(TG/NR) Double Lab, Joint Blog Post!

FIRST LAB: Iron Nails and Copper Reaction

In this lab we tried to find the ratio of moles of copper formed to moles of iron reacted in the chemical reaction we did today for the lab.

Necessary 411:

When an iron nail is placed in a solution of Cupric Chloride the Iron will slowly dissolve. The dissolved iron will combine with chloride ions to form ferric chloride. In this activity you will determine the mass of iron consumed by comparing the mass of the nail before and after the reaction. Determine the number of moles of Iron this corresponds to. From this you will predict the number of moles of copper (ii) chloride that should be formed. By measuring the change in mass of the chloride solution in the beaker you can determine the mass of ferric chloride formed.

Procedure: 

1.  Using a 10 mL graduated cylinder, measure 70 mL of Copper (ii) chloride
2.  Transfer the solution from the graduated cylinder to the 100ml beaker
3.  Record the mass of the beaker and solution
4.   Record the mass of the nail
5. Place the nail in the beaker
6. Wait between 2 and 5 minutes
7. Remove the nail and record

When doing this lab make sure you are extremely with your measurements as well as procedures. Be careful when pouring or dumping and especially when cleaning.

Wear safety goggles and lab coats, as well as tie back and long hair and roll up any sleeves that are impeding your scientific process.

SECOND LAB: Molar Volume Lab

This experiment was designed as a way to test whether the molar volume of a gas at "around" STP is actually 22.4L

• In this experiment you need a lighter, a sink, a tap, a graduated cylinder and a scale
• First measure the mass of the lighter: record it
• Then fill your sink up with water and place your graduated cylinder upside down in it
• Then put your lighter underneath the cylinder and press the gas button
• As you press the gas button you will see little bubbles of coming from the lighter, trap these inside the cylinder
• Stop as your approach the 40mL mark
• Then dry off your lighter completely
• When your lighter is dry weight it again, and note the difference

This is the physical portion of the lab, for the rest of the lab you do theoretical speculations with the data you received and recorded whilst doing the experiment

•  Next take the mass you got in the difference of your lighters and change that into moles( g -----mol/g)  
• Now take your 40mL, convert it to L (.04L) 
• Write a new equation for molar volume
• Volume over Mol
• 0.04L/Step1 = a number around 22.4

Careful with all your steps, and remember that because you're probable not operating at STP, your numbers WILL be a little bit off.

 

  

(DA) Dec. 10, 2010: Density & Moles

Density

• Density is a measure of mass per volume
                     d = m
                            v
• Measured in g/L or g/mL (not only units used to measure, but the most common)

*refer to previous posts for unit conversions.

(NR) December 8th, CONVERSIONS multi

need help to convert?!

MASS to VOLUME

__g x mol/g x 22.4L/mol = ____L


VOLUME to MASS

__L x mol/22.4L x g/mol = ____g



MASS to MOLECULES

__g x mol/g x 6.02x10^23molecules/mol = ____molecules


MOLECULES to MASS

__molecules x mol/6.02x10^23molecules x g/mol = ____g


VOLUME to MOLECULES

__L x mol/22.4L x 6.o2x10^23molecules/mol = ____molecules



MOLECULES to VOLUME

__molecules x mol/6.02x10^23molecules x 22.4L/mol = ____L



I HOPE THIS HELPS!!! :)

(TG) Moles, Molechules, and Atoms; Dec6

MASS <--->MOLES<---->VOLUME

^

MOLECULES

v

ATOMS

Today in class we learned about 2 new conversions, going from moles to molechules, and then as well as an extra step to atoms.

The conversion factor between moles and molechules is avogadro's number as shown below.

What the pictures doesnt show though is that for atoms we have to multiply by the subscripts. (an extra step)

Some Practice: 

O(little2) <--- 2 atoms of Oxygen per molechule

Al <---- 1 atom of aluminum per molecule (the extra step isn't necessary)

So now for some actual conversions...

