Dec.9--Lab 4C [Formula of a Hydrate]

Equipments:

- pipestem triangle

- iron ring & stand

- Bunsen burner

- dry crucible

  

Procedure:

1.      Put on lab apron and safety goggles.

2.      Set up the pipestem triangle, iron ring, stand and Bunsen burner. The ring should be 5-6 cm above the burner. Place a clean and dry crucible on the triangle, and heat for 3 minutes to make sure it is dry.

3.      Weigh the crucible and record the mass in Table 1.

4.      Place hydrate into it until it is 1/3 full. Record the total mass of crucible and hydrate.

5.      Place the crucible on the pipestem triangle, and begin heating. Gradually increase the heat until the bottom of the crucible is a dull red. Maintain the temperature for 5 minutes.

6.      Turn off the burner and allow the crucible to cool for 5 minutes. Weigh and record the total mass.

7.      Reheat the crucible for 5 minutes to make sure there is no more water left. Cool it down and record the mass. Use the lower mass number between the first and second heat. (The difference should be no more than 0.03 g.)

8.      Add a few drops of water to the contents of the crucible. Note any changes that occur.

Dec 5--Calculating the empirical formula of organic compounds

The empirical formula of an organic compounds can be found by:

-buring the compound (reacting with oxygen)

-collecting and weighing the product

-calculate the mass and then mole

*The mass of the product = the mass of the reaction

Ex.

What is the empirical formula of a compound that when a 10.0g sample is burned produces 20.0g of CO2 and  8.0g of H2O?

CxHy + z O2   TO   x CO2 +  y/2 H2O

1. Calculate the moles of CO2 and H2O

CO2   20.0g x (1mole/44g) =0.455 mol

H2O   8.0g x ( 1 mole/18g)=0.444 mol

2. find the mole of C and mole of H in the H2O and CO2

mol C =0.455 mol x (1mol C / 1mol CO2) = 0.455 mol  

mol H=0.444 mol x( 2mol H / 2mol H2O)=0.888 mol 

3. Divide the smallest molar amount

C  0.455/0.455 = 1

 H  0.888/0.455 = 2

Therefore ,the answer is CH2 .

Dec 3--Empirical + Molecular Formula

Empirical Formula:

gives the lowest term ratio of atoms (or moles) in the formula. (All ionic compounds are empirical formulas.)

Ex. C4H10—molecular formula     C2H5—empirical formula
 

Example#1

Consider that we have 10.87g of Fe and 4.66g of O, what is the empirical formula?

1.Convert g into moles                              
 Fe: 10.87g of Fe* （1 mole/55.8g）=0.195 moles
 O: 4.66g of O * (1 mole/16g) = 0.291 moles

2.Divide both of them by the smallest molar amount
Fe: 0.195/0.195=1       
O: 0.291/0.195 = 1.5

3.The last step is to scale ratios to whole numbers
Fe: 1 * 2 = 2        
O: 1.5* 2 = 3
 

Ans:Fe2O3

Example#2

A compound contains 31.9% k, 28.9%Cl, and 39.2% O. What is the empirical formula?
 
K: 31.9g * (1 mole/ 39.1g) = 0.816 mol
Cl: 28.9g * (1 mole/35.5g) = 0.814 mol
O: 39.2g * (1 mole/16g) = 2.45 mol
Repeat the second and third step.

Molecular Formula: 

is a multiple of the empirical formula and shows the actual number of atoms that combine to form a molecule.
 
n= (molar mass of the compound/molar mass of the empirical formula) 

Example#1

A molecule has an empirical formula of C2H5 and a molar mass of 58g/mol what is the molecular formula?

MM C2H5 = 29g/mol
n = (58g/mol / 29g/mol) = 2
MF = 2(C2H5) = C4H10 = BUTANE

Example#2

The empirical formula of a gas is CH2 what is the molecular formula if the molar mass  is 42g/mol

MM of CH2 = 14g/mol
n= (42g/mol / 14g/mol) = 3
MF = 3(CH2) = C3H6

Example#3

 A compound contains 7.44g C, 1.24g H, and 9.92g O , the molar mass of the compound is 180g what is the molecular formula?

