Formula of a Hydrate

Today we did a laboratory experiment to determine the empirical formula of a hydrate.
The lab can be found on page 45 of the Essential Experiments for Chemistry lab book.

Glossary:
Hydrate: A compound containing water (H20) in it's crystal structure. The general formula is AB xH2O
Anhydrous: The form of a compound without water.
Carbohydrate: Organic compound with the general formula Cx(H2O)y

The sutructure of a hydrate

Objectives:
To determine the percentage of water in a hydrate, the moles of water present in each mole of the hydrate and to write the empirical formula of the hydrate.

Procedure:
Summarized for full procedure refer to lab book

• Put on safety equipment and set up bunsen burner.
• Heat a crucible for 3 min, then let it cool
• Determine the mass of the crucible
• Place the hydrate in the crucible, determine the mass, and then heat for 5 min once the bottom of the crucible turns red.
• Let it cool, then determine it's mass
• Repeat the previous two steps to verify your results.
• Add a few drops of water and record the changes that appear

Results:

1. Determine the mass of the water by subtracting the after heating mass from the before heating.

Then divide the mass of the water by the mass of the hydrate. Multiply your answer by 100 to get a percent.

2. Calculate the mass of the salt by subtracting the mass of the water from the hydrate. Then divide the mass of the salt by the molar mass of the hydrate which is 120.4g. The will give you the number of moles of the salt.

3. Calculate the # of moles of water present in the hydrate by dividing the mass of the water by the molar mass of water, which it 18.0g. 

4. Divide both molar amounts by the smaller molar amount. This should equal one for the salt and seven for water

5. Substitute these numbers into the general empirical formula AB xH2O.

Answer:

AB 7H2O

Here is a video summarizing the above lab
http://www.youtube.com/watch?v=wNqzkNW72rM&feature=related

Empirical Formula of Organic Compounds

salutations drowsy readers!

BE AWAKE AS ARSENIC BASED LIFE FORMS EXIST! wait.

Before we begin what the title implies, what must brush up on our limited knowledge of what we call organic compounds!

(because embedding is disabled, i must give a link)

if the point was vague,organic compounds are basically anything that contains carbon

AND TO BEGIN NOW OUR LESSON!

We can find the empirical formula of an organic compound by BURNING it. When we BURN the organic compound (reacting with oxygen), this leaves us with the BURNT(okay i'll stop) product. From the mass of the products, the moles of each element in the original UNBURNT(last one, i swear!) reactants can be calculated. 

we can do this because of the wonderful law of conservation of mass!
-the law states that the mass of the product is the same as the mass of the reactants-
it is also assumed that all carbon and all oxygen are used up during the combustion.

Sounds confusing? PERHAPS! Let us try an example!

A 10g sample is burned, producing 20g of CO2 and 8.0g of H2O. What is the empirical formula?

• convert the grams into moles
• Mol CO2 = 20g x 1mol CO2 / 44.0g CO2 = 0.455 mol CO2
• Mol H2O = 8g x 1mol H2O / 18.0g H2O = 0.444 mol H2O

• find how many moles of C and H, these are the elements that make up the organic compound
• Mol C = (0.455 mol CO2 x 1 mol C) / 1 mol CO2 = 0.455 mol C
• Mol H = (0.444 mol H2O x 2 mol H) / 1 mol H2O = 0.888 mol H

• find the empirical mass
• divide both moles by the smallest molar amount
• C = 0.455 / 0.455= 1
• H = 0.888 / 0.455 = 2

• SUM IT ALL UP!
• the empirical formula is CH2

• don't forget to check your answer!
• convert from moles to grams, it should add up to 10g
• IF your answer does not add up, remember that oxygen could be a component of the compound
• mass of O = mass of compound - mass of c + mass of H

and as this concludes the BURNT lesson, worry not readers with a BURNING passion to learn, as I leave you a feeling of toasty(wait, what?) security!

