Notes+and+Discussions

__May 12, 2010 Balancing Redox Reactions__
Here are the handouts from class today: [|Balancing Redox Reactions-Handout.doc]

Please take a look at the handout solutions for further understanding of concepts that were unclear to you in class. [|handout1.jpg] [|handout2.jpg] (Oxid. number method) [|handout3.jpg]

[|Group Questions.doc] [|Solutions001.jpg] [|Solutions002.jpg]

Here is a link to an online practice program for balancing redox reactions: []

Here is a short clip: [|100 Greatest Discoveries: Electrochemistry]

Please complete the following problem (not hard at all :P) for homework using either of the two methods from class and we will take it up next class!** Extra Practice Problem: ** When iron pyrites, waste products from coal mining are exposed to the environment, their sulfur content is oxidized to sulfuric acid. This creates the environment problem known as //acid mine drainage//. Balance the equations for two of the main reactions involved. a) FeS2(s) + O2 + H2O à Fe 2+ + SO4 2- + H+ b) FeS2(s) + Fe 3+ +H2O à Fe 2+ + SO4 2- + H+

__April 26th, 2010~ Acids and Bases Introduction__
In your category, please do the following:
 * Give a definition (1-3 sentences)
 * Give a specific AND generic example if applicable
 * Give a diagram or picture if applicable
 * Create one question for the class to solve

1. Properties of Acids and Bases A theory described by Svante Arrhenius that describes properties of acids and bases in aqueous solutions. In an aqueous solution an acid disassociates to produce hydrogen ions and a base disassociates to produce hydroxyl ions. The product of an acid and a base is a salt and water.
 * 2. Arrhenius Theory (STOP DELETING OUR STUFF!!!!)**

When HCl and NaOH combine in equal proportion the H+ and OH- ions dissociate and combine with each other to produce water. The Chlorine and sodium ions also combine to form a salt.

HCl + NaOH à H2O + NaCl

H A + B OH à H2O + AB

An acid combined with a base in an aqueous solution will neutralize and produce water and a salt.

QUESTION: What will be the products of a reaction in which hydrobromic acid and calcium hydroxide? Write a word and chemical equation to demonstrate your findings.

3. Bronsted-Lowry Theory 4. Molecular Structure and trends in acid strength (binary acids)

4. Binary Acids are certain molecular compounds where hydrogen is combined with another nonmetallic element. (examples:HCl, HBr, HI) The strengths of the acids depend on solvation of the initial acid, the H-X bond energy, the electron affinity energy of X, and the solvation energy of X. The lower the electronegativity ofn the non-metallics, the more the acid will ionize producing a stronger acid.

HA ---> H+ + A-

Image: see pg. 552 (weak acids) 534 (strong acids)

Which is the strongest acid : a)HCl, b)HI, c)HF

5. Molecular Structure and trends in acid strenght (oxy acids) Definition: > has at least one hydrogen atom bound to oxygen
 * 1) contains oxygen
 * 1) forms an ion by the loss of one or more protons.
 * 2) has at least one hydrogen atom bound to oxygen
 * 3) forms an ion by the loss of one or more protons.

The electronegativity of the central atom (E) and the number of O atoms determine oxoacid acidity. With the same central atom E, acid strength increases as the number of oxygen attached to E increases. With the same number of oxygens around E, acid strength increases with the electronegativity of E. Specific Example: HNO3 (Nitric Acid) Generic Example: HaXbOc (where H is hydrogen, X is an element other then hydrogen or oxygen, and O is oxygen) Diagram: Question: Considering HClO3 what is it called when: a) Add an oxygen b) Nothing is changed c) Remove an oxygen d) Remove 2 oxygens

Answer Key: Relationship || General Name || Example Name || Example Formula ||
 * Relationship
 * one more oxygen atom than (root)ic || per(root)ic acid || perchloric acid || HClO4 ||
 * || (root)ic acid || chloric acid || HClO3 ||
 * one less oxygen atom than (root)ic || (root)ous acid || chlorous acid || HClO2 ||
 * two less oxygen atoms than (root)ic || hypo(root)ous acid || hypochlorous acid || HClO ||

6. Calculations with acids and bases  The pH determines how acidic/basic a substance is. Litmus paper can also be used to determine the pH depending on what colour the paper turns into.

pH = -log10[H+] [H+] = 10^-pH

Example: Calculate the pH for a specific [H+]. Calculate pH given [H+] = 1.4 x 10^-5 M pH = -log10[H+] pH = -log10(1.4 x 10^-5) pH = 4.85

Example: Calculate [H+] from a known pH. Find [H+] if pH = 8.5 [H+] = 10^-pH [H+] = 10^-8.5 [H+] = 3.2 x 10-9 M

QUESTION: What is the [H+] of a substance when the pH is 6.4??

