Amines and Amides

Amines

-have a nitrogen with carbon bonded to it

Amides

-Amides are used in nylon, kelvar, penicillin, proteins, LSD

Esters, Carboxylic Acids, Aldehydes

REVIEW from last class
Alcohols
R-OH


ketones
R=O


ethers
R-O-R1(doesn't have to be the same)


halidea
R-F, R-Cl, R-Br


NEW STUFF!!!


Aldehydes -change ending to '-al'
-attach a double bonded oxygen and a hydrogen to either end of the carbon chain


Carboxylic acids
-found in insect bites
-building blocks of fats/steroids
-change ending to '-oic'
-attach a double bonded oxygen and a hydroxide to either end of the carbon chain


Esters
-used as artificial flavoring for purfumes, synthetic fibers
-source of odour/flavour in fruits
-name the primary R with a -yl ending
-name the secondary R with -ate ending
-esters are formed through esterfication
- react a carboxylic acid with an alcohol

Functional Groups

Today we had a quiz! Also we learnt about functional groups.

1) Halides

- Group 17 elements

-bromine(bromo), chlorine(chloro), fluorine(floro)

-you use this when you see bromine, chlorine or fluorine as a side chain to your carbon chain

2) Alcohols

- have an OH(hydroxyl) function group

-same rules as usual but change the ending to '-ol'

3) Ethers

-there is an oxygen atom between two carbon chains

-use side chain endings(-yl) for both chains followed by ether

4) Ketones

-double bonded oxygen

-change the ending to '-one'

Alkynes/ Cis and Trans/ Alycyclics/ Aromatics

Alkynes (triple bonds)
-use -yne ending after parent chain
-number using lowest possible number

Cis and Trans
-side chains cannot rotate around double bonds
-cis: same side of side chain
-trans: opposite side of side chain

Alycyclics
-all the same naming rules apply
-add cyclo- to the name of the parent chain

Aromatics
-cyclohexene can have double bonds in its ring



wow okay this guy is kinda weird but he does a prett good job of explaining it so ill let him do so!

Continutation of Organic Chemistry

B. ALKENES (double bonds)
1. All the same rules for naming appy
2. Use the correct prefix and -ENE ending
3. Identify the location of the double bond using carbon numbers

Here is a website you can use to practice, the answer are here too!
http://www.sciencegeek.net/APchemistry/APtaters/alkanes.htm

C. TRIPLE BONDS
1. Use the -YNE ending

For some more detailed notes, here's a link to an online tutor!!

ORGANIC CHEMISTRY

- Organic Chemistry is the study of the carbon atom and its compounds.

- carbon compounds outnumber all other compounds combined.....

> carbon can form many different bonds

> the can have different arrangements

> can form long chains of carbon atoms

HydroCarbon Hierarchy - click here for a more in depth look

There are 3 types of chemical formulas

1) Molecular - eg C3H₈
2) condensed structural - eg

CH3-CH2-CH3

3) structural

Isomers

One molecular formula can represent several different compounds. These are called isomers

EG Alkanes

1. circle the largest carbon chain. name the chain with the appropriate prefix ane ending

2. locate any side chains by numbering the carbon chain. use lowest possible #s

3. name the side chain with appropriate prefix and yl ending

4. if there are more than one of the same alkyl side chain add the multiplier in front of the branch name

EXAMPLE

Hydrogen Bonding

k so last class we started learning about molecular polarity and intermolecular bonds, but we didn't get the chance to get to hydrogen bonds so thats what we talked about the most today

• hydrogen bonding is a special type of dipole-dipole bond
• it occurs between hydrogen and nitrogen, oxygen and flourine

H-N, H-O, H-F

now just to recap all the the types of bonding forces:

• London Dispursion Force (L.D.F) is experienced by all molecules but is the weakest of bonds
• Dipole-Dipole is only present in some molecules, and is stronger than LDF
• Hydrogen Bonds are the strongest of all the molecular bonds and are only found in a couple of different molecules

one example we went through was putting some molecules in order of the strongest bond to the weakest bond. the ones we used were:

1. C2H5OH: 26 electrons and a bp of 78 degrees Celsius

2. CH3OH: 18 elestrons with a bp of 65 degrees Celsius

3. C2H6: 18 elesctrons with a bp of -89 degrees Celsius

4. CH4: 10 electrons with a bp of -161 degrees Celsius

now that we had finished talking about molecular bonds, we discussed ions in solutions for the rest of class

• dissociation is the splitting of ionic solids into ions
• there are multiple steps you must follow in writing a dissociation equation
• determine the ions that are in your ionic solid and plit them up on the other side of the equation from your original solid
• write (aq) next to the ions to show that they are now aqueous due to the fact they are dissovled in water
• find the charge of each of the ions and write it
• balance the equation using the charge of each ions to determine how many of them there should be

we went through some concentration examples that go back to our stoichiometry section, so its basically using our old knowledge to kind of figure out new examples. I won't write out the examples bacause they would just be WAY too confusing using blogger. but instead, i have an awesome video! YAY!

it says 20 second but its more like a minute and a half. still, this does a really good job of explaining hydrogen bonding!

thats all for now!

Molecular Polarity & Intermolecular Bonds

INTERMOLECULAR: intermolecular means the bonds are outside of the molecule

A. Polarity is the result of intermolecular bonds.
1) London Dispersion Force

• experience by all molecules
• results of electrons pushing on each other
• the weakest of all forces
• As the number of electrons increase LDF increases also

2)Dipole-Dipole

• Dipoles are partial separation of charges
• LDF is a type of temporary dipole
• Some molecules have a permanent dipole
• These are polar molecules
• Polarity is determined by electron affinity(how much an atom wants electrons)
• Electron affinity is called electron negativity
• Electron negativity is highest on the top right and lowest on the bottom left

• A bond between two atoms or molecules with different electron negativites result in a dipole-dipole bond.
• Dipole-dipole forces are weak versions of ionic bonds

EXAMPLES:
CHCl3 is polar because: its structural model is not symmetrical so that means it is polar.
C2H2Cl2 is both polar and non-polar depending on how you draw the structural model

Polar and Non Polar Solvents

Polar Substances have an unequal distribution of charges:


(if you were to draw a line down the middle of this diagram it would be symetrical)
Non-polar substances have unequal charge distribution (if symmetrical)


We did a lab that involved polarity and non-polarity. We mixed sugar, salt, and iodine with paint thinner and water and came to the conclusion that polar substances can combine with polar substances and non-polar substances can combine with non-polar substances but polar substances cannot combine with non polar substances.

Electrical conductivity in solutions requires charged ions to be transferred.
Iionic solutions dissociate or form ions so they always conduct.



Molecular substances do not usually dissociate.

follow the chart to determine conductivity:

Examples
element - yes/no - reason
Mercury - yes - metal
Carbon - no - solid non-metal
NaCl - yes - ionic
acetic acid - yes - acid
Iodine - no - not ionic/acid/base/metal

covalent bonds

- electrons shared between non-metals


to draw lewis dot diagrams:
- total all valence electrons in atoms
- identify the element that can form the most bonds. this will be the central atom
- draw bonds between atoms as a line (represents 2 electrons)
- any electrons not part of a bond are lone pairs around the atom
- check to make sure each atom has a full octet

example
ammonia water and ethane


C2H4


heres a extremely educational video!!

Atoms and Ions

• atoms are electrically neutral
• number of protons= number of electrons
• ions ahve different number of protons and electrons
• ions can be either positive of negative
• cation: positive ion
• anion: negative ion

Example: how many electrons do these electrons have? what type of ions are they?


