Balancing Equations

Today, we learned about chemical equations!

- Formulas must be correct or nothing will be right

there are different types of reactions:


combustion: when oxygen combines with another compound to form carbon dioxide and water
C₃H₄ + O₃ ---> CO₂ + H₂O

synthesis: when two compounds combine to form a more complicated one
A + B --> AB

decomposition: when a complex molecule breaks down to make simpler ones
AB --> A + B

single replacement: When one element trades places with another in a compound.

A + BC --> AC + B (if A is a metal)
A + BC --> BA + C (if A is non-metal)

double replacement: When metals and non metals of an element switch places
AB + CD --> AD + CB

Acid-Base Neutralization: When an acid and base react to make a salt and water
HBr + NaOH ---> NaBr + H₂O

We balanced equations.... here's how!!
1. make sure all formulas are correct with correct subscripts
2. count how many of each element there is on each side
3. place coefficients beside compounds (and individual elements, if any) to make each element have the same number of each side
4. if oxygen is one of the elements present, do it last
5. check your answers!!!

here are a few examples

___ NaNO₃ + ___ PbO ---> ___ Pb(NO₃)₂ + ___ Na₂O
becomes
2 NaNO₃ + PbO ---> Pb(NO₃)₂ + Na₂O

AgI + Fe₂(CO₃)₃ --> FeI₃ + Ag₂CO­₃

becomes

6 AgI + Fe₂(CO₃)₃ --> 2 FeI₃ + 3 Ag₂CO­₃

C₂H₄O₂ + O₂ --> CO₂ + H₂O

becomes
C₂H₄O₂ + 2 O₂ --> 2 CO₂ + 2 H₂O

heres a worksheet!

Intro to Chemical Reactions

well firstly we went over our test from last class before we started our new unit today which is unit four- types of chemical reactions

we looked at some eqamples of when acids and bases neutralize each other


we talked about some supporting evidence that shows us when a chemical reaction has taken place:

• change in colour
• change in smell
• percipitate is formed

remember that changes in state do not always mean chemical reaction has occured, although it may seem like it ( think ice to water to water vapour)

we also saw a demonstration from mr. doktor about magnesium and how when it burns, it gives of a lot of radiant light and thats why its used in fireworks. this video clip uses a lot more magnesium than what we used in class but you get the idea :P

thats pretty much all we did today!

Copper (II) Chloride Lab

This class all we did was learn about dilution. For the next class we will be doing a cupric chloride lab. What we're will be doing is diluting a solution of copper II chloride. The lab is individual but you will be work in a group. The procedure is simple but we need to find the mass of copper II chloride we will be using. To find the mass here are a few examples we did in class.

1) Peter is asked to make a 250mL solution of K2SO4 with a concerntration of 0.55M. What steps will he follow? The conversions used here are: concentration - mole - mass
(0.55/1L)(0.25L)= make sure to convert mL into Litres!
(.1375mol)(174.3g/1mol)=
23.97grams

Dilution of Solutions:
When you add water concentration decreases. If volume is doubled, concentration is halved.
6.0L 2.0mol/L 12.0mol
12.0L 1.0mol/L 12.0mol

For the lab, the problem is how you can make a 0.300M Solution of Copper II Chloride?
In my solution I will be using 60mL of water. To know how much mass of Cooper II Chloride I will need for my solution I will use the conversions: concentration - mole - mass
(0.300mol/L)(.06L)=.018mol
(.018mol)(134.5g/mol)=2.421grams


for a dilutions worksheet, click here

Concentration

Today, we learned about the wonderful world of.....................concentration!!!!

Solution: A homogeneous mixture
Solute: The component present in the smaller amount
Solvent: The component present in the larger amount
Concentration: Ratio of solute to solvent and can be determined by the following equation:

Amount of Solute
Amount of Solvent

Possible units:
g/mL, g/L, mg/mL, micrograms/mL, m/L, % by mass, % by volume

con'c = concentration

The most useful units for con'c is mol/L which is also known as molarity (M)

molarity=

moles of solute
litres of solution

C = n/v
n = cv
v = n/c

n= number of moles
c= concentration
v= volume

***********THOSE 3 FORMULAS LISTED ABOVE CAN ONLY BE USED FOR LIQUID SOLUTIONS

if a compound formula is in bracket, [like this], it means it is the concentration of the substance.
e.g. [HCl] = 6.0 mol/L


The concentration of HCl (Hydrochloric acid) is 6.0 mol/L



EXAMPLES
What is the molarity of a solution in which 0.45 grams of sodium nitrate are dissolved in 265 mL of solution?