Acid-Base Titrations Key Concepts:

acid-base reactions involve a proton transfer

the acid donates a proton to the base

acid-base reactions are also known as neutralisation reactions

acid + base -> salt + water

acid A + base B -> conjugate acid of base B + conjugate base of acid A(Bronsted-Lowry Theory)

H+ + OH- --> H2O is the most general neutralisation reaction

Equivalence point is the point at which the moles of H+ is equal to the moles of OH- An indicator is used to show the equivalence point during a titration

A titration involves the progressive addition of one reactant from a burette(usually the acid), to a known volume of the other reactant in a conical flask(usually the base)

Calculations
Example:
 * 1) Write the balanced chemical equation for the reaction
 * 2) Extract all the relevant information from the question
 * 3) Check that data for consistency, for example, concentrations are usually given in M or mol L-1 but volumes are often given in mL. You will need to convert the mL to L for consistency. The easiest way to do this is to multiply the volume in mL x 10-3
 * 4) Calculate the moles of reactant (n) for which you have both the volume(V) and concentration(M) : n = M x V
 * 5) From the balanced chemical equation find the mole ratio known reactant : unknown reactant
 * 6) Use the mole ratio to calculate the moles of the unknown reactant
 * 7) From the volume(V) of unknown reactant and its previously calculated moles(n), calculate its concentration(M): M = n ÷ V

30 mL of 0.10M NaOH neutralised 25.0mL of hydrochloric acid. Determine the concentration of the acid.

1. 30 mL of 0.10M NaOH neutralised 25.0mL of hydrochloric acid. Determine the concentration of the acid >> NaOH(aq) + HCl(aq) -> NaCl(aq) + H2O(l) >> **NaOH** V = 30mL, M = 0.10M **HCl** V = 25.0mL, M = ? >> **NaOH** V = 30 x 10-3L, M = 0.10M **HCl** V = 25.0 x 10-3L, M = ? >> n(NaOH) = M x V = 0.10 x 30 x 10-3 = 3 x 10-3 moles >> NaOH:HCl >> 1:1 >> NaOH: HCl is 1:1 >> So n(NaOH) = n(HCl) = 3 x 10-3 moles at the equivalence point >> n = 3 x 10-3 mol, V = 25.0 x 10-3L >> M(HCl) = 3 x 10-3 ÷ 25.0 x 10-3 = 0.12M or 0.12 mol L-1 Question:
 * 1) Write the balanced chemical equation for the reaction
 * 1) Extract the relevant information from the question:
 * 1) Check the data for consistency
 * 1) Calculate moles NaOH
 * 1) From the balanced chemical equation find the mole ratio
 * 1) Find moles HCl
 * 1) Calculate concentration of HCl: M = n ÷ V

50mL of 0.2mol L-1 NaOH neutralised 20mL of sulfuric acid. Determine the concentration of the acid.

Answer: 0.25M or 0.25 mol L-1

7. Lewis Theory

50mL of 0.2mol L-1 NaOH neutralised 20mL of sulfuric acid. Determine the concentration of the acid.

Answer: 0.25M or 0.25 mol L-1

7. Lewis Theory

Acid:
 * Corrosive ('burns' your skin)
 * Sour taste (e.g. lemons, vinegar)
 * Contains hydrogen ions (H+) when dissolved in water
 * Has a pH less than 7
 * Turns blue litmus paper to a red colour
 * Reacts with bases to form salt and water
 * Reacts with metals to form hydrogen gas
 * Reacts with carbonates to form carbon dioxide, water and a salt

Base: > > Diagram: > > > Question: > What's a neutralization reaction? (also: what's your favourite colour?) > > > 2. Arrhenius Theory > 3. Bronsted-Lowry Theory > 4. Molecular Structure and trends in acid strength (binary acids) > 5. Molecular Structure and trends in acid strenght (oxy acids) > 6. Calculations with acids and bases > 7. Lewis Theory > > 7. Lewis Theory > In the early 1920's, G.N. Lewis expanded the Bronsted-Lowry model to encompass a number of substances that would not normally be classified as Bronsted-Lowry acids or bases. A Lewis acid is an electorn-pair acceptor and a Lewis base is an electron-pair donor. In order to act as a Lewis base, a substance must possess a non-bonded pair of electrons in one of its orbitals. In order to act as a Lewis acid, a substance must possess an empty valence orbital that may accept (share) a pair of non-bonding valence electrons from a Lewis base. > > > A + B = AB <--- adduct > || external image =Untitled.jpg || > > > For example: > || external image ncontent || > > Question: In the following reaction, which is the lewis acid and which is the lewis base? HINT: Draw the VSPER diagrams. > > BF3 + NH3 ---> H3NBF3 > > > >
 * Hydrochloric acid (HCl) in gastric juice
 * Sulphuric acid (H2SO4)
 * Nitric acid (HNO3)
 * Carbonic acid in softdrink (H2CO3)
 * Uric acid in urine
 * Ascorbic acid (Vitamin C) in fruit
 * Citric acid in oranges and lemons
 * Acetic acid in vinegar
 * Tannic acid (in tea and wine)
 * Tartaric acid (in grapes)
 * [[image:http://chemwiki.ucdavis.edu/@api/deki/files/3756/=Untitled.jpg caption="external image =Untitled.jpg"]] ||
 * [[image:http://www3.interscience.wiley.com/tmp/graphtoc/26737/123222008/123218217/ncontent caption="external image ncontent"]] ||

