions 1-3 on pg 15

-Roxanne.S-

NEW BLOG

New official blog for Chemistry 11 Block A

Joanne Ho, Roxanne Sobieski, and Christian Bondoc