C 7.44g * (1 mole/ 12.0g) = 0.62mole     1
H 1.24g * (1mole/ 1.0g) =1.24 moles      2
O 9.92g * (1 mole/ 16.0g) = 0.62 mole    1
CH2O = 12.0+1.0*2 +16.0=30g/mol
n = (180g/mol / 30g/mol) = 6
MF = 6(CH2O) = C6H12O6

Nov,22--Harder Mole Conversions

Mole Map





Nov.17--The Mole

Equal volumes of different gases at the same temperature and pressure have the same number of particles.

↓

Avogadro’s Hypothesis: If they have the same number of particles, the mass ratio is due to the mass of the particles.

Atomic Mass: mass of one atom of the element in atomic mass units

Ex. Fluorine = 19.0 amu

Formula Mass: mass of all the atoms of a formula of an ionic compound (in amu)

Ex. Potassium Fluorine = K (39.1) + F (19.0) = 58.1 amu

Molecular Mass: mass of all the atoms of a formula of a covalent compound (in amu)

Ex. Carbon Dioxide = C (12.0) + O₂ (16.0×2) = 44.0 amu

   express in grams per mole: 44.0 g/mol

1 mole of oxygen = 16.0 g/mol

1 mole of carbon = 12.0 g/mol

• They have the same number of particles!
 

 

Youtube Video: introduction to the idea of a mole as a number
http://www.youtube.com/watch?v=AsqEkF7hcII 

Avogadro’s Number: The number of particles in 1 mole of any amount of substance is:

  particles/mol

Nov.5--(Lab 2E) Determining Aluminum Foil Thickness

In order to find out the thickness of aluminum foil, you must remember the formula:

 D= M/ V

(D: density, M: mass, V: volume )



Equipment:
3 rectangular pieces of aluminum foil

metric ruler

centigram balance

Procedure:
1.Use a metric ruler to measure the length and width of each piece of foil.
2.Use a centigram balance to find the mass of each piece of foil
(Using significant figures to record the quantities)
3.Density of aluminum foil is 2.70g/cm’3
4.Calculate the volume use: V= M/D
5.Then H=V/LW (should be expressed in scientific notation)



Experimental Error:
(your measurement - Accepted value) / Accepted value*100  %

Nov,1--Density

Density = Mass/Volume *(D=M/V)

- unit for liquid is g/mL

  unit for solid is g/cm3

How to calculate density ? There is a Youtube Video can help you understand better.

*The density of water is 1g/mL

-If the density of an object is greater than the density of the liquid---- sink

 If the density of an object is smaller than the density of the liquid---- float

A Youtube Video about sinking and floating in liquid

Oct,31--Accuracy and Precision

Precision: is how reproducible a measurement is comparing to other similar measurements (approximately to other numbers).


Accuracy: is how close the measurement (or average measurement) is to the actual value.

Measurements and Uncertainty

- Measurements are only best estimates. Every measurement has some degree of uncertainty.
(Except for numbers like 5 people, three bags, etc.)

Absolute Uncertainty
-The uncertainty expressed in the units of measurement, not as a ratio.

How to calculate the absolute uncertainty?

♥Method 1: Measure for more than 3 times, then calculate the average.(You may omit the measurement that is most different from others.) The absolute uncertainty is the largest difference between the average and the lowest or highest reasonable measurement.

♥Method 2: (determine the uncertainty of each instrument) When making a measurement, always measure the best precision, which means you should estimate to a fraction 0.1 of the smallest segment on the instrument scale.

Relative uncertainty and significant figures

*Relative uncertainty=absolute uncertainty/estimated measurement

(The number of significant figures indicates the relative uncertainty)

Oct,31--Significant Figures

Significant digits are meaningful digits. More significant digits means more precise.
-The last digit of the measurement is uncertain (because that digit could be one digit higher or one digit lower
-ex. 5.23 then become 5.2, 5.26 then become 5.3)
Ex 1: 9.34 9 and 3 are certain, 4 is uncertain
-The significant figures include all of the certain digits and the first uncertain digitEx 2: 9.34 have 3 significant digits

Rules of counted significant digits in the measurement:
-leading zeros aren’t counted
-Trailing zeros after the decimal point are counted
-Trailing zeros without a decimal point are not counted

Exact numbers
-quantities that are defined as exactly a certain amount (we don’t need to round it)
Rounding rules:-If that digit is ＞5，round up
-if that digit is ＜5，keep it the same
-if that digit is =5, and there are more non-zero digits after the 5 round up
-if that digit is =5, and it ends at the 5 round to make the last digit “Even” (0, 2, 4, 6, 8)

Math Rules (+ and -)(x and /):
 round to the fewest number of decimal places (by position)
*: round to the fewest number significant digits

Oct.26--Significant Figures

Significant figures are measured or meaningful digits. The more significant digits there are in a number, the more precise the number is.