-JY

Empirical and Molecular Formula

The Empirical Formula gives the lowest term ration of atoms/ moles in the formula
**all ionic compounds are given in the empirical formula


Ex. C8H12-------> C2H3
     Molecular to empirical
General case of converting grams of substances to empirical formula:
As we have already learned to do convert from grams to moles for each measurement, by divide it by it's atomic mass.
Then dived each molar amount by the smallest molar amount to receive a ration. Round these ratios to whole numbers and the make them your subscripts of your equation.
An example is provided in the video below.

General Case of converting percentage ratios into empirical formula:
Assume you have 100g of the substance. Convert each percentage into grams. Divide the amounts by their atomic mass. It will give you the ratio's which you'll need to round off and then use as your subscripts.



Molecular Formula: Is a multiple of the empirical formula which shows the actual number of atoms that form  each specific compound

Molar Mass / Molar Mass of empirical formula = Molecular Formula

Examples are in the video below!

This video does a far better job of explaining than I ever could so please actually watch it! It's worth it trust me!

http://www.youtube.com/watch?v=gfBcM3uvWfs

Percent Composition

The percent composition of a compound is a relative measure of the mass of each different element present in the compound.

These are the rules on how to calculate the percent composition;

Calculate the molecular mass of the compound

Calculate the total mass of each element present in the formula of the compound

Calculate the percent compositon (percentage composition): % by weight (mass) of element
= (total mass of element present ÷ molecular mass) x 100

Examples.
1.Calculate the percent by weight of sodium (Na) and chlorine (Cl) in sodium chloride (NaCl)
Calculate the molecular mass (MM):
MM = 22.99 + 35.45 = 58.44


Calculate the total mass of Na present:
1 Na is present in the formula, mass = 22.99


Calculate the percent by weight of Na in NaCl:
%Na = (mass Na ÷ MM) x 100 = (22.99 ÷ 58.44) x 100 = 39.34%


Calculate the total mass of Cl present:
1 Cl is present in the formula, mass = 35.45


Calculate the percent by weight of Cl in NaCl:
%Cl = (mass Cl ÷ MM) x 100 = (35.45 ÷ 58.44) x 100 = 60.66%


Note: If the answers add up to 100 they are probably correct.

2.A compund has a total molar mass of 18g/mol and a % composition of oxygen and 11.1% hydrogen. How many oxygen atoms are there in the compound?
11.1/100 x 18g/mol = 2.0g/mol = 2 hydrogen

100-11.1=88.9/100 x 18g/mol = 16 = 1 oxygen

There is one oxygen atom in the compound

3.A school boy buys a bag of candy. Each piece of candy contains 5g of glucose (C6 H12 O6). A) find the % composition of glucose B) if the bag contains 150g of candy, does the percent composition of the glucose change? Explain.

a)MMg=180g/mol

C6 = 72/180 x 100 =40%

H12 = 12/180 x 100 = 7%

O6 =96/180 x 100 = 53%

b)No, because the % composition is based on each individual compound and will not affect the candy as a whole. The % composition will not change because if each piece of candy is the same than the composition will stay the same because the numbers will be larger but will have the same ratio.

4.A molecule has a mass of 142g. It contains 5 atoms of one element and 2 atoms of another. What is the other element if one of them is oxygen and oxgyen has five atoms? What is the percentage composition of both elements?

Oxygen: 16 X 5 = 80g

142g-80g=62g

62g/2=31g

**Check for a component with the atomic mass of 31g. 

Phosphorus=31g. 

Check by firguring out the percent composition. 

62/142=43.7%            43.7+56.3=100g.

80/142=56.3%

Answer= P2O5

5.Calculate the molecular mass (MM) of (NH4)3PO4:


MM = 3x[14.01 + (4 x 1.008)] + 30.97 + (4 x 16.00) = 3 x [14.01 + 4.032] + 30.97 + 64.00 = (3 x 18.042) + 30.97 + 64.00 = 54.126 + 30.97 + 64.00 = 149.096


Calculate total mass of N and then the percent:
3 N are present, mass = 3 x 14.01 = 42.03
%N = (mass N ÷ MM) x 100 = (42.03 ÷ 149.096) x 100 = 28.19%


Calculate total mass of H, and then the percent:
12 H are present in the formula, mass = 12 x 1.008 = 12.096
%H = (mass H ÷ MM) x 100 = (12.096 ÷ 149.096) x 100 = 8.11%