Sodium Nitrate = NaNO₃, molar mass = 84.99 g/mol

0.45 g x 1 mol/84.99 g = 0.005 mol
265 mL x 1L/1000mL = 0.265 L


molarity = 0.005 / 0.265 = 0.01998 mol

the molarity of NaNO_{3 in 265 mL of solution is 0.020}


What will the volume of a 0.50 M solution be if it contains 25 grams of calcium hydroxide?
<sub>calcium hydroxide = Ca(OH)2
molar mass = 74.093 g/mol

25 g x 1 mol/74.093 = 0.3329 mol

v = n/c
= 0.3329 / 0.50 M
= 0.6658 L
=0.67 L</sub>

How many grams of ammonia are present in 5.0 L of a 0.050 M solution?

Ammonia = NH₃
molar mass = 17.031 g/mol

n = cv
= (0.050)(5.0)
=0.25 moles

0.25 moles x 17.031 g/1mol = 4.26 g
= 4.3g

for a worksheet click here

Exam Time!!

So basically we reviewed the homework and got the mid-term review sheets.

here are a few sites if yaaa need extra practice/study material!!
-molar mass
-calculating percentage
-molar equations

Empirical Formulas

empirical formulas are pretty much just the simplest version of any chemical formula.

Molecular Empirical
P4H10 - - - - - P2H5
C10H22 - - - - - C5H11
C6H18O3 - - - - - C2H6O
N2O4 - - - - - - NO2

empirical formulas the simplest whole number ratios in a compound

molecular atoms show the actual atoms and bonds in a compound

finding empirical formulas:
a sample of an unknown compound in analyzed and found to contain 8.4g of carbon, 21g of hydrogen and 5.1g or oxygen. find the empirical formula

element - atomic mass- mass - moles - smallest divided - smallest ratio
carbon 12 8.4 .7 2.187 2
hydrogen 1 2.1 2.1 6.56 7
oxygen 16 5.1 .32 1 1

the empirical formula = C2H7O

comprende? thats pretty much all we did, it was a short class today

for a fun worksheet, click here!

Percentage Mass of Elements in Compounds

The only formulas you need are:
Molar Mass: G/Mol and Molar Volume: L/Mol
That is basically it and now here are some examples to practice.

EXAMPLE 1:

Calculate the percentage of each element in C3H8.
Step 1: Formula used should be molar mass: G/Mol.
Step 2: Find the molar mass of each element.
Carbon: 3(12.0) + Hydrogen: 8(1)= 44g/mol
To find the percentage divide one of the element's molar mass by the total molar mass of the compound.
Step 3: Carbon: 36g/mol/44g/mol= .818 x 100= 82.0%. Hydrogen: 8g/mol/44g/mol= .181 x 100=18%

EXAMPLE 2:
Find the percent mass of XeF2
Step 1: Use formula molar mass: G/Mol
Step 2: Find molar mass of each element.
Xenon: 131.1 + Flourine: 2(19.0)= 38 = 169.1
Step 3: Divide individual molar mass by total molar mass.
Xenon: 131.1/169.1= .775 = 78% Flourine: 38/169.1=.2247 = 22%

Some More Equations

remember:

Example:

The density of water is 1.00 g/l. Find the volume of water occupied by 1.6 x 10²⁴ molecules of H₂O

we will have to convert molecules--->moles---->mass.

A 4.00 mL sample of an unknown substance is known to contain 0.228 mols. Its molar mass is 196.967 g/mol. What is it's density?

we better get a good mark on this, it took a long time.

Density and Moles

Density --> MASS per unit VOLUME

• D=M/V
• M=D/V
• V=M/D

Density of Gases at STP

Density of Solids and Liquids

& remember...

Molar Volume Lab

today we did the molar volume lab!

for homework we made a procedure for the experiment that mr doktor showed us last class and today we carried it out, after going over some questions we had for homework.

in this lab we start by filling the sink with water and measuring the mass of the lighter. once we do that, we place the 100mL graduated cylinder under the water and let all the air bubbles flow out so it is full of water (upside down). then we hold the cheap, NONSEALED (mr doktor didnt spend much on the lighters :P) lighters under the opening in the graduated cylinder and hold the button down so the butane flows out into it up to 10mL. we then measure the mass of the lighter after the butane is released. this is where most problems occured due to the fact that when we released the butane, the water flowed into the lighter, throwing the mass off. using the difference in mass we calculated the molar volume of butane. even though it should have come out to 22.4 L per mole, we came out with 2.9 L. Yeah, just a little bit off. but hey we got to light stuff and see bubbles rise so it was fun and we got to practice some important calculations.

