MultiStep Conversions

MASS    <-------->      MOLES  <-------->    VOLUME(@STP)
               molar mass          |         mass volume
                                          |
                                          |
                                MOLECULES   ( Avogardo's number)
                                          |
                                          |
                                    ATOMS     ( subscripts)

 Remember : when you are converting, you must always go back to moles first!
                          ** with the exception of molecules to atoms or atoms to molecules

EXAMPLES:

- 11.5g of H2 gas are placed in a balloon at STP. Determine the volume of the balloon.

(there is 2 ways to do this)

1.)   step 1 : 11.5g x 1mol    = 5.75 mol
                    2.0g 
   
        step 2 : 5.75 mol x 22.4 L    = 128.8 = 129 L    (*dont forget your significant digits)
                                      1 mol
                                     
                                       OR
2.)   11.5g x 1 mol      x   22.4 L   = 129 L
                    22.4 L         1 mol



REMEMBER TO STUDY FOR YOUR MIDTERM !!!

Nicole H

The Day Mr.Doktor Was Away

We had a sub for this class, and Mr.Doktor left us some with a worksheet and some homework. Here is a few examples and an overview of what was assigned:

Moles to Atoms/Molecoles 
Molecules/Atoms to Moles


EX. How many atoms are there in 1.5mol or Iron?

1.5 mol x 6.02E32 atoms = 9.03E23 atoms
                        1mol

How many moles of magnesium bromide (MgBr2) contain 5.28E24 formula units?

5.38E24 FU x 6.02E23 FU = 8.94 mols
                          1 mol


Determine the number of atoms that are in 0.58 mol of Se.

0.58 mol x 6.02E23 = 3.5E23 atoms
                     1 mol

Janine R

Molar Volume Lab

In the lab we did last class, it helped us work towards what we're learning in chemistry right now which is converting L/mol.

During this lab, we filled a sink full of water up to the point where we could fully submerge a gradulated cylinder. We made sure that there were no air bubbles within the cylinder with water in it. Before putting the lighter in the water, we weighed it. While keeping the open end in the water and the bottom sticking out at an angle, we took the lighter and put it under the water, under the opening of the gradulated cylinder and held the lighter button so that only butane would come out. Until the cylinder was filled up to 100mL of butane, we took out the lighter and shook it, and then put it in the warm drying area so that we could weigh the lighter afterwards. There was a bit of a difference in the weight, but there shouldn't be much of a difference in the weight.

The purpose of this lab was to experimentally determine the molar volume of a gas.

Sofia Nguyen

Moles to Volume Conversion

- At specific pressure and temperature one mole of any gas occupies the same volume.
- At 0 degrees C and 101.3 kPA
1 mole = 22.4 L
- This temperature is called STP

-- 22.4 L/mol is the molar volume as STP

Examples:

- How many litres will 3.5 Cl(g) occupy at STP?

3.5 mol x  22.4 L = 78.4 L
                1 mol
( because there is only 2 significant digits the answer becomes 78 L )

- A certain gas is found to occupy 12.3 L at STP. How many moles of gas is there?

12.3 L x  1 mol    =   0.549 mol
               22.4 L

- At STP an unknown gas is found to occupy 150mL. How many moles of gas must be there?
    1L = 1000mL
150 mL x   1 mol     = 6.6964 mol
                 22.4 mL  
6.6964 mL             
  1000                          = 0.0670 mol

*** remember you must convert your mL to L. Doesn't matter if you do it first or last!
      
Nicole H

Molar Conversion

Converting Between Moles and Mass
- to convert between moles and mass we use molar mass as the conversion factors
- BE SURE TO CANCEL THE APPROPRIATE UNITS

Examples:

How many grams are there in 2.5 mole of P4  (  F2 : 19.0 x 2 = 38)

Step 1: 2.5 mol                         g   =         g
                                             mol
Step 2: 2.5mol          38.0          g  =    95 g  
                                  1         mol

Remember!!!!!
            -significant digits
            -0 is NEVER a significant digit
            - balance chemical equations first