Calculate the total mass of P and the percent:
1 P is present in the formula, mass = 30.97
%P = (mass P ÷ MM) x 100 = (30.97 ÷ 149.096) x 100 = 20.77%


Calculate the total mass of O and the percent:
4 O are present in the formula, mass = 4 x 16.00 = 64.00
%O = (mass O ÷ MM) x 100 = (64.00 ÷ 149.096) x 100 = 42.93%


The answers above are probably correct if %N + %H + %P + %O =100, that is,
28.19 + 8.11 + 20.77 + 42.93 = 100

Check out this video!
http://www.youtube.com/watch?v=xbEeyT8nK84
Search up some problems to help you through this!

Quiz Day

WE HAD A QUIZ TODAY.

IT ALSO SNOWED

AT LEAST 5CM!

WE'RE DROWNING

-gargle gargle-

the quiz was on mole conversions, just in case you forgot/ missed it

i have no cookies for you

....

i have a mole game?
while i'm at it, this blog has been lacking puns..

-JY

Harder Mole Conversions

Begin with a mole map:

Conversion from particles to mass (g)
Ex. What is the mass of 1.27x10^22 Co atoms?
1.27x10^22 atoms x (1mol/ 6.022x10^23atoms) x (58.9g / 1 mol) = 1.24g of Co

Conversion from grams to particles

If you have 10.0g of Iron how many atoms do you have?

10.0g x (1mol/55.8g) x (6.022 x 10^23/ 1 mol) = 1.08 x 10^23

Basically just follow the steps on the mole map and make sure you cancel out the units using the fractions method depicted above.

If your doing some work at home you can use this website to check some of you answers:
http://www.convertunits.com/from/grams/to/moles

Video of grams to mole conversion:
http://www.youtube.com/watch?v=ehepBBtSbDc

Video of mole to grams conversion:
http://www.youtube.com/watch?v=NMdN1LtHuDA&feature=related

Thank you!
KP

Mole Conversions: Part I

Good Evening my vaguely hostile readers!

today we will be learning how to convert

 into  .

"is this a joke?" asked the skeptical reader.

Well my dear reader, of course it is! We all know that moles don't like lotus roots!

(if anyone gets that reference, you will be forever endowed with cookies, love and my respect)

And before I am burned at the stake for creating such an obscure joke, we will introduce conversions from particles and grams to moles and vice versa. (IS IT GETTING HOT IN HERE OR IS IT JUST ME?)

Before we begin, there are a few ideas/definitions that need to be known:

Atomic Mass

• in ATOMIC MASS UNITS (u)
• the relative mass of atoms compared to the Carbon-12 atom

Formula Mass

• in (u)
• total mass of all atoms in an ionic compound
• to find formula mass
1. count the number of atoms of a single element
• keep the elements separate, if that wasn't obvious enough while counting (DON'T START BURNING MY HAIR, I'M NOT DONE TEACHING. THANK YOU)
2. multiply the by the atomic mass of that element
3. add the masses of all elements with significant figures

Molecular Mass

• in (u)
• total mass of all atoms in a covalent compound, organic, or polyatomic element
• to find molecular mass
• follow steps displayed above

Molar Mass

• mass of one mole of a substance
• the same numerical value as atomic, formula, or molecular mass...
• ...but is expressed in grams/mole (g/mol)

Now that we have laid out some key ideas/definitions, we can now begin conversions!

(and if i stall long enough, maybe your fires will go out. ARGGH I KID I KID)

Conversions Between Particles and Moles

• remember the formula 6.022 x 10^23 particles/ 1 mole
• To convert from particles to moles
• we'll use 3.01 x 10^24 particles of carbon as our example
• we want to get rid of the particles in the formula stated above
• 3.01 x 10^24 particles (1 mole/ 6.022 x 10^23 particles)
• so we divide!
• to get our answer: 5.00 moles
• To convert from moles to particles/molecules/formula units/atoms
• we'll use 0.75 moles of Carbon Dioxide into molecules
• 0.75 moles x (6.022 x 10^23 molecules/ 1 mole)
• so we multiply!
• to get our answer: 4.5 x 10^23 molecules of Carbon Dioxide
• when asked to find more than one type of atom
• count the amount of atoms wanted to be found 
• multiply it by the answer found above
• (2 atoms of Oxygen/1 molecule of Carbon Dioxide) x (4.5 x 10^23 molecules of Carbon Dioxide)
• to get our answer: 9.00 x 10^23 atoms of oxygen