EXAMPLES:

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

0.58 mol  . 6.02 x 10^23 = 3.49 x 10^23 atoms

                           1 mol 

2. Determine the number of atoms that are in 1.25 mol of O(little 2) (oxygen gas)

1.25 mol . 6.02 x 10^23  (x2) = 1.5 x 10^23 

                                      1 mol

for this example we multiplied the number by 2 at the end, because there are 2 atoms in every oxygen gas molechules

**remember to include significant figures, and to use the second "EE" function on your calculator (or brackets)

 

                               

 

 

(DA) Nov. 22, 2010: Volume Conversion

• At a specific pressure and temperature one mole of any gas occupies the same volume
• At 0 Celsius and 101.3 kPa = 22.4 Litres.
• This temperature and pressure is called STP [(S)tandard (T)emperature and (P)ressure]
• 22.4 L mol is the molar volume of STP
  
  Example:

1. A container with a volume of 893 L contains how many moles of air is STP?
   893L x 1 mol = 39.9 mol
               22.4L

            Breakdown:
What you want to do is put write down 893L:

• 893L

After, figure out what the question asks for. If it asks for a mol, you need to cancel out the 
information (units) that you already have by multiplying your information with 1 mol divided by STP.

• 893L x 1 mol =
              22.4L
  After this, divide 893 by 22.4

• 893L x 1 mol = 39.8660714
             22.4L

This is not your final answer. You need to watch out for your significant digits.

• 893L x 1 mol = 40.0
              22.4L

fun video! (not a class summary) just extra**

(NR) Nov. 18, 2010: Molar Mass (Mass of Atom)

- the mass (in grams) of 1 mole of a substance is called the molar mass.
- it can be determined from the atomic mass on the periodic table.
- measured in g/mol

  
Molar Mass of Compounds 

*REMEMBER SIGNIFICANT DIGITS

Element                           work                             Molar Mass

H2O                                       2(1.0) + 16.0                                        18.0 g/mol
NO2                                       14 + 2(16.0)                                         46.0 g/mol
NaCl                                       23.0 + 35.5                                          58.5 g/mol
FeO                                         55.8 + 16.0                                         71.8 g/mol
NaNO3                              23.0 + 14 + 3(16.0)                                  85.0 g/mol



Mole Conversions (converting between grams and moles)

- To convert between moles and mass we use molar mass as the conversion factor.
- be sure to cancel the appropriate units.  

(TG) Nov 16: AVOGDRO'S NUMBER; how we count atoms

Avogadro's Number

6.02 x 10^23

• atoms and molechules are extremely small
• macroscopic objects contain too many atoms to count or weigh individually
• Amedeo Avogadro proposed that the number of atoms in 12.000000 grams of Carbon be equal to a constant (one mole of a Carbon)
•  So what is Avogadro's number? Well its 6, 020, 000, 000, 000, 000, 000, 000, 000
• 1 mole = 6.02 x 10^23 atoms
• one mole is simply a multiply of things for example:

1.  pair = 2
2. dozen = 12
3. century = 100
4. mole = 6.02 x 10^23 

The Mole in Perspective

6.02 x 10^23

• So how big is a mole? (in perspective)
• 1 mole of meters would cross the entire galaxy over 3000 times
• 1 mole of smarties would cover 250 planets similar to the size of earth a kilometer deep!
• 1 mole of seconds is 100,000 times greater than the age of the universe
• 1 mole of blood cells more than every human on the face of the earth

A mole is also used to measure the smallest unit of a quantity. For example there is a such thing as a mole of NaCl ions. You do not thing of Na and Cl as separately, but rather as ONE UNIT. 

EXAMPLES:

A sample of  carbon contains 2.4 x 10^25 atoms. How many moles is this?

2.4 x 10^25     x       1 mole

                                    6.02 x 10^23               = 39.9 moles

THINGS TO REMEMBER:

significant figures

the units you want to cancel are always opposite each other (ie. top and bottom)

(NR) Nov. 5, 2010: Hydrate Lab

Yesterday, we did a hydrate lab. It was our first class lab and it was fun! like OMG!!!


