Download Second Semester Notes 09-10

Second Semester Notes 14-15
UNIT 7 – Periodic Table
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Arrangement of Periodic Table
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Periodic law – the properties of elements are periodic functions of their atomic number
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Periods – horizontal rows; 7; correspond to energy levels
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Groups/Families – vertical columns; “group A” Roman numerals correspond to the number of valence electrons
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Group I – Alkali metals, Group II – Alkaline earth metals, Group VII – Halogens, Group VIII – Noble gases
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Metals – everything to the left of the stairstep; including aluminum; does not include hydrogen
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Properties: Have luster (shiny), good conductors of heat and electricity, malleable ( able to be pounded into sheets), ductile (able
to be pulled into a wire), tend to lose electrons in chemical reactions, most are solids
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Transition metals – middle block over to stairstep
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Inner transition metals – bottom 2 rows; sometimes called “lanthanide series” and “actinide series”
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Nonmetals – everything to the right of the stairstep; includes hydrogen
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Properties – Dull, poor conductors, brittle, tend to gain or share electrons in chemical reactions, most are gases
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Metalloids – either side of the stairstep; does not include aluminum
oxidation number – “+” or “-“ number which indicates how many electrons will be gained (-) or lost (+) when an atom forms a compound
UNIT 8 EMPIRICAL AND MOLECULAR FORMULAS
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Mole
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In Chemistry, we use the mole to represent the # of particles in a substance
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One mole represents 6.02 x 1023 things, which is called Avogadro's number
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One mole of most elements contains 6.02 x 1023 atoms
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However, some elements naturally exist as diatomic molecules (the “magic 7”) - These contain Avogadro’s number of molecules
Molar mass - is the mass (think grams) of one mole of a substance
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Atomic masses of atoms are relative masses based on the mass of carbon-12
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To calculate the molar mass of a compound, you add up the molar masses of all the elements in that compound
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When you see 1.00 mole = _?_ g, think “g means go to the Periodic table” to find the molar mass
Molar Volume
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The volume of a gas is usually measured at standard temperature and pressure (STP)
1.
Standard temp = 0° C
2.
Standard pressure = 1 atmosphere (atm)
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1 mole of any gas occupies 22.4 L of space at STP
Putting it all together:
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1.0 mole = 6.02 x 1023 atoms or molecules = ? g = 22.4 L (at STP)
Percent Composition - the percentage by mass of each element in a compound
A.
The percent comp. is found by using the following formula:
% mass = mass of 1 element x 100
mass of compound
Empirical Formula
This is the lowest whole number ratio of the elements in a compound. For example, the empirical formula for C6H6 is C H
Example: C6H12O6
CH2O
To calculate the empirical formula, given the mass or percent of elements in compound, follow these steps:
1.
If given a percent sign, remove the sign & change to grams. You are assuming you have 100 g of the compound.
2.
Convert the grams to moles.
3.
Decide which number of moles is the lowest, then divide each number of moles by this number.
4.
If the number divides out evenly, these are the subscripts of the elements in the compound. If any of the numbers have a .5,
multiply them all by two & then place these numbers as the subscripts.
MOLECULAR FORMULA -
It is a larger version or a multiple of an empirical formula.
Example:
The molar mass of a molecular formula is 283.88 g/mole. Determine the molecular
formula if the empirical formula is P2O5.
Step 1: First find the molar mass of the empirical formula.
P2O5 = 141.94 g/mole
Step 2: Divide the molar mass of the molecular formula by the molar mass of the empirical formula.
283.88 g/mole/141.94 g/mole = 2
This tells us that the molecular formula is two times the empirical formula
Step 3: Multiply the subscripts of the empirical formula by the number you calculated in Step 2.
Multiply subscripts by two and we have P4O10.
UNIT 9 – IONIC BONDING
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Chemical Bonds - Force that holds 2 atoms together
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Valence Electrons and electronegativity determine the type of bonding.
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Ionic bonding- Occurs when electrons are completely transferred from one atom to another. Held together by electrostatic force.
This is the strongest type of bond. Occurs between metals & nonmetals
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Polyatomic ion -more than one element attached to the charge.