Conversions Between Grams and Moles

• you will need your trusty periodic table for these conversions
• Is similar to unit conversions
• To convert from gram to moles
• we'll use 3.45g of C as our example
• 3.45g x 1 mole/ 12.0g (molar mass)
• so we divide!
• to get our answer: 0.288 moles of Carbon
• To convert from moles to grams
• we'll use 2.04 moles of Carbon as our example
• 2.04 moles x (12.0g/mol / 1 mole)
• so we multiply!
• to get our answer: 24.5 g of C

So as I continue to burn from the fury of my readers, I leave you a link to a mole to molecule calculator, a link to a grams to moles calculator, and 

(and as the wise mole cricket once said “Take a turn at every turn you encounter on your way.")

-JY

The Mole

What is a Mole?
A mole is simply a unit of measurement. Units are invented when existing units are inadequate. A pure substance that containg the same number of chemical usints as there are in atoms. (i.e., 6.022 X 10^23)


Different gases have a constant ratio
-oxygen : hydrogen      16:1     <---these numbers are the atomic weight

Avogadro's Number!

 Avogadro's number, also known as Avogadro's constant, is defined as the quantity of atoms in precisely 12 grams of 12C. The designation is a recognition of Amedeo Avogadro, who was the first to state that a gas' volume is proportional to how many atoms it has. Avogadro's number is given as 6.02214179 x 10^23. Equal volumes of differest gases at the same temperature and pressure have the same number of particle. 

Relative masses are measured in Atomic Mass Unit, also written like 'amu.'

Formula Mass

Ionic- Add amu's

All atoms of a formula of an ionic comound (amu). When figuring out the atomic mass units, be sure to add the mass of each atom.

Ex.,  Potassium Fluoride
             K           F
           39.1       19
     39.1+19= 58.1amu

Molecular Mass
Atoms of a formula in covalent compound (amu)
Ex. Carbon Dioxide
         C         O2
      12.0   +    (16.2 X 2)
           =44amu

Molar Mass
Atomic/molecular/formula mass of any pure substanse (g/mol)

Ex. What is 1 mole of oxygen?
       Answer: 16g/mol


Check out this video guyss ! Thanksss!
http://www.youtube.com/watch?v=zWmyLsGBEDE&feature=related

Graphing the Density of Water

Today we had a short at the beginning of class on the lab that we preformed last day.
Then we went into the computer room and started graphing data that ms. Chen collected during our quiz.

How to Graph Data
Open excel and create a two columns, one for mass and the other for volume.
Then adjust the values so they are correct.
Then highlight the values for your graph and click create graph.
Select scatter (x,y) graph
Edit the special options such as the title, axis and legend (you can also change the colours!)
Right-click one of the data points and select create trend line. A line of best fit will appear
Now highlight the trend line and select show formula. The formula for the trend line will appear.

Our Results:
Using this method we compared the density of cold water to the density of hot water. We found that cold water had a density of 1.34g/cm3 and that hot water had a density of 1.04g/cm3.
Hot water has a lesser density than cold water because the particles are more spread out due to more energy in the substance.


The blue is cold water which is more dense so it stays at the bottom. The red is hot water which is less dense so it floats near the top


Here is a really good video explaining water density
http://www.youtube.com/watch?v=Ak9CBB1bTcc


KP

Lab: Determining Aluminum Foil Thickness

Worry not friends! The aluminum foil will always protect us from aliens, no matter the thickness!

In this lab we apply the formulas learned last post into 'real life situations.' I mean, who doesn't worry about the possible threat of mind control and mind-reading when the thickness of their aluminum foil is --?