A compound is made of phosphorous and chlorine. It is found to contain 0.200 mol and has a mass of 27.5
- determine molar mass of the compound
- suggest possible formula

27.5 g         = 137.5 g/mol
0.200 mol 

PxCly:       X              Y       MOLAR MASS
                  1              2                  102
                  2              1                 97.5
                  2              2                 133
                  1              3                 137.5                  
Phosphorous and chlorine = PCl3

Nicole H

Molar Mass

molar mass - the mass (g) of 1 mole of a substance
-the molar mass is the atomic mass
-molar mass is measured in g/mol

Molar Mass of Compounds
-to determine the molar mass of a compound, you add the mass of all the atoms together
EX.
C12H22O11 --> 12.0(12)+1.0(22)+16(11) = 342.30g/mol
KCl --> 19+17 = 36g/mol
NaBr4 --> 23.0 + 79.9(4) = 342.6
(NH4)2S --> (( 14+1 (4) )) x2 + 32.1 = 54.1g/mol

Janine Roldan

Chapter 2 Test

A review of chapter 2:

Atomic theories

JJ Thomson - raisin bun theory
                    - he beleived there was a solid, postive sphere with negative
                       particles embedded



Bohr and Lewis Diagrams



Quantum Mechanics


S orbital - each orbital holds 2 electrons
P orbital - 3 suborbitals: contain 2 electrons each
D orbital - 5 suborbitals: contain 2 electrons each
F orbital - 7 suborbitals: contain 2 electrons each

Naming Compunds



IUPAC= using roman numerals in pararanthesis
Classical System= use latin names of the element
ex. Aunn: gold
      Ferr: iron
- when using the classical system the larger charge you use ic and the smaller charge ous

Nicole H

Hydrate Lab (November 4)

The purpose of this experiment was to determine the empirical formula of hydrate. In this lab we determined the anhydrous(w/out water) mass of the hydrate. We also compared this with the actual mass of water that should have been present.


Janine Roldan

Molecular Compounds

Molecular Compounds

• molecules are diatomic 
• 2 of the same elements: H2, N2, O2, F2, Cl2, Br2, I2
• 2 molecules are polyatomic: S8, P4

Naming Molecular Compounds

• use the name of the first element
• second element ends in -ide
• 1st atom usually doesn't have a prefix EX. CO2 -> carbon dioxide
• hydrogen doesnt have a prefix EX. H2S -> hydrogen sulfide

More examples:

HBr -> hydrogen bromide

KI -> potassium iodide

H20 -> dihydrogen oxide

Naming Acids and Bases

• hydrogen compounds are acids

EX. 

HCl -> hydrochloric acid

H2SO4 -> sulfuric acid

Naming Bases

• cation and OH

EX.

Na(OH) -> sodium hydroxide

Ba(OH)2 -> barium hydroxide

Sofia Nguyen

Naming Compounds

Chemical Nomenclature

Todays most common system is IUPAC for most chemical reactions like:
- ions
- binary ionic
- polyatomic ions
- hydrates
-molecular compounds
- acids/ bases

Chemical Formulas
- be aware of the difference between and ion and a compound.
ion= charge
compound= no charge



Note: -the top charge on a multivalent is the most commonly used
           -IUPAC use roman numerals in parenthesis
           - classical systems use latin names of the elements

suffixes ic is the larger charge and ous is for the smaller charge
ex. Fe0 - ferrous oxide

Classical names are a bit different for some elements like:
Ferr : iron
Cupp: copper
Stann: tin
Plumb: lead

Hydrates
some compounds can form lattices that bond to water molecules by hydrates

-watch this video to learn how to make/ name hydrates
http://www.youtube.com/watch?v=HM2C5FEvR0g

Nicole Haughian

How to Draw Electron Dot Diagrams

Electron Structure: Electron Dot Diagrams

Drawing Electron Dot Diagrams


follow this link --> http://www.youtube.com/watch?v=y6QZRBIO0-o (:


Lewis Diagrams for Compounds and Ions

·                     in covalent compunds electrons are shared

·                     one atom gives up an electron to the other, forming positive and negative ions

1.                   First determine the number of valence electrons for each atom in the molecule

2.                  Then, place atoms so that valence electrons are shared to fill each orbital

EX.

carbon tetrafluoride

Double and Triple Bonds

EX.

oxygen: it's outer shell contains 6 electrons so it needs another 2 electrons in this case shared to get a stable configuration

Ionic Compounds

·                     the formation of an ionic bond is the result of the transfer of one or more electrons from a metal to a non-metal

·                      draw [brackets] around the metal and non-metal

·                     for metals: Energy + Metal Atom ---> Metal (+) ion + e-

·                     for non-metals: Non-metal Atom + e- --- Non-metal (-) ion + energy

EX.

JANINE ROLDAN

Trends on the Periodic Table.

 There are some things on the periodic table that you wouldn't know about just by looking at it, but by actually studying and researching about it.

Did you know that elements close to each other on the periodic table display similar characteristics? 

TRENDS

There are 7 important periodic trends that you'd have to know;

1. Reactivity

2. Ion Charge

3. Melting Point

4. Atomic Radius

5. Ionization Energy

6. Electronegativity

7. *Density 

You actually don't need to know much about density.

REACTIVITY

• Metals and non-metals show different trends
• The Most reactive metal is Francium; the most reactive non-metal is fluorine.

ION CHARGE

• Elements that that contain ion charges depend on their group (column)

MELTING POINT

• Elements in the center of the table have the highest melting point
• Noble gases have the lowest melting points
• starting from the left and moving right, melting points increase until the middle of the table

ATOMIC RADIUS

• Radius decreases up the right 
• Helium has the smallest atomic radius 
• francium has the largest atomic radius

IONIZATION ENERGY

• Ionization energy is the energy needed to completely remove an electron from an atom
• it increases going up and to the right of the periodic table
• all noble gases have high ionization energy
• Helium has the highest ionization energy and Francium has the lowest ionization energy

ELECTRONEGATIVETY

• Electronegativity refers to how much atoms want to gain electrons
• some trend such as ionization energy

IF YOU WANT TO LEARN MORE PLEASE REFER TO THE FOLLOWING LINKS! :)

http://chemistry.about.com/od/periodictableelements/a/periodictrends.htm

http://dl.clackamas.cc.or.us/ch104-06/periodic.htm

A SONG THAT COULD HELP YOU REMEMBER THEM

http://www.youtube.com/watch?v=dHcD_But7dU

http://www.youtube.com/watch?v=dHcD_But7dU

SOFIA NGUYEN

Isotopes and Atoms

^atomic number: number of protons

-atomic mass - atomic number = number of neutrons
 (p+n)           -    (p)                =  (n)

Example:
Isotopes    Mass   Atomic #    #of protons   # of neutrons
C                 14        6                     6                    8
K                 39       19                  19                   20
Na                23       11                 11                    12

Mass spectrometers: is used to determine the abundance and mass of the isotopes of elements

mass spectrom for boron:

Nicole H

Quantum Mechanics

-The difference between the bohr theory and quantum theory
Bohr Theory: the electron is a particle theat must be in orbital in the atom
Quantum Theory: the electron is a cloud of negative charge or a wave function



S orbital - each orbital holds 2 electrons
P orbital - 3 suborbitals: contain 2 electrons each
D orbital - 5 suborbitals: contain 2 electrons each
F orbital - 7 suborbitals: contain 2 electrons each


Nicole H

How to Draw Bohr Diagrams

Follow this link to learn how to draw Bohr Diagrams:
http://www.slideshare.net/profpaul/bohr-diagrams

Remember:

Electrons occupy shells which are divided into orbitals, the first shell holds 2 electrons, the second one holds 8, and the third one 8, and the fourth one 16.