Uncertain Digit: the last digit in a measurement is called the uncertain digit as it could be one digit lower or higher depending on the digit after it.

The significant digits in the measurement include ALL of the certain digits and the first uncertain digit.

Ex. 2.56g : 2 & 5 are certain, 6 is uncertain

   2.56 has three significant digits

Zeros

- Leading zeros are not counted

- Trailing zeros after the decimal point are counted

- Trailing zeros without a decimal point are not counted

Exact Numbers

Some quantities are defined as exactly a certain amount. You do not need to round those numbers.

Ex. a pair of gloves, five people, etc.

Rounding Rules: how to round measurements

1.      Look at the digit after the position of rounding.

2.      If the digit is greater than 5, round up.

3.      If the digit is less than 5, keep the same.

4.      If the digit is equal to 5, look at the digit(s) after it:

- more non-zero digits, round up

- it ends at 5, round to make the last digit even

* When rounding, do not round until the final answer!

Math Rules

Adding or subtracting: round to the fewest number of decimal places

Multiplying or dividing: round to the fewest number of significant digits.

Finally, there is a Youtube Video that can help understand better.

Lab: Separation of a Mixture by Paper Chromatography

In this lab, we learned how to use paper chromatography to separate different components of a mixture.

The mixture we used was food coloring.

[Procedure Flowchart]

Part One

l        Prepare three large test tubes and place them in three Erlenmeyer flasks. Label the test tubes A, B, and C.

l        Cut a 66cm chromatography paper into three (each 22 cm). Use a pencil to draw a lien across each strip 4 cm from one end. Cut this end into a point.

l        Place 2 cm deep water in each test tube.

Part Two

l        Use a glass stirring rod and spot the strip of paper with the color assigned. (The spot should be smaller than 0.5 cm in diameter)

l        Insert the strip in test tube A. Observe what happens to the sample spot as the water slowly moves up.

l        Identify the solute front and solvent front. After about 20 minutes, when no further separation of color will occur, remove the stripe. Immediately draw a pencil line across the top edge of the solvent front before it disappears.

l        Measure d1 and d2, record in Tables 1 and 2. Calculate the Rf value and record it as well.

l        Clean up.

Part Three

l        Take the second strip of paper and spot it with a sample of green food coloring. Spot the third strip with a sample of the unknown mixture of food colorings. Label them at the top.

l        Insert the strips in test tubes B and C. Follow the same procedures as in Part Two.

l        Record data in Table 3

l        Clean up.

Oct 17--Separation Techniques

Separating Mixture

How do you separate different components in a mixture?

- by discriminating between components with different properties.

l        high / low density

l        volatile / nonvolatile

l        reactive/ inert

l        soluble / insoluble

l        magnetic / non magnetic

Note: Separation works because components in a mixture retain their identities. However, it would be more difficult to separate if the properties are similar.

²       Filtration: select components by particle size

²       Floatation: select components by density

²       Crystallization and Extraction: select components by solubility

²       Distillation: select components by boiling point

²       Chromatography: select components by

Hand Separation (solids + solids)

- a mechanical mixture or heterogeneous mixture

- can be separated by using a magnet or sieve

Evaporation (solid dissolved in liquid solution)

- boil until the liquid evaporates

- solid part remains

Filtration [solids(not dissolved) and liquids]

- separate by passing the mixture through a porous filter

- use filter paper, the residue is left in the paper while filtrate goes through

Cystallization (solid in liquid)

- precipitation is the conversion of a solute to solid form by physical or   chemical change

-solid separated by 1. filtration

            2. floatation

-saturated solution of a desired solid

-solid come out as pure crystals, then crystals are separated from the remaining  solvent by filtration

Gravity separation

-solid based on different density

-a centrifuge whirls the test tube around at  high speed, which forces the denser  materials to the bottom

*small volumes works better

Solvent extration

-a component moves into a solvent shaken with the mixture

-solvents that dissolve only one component works better

-mechanical mixture (solid in solid) : use liquid to dissolve one solid , then you get the other one

-solution: solvent dissolves one and leaves unwanted solid behind

Distillation (liquid in liquid)

- heating cause low-boiling component to vaporize first

-collecting and condensing  volatilized component 

(Vapour ascents to distillation flask and enters condenser)

Chromatography
-The material of the mixture retains some components, different components flow over the material have different speeds.

