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


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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

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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








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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

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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

***

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