thats pretty much all we did today, as the class was shortened and the lab took up the whole time so there was no lesson otherwise. till next time "show your work and prosper"

wow sorry that was a painfully cheesy ending

ATOMS AND MOLES

1) For monoatomic elements:
a molecule= an atom

2) Diatomic elements
A MOLECULE: AN ATOM:
Cl2. Cl

3) Molecules of Compounds
2 Hydrogen atoms and 1 oxygen atom create water which is one molecule.

FORMULA FOR ATOMS AMD MOLES
6.02x10^23molecules/ 1 mol OR 1 mol/ 6.02x10^23

Examples:
how many molecules in 0.25mol of CO2?
0.25mol x 6.02x10^23 molecules= 1.51x10^23

How many moles in 5.1772x10^24 molecules of H2O?
5.1772x10^24 molecules/ 6.02x10^23= 8.6mol

For more practice here is a printable practice worksheet.
www.emp.byui.edu/PECKK/chem105/molsheet.htm

Mole Ratio Lab!

Today we did a mole ratio lab, which included the following:

+ ---->+


2Fe + 3CuCl₂ ---> 3Cu + 2FeCl₃

We:
-wanted to discover the amount of moles of copper made in the reaction of iron and copper ii chloride
- wanted to discover the amount of moles of iron was used in the reaction
- wanted to figure out the ratio of moles of iron to moles of copper
-wanted to figure out the amount of atoms and formula units.

Please refer to lab report for materials and procedures, in-class.

To sum it up:
We mixed the CuCl₂ with water, then let the iron nails sit in it for a number of minutes. Before we put in the nails, the solution was bright blue, but with the addition of iron it produced a brownish color, and copper formed on the nails. We scraped the copper off the nails, decanted the copper, and found that 2.25g of copper was produced after 8.19g of CuCl₂ was added.

Here is a youtube video very similar to our experiment, except the iron was kept in the solution for a longer time.


(apologies for the spelling errors, this is the most decent video we could find, and it shows all the steps!)

Gases and Moles

- the volume occupied by a certain gas depends on the temperature and pressure

..these are 2 very important things to remember..

1. STANDARD TEMP. = (t=0 degrees Celsius) and PRESSURE = (P=101.3kPa)
2. the volume of any gas at STP is 22.4L/1 mol or 1 mol/22.4L

examples:

1) find the volume occupied by 0.060mol of CO2 at STP

0.060mol x 22.4L/1mol = 1.34L = 1.3L

2) find the mass of a 100.0mL sample of NO2 at STP

100.0mL x 1L/1000mL x 1mol/22.4L x 46.0g/1mol = 0.41g

This is a quick review about moles(refer to the bottom of the link's page to what we learned about gases!)

the MOLE con't

watched the most loveable austin powers today to set the mood :P

for all those people who werent here last class and are too lazy to scroll down one entry, ill go over the mole again.

1mole=6.02 x 10^23

if we had a mole of your everyday pea we would have enough to cover earth 200+ times! thats how freakin small an atom is.

atomic mass is equal to the mass of one mole of the element it is the mass of. that was confusing sentence so let me make it easier with an example

the atomic mass of carbon is 12.01 right? so one mole of carbon is 12.01g of carbon! comprende?

molar mass (or molecular mass) is the mass in grams of 1 mole of the molecules of an element or compound.

so lets say you gots your typical water molecule. thats H2O is it escaped your memory. that means we have 2 hydrogen atoms and 1 oxygen atom. so the the molecular mass of one hydrogen atom is 1 gram so 2 must be 2 grams. now all we do is add that to the mass of the oxygen atom (which is 16 grams) and voila! 18g/mol is the molar mass of an H2O molecule.

need me to spell it out how we did that? here i will

2Hydrogen= 2(1.0) + 1Oxygen= 16

thus... 2.0+16=18.0

ahhhhhh make more sense now?

we did some other examples in class and mr doktor started to get tricky with us by making us convert kg and mg and such, but thats all stuff we've learned before so yeah!

haha "that was easy"

here is today's daily double to see if YOU know what you've learned!!!"

What Is The Mole?!