For this lab let's say that we do not have a micrometer available, and only a centigram balance, a ruler and 3 pieces of aluminum foil (of relatively the same size) at our disposal. 

So skeptical tin -foil hat wearing reader? How DO we find the protection level of your hat?

Let's start off with information that can possibly given to you:

• the density of aluminum foil is 2.70g/cm^3
• you have a ruler and 3 pieces of aluminum foil

Do not be so pensive, readers! 

Measuring the length and width of your aluminum foil, you can now start to feel much safer!

Remember volume can be found by:

• length x width x height (but in this situation, it is thickness)

OR 

• by mass/density ( which is given to you 2.70 g/cm^3)

And as a concerned citizen of mankind, I feel it is my duty to guide you to a feeling of safety and assurance with your tin foil hats.

You can now find the thickness of your protective headgear with simple math! As you now have:

• mass
• density
• volume
• width and length

Remember, as a proper citizen of mankind, we must always use the proper amount of significant figures when speaking about our tin foil hat thickness.

Also as a hint to my brothers, the accepted value of aluminum foil thickness is 1.55 cm.

Ah! But wait, you ask about errors within your experiment?

Well my troubled reader, worry not! You can find your experimental error with this simple formula:

experimental error = average measurement - accepted value x 100

                       accepted value

With all this information, I believe that I can safely assume my tin foil wearing, fellow humans will live on and prosper!

and for those who lack proper protection from the supernatural, i give you this! a magical link that will not allow me to read your mind!


-JY (I am sad to say, I can no longer read the minds of my readers)

Density

What is density? Well density is the amount per unit size. Density is a physical property of matter
You can use density in everyday life, like swimming for example.

Density= Mass/Volume <-- This formula is interchangeable, we can also use this formula to find the mass and or volume.
For a solid the units for density would be g/cm3.
For a liquid the units for density would be g/mL.
Example.
Volumer=Mass/Density and Mass=(Density)(Volume)
**Remember, if the density of an object is greater than the density of the liquid, it would sink. If the density of an object is lesser than the density of the liquid, it would float.

Watch this! Good song on how to remember the density formula!
http://www.youtube.com/watch?v=TRkCz3zG7w0&feature=related

Precision, Accuracy and Uncertainity

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

Accuracy: is how close the measurement comes to the accepted value

Low accuracy
High Precision

High accuracy
Low precision

High accuracy
High precision

Uncertainty
No measurement is exact; it is an estimate. Only when a set of whole objects are counted can we get an exact number ie. There are 25 cars in the parking lot.

Absolute Uncertainty
Is expressed in the units of a measurement instead of a ratio.

Method 1
The average of three measurements is calculates. (Remove unreasonable data)
The absolute uncertainty is the largest difference between the average and the largest or smallest reasonable measurement.
The answer is written as: The average plus or minus the biggest differnce
ie. 76.5 +- .2g

Method 2
Determind the uncertainty of of the innstruments used.
estimate to 0.1 of the smallest fraction of the measurment. ie. rulers smallest fraction is 1mm therefore we estimate to 0.1mm


Relative Uncertainty= absolute uncertainty/ estimated measurement
This ratio can be expressed as a percent

Significant figures tell us the relative uncertainty, since the last digit is uncertain!

http://www.mathsisfun.com/accuracy-precision.html

KP

Significant Numbers

Greetings Significant People! 

Happy Independence Day for Turkmenistan from the USSR!
!warning! excessive links ahead

When creating measurements, there is always uncertainty. To minimize the importance or meaning of the error; we write down digits of a measurement that will be meaningful. 

To better understand the meaning of significant digits/figures, we'll look at some examples:

• We'll use 3.29 grams as an example
• 3 and 2 are certain digits
• when measuring we are sure that these digits are correct
• 9 is uncertain
• the last digit of any measurement is always uncertain as it is the last number that is measured

Significant digits in the measurement include all of the certain digits plus the first uncertain digit for a given measurement. To further understand these important figures of significance, let us look at how we figure out significant figures! 

• leading zeros are not counted
• 0.000004
• there is only one sig fig
• trailing zeros are counted
• 50.0540
• there are six sig figs
• trailing zeros without a decimal point are not counted
• 1 000 000 000
• there is only one sig fig

Exact Numbers

Some quantities are defined as a certain amount.