from March 8, 2010
Atomic radius decreases as the number of protons increase. The radius increases when new layers are added. The radius increases a lot when going from the p orbital to the next s orbital. The radius size of the atoms arranged by the periodic table looks like a snowman. || ** Ionization Energy ** A + I.E. = A+ + e- The ioniztion energy is the energy needed to remove an electron from a neutral atom. Ionization energy increases when moving right on the periodic table and when moving up. || The energy released when adding an electron to an atom. Represented by the equation A + electron = A(negative) + E.A. || ** Electronegativity ** Chemical property describing an atom's ability to attract electrons We can classify the electronegativity of atoms and and determine whether the molecule is polar covalent, covalent, or ionic. (As described below) Covalent if electronegativity < 0.4, polar covalent if 0.4  1.7. || -net positive charge on electrons -the term "effective" explains the shielding of the outer electrons by the inner electrons -In the hydrogen atom, sole electron recieves full positive charge of nucleus -In all other atoms, effective nuclear charge can be calculated with the following equation: Zeff = Z-S where: Z is the number of protons and S is the number of nonvalence electrons
 * ** Atomic Radius **
 * ** Electron Affinity **
 * ** Effective Nuclear Charge **
 * ) || **Second and Third Ionization Energy

- The amount of energy spikes when passing a noble-gas configuration - Can be referred to as "electron-binding energy" - The smaller the atomic radius, the harder it is to remove an electron, and the more energy is required to remove it. (Caused by stronger attraction to positively charged nucleus) - As the distance from the nucleus increases, the amount of energy necessary to remove the electron decreases. - The energy required to remove a second or third electron from an atom is higher than that required to remove the first. - Usually applies to removing electrons from singly or doubly charged ions. ** ||

March 4th ~ 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.
 * Intramolecular Forces: **
 * A) Ionic Bonding **
 * 1) Why do ionic bonds form between metal and non-metal atoms? Manuel  - Ionic bonds between metals and non-metals occur because the metals have a surplus of electrons in terms of a full valence shell and non-metals are lacking electrons in terms of a valence shell. So metals give electrons to the non-metals, changing them both into ions with identical opposite charges. The now positive metal and negative non-metal are attracted due to the opposing charges and thus an ionic bond is formed.
 * 2) Which blocks in the periodic table do the atoms in an ionic bond typically come from? Simone- The s block and the p block.
 * 3) What are the properties of ionic solids? Ionic solids tend ts be a compound of non-metals and metals. Ionic solids are made up in a checkerboard fasion alternating between possitively and negatively charged particles. Ionic substances are more stable than molecular substances and are harder to break down.
 * 4) Why do ionic compounds generally have high melting and boiling points? Ionic compounds are made up of two atoms of with opposite charges; negative and positive. Since opposite charges attract each other, the ionic compounds have a strong bond. Thus making it harder to destroy these bonds because it takes up more energy to break them up.
 * 5) How is the **lattice energy** related to the formation of an ionic bond? Tiana Lattice energy is the energy given off in the formation of an ionic compound. In crystalline compounds, the net balence of forces (positive metal ions and negative nonmetal ions) is also called the LATTICE ENERGY.

** X+(g) + Y-(g) → XY(s) + Lattice Energy ** >>
 * 1) According to Coulomb’s Law, the force of attraction, **//F//**, between two oppositely charged particles of charge magnitude **//q//** is related the distance between them, **//d//**, by the following equation:

How would the magnitude of the charge on the oppositely charged ions, and the internuclear distance between in ions in the crystal lattice, affect the force of attraction between the ions? Tate > Covalent bonds are bonds that form between two or more atoms in which the valence electrons are shared between the atoms. There are two types of covalent bonds, covalent, and polar-covalent. Covalent bonds only occur between two or more non-metals. Covalent bonds are stronger than ionic bonds as the atoms want to remain stable and if the bonds are broken then the atoms lose their full valence shell which they do not like to do. The forces holding covalent compounds together are much weaker than the forces (intermolecular) between ionic compounds. This means that much less energy is required to seperate the atoms of a covalent compound, leading to lower melting and boiling points. 7. What is a polar covalent bond? Andy A polar covalent bond is a covalent bond where the electrons are held closer to one atom than the other (to the one with the highest electronegativity). This kind of bond will create a positively charged end and a negatively charged end to the molecule. These bonds generally happen between molecules where the difference in electronegativity is less that 1.7 (where it's still under 50% ionic), but greater than 0.
 * Covalent Bonding **
 * 1) “Covalent bonding involves a balance between the forces of attraction and repulsion that act between the nuclei and electrons of atoms”. Explain this statement. Conrad  -for a covalent bond to form, the repulsion of the electrons of the different atoms must not be stronger than the attraction of the protons and electrons.
 * 2) How is a covalent bond formed according to the hybridized orbital theory? Sakina  When hybrid orbitals are formed, it becomes easier for the atoms to form covalent bonds. The s orbitals and p orbitals combine to form the sp orbitals which contain large lobes. These lobes are able to overlap with other atoms.
 * 3) What are the two types of possible bonds and what is the difference between them? Alissa Sigma Bond is a bond created by the end -to-end overlap of atomic orbitals, while a Pi bond is a bond created by the side -by-side (or parallel) overlap of atomic orbitals, usually p orbitals.
 * 4) Which is stronger? A sigma bond or a pi bond? Why? Qingda  - Asigma bond is stronger, because it is head on, while the pi bond isn't.
 * 5) Rank the bonds in terms of bond strength (stored potential energy) //single, double, triple.// Reilly
 * Bond || # of Electons || Strength of Bond || Bond Order || Bond Length || Stored Potential Energy ||
 * Single || 2 || Weak || 1 || Long || Little ||
 * Double || 4 ||  || 2 ||   ||   ||
 * Triple || 6 || Strong || 3 || Short || Lot ||
 * 1) What are the properties of covalent bonds? Vlad
 * 1) Why are the melting and boiling points of covalent compounds lower than those of ionic compounds? Danielle
 * Metallic Bonding **
 * 1) How do metal atoms bond together? The outer-most orbital of the each atom overlaps with orbitals surrounding it. Electrons then become detached from their original atom and are then said to be "delocalised". The seperate atoms are held together because of the attraction between the delocalised electrons and the positive particles within each atom's nucleus.
 * 2) Why is metallic bonding often described as positive ions in a “sea” of electrons? Nura- //A metal's electrons delocalize from the individual atoms and occupy a single "orbital" that belongs to the metallic crystal as a whole. An analogy for this would be protons within this 'sea' of delocalized electrons. This theory is supported by the fact that metals are conductors of electricity//.
 * 3) Why are the electrons free to move? Please explain using orbital theory. David

a) Metals are excellent conductors of heat and electricity. b) Metals tend to bend rather than break When a piece of metal is hammered or stretched, the atoms slide or roll past one another. The metal atoms rearrange themselves within the “electron sea” to assume the new shape without breaking their metallic bonds. c) Almost all metals are solids at room temperature. The hardness and melting point of a metal depend to a large extent on how many valence electrons the metal has to participate in metallic bonding. The greater the number of valence electrons, the more firmly the metal ions are held in place. Moving left to right across a period, the hardness and melting points of metals generally increases as the number of valence electrons available for metallic bonding increases. So most metals are solid at room temperature.
 * 1) How does the free-electron model explain the following physical properties? Jia Hua
 * 1) What is an alloy? Alloys are solid solutions made of one or more metallic elements.
 * 2) Why do alloys have different properties than pure metals? Please give one example. Tania


 * Intermolecular Forces **
 * 1) What are the major types of intermolecular forces? Hasha: hydrogen bonds, dipole-dipole interactions, dispersion forces (london forces, weak intermolecular forces, van der waal's forces)
 * 2) Rank these forces from //strongest// to //weakest//. Claire

The molecule can not have a centre of symmetry and the non-symmetric atoms must have different electronegativity.
 * 1) What are the two requirements for a molecule to have a **//dipole//**? Ben
 * 1) List three molecules that have a dipole and three molecules that do NOT have a dipole. Scott

 ** EVERYBODY DISCUS S So…..What is the difference between an intramolecular and an intermolecular force???? ** The blips (drops) in the ionization energy trend
 * 1) Look at Figure 11 on page 263 of your textbook. Write a short paragraph using at least the following words plus any others you may want to add. Peter
 * …hydrogen bonding, crystal lattice, interstitial spaces, methane molecule, global warming, greenhouse gas, arctic ice… **
 * Explaining exceptions to the general trends in the Periodic Table **

***Formal Charge***
[] chemical bonding practice
 * Some good practice for chemical bonding!**