Esters

• The functional group:

Ethyl methanoate is one of the simplest esters

Examples:

ethyl pentanoate

butyl propanoate

Esterfication

• Esters are formed by the reaction of a carboxylic acid and an alcohol

Alicycles and Aromatics

-Carbon chains can form 2 types of closed loops
-Alicyclics are loops usually made with single bonds
-If the parent chain is a loop standard naming rules apply with one addition: "cyclo" is added infront of the parent chain

There are 3 different ways to draw organic compounds:

1. Complete structural diagram
2. Condensed structural diagram
3. Line Diagram (mostly used because its way easier to draw and you dont have to include all the hydrogens) :)

*Numbering can start anywhere and go clockwise or counterclockwise on the loop but side chain numbers MUST be the lowest possible!*

Examples:

Amines and Amides

Amines
-are funtion groups that contain a Nitrogen compound bonded to either Hydrogens or Carbons


- Primary amines have 1 carbon chain
- Secondary Amines have 2 carbon chains
- Tertiary amines have 3 carbon chains

Examples:
ethyl methyl amine

trimethyl amine

Amides
- are functional groups with CONH3

-name the alkyl (carbon) chain and add -amide ending
-The simplest amide is ethanamide 



Example
Name the following amides:
Propanamide

Ethers

Ethers

An Ether contains an oxygen group connected to two alkyl (carbon) chain

•  Name the smaller alkyl group first, then the second alkyl group followed by ether

example:

•  draw the chemical formula for dimethyl ether.

Alkenes and Alkynes

Alkenes have double carbon bonds and Alkynes have triple carbon bonds. Multiple bonds form fewer hydrogens are attached to the carbon atoms. The naming rules are almost the same as with alkanes. ( The position of the double/triple bonds always has the lower possible number and is put in front of the parent chain.)

• Double bonds (Alkenes) has the suffix -ene
• Triple bonds (Alkynes) has the suffix -yne

2 Butene

This is called 2 Butene because there is one double bond so the suffix ends in -ene, and there is also 4 carbons therefore the prefix is but-. The number 2 determines which carbon the double bond takes place. (bonds are in between 2 carbons so we take the lowest possible number)

4  methyl 1 pentyne

This is called 2 methyl 1 pentyne because there is one triple bond so the suffix ends in -yne, and there is also 5 carbons (longest chain) therefore the prefix is pent-. The number 1 determines which carbon the triple bond takes place. (bonds are in between 2 carbons so we take the lowest possible number) Since the carbon with the triple bond is at 1, counting to the left, methyl should be at 4. The name would be 4 methyl 1 pentyne.

Multiple bonds

 1,2 propadiene

This has 2 triple bonds so determine where the bonds take place. 1 and 2. prefix is propa for 3 carbons, and suffix in -ene because its a double bond. Since there is 2 double bonds we need a multiplier in between. di. 

Trans and Cis

 If two carbon bonds are bonded by a double bond and have side chains on them, two possible compounds are possible.

Carboxylic Acids

Few rules to naming Carboxylic Acids:

1. Standing naming rules for these compounds
2. Carboxylic Acids are formed by the fuction group 
3. Change the ending to the compound to ac instead of the regular ane, ene, or yne 

Aldehydes

Things to remember for Aldehydes 

1. An aldehyde is a compound that has a double bonded oxygen at the end of a chain 
2. The simplest aldehyde is methanal also called formaldehyde 
3. Follow the standard rules 
4. Change the ending to Al instead of the regular ane,ene, or yne 
5. Be careful when naming Aldehydes do not get them confused with alcohols 

Ketones

Few things to remember when finding out a Ketone

1. There is a double bonded oxygen that is NOT on either end meaning that the double bonded O will not be bonded to the last CH it will bonded to a CH in the middle. 
2. A ketone will have a one ending 

Halides

Few steps to naming a Halide and knowing which compound is a Halide:

1. Same nomenclature 
2. Still use di, tri, tetra and so on. 
3. The endings for each of the Halides will still be the same, such as ane, ene, or yne. Depending on what kind of compound you get. 
4. Cl will not be chlorine but will be changed to chloro
5. Br will not be bromine but will be changed to bromo
6. I will not iodine but will be changed to iodo
7. Prefixes will also be the same as well 

Alcohols

3 important rules in naming Alcohols 

1. Same nomenclature 
2. Ends in ol   instead of ending in ane, ene, or yne
3. An alcohol is a hydrocarbon with a OH bonded to it 

You will know if this is an alcohol if you see that it is bonded to an OH, this is a big hint on naming compounds. 
The following compounds below are methanol and ethanol. 