Hydrates are ionic compounds that contain an inorganic salt compound loosely bound to water. The purpose of this experiment is to determine the empirical formula of a hydrate. In the lab we determined the anhydrous (without water) mass of the hydrate. We compared it with the actual mass of the water that should be presented.

the materials we used were:
- Bunsen burner***
- test tube
- test tube rack
-test tube clamp
-weight scale
- Cobaltous chloride hexahydrate

***REMEMBER: BE AWARE OF BUNSEN BURNERS! You can't see the hot light blue flame... and if you accidentally touch it... PEACE TO YOU!

note for Mr. Doktor:
We should do more outdoor experiments with chemicals... We students want to see something big explode! lol

(DA) Nov. 3, 2010: Naming Compounds

Chemical Nomenclature

• Today the most common system IUPAC for most chemicals
  - Ions
  - Binary Ionic
  - Polyatomic ions
  - Molecular Compounds
  - Hydrates
  - Acids / Bases

Chemical Formulas

Be aware of the differences between ion and compound formulas

 - Zn^2+ <------------------------- Ion Charge

 - BaCl2 <------------------------ Number of Ions

Multivalent Ions

• Some elements can form more than one ion
  eg. Iron -> Fe^3+ or Fe^2+
        Copper -> Cu^2+ or Cu^1+
• IUPAC uses Roman Numerals in parenthesis to show the charge
• Classical systems use latin names of elements ans suffixes
  - ic (larger charge) and -ous (smaller charge)
  
  Example:
  - Ferric Oxide -------------> Iron (Fe)
   [ -ic refers to larger charge
     -ous refers to smaller charge]

• Ferr - Iron
• Cupp - Copper
• Mercur - Mercury
• Stann - Tin
• Aunn - Gold
• Plumb - Lead

Complex Ions

• Complex ions are larger groups of atoms that stay together during a chemical reactions
• Almost all are anions
• Write the metal name and the polyatomic ion

Hydrates

• Some compounds can form latices that bound to water molecules
  - Copper Sulphate
  - Sodium Sulfate
• These crystals contain water inside them which can be released by heating.

To name hydrates

1. Write the name of the chemical formula
2. Add a prefix indicating the number of water molecules (mono=1, di=2, tri=3 etc.)
3. Add hydrate after the prefix

ie. CuSO₄·5H₂O          Copper (II) Sulphate Penta Hydrate

      LiClO₄·3H₂O          Lithium Perchlorate Tri Hydrate

Naming Acids And Bases

• Hydrogen Compounds are acids
  - HCl ---> Hydrochloric Acid
  - H₂SO₄ --> Sulfuric Acid
• Hydrogen appears first in the formula unless it is part of a polyatomic group
  - CH₃COOH --> Acetic Acid

(TG) Nov: ELECTRONIC STRUCTURE

ELECTRONIC STRUCTURE

ELECTRON DOT DIAGRAMS: 

• the nucleus is represented by the atomic symbol
• for individual elements determine the number of valence electrons
• electrons are represented by dots around the symbol
• four orbitals (one on each side of the nucleus) ea holding a maximum of 2 electrons
• Each orbital gets one electron before they begin to pair up

examples:

CARBON

LEWIS DIAGRAMS FOR COMPOUNDS AND IONS

• In covalent compounds electrons are shared

1. Determine the number of valence electrons for each atom
2. Place atoms do the valence electron are shared to fill each orbital

example:

NF(little3)

DOUBLE AND TRIPLE BONDS:

• Sometimes the only way covalent compounds can fit all their valence levels is if they share more than one electron  

example: 

CO(little2)

IONIC COMPOUNDS

• in ionic compounds electron transfer from one element to another
• determine the number of valence electrons on the cation. Move these to the anion
• Draw [ ] around the metal and the nonmetal
• Write the charges outside the brackets

example:

(NR) Oct. 28, 2010: Trends on the Periodic Table

- Elements close to each other on the periodic table display similar characteristics.
- There are SEVEN important periodic trends:
1) Reactivity
2) Ion Charge
3) Melting point
4) Atomic Radius
5) Ionization Energy
6) Electronegativity
7) 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 ion charges depend on their group (column).

MELTING POINT:
- elements in the center of the table of the highest melting point.
- noble gases have the lowest melting points.
- starting from the left to right, melting point increases (until the middle)
> carbonis an exception!

ATOMIC RADIUS:
- radius decreases to the up and 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.
- all noble gases have high ionization energy.
- helium has the highest ionization energy.
- francium has the lowest ionization energy.
- opposite trend from atomic radius.

