DECEMBER 14

Today in class we started the section on balancing equations. I shall finish this tomorrow once I bring my notes home.

HOMEWORK:Pg. 284 #1 & 2 , Pg 281#1-4

DECEMBER 7

Today in class we were interrupted by an earthquake!!!!.... a fake one ofcourse !
After the class got settled we quickly started our LAB :) -->DILUTION AND CREATING SOLUTIONS

Mr. Doktor gave us our last set of notes before our exam.

FINDING MOLECULAR FORMULAS
-you need the empirical formula and molar mass

Example: The empirical formula for a compound is CH2O and the molar mass is 60.0 g/mol. Find the molecular formula (work is shown below)



Example: The empirical formula for a compound is C2H6O and the molar mass is 138g/mol. Find the molecular formula. (work is shown below)

STOCK SOLUTION

--are con'c solutions from suppliers

12 M HCl

15M HNO3

18M H2SO4

Homework: STUDY FOR TEST (thursday, December 10)

WHAT DOES THE TEST CONSIST OF?

-all mole conversions<---avogadro's number

-empirical formulas

-percent mass & percent composition

-concentration

-molar volume

-dilution

-molecular formulas

December 3, 2009

Giving Directions

Making solutions ----> figure out mass to add

EXAMPLES:

A.Kristina is asked to make a 350mL solution of K2SO4 with a concentration of 0.65M. What STEPS will he follow?
V= 350mL = .350L

Molar Mass: K2SO4 = 174.3g/mol

0.65mol/L x (0.350L) = 0.2275

(o.2275mol) x 174.3g/mol = 39.65g

These are the STEPS we get when we refer to the work above:
1. Weigh 39.65 of K2SO4
2. Measure 350mL of water
3. Stir your solution


B. Give directions to make 3.00L of 5.0M NaOH.
Molar Mass: NaOH = 40g

5.0mol/L x (3.00L) = 15mol

(15mol) x 40g/mol = 600g

These are the STEPS we get when we refer to the work above:
1. Weigh 600g of NaOH
2. Measure 3.00L of water
3. Stir your solution

Dilution of Solution

When you add water the concentration decreases.
If volume is doubled, concentration is halved.



EXAMPLES:

A.Heidi adds 200.0mL of water to 60.0mL of 0.80M HCl. Find the final con'c of HCl.
C1 = 0.80M
V1 = 60.0mL
C2 = ?
V2 = 260mL

0.80M/L x (0.060L) = 0.048 mol

C1V1/V2 = C2 = (0.80)(60.0mL)/260mL = 0.18M


B. Alanna adds water to 200mL of 0.45M HF. He gets a final volume of 500mL. Find the final[HF].
C1 = 0.45M
C2 = ?
V1 = 200mL
V2 = 500mL

C1V1 = C2V2
(0.45)(0.200) = C2(0.500L)

(0.45)(0.200)/0.500 = C2

C2 = 0.10M

C. Sara dilutes 60.0mL of 0.40M HNO3 to 0.15M. What is the final volunme?

C1 =0.40M
V1 =60.0mL
C2 =0.15M
V2 =?

C1V1 = C2V2

(0.40M)(0.060L)/0.15M = V2
V2 = 0.160L = 160mL

-->How much water was added?
160mL - 60mL = 100mL

Today in class we got our mid-term results back. Afterwards we learnt about CONCENTRATION

WHAT IS A SOLUTION?

a solution is a homogeneous mixture.--> not being able to differ the different substances mixed

WHAT IS A SOLVENT?

component presented in larger component

WHAT IS SOLUTE?

the component present in the SMALLER amount.

for example: SALT in WATER. water= SOLVENT salt= SOLUTE :)

WHAT IS CONCENTRATION?

the amount of solute over the amount of solvent

(i do not have my notes with me right now, but i'll complete it tomorrow!)

