May 10

Cyclo Alkanes

- Carbon compounds can form rings
- Follow the same naming rules and add cyclo- in front of the parent chain
- When numbering, you can count clockwise or counterclockwise, but where you start (#1) should be on one of the side chains because you are using the lowest numbering system

Examples of Cyclo Alkanes:

Aromatics

- When a cyclic 6 carbon chain forms, it can create a resonance structure called Benzene

Organic chem

Alkenes
- Compounds with double bonds end in -ene

- Put a # in front of the parent chain that indicates where the double bond is

- More thn one double bond changes the parent chain slightly

- Double bond always has priority (when choosing the direction of numbering)

- Any rime more than one double bond occurs, you add adi, atri, atetra...

- The longest chain has to include the double bond




Alkynes
- For compounds with triple bonds use -yne ending

- Follow all the same alkene rules

- The longest chain has to include the triple bond




ORGANIC CHEMISTRY
- There are more carbon compounds than all ionic compounds combined
- The study of carbon compounds is called organic chemistry
- Carbon can have multiple bonds and form many different shapes
- Hydrocarbons have three types of formulas:

1) Molecular formulas
C6H14

2) Condensed Structural Formula
CH3-CH2-CH2-CH2-CH2-CH3

3) Structural Formula



Nomenclature of Hydrocarbons
- One molecular formula can have a number of different structures
- Isomers are compounds that can be drawn in more than one way


Naming Alkanes
1) Name the longest chain by using the correct suffix and adding "ane"
2) Locate any branches by number carbon atoms (use the lowest possible number system)
3) Name branches by using appropriate suffix and -yl ending (Alkyl branches)
4) If there are more than one of the same alkyl group, number each one and add the multiplier number in front of the branch name

April 20, 2010

Solvents and solutes are either Polar or Non Polar

Polar substances have an unequal charge distribution (asymmetrical)

H2O = polar:

CH4 = Nonpolar

LIKE DISSOLVES LIKE!

Polar solutes dissolve in polar substances and non-polar solutes dissolved in non-polar substances

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INTERMOLECULAR BONDS

- bonds between molecules
- 3 types

1.) London Dispersion Force (L.D.F)

- Results from temporary electron dipoles
- Weakest intermolecular force
- Increases as the $ e- increases
- Occurs in any compound that has e- (ie: everything)

2.) Dipole Dipole

- Results from a permanent dipole in molecules
Polar molecules experience this force
Polarity depends how much elements want e- (electronegativity)
- Electronegativity increases to the right and up
- The strength of a dipole- dipole bond depends on the difference in electronegativity between the two atoms
- Only polar molecules experience this

Substance Boiling Point # of e-
N2 -196 degrees C 14
O2 -183 degrees C 16
NO - 152 degrees C 15
ICl 97 degrees C 70
Br2 59 degrees C 70

More electrons = higher the boiling point

3.) Hydrogen Bonds (H-Bonds)
-This is a special type of dipole- dipole bond between H and O, F, or N
- Any molecule that: H-F, H-O or H-N

Identify the substances with H-Bonds:
1) CH4
2) CH3OH
3) H2S
4) CH3-NH2
5) HCl
6) CH2-OH-OH2
/ / /
OH OH OH

Answer: Number 2, 4, and 6

Compare the boiling points of:
- Ethanol (C2H5OH)
- Ehtane (C2H6)
- Methanol (CH3OH)
- Methane (CH4)

> London Forces are the weakest intermolecular force and hydrogen bonds are the strongest.

> More electrons, the higher the boiling point.

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IONS IN SOLUTIONS
- The formation of a solution depends on the ability of the solute to dissolve in the solvent
- Solvation is the interaction between solutes and solvents
- Ionic solids (salts) are cyrstals made up of ions
- Molecular solids are crystals made up of neutral molecules
- Dissolving ionic solutions produces ions in a process called disssociation (remember?)

Ionizaiton is the break up of a neutral molecule into charged particles
Examples:
1) FeCl3 (s) ----->Fe 3+ (aq) + 3 Cl -1 (aq)
2) Ag2O (s) -----> 2 Ag + (aq) = O 2- (aq)

3) Na3PO4(s) -----> 3 Na + (aq) + PO4 (aq)

4) (NH4)2SO4 (s) -----> 2NH4 + (aq) + SO4 2- (aq)

Determining concentrations is relatively easy.

Examples:

What is the [Cl-] in a solution of 0.50 M AgCl3?