- A mobile phase sweeps the sample over a stationary phase.

-This can separate very complex mixtures and  form of separation uses very small sample sizes and analysis is highly accurate and precise.

Sheet Chromatography

Paper Chromatography  (PC) :

-Stationary phase (absorbent) is a liquid soaked into a stripe of paper, mobile phase is a liquid solvent.
-different components move at a different speeds. Some components spend more time in the stationary phase. The components appear as separated spots on the paper.

Thin Layer Chromatography (TLC) :

-stationary phase (ex. Al2O3, SiO2) is a thin layer of absorbent.
-some components bond to the absorbent strongly, some weakly. ( the speeds of the components)

Oct.13--Naming Acids

Compounds with hydrogen ions and a negatively charged ion can form acids when dissolved in water. At the same time, ions are separated from each other.

* H⁺ ion joins with H₂O to from H₃O⁺ (hydronium ion)

 For example, HCl (g) + H₂O (l) à H₃O⁺ (aq) + Cl^- (aq)

How To Name Simple Acids:

1. Use “hydro” as the beginning.

2. Last syllable of the non metal is dropped and replaced with “-ic”.

3. Add “acid” at the end.

( _ide = hydro_ ic acid )

For example: HF = hydrofluoric acid

                        H2S = hydrosulphuric acid

How To Name Complex Acids:

1. change “-ate” into “-ic

change “-ite” into “-ious”

2. Add “acid” at the end of the name

Still confused? Here's a  Youtube video for you!

 

Oct.7--Writing and Naming Ionic and Covalent Compounds

Ionic Compounds

- composed of 2 or more ions

- held together by electrostatic forces

- metal transfers electrons to non metal

Ex. Ca (2+) & O (2-)

   +2 + (-2) = 0

        CaO

     Calcium Oxide (always metal first!)

Ex. metals that have 2 or more charges

   Manganese (II) Chloride

    +2 + (-1) 2 = 0

     MnCl₂

Ex. Complex Ions: a group of atoms that behave as one atom

Calcium Hydroxide

+2 + (-1)2 = 0

     Ca(OH)₂

Covalent Compounds

-share electrons

- non metal & non metal

- Diatomic Molecules: N₂, O₂, F₂, Cl₂, Br₂, I₂, H₂

- use Greek prefixes to indicate the number of atoms

Ex. CO₂ = Carbon Dioxide

   N₂O₄ = Dinitrogen Tetroxide

Greek prefixes to indicate the number of atoms
Mono-	1
Di-	2
Tri-	3
Tetra-	4
Penta-	5
Hexa-	6
Hepta-	7
Octa-	8
Nona-	9
Deca-	10

Summary Pg 25-34, 36-39

2-1 What You Know About Matter

-Identify matter and distinguish one kind from another: different look, taste, densities, colors, boiling points.

2-2 Purifying Matter

1. Mixture - two or more kinds of matter that have separate identities

2. Solutions - mixtures like salt water or sugar water that look uniform throughout and do not scatter light

3. Mixtures are matters that can be separated into component parts and have different identities.

4. Distillation - solutions can be separated into component parts by boiling.

6. Matter can be separated into mixtures and pure substances.

2-3 Characteristics of Pure Substances

- Pure substances have a constant boiling point, mixtures do not.

- The difference between pure substances and mixtures is solidify

- The similar differences between mixtures and pure is freeze

- Freezing point: temperature at which liquid change to solid

- Melting point: temperature at which solid change to liquid

2-4 Chemical and Physical Changes

- Density: a property of matter that describes its mass per unit volume
- Chemical changes: changes that produce a new kind of matter with different properties.
  Decomposition is one type of chemical change. One kind of matter comes apart to form two or 
  more kidns of matter during decomposition.
- Melting: the change of a solid to a liquid without the formation of any new kidn of matter
- Physical changes: changes that are easily reversed to get the original material back again, 
  and do not produce new substances.
Ex. boiling


2-5 Compounds and Elements

- Electrical conductivity tester is used to test electrical conductivity. If the two wires are connected by a piece of metal or any other matter that conducts electricity, the light bulb will glow because electricity will flow through.
- Electrolysis involves passing an electric current through a substance, causing it to decompose into new kinds of matter.
- The electrolysis of sodium chloride and water are chemical changes.