S^2-: anion, sulphide ion I^1- : anion, iodide ion


chemical bonds

• a bond is an electrostatic attraction between particles
• bonds occur as elements try to acheieve noble gas electron configuration
• noble gases (usually) do not form compounds
• in noble gases the outermost energy level have stable octets
• metals lose electrons (oxidize)
• non metals gain electrons (reduced)

Lewis Dot Structure

• atoms can be represented by dot diagrams where dots represent electrons and only the valence level electrons are shown
• write the atomic symbol for the atom. this represents the nucleus and filled inner electron levels
• one dot is used to represent outer energy level electrons. one e- placed in each orbital before and pairing occurs. beginning with the 5th e-, pairing can occur up to a maximum of 8 e-

we practiced drawing lewis dot diagrams

Ionic Bonds

• electrons are transferred from metal to nonmetal. no dots are shown on the metal
• "charged" specie is written in brackets

so thats pretty much all we did today! lots of notes and lots of drawing :P

Atomic Weight

-On your periodic table silver has a weight of 107.9. The atomic mass of silver is 108. There is more than one isotope of silver and some are lighter than others.
a. 15.839% has a Mass Number of 107
b. 48.1161% has a Mass Number of 109
The element isn't split into half of the mass being 107 and 109 but it is intertwined together so it is difficult to separate. Uranium-238 is also difficult to separate.

To find the atomic weight for an element it is listed on your periodic table

24.305 would be the atomic mass for magnesium


54.938 would be the atomic weight for manganese

Emission Spectra, Atomic Structure and Isotopes

Emission Spectra

-Each element gives off a specific colour of light

-these are known as emission spectra(unique to each element)

-if electrons absorb energy they can be bumped to a higher level

-when they fall to a lower lever they release that energy as light

Atomic Structure

-atoms are made up of parts called subatomic particles

-protons(positive), electrons(negative), neutrons(neutral)

-atomic number = # of protons

Isotopes

-the # of protons determine the type of element

-changing the # of neutrons change teh isotope of the element

-all isotopes ahve the same chemical properties

Mass Number

-mass # is total of protons and neutrons

-symbol givven to mass is A

-different isotopes have different masses

-mass number = atomic # + # of neutrons(A=Z+N)

Bohr Model
- atoms are electrically neutral
-2 different models can be used to describe electron configuration : energy level model and bohr model
-electrons occupy shells which are divided into orbitals (orbitals always exist in pairs)
2e in 1st orbital, 8e in 2nd and 3rb orbital (octet)

EXAMPLES

Gold
bohr model:

energy level model:

6e-
2e-
8p+
O


aluminum
bohr model

energy level model:

3e-
8e-
2e-
13p+
Al


Orbital Shapes
atomic orbitals have a specific name and shape
1s / 2s look like this:

2px:


hybridized orbitals:
the 1st of the bohr levels is the 1st orbital and it holds 2e.
- the 2nd level contains 2s, 2px, 2py, 2p² orbitals. they hybridize to form one 2sp³


here is an extremely educational video that will enlighten 13 minutes of your life about orbitals:

Early Atomic Theory

Greeks

• in 300bc Democritus said atoms were invisible particles
• first mention of atoms (atomos)
• not a testable theory, only concept
• no mention of nucleus or sub atomic particles
• cannot explain chemical reactions
• this theory was the most accepted veiw for over 2000 years

Lavoisier (later 1700s)

• law of conservation of mass
• law of definite proportions
• wasn't a true atomic theory because it didnt discuss what atoms were or how they were arranged

Proust (1799)

• if a compound is broken down into its constituents, the products exist in the same ratio as the compound
• experimentally proved Lavoisier Laws

Dalton (early 1800s)

• atoms are solid, indestructable spheres (like billiard balls)
• provides for different elements
• no mention of subatomic particles
• cannot explain isotopes
• no mention of the nucleus

J.J. Thompson (1850s)