AgCl3 -----> Ag + (aq) + 3Cl- (aq)
Cl = 3x as many moles = (0.5 M) x 3 = 1.5 M

What is the [NO3-] in a solution of 0.82 M Fe(NO3)2?
Fe(NO3)2 -----> Fe 2+ (aq) + 2 NO 3- (aq)
(0.82 M) x 2 = 1.64 M

What is the [Cr2O7 2-] and [K+] when 3.5 g of K2Cr2O7 dissolved in 40 mL of water?
K2CrO7 -----> CrO7 2- (aq) + 2K + (aq)
3.5 g x 1 mol/294.2 g = 0.0

Introduction to 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)

- The solubility of Ba (NO3)2in water is 63 g/100 mL @ 25 degrees Celcius while the solubilty of Ba(NO3)2 in alcohol is is 1.6 g/ 100 mL @ 25 degrees Celcius

Solubility is affected by:
1. heat
2. changing the solvent
3. changing the solute

- Measuring the conductivity of a solution

APRIL 14 CLASS NOTES

- Distilled water is non-conductive

- By adding salts, we increase the conductivity

- Electrical conduction in solutions requires charged ions to be present.
- Ionic solutions dissociate (break apart) when placed in water. Molecular solutions do not usually split into ions

Ex: Dissociation of sodium chloride:

NaCl --> Na + Cl

Steps to determine conductivity:

Follow these steps to determine if the solution is conductive:

Ask...

Is it a metal?
Yes = conductive. No ask...
Is it a solid non-metal?
Yes = non-conductive. No ask...
Is it an acid or base?
Yes = conductive. If noask...
Is it ionic?
Yes = conductive. no = non-conductive.

A chemical demonstration showing that ions must be present in solution for electrical conductivity

Electronegativity [short lesson]

- Atoms affinity for electrons
- Electronegativity increase from left to right and from bottom to top

http://www.chemguide.co.uk/atoms/bonding/electroneg.html

- Polaritiy is the separation of charge inside something that is neutral

Covalent bonding

Covalent Bonding
- Electrons are shared between non-metals
- Drawing the Lewis Dot Diagrams:

1. Total all valence electrons in all atoms
2. Identify the element that can form the most bonds. This will be the central atom.
   
3. Draw bonds between atoms as a line. This represents 2 e-
4. Any e- not part of a bond are lone pairs around the atom
5. Check to make sure each atom has a full octet

Double & Triple Bonds
- Some compounds form more than one bond between two elements
- Some compounds form more than one bond between two elements

Okay, YEESh.
Timmme for a much needed break from chem with some soothing tunes from the Beatles!

Atoms & Ions

- Atoms are electrically neutral
- # of protons = # of electrons
- Ions have different # of protons and electrons
- Ions can be either positive (lose e) or negative (gain e)
- cation = positive ion
- anion = negative ion

Example:
C4-
Carbon gains 4 electrons and it is an anion.

CHEMICAL BONDS
- A bond is an electrostatic attraction between particles.
- Bonds occur as elements try to achieve Noble gas electron configuration
- Noble gases (usually) don't form compounds or bonds
- In Noble gases, the outermost energy level have stable octects.
- Metals lose electrons (oxidize)
- Non-metals gain electrons (reduced)

LEWIS DOT STRUCTURE
- Atoms can be represented by dot diagrams
- Dots represent electrons
- Only valeance level electrons are shown
- The atomic symbol for the atom represents the nucleus and filled inner electron levels
- One dot is used to represent outer energy level electrons

** Some helpful info:
http://www.ausetute.com.au/lewisstr.html



IONIC BONDS
- Electrons are transferred from metal to non-metal
- No dots are shown on metal
- 'Charged' species is written in brackets

Na + Cl → Na⁺ + Cl⁻ → NaCl

Group projects / presentation topics

• Mendeleev's Periodic Table
• Metals
• Non-metals
• Metalloids
• Trends n Physical Properties of Elements on the Periodic Table
• Trends in chemical properties of elements on the periodic table (ion charge, chemical reactivity, ionization energy)
• Properties of Alkalis, Alkaline Earth Metals, Halogens, Noble Gases and Transition Metals

Atomic Weight [ jojojoojojoanneanneanne]

Example: Silver



The atomic weight for silver is 107.9. in the last example we saw the atomic number of silver was 108. This is because there are more than 1 isotope of silver and some are lighter than others

Atomic Structure / Isotopes [Joanne]

Emission Spectra
- Each element guves 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 level they release that energy as light

Atomic Structure
- Atoms are made up of parts called subatomic particles
- Protons (positive), Neutrons (neutral), and Electrons (negative)

Atomic Number
= number of protons

A = atomic number
B = Ion charge
C = Symbol

D = Element name
E = Atomic Mass

Isotopes
- The number of protons determine the type of element
- Changing the number of neutrons changes the isotops of the elements
- All isotopes have the same chemical properties