Janine R

October 12 class (Bohr Notes)

Bohr (1920's)

-Rutherfords first tried to make a model before Bohr, his model was very unstable

-Bohr based his model on the energy (light) emitted by different elements

-Bohr believed each atom has a specific spectra of light 

Bohr's Theory

-electrons exist in orbitals

-when they absorb energy they move to a higher orbital

-as they fall from a higher orbital to a lower one they release energy as a photon light

Sofia N

THE TEST

This class we did the test :(

A Review of what was on it :
1. Lab safety
     -Do not run in the lab
     -Wear safety equipment
     -Tie your hair back

2. Balancing Equations
     -Refer to link below on a review of how to balance equations
                 http://www.youtube.com/watch?v=8KH3laR2iR4

3. Dimensional Analysis
     -How many miles are equal to 120km?
                1.6km =1 mi
            120km  x  1mi       =75 mi
                           1.6km
4. Scientific Notation

Nicole H

(September 22 Class) Units

In class we learned all about the different SI units, which include;

SI BASE UNITS 


Base Quantity                          Base Unit                                     Symbol
Length-----------------------------meter------------------------------m
Mass------------------------------kilogram---------------------------kg
Time------------------------------second-----------------------------s
Temperature----------------------kelvin------------------------------K
Amount of a substance------------mole------------------------------mol
Electric Current-------------------ampere----------------------------A
Luminous intensity----------------candela-----------------------------cd

COMMON METRIC PREFIXES 

Prefix	Abbreviation	Meaning
Mega-	M	106 or 1000000
Kilo-	k	103 or 1000
Deci-	d	10-1 or 0.1
Centi-	c	10-2 or 0.01
Milli-	m	10-3 or 0.001
Micro-	µ	10-6 or 0.000001
Nano-	n	10-9 or 0.000000001
Pico-	p	10-12 or 0.000000000001

IF YOU HAVE ANY OTHER QUESTIONS ABOUT UNITS PLEASE REFER TO THE LINK BELOW THANKS!
http://en.wikipedia.org/wiki/SI_derived_unit


SOFIA NGUYEN

Atomic Theories

Democritus 

-called these particles atoms
-in 300 BC said atoms were indivisible particles






















John Dalton 

-each element is composed of extremely small particles called atoms
-identical, but differ from other elements
-neither created or destroyed in any chemical reactions
-always same relative numbers












Lavoisier

-law of conversation of mass














Proust

-if a compound is broken down into its constituents, the producers exist in the same ratio as in the compound
















JJ Thompson

-raisin bun model 
-said, positive, spheres with negative particles embedded in them
-1st atomic theory to have positive (protons) and negative (electrons) changes












Rutherford

-showed that atoms have a positive, dense centre with electrons outside it














Janine R

Dimensional Analysis

Steps to Converting:
1. Identify what units you want to end up with
2. Find the conversion factor
3. Place units in appropriate places
4. Cancel units

Examples:
HOW MANY MILES ARE EQUAL TO 120KM?
-1.6 Km = 1 Mi
120 km x 1 mi     = 75 mi
              1.6 km

HOW MANY SECONDS ARE IN 1.4 HRS?
-1.4h x 60 min x 60s     = 5040s
              1hr       1min

Nicole H

Scientific Notation

Scientific notation is a way of writing numbers that accommodates values too large or small to be conveniently written in standard decimal notation. Scientific notation has a number of useful properties and is commonly used in calculators, and by scientists, mathematicians, doctors, and engineers.
In scientific notation all numbers are written like this:

When are Digits Significant?

Non-zero digits are always significant. Thus, 22 has two significant digits, and 22.3 has three significant digits.