- kristina

ALSO!!** don't forget.....in order to figure out the concentration or anything else, we need to know how to get MOLES. * don't forget about moles :)

ENJOY THIS VIDEO :)

homework: unit 2 homework questions 25-28

NOVEMBER 17 2009

(Heidi)

NOVEMBER 13 2009

EMPIRACAL FORMULAS

Find the total mass of Carbon in a 3.0kg sample of ethanol.
(C2H5OH)
^ ^ ^ ^
2(12) 5(1) 1(16) 1(1)=46.0g/mol

C2H6O<--Empiracal Formula Empiracal Formulas are simplest formulas Molecular Empiracal
P4H10 ------> P2H5
C1oH22 ------> C5H11
C6H18O3 ------> C2H6O
N2O4 ------> NO2
(*Think lowest common denomenator*)

Empiracal Formulas the simplest whole number ratios in a compound
Molecular formulas show the actual atoms/ and ...(guys finish the sentence.. i can't read my writing)

FINDING EMPIRACAL FORMULAS:
A sample of an unknown compound is analyzed and found to contain 8.4g of 'C', 21g of 'H', and 5.1g of 'O'
Find the empiracal formula


Example: 0.888 grams of a compound made up of C2H7O
C, H, and O are found to contain 0.576g of C and 0.120g of H
a) Find the mass of oxygen
b) Find the empiracal formula

0.120g

+ 0.57g

-----------

.696

*.696

-.888

----------

.192-->oxygen

C=.192 x mol=0.48-->4

H=.120xmol=.12--->10

O= .192xmol=.012-->1

Homework: Worksheet questions # 22, 23, 24

NOVEMBER 10

In the beginning of class Mr. Doktor gave us the answers to the 10-2 worksheet # 1-30
But ofcourse right after that we got some notes on the board !

PERCENTAGE MASS OF ELEMENTS IN COMPOUNDS
Molar Mass: g/mol
Molar Volume: 1/mol

example: 0.0263 mol of gas occupies a volume of 482.0 mL
calculate the molar volume. Is this at STP?
Molar Volume=0.482L/0.0263 mol = 18.33 L/mol
--> Find the % carbon by mass in ethane (C2H6)
C2H6
2(12)+6(1)=30g/mol
%C 24g/mol /(over) 30g/mol x100= 80%
%H 6g/mol / (over) 30g/mol

--> Find the % mass of each element in K2Cr2O7
2K= 78.2g/mol
2Cr= 104g/mol
70= 112g/mol/(over) 294.2g/mol
%K= 78.2g/mol /(over) 294.2g/mol =26.6% %Cr=104g/mol /(over)294.2g/mol =35.4%
%O= 112g/mol /(over)294.2g/mol =38.1%

FINDING MASS OF AN ELEMENT IN A GIVEN SAMPLE

example: Find the mass of carbon contained in a 25.0g sample of CO2

1C= 12g/mol %carbon= 12g/mol /(over) 44g/mol =27.3%
2O=32g/mol /(over) 44g/mol (o.273)(25.0g)=6.82g

32/44= (0.727)(25g)=18.12g
25g-6.82g=18.12g

Find the mass of Potassium, carbon and oxygen contained in 450.0 of K2CO3--138.2g/mol
K=254.7g K=56.6%
C=39.2g C=8.7%
O=156.2g O=34.7%

After taking notes Mr. Doktor gave us time to do our homework while checking around if we had completed the worksheet 10-2

Homework: worksheet # 20 & 21

November 6, 2oo9

This class we did a review on moles by using the new chart that Mr. Doktor drew for us!



Here is the chart we used:

Here are a few questions that Mr. Doktor reviewed with us:

1. Liquid Mercury has a density of 13.55 g/mL. Find the volume of Mercury occupied by 1.806 x 10^27 atoms of Hg.
   molecules -> moles -> mass
   
   1.806 x 10^27 moles x 1mol/6.02 x 10^23 moles x 200.6 g/1mol = 601800 g
   
   V = m/d = 601800/13.55 g/mL = 44413 mL = 44.41 L
2. A 5.00 mL sample of lead is known to contain 0.274 mol of Pb. Calculate the density of lead.
   
   0.274 mol x 207.2 g/1mol = 56.773 g
   
   D = m/v = 56.773 g/ 5.00mL = 11.4 g/mL
3. 250 mL of a gas which is known to contain only fluoride and sulfur has a mass of 1.63 g at STP.
   a. Find the molar mass:
   250mL/1 x 1L/1000mL x 1mol/22.4L = 0.0116 mol
   
   1.63g/0.0116 = 146 g/mol
   
   b. Find the chemical formula
   S(x)F(y) = 146 g/mol
   
   x=1 & y=6

November 4, 2oo9

November 2, 2oo9

This class we worked on the

Molar Volume Lab

.

The problem of the lab was to experimentally determine the molar volume of a gas

For homework we had to do the prelab questions. What we found was that the chemical formula for butane is C4H10 and also that the molar mass of butane is 58.124 g/mol.