There is no need for rounding and have an infinite amount of significant figures.

Objects that are counted are exact numbers.

Rounding Rules

When writing answers, we must round to the appropriate number

• if the number is > than 5; round up
• if the number is < than 5; round down
• if the number is 5 with non-zero digits following it; round up
• if the number is 5; make the last digit 'even'
• even numbers to be exact (2, 4, 6, 8)

http://www.youtube.com/watch?v=5UjwJ9PIUvE


Adding and Subtracting Significant Digits
When adding or subtracting, we round to the fewest number of decimal places, as well as the first uncertain digit.
http://www.youtube.com/watch?v=U6k7VpdW_rQ
http://www.youtube.com/watch?v=cEMSHsIeKMM&feature=related


Multiplying and Dividing Significant Digits
When multiplying or dividing, we round to the fewest number of significant digits.
http://www.youtube.com/watch?v=qm2bY8tcNQ0
http://www.youtube.com/watch?v=oxqW1LFm7aw&feature=related

and for those who enjoy awkward music videos

http://www.youtube.com/watch?v=ZuVPkBb-z2I

-JY

Lab 3B: Separation Of A Mixture By Paper Chromatography

So todays class we had a lab involving one of the separation techniques we learned from the previous. The separation technique we used for todays lab was paper chromatography. We used the paper strips to make separations between various food colorings. The was very successful with mild problems. Some of the groups results were very low, but as an over all average everything was pretty up to scale. The video below is an example of paper chromatography, so a basic summary of todays lab, enjoy :)

http://www.youtube.com/watch?v=IRZ4lHEe1DI&feature=related

P.S., STUDY FOR THE TEST!

Separation Techniques

The basis of separation is that everything is made up of different components or properties. These components can be separated and then individually analysed.

Strategy: Devise a process that will separated the substance best. Use properties such as

Low/High density

(non) volatile

(non) soluble 

(un) reactive

(non) magnetic

(non) polar

Properties of Separation:

The components stay the same

The more similar the components are, they harder it is to separate them.

Examples of separation methods:

Hard Separation (two solids): mechanical or heterogeneous mixture separated by a magnet sieve, etc.

Evaporation (Solid dissolved in a liquid): Boil away the liquid to leave the solid. 

Filtration ( solid (not dissolved) in a liquid): Pour through a porous filter, solids stay behind because they are larger than the pores in the filter.

Crystallization (solid in a liquid): A precipitate is formed by a chemical or physical reaction. The solution is then filtered. Then it is evaporated to form crystals which is then filtered again to produce the remaining solvent.

Gravity (solids): A centrifuge whirls a test tube around, pushing the more denser objects to the bottom. (Works best for small volumes)

Solvent Extraction: Solvent is used to dissolve one component of the mixture. Only works when the solvent dissolves one component.

     Liquid: dissolves one solid, leaves the other behind.

     Solution: Solvent is insoluble with solvent present. The solvent dissolves one or more of the substances and leaves behind the unwanted substances. Then it is shaken in a separatory funnel which forms layers in the solution. Some of these layers can be drained to leave you with the wanted material.

Distillation(Two Liquids): Mixture is heated, one substance has a lower boiling point than the other. It evaporates up a tube and then condenses on the other side. 

                                            

Chromatography: Used to separate very complex mixtures. A mobile phase is swept over a stationary phase. Components move over stationary phase at different speeds, allowing each components to be collected individually, producing very precise and accurate information.

    Sheet Chromatography

      Paper Chromatography: The stionary phase is paper soaked in a liquid. The mobile phase is a solvent. The results appear as dots on the paper.

      Thin Layer Chromatography: The stationary phase is Al 2 O3 or Si O2 which is an absorbent on glass. The mobile phase is poured over, some components bond with the absorbent, appearing as dots on the sheet.

Here is a very good video explaining many of the above mentioned separation techniques

http://www.youtube.com/watch?v=jWdu_RVy5_A&feature=related

Law of Multiple Prportions

Elements can combine in different ratios to form compounds

Such as NO2, NO, N2O or N2O5

Naming Acids

Alright wonderful people! If you haven't guessed, we're going to be learning about how to name acids!