Organic Chemistry

Organic Chemistry is the study of carbon compounds. Carbon can form multiple covalent bonds such as chains, rings, or bracelets. There are less than 100,000 non-organic compounds while, organic compounds number more than 17,000,000. The simplest organic compounds are made of carbon and hydrogen.
Ex.

CHCCH3

CH3CH3

CH4

Saturated compounds have no double or triple bonds. Compounds with only single bonds are called Alkanes and always end in -ane.

Nomenclature (naming)
There are 3 forms of bonds in Organic Chemistry:
Straight chains, Cyclic chains, and Aromatics

1) Straight chains
To name Straight chains:

• Circle the longest continuous chain and name this as the base chain.

The base chain can either be the red or green from the drawing above because both chains are in equal length.

• Number the base chain so side chains can have the lowest possible numbers

• Name each side chain using the suffix -yl

"Meth" has 1 so theres a methyl, and "eth" has 2  so it's ethyl

• Give each side chain the "appropriate number" ( if there is more than 1 identical side chain numbers/labels are slightly different: 2 methyl 4 ethyl
• List side chains alphabetically

So the Name for this structure is 4 ethyl 2 methyl hexane

Ion Concentration

DISSOCIATION

• ionic compounds are made up of two parts
• Cation: positively charged particles
• Anion: negatively charged particles
• when ionic compounds are dissolved in water, the cation and anion separate from each other
• this process is called dissociation
• when writing dissociation equation, the atoms and charges must balance.
• the dissociation of sodium chloride is:

NaCl -> Na+ + Cl-

• if the volume does not change then the concentration of individual ions depends on the ba;anced coefficient in the dissociation

ex. Determine the [Na+] and [PO4 3-] in a 1.5 M solution of Na3PO4

Na3PO4 -> 3Na+ + PO4 3-

1.5 M = [PO4 3-]

1.5 x 3 = 4.5 M = [Na+] 

Dilution

DILUTION SOLUTIONS

• when two solutions are mixed, the concentration changes
• dilution is the process of decreasing the concentration by adding a solvent (usually water)
• the amount of solute does not change

• because concentration is mol/L, we can write:

C = n/V                C1V1 = C2V2

n = CV

ex. determine the concentration when 100 ml of 0.10 M HCl is diluted to a final volume of 400 mL.

V1 = 100mL

C1 = 0.10 M

V2 = 400 mL

C2 = ?

C1V1 = C2V2

(0.1)(100) = C2(400)

C2 = 0.025 M

ex. how much water must be added to 10.0 mL of 10.0 M Na2SO4 to give a solution with a concentration of 0.50 M?

V1 = 10.0 mL

C1 = 0.1 M

V2 = ?

C2 = 0.50 M

C1V1 = C2V2                                  ΔV = 200 - 10 = +190 mL

(10.0)(10.0) = (0.5)V2

V2 = 200 mL

Bonds and Electronegativity

There are 3 types of bonds:

Ionic.
Covalent.
Metallic.

Ionic bonds are when electrons are transfered from metal to non-metal. Covalent bonds are when electrons are transfered between non-metals. Metallic bonds are when pure metals are held together by electronegativity attraction.

electronegativity - (EN) is a measure of an atoms attractions for electrons in a bond

1. Atoms with a greater EN can attract more e- 
2. polar convalent bonds form from an unequal sharing
3. non-polar covalent bonds form from an equal sharing

Electronegativity has a specific rule for each elements.

• EN > 1.7 = ionic bond
• EN < 1.7 = polar covalent bond
• EN = 0 = non-polar convalent bond

Example:

1. Ba-I
   0.89-2.66
   1.77 = ionic bond
2. Co-P
   1.88-2.19
   0.31 = polar covalent bond
3. Hg-Po
   2.0-2.0
   0 = non-polar convalent bon

 

Inter-molecular Bonds

-Inter-molecular bonds exist within a molecule such as the ionic and covalent bonds
-Inter-molecular bonds exist between molecules

• If the inter-molecular is stronger it gives you a higher a BP or MP 
• There are two types of inter-molecular bonds such as Van Der Waals bonds and Hydrogen Bonds
• They are based on electron distribution 
• And two categories 