Also for homework we had to write out a procedure for the lab before doing it, and this is what we wrote:

1. Weigh the lighter on the weigh scale and record it
2. Fill your sink with water, about three inches from the top
3. Place the 100 mL graduated cylinder under water until there are no more bubbles
4. Put the lighter under water at the opening of the graduated cylinder and carefully release butane from the lighter (about 10 mL)
5. Record the volume to the nearest mL
6. Dry the water from the light and then measure the mass of the lighter again on the weigh scale

What we observed during the lab was that after adding the butane gas, the amount of water in the beaker decreased.

Homework: Do analysis and conclusion questions from the lab! =)

November. 4,2009

In the beginning of class, Mr. doktor handed back our mole ration lab. After, we marked the unit 2 homework sheet numbers 11-15.
Mr. Doktor taught a new topic about DENSITY & MOLES:
To find either the density, volume or even mass we can use this triangle.
To use this triangle, you cover the one that you are trying to find.

Let us say that you need to find the mass. You cover the mass and your left with density times volume. Just remember to cover the one that you are trying to find and the triangle will tell you how.

DENSITY OF GASES @STP
To find the density of gases @ STP you can use this formula:

DensitySTP = mass of 1 mole/volume of 1 mole = molar mass/ 22.4 L = gram per mole/liter per mole = gram/litre

DENSITY OF GASES @ STP

ATOMS & MOLECULES

1. for monoatomic elements
a molecule= an atom

2. diatomic elements:
a molecule an atom
Cl2 Cl

Molecules of compounds:
2 hydrogen atoms}
1 oxygen atom } 1 molecule

EXAMPLE:
write the formula for ammonium carbonate:
(NH4)2 CO3
2-N
8-H
1-C
3-O
= 14 atoms

moles<--->molecules
6.02x10^23 molecules
---------------------
1 mol

**HOW MANY MOLECULES ARE IN A 0.25MOL SAMPLE OF CO2?
0.25<----> molecules
0.25molx 6.02x10^23molecules= 1.51x 10^23 molecules
----------
1 mol

1.51x10^23 carbon atoms
3.02x10^23 oxygen atoms

**5.1772x10^24 molecules of water= ??? moles
5.1772x10^24 x1mol}
------------------}= 8.6mol
6.02x10^23}

OCTOBER 27

Today in Chemistry 11 we did a lab :) YAAAY
we did a lab to determine the ratio of moles of iron consumed to copper produced during the chemical reaction between iron and copper (II) chloride.

We used Nails and blue copper sulphate solution! this is how it went



Homework: Finish the lab (questions)

OCTOBER 22

For today's class Mr.doktor went around the class and checked our unit 2 homework questions. After he took up some questions from the homework sheet.

Gases and Moles
The volume occupied by a certain gas depends on the temperature and pressure.

Standard Temperature and Pressure
T=0 degrees Celsius
P=l01.3 kPa

SATP
T=25 degrees celsius
P=100kPa

The volume of any gas at STP is 22.4L for every mole
22.4 L/1mole OR 1mole/22.4L

Some examples for use of STP
Find the volume (l) occupied by 0.060 mol of c02 at STP
0.060mole x 22.4L/1mole = 1.3L

Find the mass of a 200.0 mL sample of NO2 at SATP
mL -> L -> mol -> g
200.0mL x 1L/1000mL x 1mol/22.4L x 46 g/ 1 mole = 9200/22400 =.41g

• Times the top and bottoms then divide

HOMEWORK: Mole ratio lab -> do up to summary of procedure

Questions from 5-10 on the unit 2 homework questions

OCTOBER 20

To start off class and continue our talk about moles Mr.doktor showed a clip from the Austin powers 3 movie.


Atomic Mass
-The mass of 1 mole of atoms of an element
The mass of 1 mole of "c" atom is 12.01g
The mass of 1 mole of "ca" atom is 40.1g

Molar mass
The mass if grams of 1 mole of molecules of an element or compound is the molar mass

Diatomic numbers
H2, O2, N2, Cl2, Br2, I2, F2

Polyatomic numbers
P4,S8

We did a chart to give us some practice how molar mass is calculated.
Element Symbol Formula Atomic mass Molar mass
Bromine Br Br2 79.9 159.8
Neon Ne Ne 20.2 20.2
Silicon Si Si 28.1 28.1
Hydrogen H H2 1 2
Iron Fe Fe 55.8 55.8

To find the molar mass it is basically like this
Equation - >Ca(NO3)2
1.find out how many elements their are in the equation
Ca=1
N=2
O=6
2.Look on the periodic table and get the atomic mass of the element and times it by how many elements their are in the equation then add
Ca=1x40.1 -> 40.1
N=2x14 - > 28
O=6x16 - > 96
40.1 + 28 + 96 = 164.1g/mol

We also learnt how to convert grams to moles and moles to grams.
Mr.Doktor also mentioned that for how air balloons it is filled with helium instead of hydrogen.If hydrogen is used instead the balloon will burst when coming close to fire.