Differences between decomposition and distillation:
Decomposition
- a single, pure substance is changed into new substances with different properties
- chemical change


Distillation
- the separated components exist in the original mixture as separate substances
- physical change

2-6 Compounds Have a Definite Composition

- Law of definite composition: the compounds of elements always have a definite composition.

Ex. The definite composition of water was described in terms of volume: The volume of hydrogen gas obtained from water is always twice the volume of oxygen gas obtained.

- Law of multiple proportions: two or more compounds with different proportions of the same elements can be made.

- Any sample of a given compound contains the same percentage by mass of each element. Samples of compounds which contain the same elements but in different percentages by mass are different.

2-7 Matter Is Made Of Atoms

- Macroscopic observations: observations of an experiment that are large enough to be visible to the naked eye

- Macroscopic properties: melting point, boiling point, heat of fusion, temperature, mass

- Mixture differs from a pure substance: A pure substance has a definite boiling point, a mixture does not. A mixture can be broken down, but a pure substance cannot.

- Explain: a word that talks about the things as they might be.

- Scientific explanation: off a way to observe why things happen (can be intelligent guesses that are based on observation)

- Microscopic models are too small to be seen except under a microscope. When used with SI units, means one millionth) and use to explain the behavior of matter.

- Matter is composed of some kind of small pieces-atoms, molecules

- Atoms: the smallest component of an element which have the chemical properties of the element.

- Atoms are represented with spheres of various sizes and colors

- Size of atoms: the relative size of the actual atom, but not the representations of real atoms

- Elements are pure substances and cannot be broken down.

(If each element has different kind of atoms, there are 109 atoms because there are 109 elements)

1-109: atomic numbers

Solid: atoms are stuck together

Liquid: atoms are still close together, but also can move past one another

Gas: atoms are far apart, moving in a straight line until they collide with other gas atoms or the wills of container)

As temperature increases, atoms in solid vibrate move. Heat is absorbed to overcome the forces of attraction that hold the particles together.

The force of gravity pulls the atoms down, so the liquid flow to take the shape of its container.

Temperature of a liquid is raised to its boiling point. Then kinetic energy increases, and atoms move faster. They eventually escape from the liquid to move farther apart and become gas.

Molecules: particles made of more than one atom

(Some elements exist in various combinations of their atoms forming molecules.)

Compound: a substance formed by chemical union of two or more elements in definite proportion. (Because elements contain only one kind of atom, so compounds are made of two or more kind of atoms.)

When a compound decomposes, the different atoms are separated (need enough energy-heat, electricity)

Compounds: solids, liquids, gases

If the atoms within a water molecule came apart, new substances are formed (hydrogen and oxygen) with new properties

Molecules can separate, atoms cannot. More energy is required to break the bonds which hold atoms together in a molecule.

Ions: particles that have an electrical charge

The only way to know which compound are ionic (melt to form ions)and which compounds are molecular(melt as molecules)is to check them for conductivity.

Oct.4--Heating/Cooling Curve Of a Pure Substance

 

A: (solid state at any temperature below its melting point) Particles are very closely packed together in an orderly manner. Forces between the particles are very strong and can only vibrate at a fixed position.

A-B: As it is heated, heat energy is converted to kinetic energy. Kinetic energy increases and the molecules vibrate faster about their fixed positions, and the temperature increases.

B: (still solid) Melting has begun. Temperature remains the same, which is called the melting point. Solid begins to change into liquid.

B-C: (exists in both solid and liquid state) The temperature remains constant because the heat that is supplied to it is used to overcome the forces of attraction that hold the particles together. The heat energy that absorbed to overcome intermolecular forces is called the latent heat of fusion.

C: All has completely melted. Solid has turned into liquid.

C-D: As the liquid is heated, the molecules gain more heat energy and the temperature increase. The particles move faster because their kinetic energy is increasing. (in liquid state)

D: (still in liquid state) Molecules have received enough energy to overcome the forces of attraction between the particles. Some of the molecules start to move freely and liquid begin to change into gas.

D-E: (exists in both liquid and gaseous states) Temperature remains unchanged. The heat energy absorbed is used to overcome the intermolecular forces between the particles of the liquid rather than increase the temperature. This constant temperature is the boiling point.

E: All liquids have turned into gas.

E-F: The gas particles continue to absorb more energy and move faster. Temperature increases.