With zeroes, the situation is more complicated:

1. Zeroes placed before other digits are not significant; 0.046 has two significant digits.
2. Zeroes placed between other digits are always significant; 4009 kg has four significant digits.
3. Zeroes placed after other digits but behind a decimal point are significant; 7.90 has three significant digits.
4. Zeroes at the end of a number are significant only if they are behind a decimal point as in (c). Otherwise, it is impossible to tell if they are significant. For example, in the number 8200, it is not clear if the zeroes are significant or not. The number of significant digits in 8200 is at least two, but could be three or four. To avoid uncertainty, use scientific notation to place significant zeroes behind a decimal point:

8.200 ´ 10³ has four significant digits
8.20 ´ 10³ has three significant digits
8.2 ´ 10³ has two significant digits


Janine R

Classification of Matter

Divide Matter in 2 types:

1. Homogeneous: consist of only one visible component.
     ex. distilled water, oxygen, graphite
2. Hetergeneous: contains more than one visible component.
    ex. chocloate chip cookies, granite

                                                                            MATTER
                                       HOMOGENEOUS                                    HETEROGENEOUS
                   PURE SUBSTANCE       HOMOGENEOUS MIX               MECHANICAL MIXTURE
                        Element                                    Solution
                      Compound

Pure Substances
    Element: Substance cannot be broken into simpler substances by chemical reaction
                                         EX. oxygen, Iron, Magnesium
Compound: Substances made up of 2 or more elements and can be changed into elements by   chemical reactions
                                         EX. water, sugar

Mixtures
- many mixtures are easy to identify (chocolate chip cookies) BUT others are easily confused as pure substances.
Heterogeneous mixtures- different parts are clearly visible (granite, sand)
Homogeneous mixtures- different parts are not visible (salt, water, air)

Seperating Mixtures
- there are many methods to seperate mixtures, depending on the type of mixtures
           -by hand
           -filtration
           -distillation                                           Physical Changes
           -crystallization
           -chromatography

Nicole H

Chemical and Physical Changes

Chemical Changes

• a chemical change produces a new substance 
• the change cannot be undone
• the old matter is no longer present
• the original matter cannot be recovered

Rust forming on a nail

Physical Changes

• a physical change does not produce a new substance
• changes in state or phase such as melting, freezing, condensation and sublimation are physical changes
• you can return the substance back into its original state and vice versa

Ice melting can be turned back to ice

How to tell the difference ?
A chemical change makes a substance that wasn't there before. There may be clues that a chemical reaction took place, such as light, heat, colour change, gas production, odor, or sound. The starting and ending materials of a physical change are the same, even though they may look different.


Janine R.

Chemical and Word Equations

Today in Chemistry we learned how to form a chemical equation, going from word form -----> chemical form


For Example: 
      
        A solution of barium phosphate is mixed with aqueous sodium sulphate to yield solid barium sulphate and aqueous sodium phosphate.

*Notice the bolded words such as aqueous and solid. Those are the phases that those certain chemicals are in.

       Ba3(PO4)2 (aq) + Na2 SO4(aq) --->  Ba SO4 (s) + Na3 PO4 (aq)
*This is what the word equation should look like turning into a chemical equation 


       Ba3(PO4)2 (aq) +3Na2 SO4(aq) --->3Ba SO4 (s) +2Na3 PO4 (aq)
*This is what you would call a balanced chemical equation. It's so that on both sides of the equation (the arrow), there are a balanced number chemicals on each side. Now on both sides there are:

3 Bariums
3 Sulphates
2 Phosphates
6 Sodidums

*IF YOU STILL DON'T HOW I GOT THAT ANSWER PLEASE REFER TO THE FOLLOWING LINKS! THANK YOU FOR READING!  :)


http://www.youtube.com/watch?v=8KH3laR2iR4
http://www.standnes.no/chemix/examples/chemical-word-equations.htm

Sofia 

Lab Safety

Here are 10 rules for keeping safe while performing in the lab:

1. Wait for instructions before starting your lab
2. Long hair must be tied back at all times
3. Always wear safety equipment
4. Keep lab clean
5. No running in the lab
6. No fooling around in the lab
7. If a substance spills, alert Mr. Doktor immediately
8. Waft chemicals towards you, do not smell directly
9. Don't consume the substances
10. No food or drinks in the lab

Lab Safety Rap:
http://www.youtube.com/watch?v=yclOrqEv7kw

Nicole H

Welcome to our Chem 11 blog ... enjoy :)