HOMEWORK:

Unit 2 homework questions # 1-4

OCTOBER 16

Today we started chapter 3, The Mole.

One mole = 602 000 000 000 000 000 000 000
or also written as: 6.02 x 10^23 <-Avogadro's Number

2H2 + O2 --> 2H20

How big is Avogadro's number?

• distance from earth to the moon: 3.84 x 10^5 km
• distance from earth to pluto: 6 x 10^9 km

How gases combine:

1. John Dalton

• looked at the masses or gases:
  -11.1g of H2 reacts with 88.9g of O2
  -46.7g of N2 reacts with 53.3g of O2
  -42.9g of C reacts with 57.1g of O2
• *No pattern*

2. Joseph Gay-Lussac

• 1L of hydrogen reacts with 1L of Cl2 -> 2L of HCl
  1L of N2 reacts with 3L of H2 0> 2L NH3
  2L of CO reacts with 1L of O2 -> 2L of CO2
• reactions occur in simple rations

3. Avogadro's Hypothesis:

• Equal volumes of any gas at the same temperature and pressure contain equal numbers of molecules

H2 + O2 -> H2O

*Our homework for today is: numbers 1, 2, 3, and 5 on page 322*

OCTOBER 14, 2009

Today's chemistry block went by quickly, we had our unit 2 test that covered nomenclature, the different types of seperation of mixtures & a chart on matter. This took up all our time in class, but we were assigned homework.- read page 313-319.

October 9,2009

Today in class we did a lab where we had to turn a solid into a liquid. During the process, we had to evaporate the water and record the weight of it before and after.When it was a solid, the substance was red and had a crystal look. While turning it to a liquid, the colour changed from red to blue then light blue.

OCTOBER 7, 2009

In the beginning of our chemistry class Mr. Doktor taught us how to use a bunsen burner!

In this picture you'll see the different parts of the bunsen burner (click to enlarge)

In this picture it shows the different types of flames the bunsen burner can create.

Of course we also did notes about ACIDS AND BASES

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
example HI(aq)--> Hydro Iodic Acid
-classical rules use the suffix IC and/or the prefix HYDRO-
example sulfuric acid
hydrochloric acid
~IUPAC system uses the aqueous hydrogen compound
example 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


We took a break from all the note taking and Mr. Doktor did another experiment :):)
He mixed sulfuric acid with sugar. Here is a video... and you'll see what happens!

Just incase you wanted to see how the bunsen burner worked :)

ps: instead of the striker, we used the stick and match. Just because we're cool.


HOMEWORK: read pages 252-253
do questions 1-9, 28, 34

October 5, 2oo9

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
  *KNOW ALL TEN*

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

IUPAC NAME/FORMULA

• 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

*Notes were taken down from Mr. Doktors powerpoint, we were given 2 pages of homework, 7-3 Practice Problems & 7-3 Apply worksheets. He's going to be collecting them next class!*

October 1, 2oo9

A. SEPARATING MIXTURES

1. there are many methods to separate mixtures, depending on the type of mixture

• by hand-----
  --> Heterogeneous Mixture


• filtration-----


• distillation:


• crystallization


• chromatography
  
  *All are physical changes

B. 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

C. 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

D. 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

E. 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

F. 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)

G. 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

H. 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

I. NAMING BINARY IONIC COMPOUNDS

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

J. STEPS

• write formula


• criss-cross charges


• reduce ion numbers to lowest common multiples

September 29, 2oo9

A. CLASSIFICATION OF MATTER

• understanding matter begins with how we name it. 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)
  

C. TELLING THE DIFFERENCE

• it is often very difficult to know if something is an element or a compound
  -only 'visible' on the atomic level
• one method is to connect the substance to an electric current. This technique, called electrolysis, can split the compound apart into its constituent elements

September 24, 2oo9

A. WHAT MATTERS?

• anything that has mass and occupies space


• matter can exist in many different states, the most common are:
  -solid, liquid, gas
  -plasma, aqueous, amorphous


• SOLIDS: holds one shape and has a definite volume


• LIQUID: can change shape, but has a definite volume


• GAS: can change shape and volume


• AQUEOUS: something dissolved in water


• Solid - strong bonds


• Liquid -weak bonds


• Gas - no bonds


• Plasma - ionization

B. CHANGES IN MATTER:

• matter can undergo any changes


• nearly all changes can be broken down into 3 categories:
  -physical changes, chemical changes, and nuclear changes

PHYSICL CHANGES:

• involves changing shape or state of matter
  -crushing, tearing, ect.