ELECTRONEGATIVITY:
- refers to how much atoms want to gain elections.
- same trend as ionization energy.

(DA) Oct. 26, 2010: Isotopes & Atoms

Atomic Number

• Atomic Number: Number of protons

Atomic Number = 22

Symbol = Ti

Atomic Mass = 47.87

Atomic mass - Atomic Number = # neutrons

   (p+n)         -           (p)            =        (n)

Isotopes - Same atomic number but different mass

FOR EXAMPLE:

Isotope     Mass#     Atomic#    # of Protons    #of Neutrons

 54Fe          54             26                26                   28

 56Mn         56             25                 25                   31

237Np        237            93                 93                  144

  14C          14              6                   6                     8

(TG) Oct. 21, 2010: Quantum Mechanics

CHEMISTRY NOTES

QUANTUM MECHANICS

·         Different levels of electrons: s, p, d, and f orbitals

·         As you increase levels, the orbitals get larger

BOHRS THEORY:

·         The electron is a particle that must be in orbital in the atom

QUANTUM THEORY:

·         The electron is like a cloud of negative energy or a wave

·         Orbitals are areas in 3D spaces where the electrons most probably are

·         The energy of the electron is in its vibrational modes – like notes on a guitar string

·         Photons are produced when high energy modes change to lower energy modes

S ORBITALS:

·         Each orbital holds 2 electrons

(hydrogen and helium)

                         P ORBITALS

·         There are 3 suborbitals

·         Each contains 2 electrons

·         Total of 6 electrons (3 x 2 = 6)

D ORBITALS

·         There are 5 suborbitals

·         Each contains 2 electrons

·         Total electrons: (2 x 5 = 10)

F ORBITALS

·         There are 7 suborbitals

·         Each contains 2 electrons

·         Total electrons: (7 x 2 = 14)

·         Can be made into nuclear bombs

Examples:

1.       How many and what type of electrons does the atom calcium have?

1s2 2s2 2p6 3s2 3p6 4s2

(NR) Oct. 19, 2010: Bohr Diagram

 

these are examples of the energy level model and the Bohr Model... by Nikko Rey.

Bohr Model

- Atoms are electrically neutral.

- Two different models can be used to describe electron configuration.

        > Energy level Model

        > Bohr Model

- Electrons occupy shells which are divided into orbitals.

        > 2e in the first orbital

        > 8e in the second orbital (OCTET)

        > 8e in the third orbital (OCTET)

 

(DA) Oct. 15, 2010: Neils Bohr

Bohr (1920s) 

• Rutherford's model was inherently unstable
  - Protons and electrons should attract eachother
• Matter emits light when it is heated (black body radiation)
• Light travels as photons
  
  
  PHOTONS
• The energy photons carry depends on their wavelengths

Bohr based his model on the energy (light emitted by different atoms. Each atom has a spectra of light. To explain this emission spectra, Bohr suggested that electrons occupy shells or orbitals

 

If given enough energy, an atom can move down or up levels of the orbitals.

BOHRS THEORY:

• Electrons exist in ortibals
• 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 of light.

(TG) October 13, 2010: Atomic Theory

THE ATOMIC THEORY

ARISTOTLE:

• One of the first theories ever created was called the "Four Elements Theory"
• Consisted of: Water, Earth, Wind, and Fire
• Lasted about 2000 years, but is not currently a valid theory because it cannot be proven 

DEMOCRITUS: 

• In 300 BC Democritus said atoms were indivisible atoms
• It was the first mention of atoms: atomos
• Was only a conceptual model 
• No mention of a nucleus
• Does not explain Nuclear Reactions

LAVOISIER (LATE 1700'S)

• Created the law of the conservation of mass: "in a chemical reaction no matter is destroyed or created"
• Also had a law of definite proportions: (e.g. in a H20 molecule there is always 11% H and 89% O)

PROUST

• In large proportions atoms have the same ratios as they do in molecules or compounds
• Proved Lavoisiers Laws correct 