1. Dipole- Dipole Bonds
-If a molecule is POLAR the positive end of one molecule will be attracted to the negative end of another molecule (One must be negative and one must be positive)

2. London Dispersion Forces (LDF) (Count the # of electrons)
-LDF is present in all molecules
-they create the weakest bonds
-If any substance is NON-POLAR you do not need to check if the substance is Dipole Dipole because it will not exist.
-electrons are free to move around and will randomly be grouped on one side of the molecule
-this can create a temporary dipole and can cause a weak bond to form
-the more electrons in the molecule the stronger the LDF can be

3. Hydrogen Bonding
-If hydrogen is bonded to a certain element such as Flourine, Oxygen or Nitrogen the bond is highly polar
-This forms a very strong intermolecular bond

Solution Stoichiometry (cont)

Ex. 100 ml of 0.250 M of Iron (II) Chlorine reacts with excess copper. How many grams of Iron are produced?

FeCl2 + Cu -> Fe + Cu Cl2

0.25 M x 0.1L x 1/1 x 55.8g = 1.40 g

• How many moles of Copper Chloride are produced?

0.25 M x 0.1L x 1/1 = 0.0250 mol of copper chloride

• Determine [CuCl2]

0.025 mol x 1/ 0.1 = 0.25 M of CuCl2

Ex. A beaker contains 100 ml of 1.5M HCl. Excess Zinc is added to the beaker. Determine how many leaders of hydrogen gas should be produced.

2 HCl + Zn -> ZnCl2 = H2

1.5M x 0.1L x 1/2 x 22.4L/mol = 1.7L of H2

• If 1.40L Hydrogen gas are actually produced what is the present yield?

1.4/1.7 x 100 = 82%

Ex. A 0.330 M solution of Ba(OH)2 reacts with 25 ml of 2.50M HCl. What volume of Ba(OH)2 is required for a complete reaction.

Ba(OH)2 + 2 HCl -> BaCl2 + 2 H20

2.50 M x 0.025 x 1/2 = 0.0313 mol of Ba(OH)2

 0.330M x 0.0313 mol = 0.095L or 95 ml

Ex. A 3.00g of piece of Iron is added to a beaker containing 100 ml of 0.750 M AgNO3, Determine the L.R.

Fe+ 3 AgNO3 - > Fe(NO3)3 + 3 Ag

0.750 M x 0.1 L x1/3 x 55.8 = 1.40g of Iron       Ag is L.R.

• How many grams of Ag are produced?

0.750 M x 0.1L x 3/3 x 107.9 = 8.09g

example of Titration

- In limiting reactions, usually one chemical gets used up before the other.

• The chemical used up first is called limiting reactant
• Once it is used up it stops
• LR determines the quantity of products formed

- To find the LR, assume one reactant is used up. determine how much of this reactant is required.

Other Conversions (Volume and Heat)

• Volume @ STP can be found using the conversion factor 22.4 L/mol.
• Heat can be included as a seperate term in chemical reactions (This is called Enthalpy)
• Rxns that release heat are exothermic.
• Rxns that absorb heat are endothermic.
• Both can be used in stoichiometry.

Mole to Mole

• Mass to Mass Problems involve one additional conversion
• Start from mass A convert to moles A then convert to Moles B to convert to Mass B
• EX. Lead IV Nitrate reacts with 5.0g of Potassium iodide. How many grams of Lead (IV) nitrate are required for a complete reaction?

Pb(NO3)4 + 4KI -> PbI4 + 4 KNO3

5.0g (mass A)   x   mol/166g (Mole A)  x 1/4 (Mole B)  x 455.2/mol (Mass B) = 3.4g

• How many grams of O2 are produced from the decomposition of 3.0g of Potassium Chlorate?

2 KClO3 -> 2 K + Cl2 + 3 O2

3.0g x mol/290.7g x 3/2 x 32/mol = 1.0g

• When Solid Zinc reacts with hydrochloric acid what mass of hydrogen gas is produced when 2.5g of Zinc react?

Zn + 2 HCl -> ZnCl2  + H2

2.5g x mol/65.4 x 1/1 x 2/mol = 0.076g of H2

• If 100g of Octane are burnt in a car engine what mass of oxygen is needed?

