Class+Notes+and+PowerPoints

=__Electrochemistry Study Questions__=
 * //Please add your questions next to your topic. And include answers if not in the book.// **

Definition of Oxidation and Reduction :
 * what is the difference between oxidation and reduction?
 * p 657 #9

Oxidation States:
 * differentiate between oxidation number and electric charges
 * p 659 # 12f

Galvanic and Voltaic Cells What is a galvanic cell? In what direction does the electrical current travel? (cathode-anode or anode-cathode)

In a silver-copper cell, there is evidence that the copper electrode decreases in mass and the intensity of the blue colour of the electrolyte increases. Therefore:

a) oxidation of copper metal is occuring Cu (s) -> Cu^2+ (blue) + 2e- b) Recudtion of silver ions is occuring Ag+ (aq) + e- -> Ag (s) c) Electrons are moving from the silver electrode to the copper electrode d) Reduction of copper metal is occuring

Write equations for the half-reaction and the overall reaction that occur in the following cell: Cu (s) | Cr2O7 2- (aq), H+ (aq) || Cu2+ (aq) | Cu (s) Cathode Cr2O7 2- + 14H+ + 6e- -> 2Cr 3+ + 7H2O Anode 3[Cu(s) -> Cu2+ + 2e-

Net Cr2O7 2- + 14H+ 3Cu(s) -> 3Cu 2+ + 2Cr 3+ + 7H2O

Balancing Redox Reactions Electrolysis

Faraday's Constant What are the units for Faraday's constant? Answer: C/mol How long will it take a current of 3.50A to transfer 0.100 mol of e-? Answer: 46min, pg 749

a) For this system write the shorthand notation. b) Indicate the cathode and anode. Also include if the reduction or oxidation reaction occurs at either one. c) Write the two half reactions and the Voltage.
 * Shorthand notation for Electrochemical Cells**

a) Cu(s)|CuSO4(aq)||ZnSO4(aq)|Zn(s) b)Cathode: Copper; here the reduction occurs Anode: Zinc; here the oxidation occurs c) Cu(s) --> Cu2+(aq) + 2e- V= -0.34 Zn2+(aq) + 2e- --> Zn(s) V= -0.76
 * Answers:**

Calculating Cell Potential

pg. 708 #11 pg. 709 # 2 The standard cell potential E(naught) is a) max electric potential between cathode and anode b) min electric potential between cathode and the anode c) does not affect the electric potential between cathode and the anode

Using Reduction Potentials Table

To help you study, I suggest doing
 * the self-quiz on p. 764 (lots of m.c. questions)
 * Unit 5 Review questions pp. 766-769 ~ All questions except for #14, 15,

Considering the reduction of a metal cation to form its elemental form (i.e. Zn 2+ + 2 e- ---> Zn), which of the following is FALSE?

a) This usually occurs spontaneously in water. b) High voltages will favour the reduction of the cation over the reduction of water. c) This type of reaction usually occurs at temperatures in excess of 500C. d) The reduction potentials of different cations can be used to determine the voltage needed to create their elemental forms. e) None of the above ANSWER: (a)

Pg 679 #19

How does one read the reduction potential chart? The chart is a assembled in a list of half-reactions of increasing strength of the reducing agent. The chart also lists their corresponding reduction potentials in volts. It should be noted that all of the equations are written with the oxidizing agents on the reactants side. =__Organic Functional Group Priority List__=

Here is a great website with the priority list. It even has pictures you can manipulate with your mouse! Some of these, the anhydrides, acid halides and sulphides, will not be encountered in this course but it's still pretty great. You will get this list on the exam. Thank you Prof. Hunt from the University of Calgary. [|functional groups page]

Here is that powerpoint...
 * __Balancing Using Half-Reactions__**


 * __Introduction to Acids and Bases__**
 * __A) Bronsted-Lowry Theory

__** Summary : Which elements are the acids and bases in a reaction is determined by the tranferral of protons, as opposed to their production. Whether something is an "acid" or a "base" can only be said with reference to a specific reaction; many substances are amphoteric, which means they can act as an acid or a base depending on the reaction. A Bronsted-Lowry acid is a proton donor. A Brondstead-Lowry base is a protein acceptor. Acid-base reactions are universally reversible and result in acid-base equilibrium. A conjugate acid-base pair consists of two substances that differ only by a proton: the acid in a conjugate acid-base pair has one more proton than the base.