Rufus the naked mole rat??? NO! We're learning about the OTHER mole....
THE MOLE
The mole is a quantity of a substance that has a mass in grams numberically equal to its formula mass (6.02 x 10 to the 23rd power). It is just another way to represent a number.
Example: triplets = 3, a dozen = 12, a mole = 6.02 x 10 to the 23rd power.
The number 6.02 x 10 to the 23rd power is also known as AVOGADRO'S NUMBER.

How big is Avogadro's Number?
Imagine a small green pea, it covers one cubic centimeter. 1 mole would cover the entire earth plus the pathway to jupiter!!

How Gases Combine
John Dalton look at the masses of gas
11.1 g of H(2) reacts with 88.9 g of O(2)
46.7 g of N(2) reacts with 53.3 g of O(2)
42.9 g of C reacts with 57.1 g of O(2)
In this observation there are no patterns
Joseph Gay-Lussac
1L of H(2) reacts with 1L of Cl(2) to create HCL
1L of N(2) reacts with 3L of H(2) to create NH(3)
2L of CO reacts with 1L of O(2) to create 2L of CO(2)
In this observation there is a pattern. Reactions occur in simple ratios

Avogadro's Hypothesis
Equal volumes of any gas at the same temperature and pressure contain equal numbers of molecules.

As for the fun parts of class we blew up more balloons. One filled with Hydrogen and the other with Hydrogen and oxygen. The balloon filled with both hydrogen and oxygen should have blown up louder than the balloon with hydrogen but a hole was burnt into the balloon. Mr. doktor brought his potato gun and the result was a big bang and the wet tissue being blasted at the wall on the floor.

This is what the potato gun looked like.

Acids and Bases

Today we started off with some fascinating experiments! One of note was when we combined sugar and sulfuric acid which produced carbon and water. here is a video summarizing what happened:

the equation is:
C12H22O11(s) → 12 C(s) + 11 H2O(aq)

We also learned about acids and bases.

Acids
- solid liquid or gas at SATP - Standard Ambient Temperature and Pressure
- form conducting aqueous solutions
- turn blue litmus red
- dissolve in water to produce H⁺
- Tastes sour
- PH of less than 7

Bases
- turn red litmus blue
- slippery
- nonconductive
-dissolve in water to produce OH^-
- PH of more than 7

here are some common acids and bases you can find around your home:


Naming Acids
- Acids are aquaeous (dissolved in water)
- Hydrogen compounds are acids
- HCl_(aq) ---> hydrochloric acid
- H₂SO_4(aq) --> Sulfuric Acid
- Hydrogen appears first in the formula unless it is part of a polyatomic group.
-CH₃COOH ----> acetic acid
(aka C₂H₄O₄ but usually appears in the first form as that is how it is structured: )

HI_(aq) Hydro Iodic Acid
notes: the "ic" replaces the "ide" ending in iodide, hydrogen is added and the 'gen' is dropped

- Classical rules use the suffix -ic and/or the prefix hydro-
- eg. sulfuric acid, hydrochloric acid

- IUPAC system uses the aquaeous hydrogen compound
- eg. HCl_(aq) ---> Aquaeous Hydrogen Chloride

Naming Bases
- For now, all bases will be aquaeous solutions of ionic hydroxides.
eg. NaOH - sodium hydroxide, BaOH₂- Barium Hydroxide
- Use the cation name followed by hydroxide (see above)

INTERESTING FACT: the "ous" in nitrous acid HNO_2(aq) means that this unit has "the smaller #) of oxygens in the polyatomic ion.

To practice naming acids and bases, Click here!!!

Hydrates and Molecular Compounds

Hydrates

• some compounds can actually form lattices that bond to water molecules. some examples of these are copper sulfate and sodium sulfate.(without water the compound is often preceded by -anhydrous)
• these crystals contain water inside them which can be released by heating.

Naming Hydrates

1. write the name if the chemical formula
2. add a prefix indicating the number of water molecules
3. add hydrate after the prefix

ex/ Cu(SO4) . 5H20(s) = copper(II) sulfate-pentahydrate

nickel(II) sulfate - hexahydrate = Ni(SO4) . 6H2O(s)

..(keep in mind that if an element ends in -ic it is referring to the larger charge and when the element ends in -ous it is referring to the smaller charge)

here is a simple website you could use to test and practice your understanding!

Molecular Compounds

• composed of 2 or more non-metals
• low melting and boiling points
• share electrons rather than exchange
• usually end in -gen(hydrogen, oxygen, nitrogen)
• 7 molecules are diatomic(H2, N2, 02, F2, Cl2, Br2, I2)
• 2 molecules are polyatomic(P4, S8)

Examples/ N204 - dinitrogen tetraoxide

SC2 - carbon disulfate

sulfur dibromide - SBr2

dihydrogen exide(water) - H2O

Here is a list of Molecular Compounds that have their own very special names..