• raisin bun model
• solid positive spheres, with negative particles embedded in them
• first atomic theory to have positive and negative charges (protons and neutrons)
• introduces idea of nucleus
• no mention of neutrons so radioactive decay cannot be explained
• does not explain how electrons can exist outside nucleus
• does not explain neutrons role in chemical bonding

Rutherford (1950)

• showed that atoms have a positive centre with electrons outside it
• resulted in planetary models
• explains why electrons spin around the nucleus
• suggests atoms are mostly empty space
• should be unstable
• no mention of neutrons
• does not explain valence level electrons role in chemical bonds

i just had to put these guys in here cuz it is a prime example of how white kids can't rap and how i should never attempt it :P they mention dalton, thompson, rutherford and some other guy i dont know. ill explain Bohr after

Bohr (1920s)

• electrons must onlt exist in specific orbitals around nucleus
• explains how valence electrons are involved in bonding
• explains the difference between ionic and covalent bonding
• resolves the problem of atomic instability
• includes the neutron (discovered 1932)
• explains atomic emission spectra

oh look more white kids doing a terrible job of rapping! hahah at least they get his story across.

k so thats all we did today, a whole bunch of notes and learning about old folks who change chemistry as we know it!

Limiting Reactants

- usually 1 reactant is used up before the other
- this reactant is called the limiting reactant (or reagent) ----> it stops or limits the reaction.
- the limiting reactant determines the amount of products produced.
- assume one reactant is the limiting reactant, and determine what quantity of the other is needed.

example

what is the limiting reactant when 25 g of p₄ reacts with 64.6 g of Cl₂ forming phosphorus trichloride?
P₄ + 6 Cl₂ ---> 4PCl₃
P₄ = 123.6 g/mol
Cl_{2 = 71 g/mol}


Since you need more than the given amount of Cl₂ to fully react the P₄, Cl₂ is the limiting reagant.


Determine the theoretical yield of the above reaction.
PCl<sub>3 = 137.4 g/ mol

when we figure out the theoretical yield, we have to calculate the amount of the product that will be produced from the limiting reagent, because that is how much is being used before the reaction is over.</sub>







83.3 g of PCl3 will be produced. (we used the limiting reagent in the equation because that is the amount that will be used in the reaction.


example.
in the formation of water, 25g of oxygen reacts with 10g of hydrogen. determine the limiting reactant, theoretical yield and amount of excess reactants.

2O2 + 2H2 -----> 2 H2O

H2 = 2 g/ mol
O2 = 32 g/mol

For 10g of hydrogen, we need 80 g of oxygen, so
O2 is the limiting reagent.



H2O = 18 g / mol
THEORETICAL YIELD
28.1 g of water will be produced.


EXCESS - we plug in the calculation so we know exactly how much hydrogen we need, then we subtract this amount from the amount that is given to know how much excess there is.

10g - 3.13g = 6.87 g will be left.

***** Notice that the amounts you need (25 g O2 + 3.13g H2 = 28.1 which is the theoretical yield!!!)



here's a worksheet

and heres a video for further help:

Stoichiometry Lab

In class we tested if stoichiometry could predict precipitate. The materials we used were:
-a lab stand
-a funnel
-a clamp
-filter paper
-beakers
-squirt bottle
To test stoichiometry we used 3 grams of Copper(II) Sulfate and 2 grams of Strontium Nitrate. We dissolve those two in seperate beakers and mixed them together. After stirring for about 7-10 minutes we carefully poured the mixture into the filter. We diluted and continued to dissolve the leftover precipitate. When the mixture was all dissolved we put the filter paper in the dryer to leave over night

Mole to Mass and Other Conversions

Examples
1. Lead(II) Nitrate reacts with 5.0g of potassium iodide. How many grams of Lead(IV) nitrate are required for a complete reaction?