Now because this is chemistry, and not computer virus class; I will not be teaching you how to corrupt files from 1992. BAD READERS.  BAD.

Acids are formed when a compound made of hydrogen ions and a negatively charged ion are dissolved in water. We call this state aqueous.

When acids are placed in water:

• Hydrogen combines with water
• Negatively charged ions dissolve in the water; separating them
• Hydrogen joins with water to form H₃o (an hydronium ion)

Here are basic guidelines to naming acids

meaning not EVERY acid is named this way, every rule has an exception

Simple Acids (group 16&17 of periodic table)

• use 'hydro' as your prefix
• this indicates that hydrogen is present
• last syllable of the non-metal is dropped
• the suffix 'ic' is used in it's place
• add the word 'acid' at the end
• _______ide --> hydro_______ic acid
• insert your ion into both blanks

Complex Acids 

• you do not use 'hydrogen' while naming complex acids
• your negative ion suffix will change according do its name
• ___ate is replaced with ____ic
• ___ite is replaced with ____ous
• add the word 'acid' at the end

-acids that end with 'ide' are considered simple-

Helpful Hint Honouring Hamish

Acronyms are extremely helpful when memorizing anything.

Here is one courtesy of a teacher:

we ate ic-y sushi and got appendic ite-ous.

chortle, chortle, it's so clever!

• the endings are together in this acronym, so it's pretty self explanatory.

If this post isn't clear enough; here is a website that can explain, and even has interactive excercises!

JY

Review: Naming and Writing Ionic and Covalent Compounds

In todays class we had a review for writing and naming ionic compounds. Listed below is a summary of the class :)!
Ionic Compound
-composed of two or more particles that must be oppositely charged
-electorstatic forces hold them together
-metal-->non-metal transfer

Example: Na^+ P^3-
=Na3P  (criss cross)
=Sodium Phosphide

Example: Fe^2+ S^2-                **use roman numerals when an element has more than one charge**
=FeO
=Iron(II)Oxide

Complex Anions
-group of atoms that act as one atom

Example:
a)zinc phosphate-->Zn3PO4
b)tin(II) permanganate--> Sn(MnO4)2
c)Pb(CO3)2--> lead(II) carbonate

Covalent Compounds
-share electrons
-non-metal with non-metal
**Diatomic molecules: H2, O2, F2, Br2, N2, Cl2, I2**
-use greek prefixes to indicate the nuber of atoms


Example:
a)CO2--> carbon dioxide
b)sulphur trioxide--> SO3

DONT FORGET TO DO YOUR HOMEWORK! ENJOY THE LONG WEEKEND! HAPPY THANKSGIVING :)!

Heating and Cooling Curves of a Pure Substance Lab 2B

Today in class we preformed an experiment in which we studied the properties of cooling and heating dodecanoic acid (C11H23COOH)

I would write down the answers to our investigations but my work in currently in the classroom, since Ms. Chen collected it after the lab. However here is a video explaining heating and cooling curves.

http://www.youtube.com/watch?v=DhZ3r9qp7Ik

Dodecanoic acid

KP

Concepts of Matter (warning: teal deer crossing)

Hello fellow chemists! To segway out of an awkward introduction and into the nitty gritty stuff; I present you a warning!

an online phrase commonly used in discussion forums as a response to previous posts that are deemed unnecessarily long and extensive; it’s short for Too Long; Didn’t Read.


Alright! So with our previous knowledge of matter, we are going to explore different concepts around matter. 

HEATH CHEMISTRY: pg 25-34, pg 36-39

Macroscopic changes in matter

• are changes we are able to see with our naked eye

Mixtures and Pure Substances

Mixtures are two (or more) kinds of matter that have visible components, they are also called impure
Solutions are substances that look uniform but are made of two (or more) different kinds of matter

• mixtures scatter light since it contains different visible components.



