Functional Groups

Halides:

Halides are the elements in group 17 of the periodic table, and the naming of them as side chains in a compound. To name halides, take the first syllable of the element, for example bromine, and add -o ending: bromo. Others that we use in chemistry class are Chlorine: Chloro, and Fluorine: Floro.

Alcohols:

If a carbon is connected to an OH (Hydroxyl), it is an alcohol. To name alcohols, you must circle the longest carbon chain containing the carbon connected to the hydroxyl, and appropriately label the chain with the proper prefix, ending with an -ol.

Ethers:

An ether is the oxygen atom that connects two carbon chains together. To name ethers, you must label each carbon chain as a side chain, with the right prefixes and ending -yl, then simply putting ether. For example, ethyl methyl ether.

Ketones:

A carbon chain with a carbon double-bonded oxygen is a ketone. To name ketones, you must cirlce the longest carbon chain and name with the right prefix, and ending with -one.

Functional

Alkenes and Alkynes

Alkenes are organic compounds with double-bonded carbons. For naming alkenes, same rules apply with Alkanes, except the ending is -ene.

Alkynes are organic compounds with a triple-bonded carbon. Same rules also apply and the ending is -yne.

Alkanes

Naming Alkanes is easy:

1. Circle the longest carbon chain
2. Name the chain with the appropriate prefix
3. -ane ending
4. Locate any side chains
5. Label side chains with appropriate prefixes
6. give side chains -yl ending

Intro to Organic Chemistry

It is the study of carbon atoms and its compounds. Carbon compounds outnumber all other compounds combined > carbon compounds can form many different bonds > bonds can have different arrangements > carbon can form long chains.

types of bonds:
> alkanes - single bonds
> alkenes - double bonds
> alkynes - triple bonds

There are 3 types of chemical formulas
1. molecular formular : C7H16
2. condensed structural formula: CH3 - CH2 - CH2 - CH2 - CH2 - CH3
3. structural formula:




-- Jael Lumba

Molecular Polarity and Intermolecular Bonds

Polarity is the result of intermolecular bonds. There are 3 kinds of intermolecular bonds: lond dispersion force (LDF), dipole-dipole, and hydrogen bond.

1. London Dispersion Force (LDF)
this bond is experienced by all molecules. in result of electrons pushing on each other. It is the weakest of all forces. **as the number of electrons increase, the LDF increases as well.

2. Dipole-dipole
It is the partial separation of charges. LDF is a type of temporary dipole-dipole. Molecules who undergo dipole force is polar. Polarity is determined by electron affinity (how much an atom wants an electron). Electronegatively is electron affinity.

>Polar have unequal distribution of charges distribution
>Non-polar molecules have equal charge distribution



3. Hydrogen Bonding
It is a special type of dipole-dipole bond and occurs between hydrogen and nitrogen, oxygen, or fluorine.

H - N , H - O , H - F



-- Jael Lumba

Solution Chemistry

A solution is a homogeneous mixture. Solvents are components present in larger amounts Solutes are componets present in smaller amounts. A solute is soluble in a solvent if it dissolves to form a homogenous mixture. A saturated solution contains as much solute as possible. An unsaturated solution can dissolve more solute. Solubility is the measure of how much solute can dissolve in a given solution (g/L, g/ml, mol/L, ppm)

Solubility is affected by heat, changing the solvent, and changing the solute.

To determine conductivity:

> if it is a metal, it is conductive, if not > if it is a solid, it is not conductive, if not > if it is an base.. > is it ionic, it is conductive, if not it isn't conductive.

-- Jael Lumba

Chemical Bonds

A bond is an electrostatic attraction between particles. They occur as elements try to achieve noble gas electron configuration. Keep in mind that noble gases have full valence electrons and do not form compounds because of their stability. Metals lose electrons, so they oxidize. Non-metals gain electrons, so they reduce.

There are two types of bonds: ionic and covalent. Ionic occurs when a metal losses electrons to a non-metal to achieve noble gas electron configuration. An example is potassium chloride, NaCl. Covalent bonding occurs between non-metals that share electrons to achieve stability. An example is water, H2O.

-- Jael Lumba

Atoms & Ions

Atoms are electrically neutral ( # of protons = # of neutrons). Where as ions have different number of protons and electrons. Ions can be either positive (loss of an electron) or negative (gain of an electrons). Cation is a positive ion and an anion is a negative ion.

ex: Bromine (Br-) has a charge of - , therefore, it gains 1 electron making it an anion.
Phosphorus (P3-) has a charge of 3-, therefore, it gains 3 electrons making it an anion.
Magnesium (Mg2+) has a charge of 2+, therefore, it losses 2 electrons making it a cation.
Potassium (K+) has a charge of +, therefore, it losses 1 electron making it a cation.