IUPAC Formula

water H20

hydrogen peroxide H2O2

ammonia NH3

glucose C6H12O6

sucrose C12H22O11

methane CH4

ethane C2H6

propane C3H8

octane C8H18

methanol CH3OH

ethanol C2H5OH

here is another website where you can read more about molecular compounds and practice your understanding!

OR if you're not into the website thing.. here's a quick video!


Enjoy! :)

Mixtures, Atoms and the Periodic Table

Mixtures

Yesterday in class we reviewed the various ways of separating mixtures. A few of the most commonly used methods to separate homogeneous mixtures are distillation filtration and chromatography.

A) B)C)

- Figure A is showing distillation, where the substance's state is changed from a liquid to a gas and back to a liquid by boiling the liquid substance.

- Figure B is showing filtration, which can be used to separate solid and liquid parts of a mixture

- Figure C is showing the results of chromatography, where a solution can be separated by allowing it to flow along a stationary substance

Atoms

- matter is made up of atoms

- molecules are groups of atoms held together by electrical bonds

- ions are atoms or molecules that have an electric charge

- (+) cations(remember, cations are PUSSYtive)

- (-) anions

- atoms are made up of 3 sub-atomic particles

- protons have a positive charge, are inside the nucleus, each element has a different number of protons and the number of protons in an element is called its atomic number

- neutrons are neutral, they are inside the nucleus, have often the same mass as protons and adding or removing neutrons doesn't change the element

- electrons have a negative charge, are located outside of the nucleus, are 1800 times smaller than protons and chemical reactions occur between electrons in different atoms/compounds

Periodic Table

.. just a few quick facts/reminders about the PT

- Families(or groups) form vertical columns, all elements of a family have similar traits/characteristics

- Periods are horizontal rows. Elements gradually change from metals to non-metals as you move from left to right.

- most elements use Latin names as well(some exceptions are copper-cuprum, gold-aurum, iron-ferrum, lead-plumbum and silver-argentum)

Naming Nomenclatur

Naming chemical compounds has been a very difficult task and different systems have been used through the centuries. Today the most common system is IUPAC for most chemicals(ions, binary ions, polyatomic ions, molecular compounds and acids)

When you are naming ions remember there are metals, non-metals and polyatomic ions. For metals, use the name of the element and add the ion(AL+3 = aluminum ion). For non-metals, remove the original ending and add -ide(F=fluorine to fluoride). And remember that polyatomic ions have special names.

When you are creating binary ionic compounds, know that they contain a metal and a non-metal. Also know that metallic and non-metallic ions bond together, that the electron is transferred from the metal to the non-metal and the net charge must always be zero.

This is a handout you can use to practice naming and writing chemical formulas

Conservation of matter

we had a short class today so we didnt get in as much as usual but hey we still learned a lot and even got around to blowing stuff up!

we learned about all the ways matter can be divided up in respect to mixtures and substances.

understanding matter begins with how we name it. we can divide matter into two types:

• HOMOGENEOUS: consists of one visible component. i.e. distilled water, oxygen or graphite
• HETEROGENEOUS:contains more than one visible component. i.e. chocolate chip cookies, granite
  

(sorry for the home-made-ness, but i couldn't find it on the web and didnt know other wise :P)

this chart we were shown in class shows us how matter can be catergorized with regards to their homogeneous or heterogeneous state

PURE SUBSTANCES

we learned there are two types of pure substnaces:

• ELEMENTS: substances that cannot be broken down into simpler substances by chenical reactions. i.e. oxygen, iron, magnesium
  
• COMPOUNDS: substnaces that are made up of two or more elements and can be changed into elements or other compounds by chenical reactions i.e. water, sugar

telling the difference...

it is often very difficult to tell the difference between elements and compounds. the differences are only visible on the atomic level. but there are ways to find out without looking at the atomic level. one way is through ELECTROLYSIS. this method is to connect the substance to an electric current. this method will split the compound into the elements it is made of, but if the substance is a pure element, no change will occur.

well this is pretty much all we learned today... oh right and i forgot to show you what we blew up!

as seen in this video, we blew up a balloon filled with hydrogen today in class. it was a great way to end the class!