-insert irrelevant picture and caption here!- 

• this would mean we are able to separate the mixture into it's component parts
• With muddy water; we are able to use alum or lime to produce that a jelly-like substance that can be filtered out. This separates the substance into its component parts
• BUT, what if we described solutions like salt water as a mixture?
• we than imply that we can separate it into different components.

Distillation is the process of heating a liquid until it boils, capturing and cooling the resultant hot vapors, and collecting the condensed vapors.

•  distillation gives us a clue about how solutions can be separated
• With salt water, we are able to boil away the water. This leaves the salt behind, therefore, separating the components.

Only when we can no longer separate substances can we call it pure, however, sometimes it is extremely difficult to separate some substances. This could mean it could take years to discover impure properties of a substance (Say hello to the future advancements in science!)

Pure Substances are matter that are made out of the same particle

• pure substances have a constant boiling point and freezing point
• with water: 0° (freezing point) and 100° (boiling point)
• mixtures usually do not have constant boiling/freezing points

Heating/Cooling Curve

The curve shows the process in which a pure substance changes from solid to gas

The cooling curve is the opposite of the heating curve; going from top left to bottom right.

• Point A
• particles are closely pack together
• Point A-B
• heat is converted to kinetic energy
• substance starts to melt
• Point B
• point where it will either melt or freeze
• Point B-C
• exists in both solid and liquid state
• temperature remains constant
• melting/freezing point
• Point C
• substance is liquid
• Point C-D
• still in liquid state
• kinetic energy increases as particles move faster
• Point D
• begins to change into gas
• Point D-E
• exists in both liquid and gas state
• boiling point
• Point E
• all liquid is now gas
• Point E-F
• gas particles continue to absorb energy and move faster
• temperature increases as heating continues

Chemical and Physical Change

Density is a property of matter that describes volume

Chemical Changes are irreversible 

• produce new matter with different properties than the original matter
• When sugar is heated, it bubbles and turns black as well as a colourless liquid forming. When the black matter and liquid are mixed together it does not change back to the original sugar
• electrolysis
• matter decomposes to form new kinds of matter
• decomposition 
• of a pure substance in which they are separated into components

Physical Change is easily reversed

• it is a change of state 
• it does not appear to create a new kind of matter
• we are able to get the original matter back again
• When moth flakes are heated, the solid material disappears and a liquid is in it's place. If we allow the liquid to cool it changes into a solid. This solid has the same properties as the moth flakes.
• distillation
• of a pure substance in which they are separated into components that already exists

Compounds

Compounds are pure substances that can be decomposed into new kinds of matter

• They appear to be put together from simpler substances
• must be made out of 2 more kinds of atoms
• separated if there is enough energy supplied to separate the components
• not all compounds are made out of molecules; some are made out of elements

Ions are particles that have an electrical charge 

• to know which compounds are ionic and which compounds are molecular is to check them for conductivity

Law of Definite Composition

Law of definite composition is an experimental fact that defines that all compounds have a definite composition

• water is always made out of 2 hydrogens and 1 oxygen
• however, mixtures can have almost any composition desired

Law of Multiple Proportions

Law of Multiple Proportions explains that two or more compounds with different proportions of the same element can be made

• they do not represent the same compound
• they are just multiples of each other
• do not have the same proportions as mixtures

Elements

Microscopic Model is a smaller representation of matter, rather than macroscopic(which is seeing, smelling, touching)

Elements are pure substances that cannot be decomposed into simpler substances

• we assume elements are made out of only one kind of atom
• each element contains a different kind of atom
• elements can exist in solid, liquid, or gaseous state
• elements change state through changes in temperature

Molecules are particles made of more than one atom

• some elements exist in larger units
• have definite shapes and composition

Atoms are the smallest unit of an element

• are usually represented in spheres, spheres are used to suggest the relative size of the atom
• atoms in solid state are closely packed together in an orderly,organized pattern
• atoms in liquid state are still close together but; are no longer in an organized pattern
• atoms in gaseous state move very far apart and move into a straight line until it collides with another atom/container

Still alive and kicking everyone? Don't worry, if you've stayed to the end you get a cookie. But don't lie. I don't give cookies to people who tl;dr. People who tl;dr make me sad.

enjoy your weekend!

-JY