-- Jael Lumba

Atomic Structure

An atom is made up of subatomic particles: protons (+), electrons (-), and neutrons (n).
- Protons: they are positive and located in the nucleus
- Neutrons: they are neutral and located in the nucleus
- Electrons: they are negative and are located outside the nucleus.



The atomic number is the number of protons in an atom.

Isotopes are different types of atoms of the same chemical element, each having different number of neutrons. Therefore, changing the number of neutrons changes the isotope of the element. All isotopes have the same chemical properties.

The mass number is the total of protons and neutrons. The symbol given is A. Different isotopes also have different masses.

Keep in mind that mass number equals the total number of protons and neutrons

A = atomic # + # of neutrons

-- Jael Lumba

Emission Spectra

Emission spectra is the specific colour of light that each element gives. Each colour is unique to each element. If electrons absorb energy, they can be bumped to a higher level, but when the fall to a lower level, they release that energy as light.

The picture to the left is a spectra of the elements of lithium, sodium, potassium, rubidium, cesium, mercury and neon.

-- Jael Lumba

Orbital Shapes

Atomic orbitals each have a specific name and shape.

- 1s:
- 2px:


Hybridized Orbitals:
The first of the Bohr levels is the 1s orbital and holds two electrons. The second level contains the 2s, 2px, 2py 2pz orbitals. They hybridize (combine) to form one 2sp3 orbital:



-- Jael Lumba

Energy Level and Bohr Models

We learned that atoms are electrically neutral. To model the atom we can use either energy level models or Bohr models. Electrons occupy shells which are divided in orbitals: 2 electron ins the first orbital, 8 in the second, 8 in the third, etc..

Here is a Bohr Model (left) and an Energy Level Model (right) of Calcium:


Bohr Models gives a picture all of the atom's valence shells and its electrons along with the number of protons and neutrons present in the nucleus. Energy Model diagrams show all of the atom's valence shells and its electrons along with the atomic mass (40) and number of protons (20).


--- Jael Lumba

The Wonders of the ATOMIC THEORY!

The developement of the early Atomic Theory first began with the Greeks in 300 BC with a man called Democritus. He said that atoms are indivisible particles. This is the first mention of atomos, also known as atoms! Democritus theory was not testable, but only a conceptual model. It had no mention of any atomic nucleus of its constituents. It could not be used to explain chemical reactions. Thus, this theory was accepted for about 3000 years.

Lavoisier: ( late 1700s)

This man came up with the law of conservation of mass and law of definite proportions. It is said his concepts was not considered to be an Atomic Theory because it didn't discuss what atoms were or how they were arranged.

Proust: (1799)

Proust said if a compound is broken down into its constituents, the products exist in the same ratio as in the compound. This theory experimentally proved Lavoisier's Law.

Dalton: (early 1800s)

Dalton considered atoms as solid indestructible spheres, like billiards. The problem with Dalton's theory is that it didn't mention subatomic particles, explain isotopes, and had no mention of the nucleus.

J.J Thomson: (1850s)

This man came up with the Raison Bun Model. It showed solid, positive spheres with negative particles embedded in them. It was the first atomic theory to have positive and negative charges! It introduced the idea of the nucleus. The problem with this theory is that it no mention of neutrons, therefore, radioactive decay cannot be explained. it also doesn't explain how electrons exist outside the nucleus and didn't explain electrons role in chemical bonding.

Rutherford: (1905)

Rutherford showed that atoms have a positive dense center with electrons outside it. this resulted in the Planetary Model which explains why electrons spin around the nucleus. it suggest atoms are mostly empty space. But, it showed instability because the negative and positive should attract and destroy the atom. Thus, it had not mention of neutrons and did not explain valence level electons role in chemical bonding.

Bohr: (1920s)

Bohr said the electrons must only exist in specific orbitals around the nucleus. It explains how valence electrons are involved in bonding and the difference between ionic and covalent bonding. It resolved the problem of atomic instability and the atomic spectra. It also includes the neutron

-- Jael Lumba

Limiting Reactants (click the title)

Usually, one reactant is used up before the other> this reactant is called the LIMITING REAGANT in other words, it stops or limits the reaction. The limiting reagant determines the amount of product produced. Assume one reactant is the limiting reactant and determine what quantity of the other is needed.