Properties of Matter

Today we began unit 2 of chemistry, and actually began learning about the subject.

we learned that matter:
- is anything that has mass and occupies space
- can exist in many different states, but the most common are
- solid, liquid and gas
- plasma (in stars, can be aqueous or amorphos)

solids: has definite shape and volume

liquids: can change shape but has definite volume

gases: can change shape and volume

aqueous: something dissolved in water
e.g. NaCl(s) ----> NaCl(aq)



We also learned that matter can undergo many changes, and nearly all of them can be classified as physical, chemical or nuclear

***a note about phase changes ***
- changing from solid to gas can often be confused as a chemical change
- the chemicals remain the same


in a physical change:
-no new substances are formed (like boiling water, cutting wood, and smashing cars)
- involves changing the shape or state of matter (crushing, tearing, etc)

in a chemical change:
- new substances are formed
- properties of matter change (conductivity, acidity, colour, etc)
eg. iron rusting, burning water, digesting

here is a video of physical changes and chemical changes:


conservation of matter:
- in physical and chemical changes matter is NEVER created or destroyed. ever.
- this is called conservation of matter

here's a diagram to help with the idea of conservation of matter:




Thats all for now!

Sodium Chloride

Today, we did our first lab of the year! We explored the following problem:

What is the maximum amount of Sodium Chloride (Table Salt) you can dissolve in 200mL of distilled water?

We measured salt on a scale, then added it bit by bit to different amounts of distilled water with a scapula. We got the following results:

Trial	Volume of Water (mL)	Mass of Salt (g)
1	15	0.65
2	25	0.84
3	40	1.44
4	50	1.17

We also learned that we must expect that errors are going to be made during an experiment. For example, one time when the salt was being transported from the scale to the beaker, some dropped onto the counter. This would have resulted in a different mass of the salt that dissolved in the water, since we subtracted the new amount from the original amount on the scale.


We then created a graph to determine how much salt can dissolve in 200mL of water by graphing the data points and drawing a line of best fit.

above: salt dissolving into water

above: a look at salt dissolving, scientifically

Graphing!

today we learned all the things we needed to remember during the year for when we make a graph. appearently we will be doing it a lot so THIS IS IMPORTANT!

5 BASIC THINGS WE NEED

1. title- short and descrptive works best here. titles such as "time verses distance" or "speed of car verses time it takes" would be accepted as titles
2. labeled axis- full names of the data being respresented... AND DONT FORGET THE UNITS!!!
3. scale- we should be using about 70 percent of the space given for the graph, dont bunch it up in the corner. and dont use those squiggily things to make your graph longer, cuz its messes up your line. if your data starts at something higher than zero, just start there!
4. correct data plot- make sure you are accurate and precise with your data plotting. even the smallest mistake can cause things to be off.
5. line of best fit- use either a straight line of a SMOOTH curve. DONT JUST CONNECT THE DOTS!

linear graphs use the slope=rise/run to find the slope.

dont use the points you plotted for your slope calculation. instead, use the line of best fit. occasionally, there may have been an error in some of the gathered data, so the line of best fit will even it out so you can make a more accurate calculation

remember your line of best fit doesn't HAVE to start at zero.

there are different types of linear graphs

• direct relationship- the graph starts at zero and has a constant slope
• y=mx+b- this graph has a y-intercept, meaning it does not start at zero, making it a different type of linear graph

thats pretty much all thats new from today! problems? questions? queries? HAHA :P

here is an activty you can do to practice the placement of your line of best fit here (scroll down to the box that says "line of best fit")

Measurement, Accuracy, and Significant Digits

Today we learnt about the metric system and that it has 7 fundamental units.

1. Meter (m)- Length
2. Kilogram (kg)- Mass
3. Seconds (s)- Time

4. Amphere (A)- Current

5. Mole (mol)- Amount

6. Kelvin (k)- Temperature

7. Candela (cd)- Luminosity Intensity

We also learnt about Prefixes used with the "Le systeme International d'Unites" = Si System.

Apple has used the word "Nano" to exaggerate the size of the ipod nano.

We also took notes on significant zero

1) If there is a decimal after the zero it is significant.

IE. 1020. 302980.

2) You must round to the least percise number

IE. 4212-3.54= 3.8812 S.D.s= 3.88

3) When you multiply and divide, round to the number with the fewest S.D.s

IE. 2.5x5.55= 13.875 S.D.s= 14

Near the end of class we watched two youtube videos. One of a bridge collapsing and the other a swing fail. http://www.youtube.com/watch?v=3q27dzRgHLo