DALTON

• Believed atoms are solid, indestructible spheres (e.g. billiard balls)
• This theory provides an explanation for different elements
• Based on the law of conservation of mass (if atoms are not destroyed, mass doesn't change)
• Having a molecule explains the law of of Constant Composition
• 2H2     + O2             =      2H2O 

J.J. THOMPSON

• Raisin Bun model
• solid, positive spheres, with negative particles embedded in them
• first atomic theory to have positives and negatives
• demonstrated the existence of electrons (cathode tube)  

 RUTHERFORD's assistant (1905) (

• showed that atoms have a positive, dense, center, with electrons outside of it
• resulted in planetary model
• explains why electrons spin around nucleus
• suggests that electrons are made up of mostly open spaces

RAISIN BUN MODEL

RUTHERFORDS MODEL (ELECTRON SPACING)

COMPARING THE MODELS
 

(DA) Oct. 4: Sodium Chloride Lab

Today we started off the class with a simple homework check.
Soon after we were given instructions on how to do our lab by Mr. Doktor (the best teacher in the world, as well as the smartest), and how to use each of the materials.


We split into groups of 3 and 4, and set off to gather our materials and safety clothing and equipment.
As we were in our groups we put our weight paper on the electronic scale and waited for Mr. Doktor to
put the Sodium Chloride (salt) onto the scale.


The Procedure we needed to follow:
1. Gather all the materials and put them on the lab bench
2. Measure 30 mL of distilled water using the graduated cylinder. Transfer this water into a 50 mL beaker
3. Weight 50g of Sodium Chloride.
4. Add Sodium Chloride to the water until it stops dissolving (the solution is saturated) and the first salt crystals begin appearing on the bottom of the beaker.
5. Measure the mass of salt remaining. Record the difference in salt as the amount added to 30 mL of water
6. Repeat steps 2-5 for 50 mL of water, 80 mL of water and 100mL of water.
7. Record all your data in the table below.
8. Create a graph of Mass of \salt vs. Volume of Water (Be sure you include a title, axis, data points, scale and a straight line of best fit). 




Electronic Scale:

Graduated Cylinder:

Beaker:

(NR) Sept 30: Density and Graphing

Density

- The density of an object is its mass divided by its volume. It is usually expressed in kg/L, kg/m^3, or g/cm^3.

example:

1) Determine the density of the hot tub that has the mass of 115kg and a volume of 70L.

d = 115 kg 
        70L

  = ____kg/L




(answer: 1.6 kg/L)

Graphing

• All graphs MUST contain 5 important things:
  1) Labelled axis
  2) Appropriate scale
  3) Title
  4) Data points
  5) Line of best fit line
• Three things can be done when working with graphs:
  1) Read the graph.
  2) Find the slope (RISE/RUN)
  3) Find the area under the graph.
  
  
  
  

m1) rise = 4                           m2) rise = 2                          m3) rise = 4                              m4) rise = 2      
       run     1                                  run    1                                 run     2                                     run     1
             
      = 4 miles/hour                 = 2 miles/hour                       = 2 miles/hour                           = 2 miles/hour

m5) rise  = 0
       run     1
    
     = 0 mile/hour

AREA:

a = 1/2bh = 1/2(1)(4)                                   d= 1/2bh = 1/2(1)(2)
               = 2                                                               = 1
                                                                     = lw       = (10)(1)
b= 1/2bh = 1/2(1)(2)                                                  = 10
              = 1                                                 = 11
  = lw     = (4)(1)
             = 4                                                 e= lw     = (12)(1)
 =  5                                                                         = 12

c= 1/2bh = 1/2(2)(4)                                 
              = 4                                               TOTAL AREA: a + b + c + d + e
  = lw     = (2)(6)                                                                 = 2 + 5 + 16 + 11 + 12
             = 12                                                                      = 46 miles x hour
= 16

(TG) Sept 28: Dimensional Analysis

Want to know what 50km/hour is in miles? Its the speed limit, so it would be pretty useful to know if you were ever in a sticky situation...