Step 1. Find balanced equation
Pb(NO3)4 + 4 KI-> PbI4 + 4 KNO3

Step 2. Convert and multiply what you need over what you're given
5.0g x 1mol/166g x 1Pb(NO3)4/4KI x 455.2g/1mol Pb(NO3)4 = 3.5g


2. How many grams of O2 are produced from the decomposition of 3.0g of potassium chlorate?

Step 1. Find balanced equation
2 KClO3 -> 2 KCl + 3 O2

Step 2. Convert and multiply what you need over what you're given
3.0g x 1mol/122.6g x 3mol of O2/2mol of KClO3 x 32g/1mol of O2 = 1.2g


3. If a 100mL solution of 2.0M H2SO4 is neutralized by sodium hydroxide. What mass water is produced?

Step 1. Find balanced equation
H2SO4 + 2 NaOH -> 2 HOH + Na2SO4

Step 2. Convert(in this case use concentration formula to find the number of mol) and multiply what you need over what you're given
0.100L x 2.0mol/1L = 0.200mol H2SO4 x 2mol HOH/1mol H2SO4 x 18.0g/1mol = 7.2g

*Remember for next class' lab*
-Theoretical yield of a reaction is the quantity of products expected(calculation)
-The amount produced in an experiment is the actual yield
-The percent yield is : %yield = actual/theoretical x 100

More Stoichiometry

Remember last class when Mr. Doktor introduced us to the wonderful work of stoichiometry?(If you don't we wrote a blog about that one too! haha) Well today we learnt even more!! In adition to moles to mass conversions, we learnt how to solve mass to mole conversions. We also learnt how to solve problems involving volume at STP(22.4L). To help you remember info about volume at STP the chart below could be very handy. You can use this chart to help you by seeing which order to convert your elements and by how much.

Here's a quick example..

What mass of water vapour is produced when 3.5L of Hydrogen is burnt(with oxygen) at STP?

Click here for a few pratice questions!!

And here's a little video of a guy with a cool accent to try to explain it for ya!!

Stoichiometry

so beginning today, we will begin to put together all the things we have learned in the last couple months so that they work together. anything from mole conversions to balancing equations will be involved and it sounds like pretty interesting stuff!

mole to mole calculations

• coefficients in balanced equations respresent the number of moles in each part
• they can also be used as conversion factors

example:

if .15 mol of methane reacts with oxygen, how many moles of water are produced?

first we find the balanced equation:

CH4 + 2O2 -> CO2 + 2H2O

then we find out how many moles of water are needed for the equation by using the formula:

(what you're given) X (what you need)/(what you know)

so... (.15 moles of methane) X (2 moles of water)/(1 mole of methane) = .30 moles of H2O

so by using the quantity we are given, we can find out any other of the factors in the equation

we can also take it one step farther by multiplying the mole ratio we found by the molar mass of the chemical we are trying to find the mass of

for example we can take the .30 moles of water and multiply that by its molar mass (18.02 g/mol) and we find that 5.4 g of water has been used in the equation.

so thats all the notes for today!

Candle Wax Lab

In class we tested how much energy a small candle can release!

The material needed are:
1) candle
2) matches
3) weigh scale
4) lab stand
5) wire mesh
6) thermometer
7) water

Our observations were:
1) volume of water used: 100mL
2) inital temperature of water: 26 degrees C
3) inital mass of candle: 10.1 g
4) final temperature : 42 degrees C
5) final mass of candle: 9.80 g

In this experiment 0.000880 mole of paraffin wax was reacted.
6.704 kJ was the amount of energy gained by the water. The molar heat of combustion of Paraffin wax is 7.6 x 10 to the power of 3 kJ/mol

Calorimetry

To know the amount of energy released three things must be known.
1) Temperature change (▲T) - measured with a thermometer
2) Amount of water (kg) ..... 1 g = 1 mL.... 1kg = 1000mL - measured with a scale
3) Specific heat capacity - how much heat a given substance can hold
water can hold 4.19 kj/kg °C

Equation
▲H= mC▲T
▲H= enthalpy change (kj)
m= mass of water (kj)
C = specific heat capacity (kg/kj °C)
▲T= change in temperature (°C)

Molar Enthalpy - Change in heat for each mole reacted

If you burn 0.315 moles of hexane (C₆H₁₄) in a bomb calorimeter containing 5.65 kg of water, what’s the molar heat of combustion of hexane is the water temperature rises 55.4⁰ C?