2 C8H18 + 25 O2 -> 16 CO2 + 18 H2O

100g x mol/114 x 25/1 x 32/mol = 352 g 

Moles to Mass & Mass to Moles

Map

Mass A to Mol A to Mol B to Mass B

• Some questions will give you the amount of moles and ask to determine the mass
• converting moles to mass only requires on additional step.
• How many grams of Bauxite (Al2 O3) are required to produce 3.5 mol of pure Aluminum?

2 Al2O3 --> 4 Al + 3 O2

3.5 mol x 2/4 x 102/mol = 178.5 = 1.8 x 10^2 g of Al2O3

• How many grams of water are produced of 0.84 mol of Phosphoric Acid is completely neutralized by Barium Hydroxide?

2 H3PO4 + 3 Ba(OH)2 -> 6 HOH + Ba3(PO4)2

0.84 mol x 6/2 x 18/mol = 45.56 = 45 g of H2O

• How many gram of Silver nitrate are needed to produce 1.02 mol of silver chloride according to the reaction.

2 AgNO3 + BaCl2 -> Ba(NO3)2 + 2 AgCl

1.02 mol x 2/2 x 169.9/mol = 173 g

• How many moles of Lead(II) nitrate are consumed when 4.5 g of sodium sulfide completely reacts

Na2S + Pb(NO3)2 -> 2 NaNO3 + PbS

4.5 g x mol/78.1 x 1/1 = 0.058 mols of Pb(NO3)2

Mole to Mole Conversions

• Coefficients in balanced equations tell us the number of moles reacted or produced.
• They can be used as conversion factor

               _X + _Y = _Z

*** WHAT YOU NEED over WHAT YOU HAVE





ex. If 0.15 mol of methane are consumed in a combustion reaction, how many moles of CO2  are produced?

          1CH4 + 2O2 à 1CO2 + 2H2O

          

          0.15mol x 1/1 = 0.15 mol of CO2

ex. How many moles of bauxite (Aluminum Oxide) are required to produce 1.8 mol of pure Aluminum?

          2Al2O3 à 4Al + 3O2

          1.8mol x 2/4 = 0.9mol off Al2O3

ex. When 1.5 mol of copper react with iron(II) chloride. How many moles of iron should be produced?

          1Cu + 1FeCl2 à 1Fe + 1CuCl2

          1.5mol x 1/1 = 1.5mol of Fe

Stoichiometry (Quantitative Chemistry)

There is Qualitative Chemistry and Quantitative Chemistry
Under Quantitative Chemistry:

• Stoichiometry is a branch of chemistry that deals with the quantitative analysis chemical reactions
• It's a generalization of mole conversions to chemical reactions
• Understanding the 6 types of Chemical reactions is the foundation of stoichiometry

6 Types of Reactions

1. synthesis (formation)
2. Decompostion
3. Single Replacement (SR)
4. Double Replacement (DR)
5. Neutralization
6. Combustion

Synthesis

A+B>AB

• usually elements> compounds
• ex. K + O2 > 2K2O

Decomposition

AB>A+B

• Reverse of synthesis
• Always assume the compounds decompose into elements during decomposition
• Ex. 4H3PO4>6H2+P4 +8O2

Single Replacement (SR)

A+BC>B+AC

• ex. Ca+2KCL>2K+CaCl2

Double Replacement

AB+CD>AD+BC

• Ex. MgCl2+K2(SO4)>Mg(SO4)+2KCl

Neutralization (All DR)

Reaction between acid and base

• Ex. H2SO4+2KOH>2HOH+ K2SO4

Combustion

• Reaction of something (usually hydrocarbon) with air 
• Hydro carbon combustion always produces CO2 and H2O
• Ex. CH4+2O2>CO2+ 2H2O
• Always balance in the order C, H, then O

http://www.youtube.com/watch?v=XS6JTr-mTWY

Molecular Formulas

• If you know an empirical formula to find the molecular formula, you need the molar mass.

ex. The empirical formula for a substance is CH2O and its molar mass is 60.0 g/mol. Determine the molecular formula.

Empirical Formulas

• Empirical Formulas are the simplest formula of a compound
• Show only the simplest ratios, not the actual atom
• Example: the empirical formula for Chlorine gas is Cl
• Molecular formulas give the actual number of atoms

• To determine the empirical formula, we need to know the ratio of each atom
• To determine the ratio, fill in the table below:

ex, A sample of an unknown compound is found to contain 8.4 g of C, 2.1 g of H, and 5.6 g of O. Determine the empirical formula.

• The simplest ratio may be decimals. For certain decimals, you need to multiply everything by a common number.