Example:

Water is amphoteric! Here, water acts as a base, accepting a proton. CH3CO2H + H2O CH3CO2- + H3O+
 * [[image:http://upload.wikimedia.org/wikipedia/commons/thumb/9/96/Equilibrium.svg/20px-Equilibrium.svg.png width="20" height="17" caption="is in equilibrium with" link="http://en.wikipedia.org/wiki/File:Equilibrium.svg"]] ||
 * ||

Here, water acts as an acid, donating a proton.

H2O + NH3 OH- + NH4+
 * [[image:http://upload.wikimedia.org/wikipedia/commons/thumb/9/96/Equilibrium.svg/20px-Equilibrium.svg.png width="20" height="17" caption="is in equilibrium with" link="http://en.wikipedia.org/wiki/File:Equilibrium.svg"]] ||
 * ||

Diagram

Suggested Problems

Page 532, Question 1 & 2. Page 632, Question 10a.

__**B) Reversible Reactions**__

- **reversible reactions are proton transfer reactions at equilibrium**, where both forward and reverse reactions involve Bronsted-Lowry acids and bases - there are always 2 acids and bases involved in the equilibrium, and they form 2 conjugate acid-base pairs - **conjugate acid-base pair**: 2 substance whose formulas differ by one H+ ion e.g. H2O and H3O+ make a conjugate pair, where H2O is the conjugate base and H3O+ is the conjugate acid - General equation: conjugate acid 1 + conjugate base 2 <--> conjugate base 1 + conjugate acid 2 - the substances on the right can just as easily react, producing the substances on the left. In the reverse reaction, the reactants of the forward reaction would become products, and the products of the forward reaction serve as the reactants.

Hydrofluoric acid is a weak acid. Its reaction with water is as follows: HF + H2O <--> H3O+ + F− Here HF and F- are conjugate acid-base pairs. HF = conjugate acid, F- = conjugate base H2O and H3O+ are also conjugate pairs. H2O = conjugate base, H3O+ = conjugate acid
 * Example**:

Acid-Base Equilibrium in Water Animations: http://www.chembio.uoguelph.ca/educmat/chm19104/chemtoons/chemtoons.htm

Question: Identify the conjugate acid-base pairs. NH3 + H2O <--> NH4+ + OH- H2SO4 + H2O <--> HSO4- + H3O+

  - certain elements have a higher affinity for protons, where as some cannot hold on to protons very well - for example, C2H3O2- has a greater ability to hold on to its proton (H+) than the ability of H2O to pull the proton away. Therefore, the percent ionization is low 1.3% - HC2H3O2(aq) + H2O2(l) <=> H3O+(aq) + C2H3O2- (aq) - The stronger an acid, the weaker its conjugate base, and conversely, the weaker an acid, the stronger its conjugate base
 * __Competition for Acidity__**

Example Question: What does it mean when the arrow in an equation is a one-way arrow? A two-way arrow?

Answer: One-way arrow means that it is a strong acid or base, because equilibrium lies in the direction of the arrow. Reversible arrow means that it is a weak acid or base.

D) Strong Acids

A strong acid is an acid that ionizes completely in an aqueous solution. This means that greater than 99% of molecules ionize, leading to the assumption of 100% ionization for calculations. Monoprotic acids contain one ionizable hydrogen atom while diprotic and triprotic acids contain 2 and 3 respectively.

Kw = [H+(aq)][OH-(aq)] This equation is used to solve problems involving strong acids. All of the hydrogen atoms come from the strong acids. The concentration of the strong acid is equal to the concentration of hydrogen ions.

Eg) HCl(g) -(100%)-> H+(aq) + Cl-(aq)

Sample Problem: page 537, question 6

__**E) Strong Bases** Base: Substance that dissociates to **increase** the hydroxide ion concentration when dissolved in water. Strong Base: Substance that dissociates **quantitatively** (completely) to increase the hydroxide ion concentration when dissolved in water.

Basically, since it dissociates completely, you can just balance the equation and then determine the mole ratio.

Equation: Kw=[H][OH]

Example: NaOH --> Na + OH For every mole of sodium hydroxide, one mole of hydroxide ion is poroduced.

Problem:
 * Calculate the hydrogen ion concentration in a 0.25 M solution of barium hydroxide. a strong base.

F) Hydrogen ion Concentration and pH**__



//**Write your answers next to and/or under each questions. The sooner the better. :) **//
 * __Chemical Bonding and Intermolecular Forces Questions__**

Answer the general questions below about chemical bonding to help understand the difference between **//intermolecular//** and //i**ntramolecular**// forces.