EXAMPLES:

1. What is the limiting reactant when 125g of P4 reacts with 325 g of Cl2 forming phosphorus trichloride?
   P4+6Cl2=4PCl3 (125g=p4/ 323g=6Cl2)
   
   125 g X1 mol/71 gX 4 mol/ 6 mol x 137.4g/ 1mol= 417 g PCl3
   

-- Michelle Haughian

Percent Yield

The theoretical yield of a reaction is the quantity of products expected (calculation). The amount produced in an experiment is the actual yield. To find the percent yield follow this calculation:

% yield= actual / theoretical x 100

Here is a short video to further explain:

http://www.youtube.com/watch?v=mOo3KuqM64c


Examples:

1. The production of Urea CO(NH2)2 is given by: 2NH3 + CO2 ---> CO(NH2)2 + H2O
If 47.7g of Urea are produced, determine

a.) theoretical yield if 1mol of CO2 reacts
1mol CO2 x 1 CO( NH2)2 \ 1 CO2 x 60.1 g= 60.1 g

b.) actual yield
47.7 g

c.) percent yield
47.7g / 60.1 g x 100= 80%

-- Krizia Umali

Stiochiometry

In balanced equations, the coefficients represents the number of moles. To find the number of moles of a compound or element in a balanced equation, we divide the given mole by the coefficient.

Example - : If 0.20 mol of methane reacts with oxygen, how many moles of each product are produced?

- first write a balanced equation, be sure that the equation is correctly balanced:
2 CH4 + 4O2 -> 2 CO2 + 4 H2O
- to find the number of moles of CO2: multiply the number of moles of methane by CO2's coefficient which is 2, and divide by the number of moles of methane.
(0.20 mol CH4 x 2 mol CO2) / (2 mol CH4) = 0.4 mol CO2

When converting moles to mass, one more step is needed (g/mol)

Example - : How many grams of water are produced if 1.0 mol of phosphoric acid is neutralized by barium hydroxide?

- first write a balanced equation:
2 H3PO4 + 3 Ba(OH)2 -> Ba(PO4) + 6 H2O
- to find the number of moles of H2O: multiply the number of moles of H3PO4 (phosphoric acid) by H2O's coefficient which is 2, and divide the number of moles of H3PO4
(1.0 mol H3PO4 x 6 H2O) / (2 mol H3PO4) = 3.0 mol H2O
- now multiply H2O's molar mass and divide by mol (g/mol)
(3.0 mol x 18 g) / (1 mol) = 54 g H20

Here's a video of Chemguy solving many moles of silver chloride forms when 2.6 mol of KCl reacts with excess silver nitrate in solution:


That is all. Mr. Doktor is the best!

-- Jael Lumba

Classifying Chemical Reactions

There are four main types of chemical reactions. They are direction combination, decomposition, single-replacement, and double-replacement.

Direct Combination Reactions
A + B -> AB
EXAMPLE: 2 Na + Cl2 -> 2 NaCl

- Combustion reactions are direct combination when the products are carbon dioxide and water:
C2H5OH + 3 O2 -> 2 CO2 + 3 H2O

Decomposition Reactions
AB -> A + B
EXAMPLE: 2 H20 -> 2 H2 + O2

Single-Replacement Reactions
A + BX -> AX + B
EXAMPLE: Mg + CuSO4 -> Mg SO4 + Cu

Double-Replacement Reactions
AX + BY -> AY + BX
EXAMPLE: CaCO2 + 2 HCl -> CaCl2 + H2CO3

That is all. Mr. Doktor is the best!

-- done by Jael Lumba

Concentration, number 9!

We learned that a solution is a homogeneous mixture. In a solution, there is a solute and solvent. A solute is the component present in smaller amount, a solvent is the component present in a smaller amount. Concentration (con'c) is amount of solute / amount of solvent.

The possible units for concentration are g/mL, g/L, mg/L, % by mass, % by volume, etc..
- mol / L = molarity

Concentration is found by using C = n / v
- C = concentration
- n = number of moles
- v = volume

- Example -: Bob dissolves 38.0 g of NaOH in enough water to make 200.0 mL of solution. Find the molarity.
- [NaOH] = ? - first find the molar mass which is 40.0g
- find the amount of moles by dividing the mass of NaOH by its molar mass: 40.0g / 38.0g = 1.05 mol
- find the molarity in mol/L: (1.05 mol x 1000mL) / (200 mL x 1 L) = 5.25 mol / L

That is all. Mr. Doktor is the best!

-- done by Jael Lumba