Mole Conversion

Recall: 1 mol = 6.022 E 23 units of substance

Conversion between particles and mole

1. Particles to mole

• multiplied particle/ formula units/ molecules by 1 mol/6.022 x 10^23
• Example:   How many moles of barium nitrate (BaNO3) contain 6.80 x 10^24 formula units?

• Well, to do this, divide 6.80 x 10^24 by Avagadro's number, 6.022 x 10^23, to get 11.29 moles. 

     2.  Mole to particle

• multiplied mole by 6.022 x 10^23/ 1 mol
• Example:  Determine the number of atoms that are in 0.58 mol of Se

• If you're thinking you multiply .58 moles by 6.022 x 10^23 atoms/mol, you're absolutely right. This yields the answer of 3.49 x 10^23 atoms of Se.

Conversion between moles and grams 

To do this type of questions, you have to know the molar mass of the substance.  Using the periodic table we can find the molar mass, or the mass of a mole of a substance.  The unit of the molar mass is grams/ mol

1. Moles to grams

•  multiply moles by _grams / 1mol
• Example:  how many grams are there in 2.04 moles of Carbon?

•  To do this question, first you have to find the molar mass of C and it is 12.0g/ mol. Then, you multiply 2.04 mole by 12.0g/ 1 mol and get the answer of 24.5g of C

     2. Grams to moles

• multiply grams by 1 mol/ _grams
• Example:  how many moles are in 3.45 g of Carbon?

• Well, the molar mass of Carbon is 12.0 g/ mol so you multiply 3.45 g by 12.0g/ mole and get the answer of 0.288 mol.

Harder mole conversion

the following question is more difficult and a mole map can help you understand it.

1. Particles to mass

• To do this type of question:
• Find the moles of the item you wish to convert
• Determine the molar mass of the item you wish to convert. (set up the conversion factor - this time moles are on the bottom)
• Set up the equation and solve  
• Example: what is the mass of 2.78 x 10^ 22 Fe atoms?

• To do this, we multiply 2.78 x 10^ 22 by 1 mol/6.022 x 10^23 and times the molar mass: 55.8 g/ 1 mole to the answer and get 2.58 g Fe

     2. Mass to Particles

• The rules for this type of questions are
• Find the mass of the item you wish to convert
• Determine the molar mass of the item you wish to convert.
• Set up the equation and solve.
• Example:  how manu atoms of Iron in 20.0 g of Iron?

• For this question, we time 20.0 g by the molar mass of the Iron : 1 mol/ 55.8 g, then we multiply the result 6.022 x 10^23/ 1 mol and get the result of the number of atoms, which is 2.16 x 10^ 23 atoms of Fe

Finally, this is a cool story about mole conversion. Watch it!

Atomic Mass, Formula Mass, and Molar Mass

Yay it was snowing last night :). Who doesnt love snow??? (besides people who drive.) Anyways, this blog is for science... So, the last 2 classes we learnt about different masses and how to find them. Here are the notes to remind you of what we know so far.

Avogadro's Hypothesis
- The volumes of different gases at the same temperature and pressure have the same amount of particles.
- If they have the same number of particles, mass ratio is caused by the mass of the particles
- This principle is used for the relative mass of all the atoms on the periodic table

Avogadro's Number (not phone number :p)
-The number of particles in 1 mole of any amount of substance is 6.022x10²³ particles/mole
- The mole allows chemists to count atoms and molecules a lot easier.
-Atoms can also be counted by weighing



Relative Mass
-It is the mass of the different elements that make up a particular formula.

Atomic Mass
- The mass of an atom of a chemical element expressed in atomic mass units. (Its found on the periodic table to make it easier.)

* amu stands for atomic mass units


Formula Mass
- All atoms of a formula of an ionic compound (in amu)
ex: Potassium has 39.1 atomic mass and Fluorine has 19.0. That would mean KF (potassium fluoride) would equal to 58.1 amu. Just need to add the atomic mass.

Molecular Mass
- All atoms of a formular in a covalent compound. (in amu)
ex: O2 would be 16.0u + 16.0u = 32.0u



Molar Mass
- The mass of 1 mole (6.022x10²³ particles) in a substance and is the same numerical value of atomic mass, formula mass, or molecular mass but expressed in GRAMS PER MOLE.
ex: The molecular mass of oxygen is 16.0u and the molar mass of oxygen is 16.0g/mole.



That's all for today :)
HAVE FUN IN THE SNOW!!

Joke of the day!

A small piece of sodium that lived in a test tube fell in love with a Bunsen burner. "Oh Bunsen, my flame," the sodium pined. "I melt whenever I see you," The Bunsen burner replied, "It's just a phase you're going through."