• This type of conversion is called Dimensional Analysis
• It is very similar to the process of converting currencies

4 Steps to Know:

1.    Find a unit equality
2.    Find the conversion factors
3.    Apply Conversion Factors 
4.    Cancel Units

EXAMPLES:

50 km/hour to miles/hour

50km/h x 1mile/1.6km

UNITS CANCEL: km's is crossed out in the equation

You now multiply the two fractions and get:

50/1.6 per hour

50/1.6 = 31.25

Final Answer: 50km/hour is = to 31.25mi/hour


A Helpful Picture to Illustrate the Process:

(NR) Sept 23: Scientific Notation & Significant Digits

TO REMEMBER!
* Accuracy and precision is VERY important in science.
* Calculators are NOT smart enough to decide what is precise and what isn't.
* Scientists have established rules for rounding off extra digits; you MUST follow them!


SIGNIFICANT DIGITS

• Non-zero digits are always significant.
• If the zero is a place keeper it is not significant.
  example: 0.0098 (2 s.d.)
                0.00043 (2 s.d)
• Any numbers to the left of a decimal point are significant.
  example: 2.00, 27.06, 38.9000
• Zeros after another number are significant.

examples:

HOW MANY SIGNIFICANT DIGITS ARE IN...

14.78? ............_____s.d.
2.04? .............._____s.d.
9.00? .............._____s.d.
0.0038?..........._____s.d.

(answers: 4 s.d., 3 s.d., 3 s.d., 2 s.d.)

• when you multiply/divide ROUND to the number wit hthe fewest S.D.'s 
  example:
  
  
  • 1.38797 x 38.234 = 53.067644 => 53.068

• when adding/subtracting ROUND to the LEAST precise number
  example:
  
  
  • 10.978 - 5.68 = 5.298 => 5.30
  • 20.3 + 18.348 = 38.648 => 38.6
  
  http://www.youtube.com/watch?v=qrpq4qU6Wgg&feature=related 

SCIENTIFIC NOTATION

• used if we need to write the number 2000 with only 2 s.d.'s
  > 2.0 x 10 ^ 3
• used if we want to write the number thirty three billion four hundred million without taking up an entire line.
• Shows really big or really small numbers easily
• makes use of power of 10 

examples:
10^5 = 100000
10^-5 = 0.00001

-convert each question into scientific notation.

13400000 => ____ x 10^__
0.00000542 => ____ x 10^__
23900000 => ____ x 10^__

(answers: 1.34 x 10^7, 5.42 x 10^-6, 2.39 x 10^7)

REMEMBER:
DO NOT USE "^" ON YOUR CALCULATOR!!!  

(DA) Sept 21: SI System & Percent Error

Prefixes Used with SI Units

• We can put a prefix in front of the unit and change the power of it

         ^{- tera (T) 10^12
         - giga (G) 10^9
         - mega (M) 10^6
         - kilo (K) 10^3
         - hecto (h) 10^2
         - deca (da) 10^1
SI Prefixes}

• The SI System uses many prefixes to represent very large or very small numbers
  - deci (d) 10^-1
  - centi (c) 10^-2
  - milli (m) 10^-3
  - micro (µ) 10^-6
  - nano (n) 10^-9
  - pico (p) 10^-12
  - fempto (fm) 10^-15

^{***Don't use scientific notation and prefixes together! IT BECOMES VERY CONFUSING.
Experimental Accuracy}

• In general, the maximum accuracy of any measurement is half of the smallest division of the measuring device
• A ruler with measurements of millimeters has a meximum accuracy of +/- 0.5mm
• The odometer in a car has a maximum accuracy of +/- 50m.

EXAMPLE:

• A graduated cylinder has units of 1.0 mL. The accuracy of the cylinder is +/-0.5mL
• Liquid in a graduated cylinder will typically form a curved top called a meniscus
• The volume is taken at the bottom of the cylinder

The curved top is called a Meniscus

Expressing Error

• Error is a fundemental part of science
• There are usually 3 reasons of error
  1) Physical errors in the measuring device
  2) 'Sloppy' measuring
  3) Changing ambient conditions
  

Calculating Error

• Two different possibilities
  
  1) Absolute Error
           - Measured value minus accepted value
           - Absolute Error = Measured - Accepted
  
  2) Percentage Error
           - Most common
           - Percent Error = Absolute Error / Accepted Value
      
        |  Measured - Accepted |
  % = |           Accepted           | x100