▲H = mC▲T

▲H = (5.65 kg)(4.19 kj/kg⁰C)(55.4⁰ C)

▲H = 1311.5 kJ

= 1312 kJ

If you burn 22.0 grams of propane (C₃H₈) in a bomb containing 3.25 kg of water, what’s the molar heat of combustion of propane if the water temperature rises 29.5⁰ C?

▲H = mC▲T

▲H = (3.25 kg)(4.19 kj/kg⁰C)(88.5⁰ C)

▲H = 1205.15 kj

▲H = 1205 kj

here is a quick worksheet to test your skills!

and here is a quick video that shows an experiment involving enthalphy. It is an exothermic reaction



and here is a how to video if youre still confused

Heat in Chemical Reactions

Here are a few things to remember...

• reactions that release heat are exothermic
• reactions that absorb heat are endothermic
• heat IS a form of energy
• all chemicals have stored energy, also known as enthalpy
• enthalpy is given the symbol 'H'
• ΔH is the change in enthalpy

Enthalpy Diagrams

Exothermic

• high to low enthalpy
• ΔH is negative
• heat is released

Endothermic

• low to high enthalpy
• ΔH is positive
• heat is absorbed

wanna learn more about enthalpy change? click here....... you know you waaaaaannaa

--
this is totally off topic but i think Mr. Doktor will find this pretty cool. well at least i think its pretty awesome..its a tv show on discovery or FSN called sport science. Hope you enjoy :)

Balancing Chemical Equations!

There are 6 types of Chemical Reactions!

1 SYNTHESIS: A + B = AB
Al + Cl = AlCl3 The balanced equation is: 2Al + 6Cl = 2AlCl3

2 DECOMPOSITION: AB = A + B
AlBr3 = Al + Br2 The balanced equation is: 2AlBr3 = 2Al + 3Br2

3 SINGLE REPLACEMENT: A + BC= B + AC
Ag(No3) + Zn = Ag + Zn(No3)2 The balanced equation is: 2Ag(No3) + Zn = 2Ag + Zn(No3)2

4 DOUBLE REPLACEMENT: AB + CD = AD + CB
MgCl2 + K2So4 = MgSo4 + KCl The balanced equation is: MgCl2 + K2So4 = 2KCl

5 NEUTRALIZATION: Acid + Base = Water + Ionic Salt
Ba(OH)2 + HCl = BaCl2 + H2O The balanced equation is: Ba(OH)2 + 2HCl = BaCl2 + 2H2O

6 COMBUSTION: Reaction with Oxygen
Mg + O2 = MgO The balanced equation is: 2Mg + O2 = 2MgO

For more practice here is a link to a worksheet!
"http://www.bishops.k12.nf.ca/science/1206/chem/reactiontypes.htm"

Balanced Equations

K so we went over all the stuff about balanced equations that we were talking about before the break. Not many notes for this class, just a lot of examples so ill use some of the notes from last year to supplememnt.

steps to balancing:

• first we must make sure the equation is correct

Mg + O2 - MgO

• balance the equation so there are the same amount of atoms on each side

2Mg + O2 - 2MgO

here are some acids that me are going to have to remember for the acid neutralization reactions:

• HCl: hydrochlroic acid
• HNO3: nitric acid
• H2SO4: sulphuric acid
• H3PO4: phosphoric acid
• CH3COOH: acetic acid

k so this video is pretty long but it gets the point across so s'all good :P

thats pretty much all we did today. ttfn!