Safiyyah Yawen In the formation of an ionic bond, lattice energy is given off when oppositely charged ions in the gas phase come together to form a solid. This can be illustrated in the following equation:  X + (g) +Y¯(g) -->XY(s) + lattice energy This also works vise versa: lattice energy is also the energy required to separate an ionic solid into gaseous ions. //F//, between two oppositely charged particles of charge magnitude **//q//** is related the distance between them, **//d//** F = d^2 -the the charge on the ions, the the force of attraction -the the internuclear distance between ions in the lattice, the  the force of attraction
 * Intramolecular Forces:**
 * A) Ionic Bonding **
 * 1) **Why do ionic bonds form between metal and non-metal atoms?** **Tessa:** Well everything is just jonesing for stability. If you have a metal and non-metal interacting to get that stability, you'll have an ionic bond. This is because metals have a low ionisation energy (their electrons are not clutched very tightly to their nucleis) and few electrons in their outer n-level. Meanwhile, non-metals have many electrons in their outer n-level, and a high electronegativity (their nucleis attract electrons most vigorously). So you have the non-metal tugging electrons away from the metal. And then you have ions!
 * 2) **Which blocks in the periodic table do the atoms in an ionic bond typically come from?** **Elisa:** Ionic bonds are typically formed by s or d block elements, where most metals are found, bonding with p block elements, where mostly non-metals are found.
 * 3) **What are the properties of ionic solids?** **Donnette**: Ionic solids are made up of metal and non-metal ions. They have high melting and boiling points. They do not conduct electricity unless they are dissolved in water (most of the time).They tend to be crystals that are hard but brittle. Their crystal lattice structures and strong ionic bonds are what cause this.
 * 1) According to Coulomb’s Law, the force of attraction,
 * 1) , by the following equation:

LucWell the covalent bond is formed with 1 electron from each of the atoms. The only difference in this theory is that now more covalent bonds can occur, instead of having separate s and p orbital’s the atom now has 4 sp orbital’s. occur between two different non-metals that share electrons unequally. (electronegativity difference < 1.7) ( 2) Non-polar covalent bonds occur when two identical non-metals equally share electrons between them. (electronegativity diff. is 0) <span style="color: rgb(255,0,6);">Mads... <span style="color: rgb(13,128,0);">Answer: Sigma bonds are stronger for two reasons 1) They have less energy than the pi bonds because they are in a hybridized state 2) They have more area for the interference to occur //single, double, triple.// Danielle <span style="font-size: 10pt; color: rgb(128,0,128); font-family: Arial,Helvetica,sans-serif;">Triple, double, single (refer to "evidence" on page 240 of the chemistry book)
 * Covalent Bonding
 * <span style="color: rgb(110,41,168);">Electrons between covalently bonded atoms are attracted to both nuclei bu repelled by each other. The attraction to the nuclei holds the atoms together and compress the orbitals. The repulsion of the electrons forces them to spread out into the shapes they form around the nuclei. Curtis
 * <span style="color: rgb(110,41,168);">Electrons between covalently bonded atoms are attracted to both nuclei bu repelled by each other. The attraction to the nuclei holds the atoms together and compress the orbitals. The repulsion of the electrons forces them to spread out into the shapes they form around the nuclei. Curtis
 * . ( <span style="font-size: 10pt; color: rgb(6,50,96); font-family: Arial;">1) Polar covalent bonds
 * <span style="font-size: 10pt; color: rgb(0,0,0); font-family: Arial,Helvetica,sans-serif;">
 * <span style="font-size: 10pt; color: rgb(0,0,0); font-family: Arial,Helvetica,sans-serif;">
 * <span style="font-size: 10pt; color: rgb(0,0,0); font-family: Arial,Helvetica,sans-serif;">
 * <span style="font-size: 10pt; color: rgb(128,0,128); font-family: Arial,Helvetica,sans-serif;">
 * <span style="font-size: 10pt; color: rgb(128,0,128); font-family: Arial,Helvetica,sans-serif;">Iza ... <span style="font-size: 10pt; color: rgb(224,11,11); font-family: Arial,Helvetica,sans-serif; background-color: rgb(255,255,255);">Low melting/boiling points (compared to ionic), gas/liquid at room temp., nonconducting, insoluable in water.
 * <span style="font-size: 10pt; color: rgb(0,0,0); font-family: Arial,Helvetica,sans-serif;">