• no newsustances are formed
  -eg. boiling water, cutting wood, smashing cars

PHASE CHANGES

• changing from a solid to a gas can often be confused as a chemical change
  -chemicals remain the same


• during the melting process chemicals usually follow this path:
  

CHEMICAL CHANGE

• new substances are formed
• properties of the matter change
  -conductivity, acidity, color, etc.
  -eg. iron rusting, burning wood, digesting food

C. CONSERVATION OF MATTER

• in physical and chemical changes, matter is neither createrd nor destroyed...EVER. Period.
• this is called the 'conservation of matter'

September 22, 2oo9

Today in class we had our Chapter 1 Unit Exam!


1. In part one of the exam we had to match the WHMIS symbols with their appropriate labels.
And here they are:






Compressed Gas








Flammable Material

Oxidizing Material

Material Causing Immediate and Serious Toxic Effects

Corrosive Material

Biohazard Infection Material

2. In part two we listed 4 fundamental units of the SI system and their abbreviations.

These units were:

Meter - M

Mole - Mol

Second - S

3. In part three we did fill in the blanks:

When measuring liquid in a chemistry lab first use a graduated cylinder. Most liquids will not form a flat horizontal line but will curve. This curve is called a meniscus. To properly measure the volume of fluid measure to the bottom of the curve. After recording the volume of fluid transfer it to a beaker or test tube.

4. In part four of the exam we had to count how many significant digits were in the following numbers:

1.500m = 4 significant digits

2.0 x 10^-3g = 2 significant digits

0.00202 mL = 3 significant digits

7.01 x 10^4 J = 3 significant digits

Pi = never ending

For the rest of our chapter 1 unit exam we went over unit conversions.

SEPT 14, 2009

The Metric System

-->Started about 300 years ago in France
-->Le Systeme International d'Unites *S1 system
-->7 Fundamental units (used for basic measurements)

*meter(m)-length
*kilogram(kg)-mass
*second(s)-time
*ampere(A)-current
*mole(mol)-amount
*kelvin(k)-temperature
*candela(cd)-Luminous Intensity

Length
-->standard unit is the metre
Mass
-->standard unit is the kilogram
Time
-->standard unit is the second
-->9, 192, 631, 770 vibrations of a cesium atom

Prefixes used with S1 Units
*we can put a prefix in front of the unit & change the power of it
*GEGA=Billion
*MEGA=Million
*kilo=thousand
*micro=millionth
*nano=small (billionth 10 to -9)
*pico=ten to the -12

Fundamental Unit 1m

deci(d) 10-1 10dm
centi(c) 10-2 100cm
milli(m)10-3 1000mm
micro(M) 10-6 1 000 000Mm
nano(n) 10-9 1 000 000 000nm
pico(p) 10-12 1 000 000 000 000pm
femto(fm) 10-15 1 000 000 000 000 000fm

254,000,000,000=2.54x10(11)
254 Gm
0.00000411 =4.11x10-6
4.11um

*DON'T use scientific notaion & prefixes together!!
Becomes very confusing

Experimental Accuracy
--> The max accuracy of any measurement is 1/2 of the smallest division of the measuring device
--> A ruler with measurements of millimeters has a max accuracy of +or- 0.5mm

ex. 20.6cm +or- o.05cm (1/2)
mm^ <-------------^

Meniscus --> curved top of liquid in a graduated cylinder
Significant zeros: *zero is the place keeper it is generally not significant.
*if a decimal is placed to the right of a zero it IS significant
*no zeros= all digits are important

Adding or Subtracting: Round to the LEAST precise number

1) 7.4212 - 3.54= 3.8812= 3.88
2)12.31 + 10.771= 23.071=23.1

Multiply or Divide: Round to the number with the FEWEST S.D.s

1)2.5 x 5.55 = 13.875-14
2)1.33421 x 19.531= 26.05845= 26.058