Mass Number

- Mass number is the total of protons and neutrons

- Symbol given is A

- Different isotopes have different masses

- Mass number = atomic number + number of neutrons

A = Z + N

Ex. Boron:
Isotope - 11
Mass - 11
Atomic # - 5
Protons - 5p
Neutrons - 6n

Ex. Nickel
Isotope - 59
Mass - 59
Atomic # - 28
Protons - 28p
Neutrons - 31n

Bohr Models [ Joanne ]

number of electrons = number of protons in a neutral atom
number of neutrons = atomic mass - the number of protons.
***

- Atoms are electrically neutral
- Two different models can be used to describe the electron configuartion:
1.) Energy Level Model
2.) Bohr Model

BOHR MODEL
Electrons occupy shells which are divided into orbitals
2e in the first Orbital
8e in the second orbital
8e in the third orbital
(the ones with the 8e are called Octet)

Ex. Aragon (Energy level Model)
8e- 3rd level
8e- 2nd level
2e- 1st level

40 <---atomic mass 18 Ar <---- atomic number / number of protons

Bohr model of Argon

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Ex. Chlorine (Energy Level Model)

7e- 3rd level
8e- 2nd level
2e- 1st level

35
1 Cl

Chlorine (Bohr Model)

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Ex. Fluorine (Energy Level Model)

7e- 2nd level
2e- 1st level

19
9 F

Flourine (Bohr Model)

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_____________________________

Hybridized orbitals
- The first of the bohr levels is the 1st orbital and it holds only 2e
the second level contains the 2s, 2px, 2py, 2pz orbitals. they combine (hybridize) to form one 2sp3 orbital

***



"OH MY GOSH,OH MY GOSH,
OHHHH MY GOSH !!!!!!"

Atomic Theory [Joanne]

Early Atomic Theory

-GREEK

- In 300 B.C., Democritus said atoms were invisible particles,

- First mention of atoms (atomes)

- Not a testable theory, only a conceptual model

- No mention of any atomic nucleus or its constituents

- Cannot be used to explain chemical reactions

- This theory was the most accepted view for over 2000 years


____________________________________________

Lavoisier (late 1700s)

- Law of Conservation of Mass

- Law of definite proportions

- Wasn't a true atomic theory because it didn't discuss what atoms were arranged

____________________________________________

Proust (1799)

- If a compound is broken down into its constituents, the products exist in the same ratio as in the compound

- Experimentally proved Lavoisier Laws

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Dalton (early 1800s)

- Atoms are solid, indestructible spheres (like Billard balls)

- Provides for different elements (these would be different spheres)

- Doesn't metntion subatomic particles

- Can't explain isotopes
- Didn't mention the nucleus

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J.J. Thomson (1850s)
- Raison Bun Model
- Solid, positive spheres, with negative particles embedded in them
- First atomic theory to have positive (protons) and negative (electrons) charges
- Introduces idea of nucleus
- Didn't mention neutrons, so radioactive decay could'nt be explained
- Doesn't explain how electrons can exist outside nucleus
- Doesn't explain electrons role in chemical bonding








_________________________________________

Rutherford (1905)
- Showed that atoms have a positive, dense center with electrons outside it
- Resulted in a planetary model
- Explains why electrons spin around the nucleus
- Suggests atoms are mostly empty space
- Didn't mention neutrons
- Doesn't explain valence level electrons role in chemical bonds

__________________________________________

Bohr (1920s)

- Electrons MUST only exist in specific orbitals around nucleus

- Explains how valence elctrons are involved in bonding

- Explains the difference between ionic and covalent bonding

- Resolve the neutron (discovered in 1932)

- Explains atomic emission spectra

***

Here's a random video of an unusual singer from Filipina's got talent....

Caloimetry and Molar Enthalpy

To measure heat absorbed/ released by water, we need to know:

• Temperature change (° C)
• Amount of Water (g, kg, mL, L)
• Specific Heat Capacity (kJ/kg ° C)- the heat needed to change 1 degree C in 1 kg

Example:
Calculate the amount of heat required to warm 400g of water from 20°C to 50°C.

ΔH = mCΔT
ΔH = (0,400kg)(4.19kJ/kg x T)(30°C)
ΔH = 50 kJ

Molar Enthalpy
- Heat absorbed / released by one mole

Example:
When a candle (C25H25) is burnt, heat is released according to the following reaction:

C25H52 + 38O₂ → 25CO₂ + 26H₂O +1100kJ

If 1.0g of wax is burnt, how much energy is released?