Marie <span style="color: rgb(232,56,255);">Polar Covalent bonds are covalent bonds between atoms where the electrons are shared unevenly due to electronegativity variations. For a covalent bond to be polar it must have a difference in electronegativity of 0.4 or greater.
 * <span style="font-size: 10pt; color: maroon; font-family: Arial,Helvetica,sans-serif;">... <span style="font-size: 10pt; color: rgb(128,0,128); font-family: Arial,Helvetica,sans-serif;">Answer: Ionic compounds have very high melting and boiling points because it takes a lot of energy to seperate the positive and negative ions in the crystal lattice structure; they have very strong forces of electrostatic attraction. Covalent compounds, however, form molecules with weaker intermolecular forces in which the ions do not interact much. Therefore, it is very easy to separte those bonds as little energy is required to do so.
 * <span style="font-size: 10pt; color: rgb(0,0,0); font-family: Arial,Helvetica,sans-serif;">
 * 1) <span style="font-size: 10pt; color: rgb(0,0,0); font-family: Arial,Helvetica,sans-serif;">

Metallic Bonding Matthew Melody Sophie Neeti <span style="font-size: 10pt; color: rgb(11,137,11); font-family: Arial,Helvetica,sans-serif;">a) Metals are excellent conductors of heat and electricity. Due to metallic bonds having free flowing electrons, when some of the electrons are excited on one end, they cn move and pass of the charge at the other end. Also due to the ions being closely packed in a metallic bonding, when heat is applied, the vibrations of the atoms pass causing them to conduct heat/ b) Metals tend to bend rather than break As the positive ions have a 'sea' of electrons around them, when pressure is applied, instead of the bond breaking, the ions slip ontop of eachother, causing them to bend instead of break c) Almost all metals are solids at room temperature. Metals are solid in room temperature because of the strong electromagnetic attraction between the electrons thus requiring a lot of energy to break them and it is due to this they are solid at room temperature. Cassie
 * Metallic bonding is the result of the electromagnetic attraction between the free-flowing, or delocalized, electrons of metals and the positively charged nuclei of these metals. In a metallically bonded lattice, the positively charged nuclei form the shape of the lattice, with the free-flowing electrons between them acting like a glue between the nuclei. The positive nuclei are electrostatically attracted to the sea of electrons between them. Since the electrons are free-flowing, there is attraction between nuclei and electrons in almost every direction in the lattice, causing metal nuclei to pack closely together and forming a strong latice structure. Note: Single metallic bonding doesn't really exist, metals usually bond in a collective.
 * Metallic bonding is the result of the electromagnetic attraction between the free-flowing, or delocalized, electrons of metals and the positively charged nuclei of these metals. In a metallically bonded lattice, the positively charged nuclei form the shape of the lattice, with the free-flowing electrons between them acting like a glue between the nuclei. The positive nuclei are electrostatically attracted to the sea of electrons between them. Since the electrons are free-flowing, there is attraction between nuclei and electrons in almost every direction in the lattice, causing metal nuclei to pack closely together and forming a strong latice structure. Note: Single metallic bonding doesn't really exist, metals usually bond in a collective.
 * Since the atoms are not very electronegative, their electrons are free to move around between atoms. The protons stay stationary while the electrons "flow" around them.
 * Since the atoms are not very electronegative, their electrons are free to move around between atoms. The protons stay stationary while the electrons "flow" around them.
 * <span style="color: rgb(98,144,65); font-family: 'Times New Roman',Times,serif;">Response: The valence electrons are mobile because metal atoms generally have low ionization energy, thus the electrons are only loosely held to them. As well, metals generally have a great number of empty valence orbitals (compared to their few electrons) into which the electrons can be easily excited. Furthermore, the close-packing structure of metal atoms causes overlaps of orbitals that allow electrons to have great freedom in movement between atoms through these orbitals(?).
 * <span style="color: rgb(98,144,65); font-family: 'Times New Roman',Times,serif;">Response: The valence electrons are mobile because metal atoms generally have low ionization energy, thus the electrons are only loosely held to them. As well, metals generally have a great number of empty valence orbitals (compared to their few electrons) into which the electrons can be easily excited. Furthermore, the close-packing structure of metal atoms causes overlaps of orbitals that allow electrons to have great freedom in movement between atoms through these orbitals(?).

Cassie

Intermolecular Forces

//strongest////weakest//
 * Grace- The three major types of intermolecular forces are Hydrogen bonds, Dipole-Dipole bonds, and Dispersion (aka London) forces.

//dipole// The molecule must be made up of covalent bond(s) and be polar.**
 * 1) from strongest to weakest: hydrogen bonds, dipole-dipole, then dispersion.