GRAPHING!! (:

So.... for 2 classes we did graphs by using the microsoft excel.

Graphs that were taken includes: Density of Water, Density of Hot water, Temperature Vs volume of gas and lastly Mass Vs Volume.

How to graph?

As an example, here's how to graph the Density of Water.

STEPS:

• open excel
• plug in all the numbers for volume(mL) and mass (g) { make sure its separated into 2 columns}
• highlight both column then click "chart"
• choose "scatter"
• lable: title, subtitle, trend line type ( highlight + click "show formula" to get the formula) ,  x and y value

A linear line looks like this:

• click "add decimal place ( or something)" if the # keeps rounding itself to 1 ( this refer to the last 2 graph that we did)
• NOW, you get to design your graph. BE creative. Do ANYTHING you want. Lol

 

Question: Which one has a higher density? Cold water or Hot water?

Answer: Cold water, because as you heat water its molecules gain kinetic energy and the water becomes much less dense

VIDEO!

.. on Water Density.

Lab 2E

Lab we did on Wednesday was about Density.  The purpose of this lab was to calculate the thickness of a sheet of aluminum foil and express the answer in terms of proper scientific notation and significant figures

The procedure of this lab:

1. Get 3 rectangular pieces of aluminum oil and label 1, 2, and 3
2. Measure the length and width (using scientific notation)
3. find the mass with centigram balance
4. compare and discuss the result 

To get the thickness of the aluminum foil, we needed two formula

1. density = mass/ volume
2. volume = width * length* height

 So, how do we know whether our result is accurate?......Well, we can calculate the Experimental Error

Experimental Error = |actual-experimental| /actual x 100%
the smaller experimental error percentage we get, the more accurate the result is

Density

What is DENSITY ?


DENSITY is a physical property of matter, as each element and compound has a unique density associated with it. Density defined in a qualitative manner as the measure of the relative "heaviness" of objects with a constant volume.Density may also refer to how closely "packed" or "crowded" the material appears to be.

To calculate:

Density =	mass =	g/mL
	volume

The units of density are typically g/cm³, but they can also have other designations that may be more convenient. The densities of various materials range from large values for heavy metals to very small values for gases.

In most cases, the density of an object can be used to predict whether it is a solid, liquid or a gas.

• Solids generally have a density greater than water (1 g/cm³). For example, aluminum has a density of 2.7 g/cm³. Yet, oak (0.75 g/cm³) is a solid and floats in water. Some exceptions are most varieties of wood, many plastics, and pumice.

• Liquids generally have a density near 1 g/cm³, with some being slightly above and some being slightly below. An exception is mercury (Hg), or quicksilver, which has a density of 13.6 g/cm³.

 

Matter with smaller density will float on top of the matter with bigger density.

The densities for some common substances are:

Substance	Density (gm/cu.cm)
Air	0.0013
Wood (oak)	0.85
Water	1.00
Ice	0.93
Aluminum	2.7
Lead	11.3
Gold	19.3
Ethanol	0.94
Methanol	0.79

Accuracy, Precision, and Uncertainty

It is true that we cannot measure things exactly (unless we're robots...) however that doesn't mean we can't give it our best. Here are a few notes we took last class, explaining the uncertainties of measurements.

- Precision: How reproducible a measurement is compared to other similar measurements.

- Accuracy: How close the measurement (or average measurement) comes to the accepted or real value.

Measurement and Uncertainty

- No measurement is exact. Only best estimates which has a some degree of uncertainty.

-Only when we count do we get an exact number (ex: There are 28 students in this class. There cannot be   

 28.5... that would be just weird.)

Absolute Uncertainty

-Uncertainty is expressed in the units of measurement not ratio

Method 1: Make 3 measurements, calculate the average (Make sure to take out the measurements that are "different" from the others. Ex: 2.3, 2.32, 2.46, 2.33). The absolute uncertainty = largest difference between the average and lowest/highest reasonable measurement.

Method 2: Determine the uncertainty of each instrument. Always measure to the best precision that you can. Therefore you should estimate to a fraction 0.1 of the smallest segment on the instrument scale. (On your ruler the smallest division = 1mm. Your best precision should be to break this into 10 equal pieces.)

Relative Uncertainty and Significant Figures

-Relative uncertainty = absolute uncertainty divided by estimated measurement

-Relative uncertainty can be expressed

• In percent %
• or using significant figures (sig. figs.)

- The number of sig. figs. = relative uncertainty: The last digit in a measurement is uncertain as it could be one digit higher or lower very easily.