1.0g (mol / 352g)
= 0.00284mol

0.00284mol × 1100kJ / 1molC25H52
= 31.2 kJ/mol

Therefore: 31.2 kJ/mol of energy are released when 1.0g of wax is burnt

HEAT AND ENTHALPY

HEAT AND ENTHALPY

Ø Reactions that release heat are exothermic

Ø Reactions that absorb heat are endothermic

Ø Heat is a form of energy

Ø ENTHALPY - Stored energy

Ø Enthalpy of gasoline > Enthalpy of water

Ø Enthalpy symbol is H and change in enthalpy is ΔH

ENTHALPY GRAPHS:

EXAMPLES:
EXOTERMIC-
2C8H18 + 25O2 -----> 16 CO2 + 18H2O + 5076 kJ
2C8H18 + 25O2 -----> 16 CO2 + 18H2O ΔH= - 5076 kJ
ENDOTHERMIC-
3.2 C + 2H2 + 52.3 kJ -----> C2H4
3.2 C + 2H2 + -----> C2H4 ΔH= - 5076 kJ

MORE Practice:
State whether each of the following are exothermic or endothermic.
a.) H + Cl ---> HCl + 432 kJ EXOTHERMIC
b.) 12CO₂ + 11H₂O ---> C1₂H₂2O11 + 12 O1₂ ΔH= 5638kJ ENDOTHERMIC

jan 8 - types of chemical reactions

1) Synthesis

A+B ----> C (Two or more substances combine)
EX: H2 + Cl2 ----> 2HCl

6) Combustion


AB---->A+B (Breaking down into simpler substances)
EX: 2Ag2O---->4Ag + O2

3) Single Replacement

A + BX---->B + AX (Compunds must have a metal and a nonmetal-replacing one atom in a compound by anohter atom)
EX: Cl2 + 2KI ----> I2 + SKCl

4) Double Replacement

AB + XY ----> AY + XB (exchange of atoms between two different compounds)
EX: 2NaCl+ H2SO4 ----> 2 HCl + Na2SO4

5) Neutralization

(Products are water and an ionic salt/ Always between acids and bases)
EX: HCl + NaOH----> NaCL + H2O

6) Combustion

There are two types: Metallic (Can also be known as a synthesis reaction/includes oxygen) and hydro-carbon (includes carbon and oxygen). Also, the productes are always CO2 and H2O

EX: C5H12 + 8 O2 ----> 5CO2 + 6 H2O

Jan 6

Jan 6 - First class back from christmas break

Balancing with C, H, & O

1.) CH4 + O2 ---> CO2 + H2O
CH4 + O2 ---> CO2 + 2H2O
2.) CHH6 + O2 ---> CO2 + H2O
2CHH6 + 7O2 ---> 4CO2 + 6H2O
3.) C8H18 + O2 ----> CO2 + H2O
2C8H18 + 25O2 ----> 16CO2 + 18H2O

ALCOHOLS:
- Octane = most important chemical used to dilate fuel gas for cars
- OH means alcohol (Ex. C2H5OH = ethane)

Examples:
1.) CH3OH + O2 ---> CO2 + H2O
2CH3OH + 3O2 ---> 2CO2 + 4H2O
2.) C2H5OH + O2 ---> CO2 +H20
C2H5OH + 3O2 ---> 2CO2 +3H20

WORDS TO BALANCED EQUATIONS:

Example 1:
Aluminum chloride is mixed with potassium carvbonate. Aluminum carbonate and potassium chloride are formed.
Write the equation and balance:
AlCl3 + K2CO3 -----> Al2(CO3)3 + KCl
2 AlCl3 + 3 K2CO3 -----> 1 Al2(CO3)3 + 6 KCl

Example 2:
Aluminum metal reacts violently with bromine to produce aluminum bromide.
Al + Br2 -----> AlBr3
2 Al + 3 Br2 -----> 2 AlBr3

Example 3:
Magnesium sulphate hepta hydratee decomposes to form water and magnesium sulphate.
MgSO4 - 7 H2O -----> MgSO4 + 7 H2O


Acids - NEED TO KNOW:
HCl - Hydrochloric Acid
HNO3 - Nitric Acid
H2SO4 - Sulphuric Acid
H3PO4 - Phosphoric Acid
CH3COOH - Acetic Acid

EXTRA Practice:
1.) NH3 + O2 ---> NO + H2O
4NH3 + 5O2 ---> 4NO + 6H2O
2.) (NH4)2C2O4 + AlCl3 ---> Al2(C2O4)2 + NH4Cl
3(NH4)2C2O4 + 2AlCl3 ---> Al2(C2O4)2 + 6NH4Cl

3.) aluminum metal reacts with bromine to form aluminum bromide
2Al + 3Br ---> 2AlBr3