Bonding Basics
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Make sure you understand valence electrons and electron configurations
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Draw Dot structures of valence electrons
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Know your oxidation numbers
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Make sure that positive and negative charges add up to zero!
NAMING BINARY IONIC COMPOUNDS
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Binary compounds are made up of only 2 different elements
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STEP 1: IDENTIFY THE TYPE OF METAL
Type I Metal: Group I, II or Al+3, Zn+2, Cd+2, Ag+1
These metals have only 1 oxidation state
TYPE II metal: all metals that are not type I metals
You must work backwards from the formula to determine the oxidation #
for the metal
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STEP 2: WRITE THE NAME
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Type I metals: name the metal and change the nonmetal ending to –ide
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Metal nonmetal-ide Ex. CaCl2 Calcium Chloride
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Type II metals:
name metal, write oxidation number as a Roman numeral and change nonmetal ending to –ide
Metal (oxidation #) nonmetal-ide
Example: FeCl3 Iron (III) Chloride
POLYATOMIC IONS (PAI)
Two or more nonmetal atoms bonded together
Act as a single unit with a net charge
UNIT 10 COVALENT BONDING
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Octet Rule - Atoms (except H & He) are most stable when they have 8 electrons in their outer energy level
Covalent Bond -the chemical bond that results from the SHARING of valence electrons
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Occurs when elements are close together on the periodic table
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Between nonmetallic elements
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Molecule-formed when two or more atoms bond covalently
Characteristics of Covalent Bonds
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Can exist as gases, liquids, or solids depending on molecular mass or polarity
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Usually have lower MP and BP than ionic compounds
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Do not usually dissolve in water
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Do not conduct electricity
Types of Bonds:
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Nonpolar covalent (also called pure covalent) – equal sharing of electrons between atoms; example: O2
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Polar covalent – unequal sharing of electrons between atoms; example: H2O
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Ionic – complete transfer of electrons; example: NaCl
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*electronegativity (EN) – indicates how strongly an atom wants to gain an electron
Diatomic Molecules - Contains only two atoms. 7 naturally occurring in nature
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Examples: H2 O2
F2
Br2
I2
N2
Cl2
To Determine Molecular Shape: - Use VSEPR (valence shell electron pair repulsion) rules:
1.
Draw the Lewis dot structure for the molecule
2. Identify the central atom
3. Count total # of electron pairs around the central atom
4. Count # of bonding pairs of electrons around the central atom
5. Count # of lone pairs of electrons around the central atom
6. Look at summary chart, identify shape
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Naming Covalent Molecules
1.
Name the first element using the entire name
2. The second element is named using the root and the suffix (ending) –ide
3. Prefixes are used to indicate the number of each type of atom.
Exception-the first element will never have the prefix mono
Prefixes:
1
_mono_________________
2
_di__________________
3
_tri__________________
4
_tetra_________________
5
__penta_____________
6 ___hexa____________________
7 ___hepta____________________
8 ___octa____________________
9 ___nona____________________
10 __deca_____________________
UNIT 11
EQUATIONS AND TYPES OF REACTIONS
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Parts of a reaction
Reactants
Products
2H 2(g)  O 2(g)  2H 2 O (l)
Coefficient
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Subscript
State of Matter
Arrow means “to yield”
Law of Conservation of Matter
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Matter cannot be created nor destroyed, it can only change forms.
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Therefore, in a chemical reaction the number of atoms, the mass, & the charge must be conserved.
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This means it is the same on both sides of the arrow.
Types of Reactions - Reactions are classified into several categories.
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Combination
A + B  AB
2. Decomposition
AB  A + B
3. Single Replacement
A + BC  AC + B
4. Double Replacement
AB + CD AD + CB
5. Combustion
CxHy + O2 CO2+ H2O
(oxygen combines with a substance & releases energy (heat & light))
Writing Equations from words
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Write the formula for each reactant and each product.
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Your formula MUST be correct. That’s why we learned the polyatomic ions!
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Remember the diatomic elements H, O, F, Br, I N, Cl. The are written H2, O2, F2, Br2, I2, N2 and Cl2
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Add the symbols for states of matter
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Take inventory of the equation and balance
Predicting Products
Steps
1.
Determine what type of reaction is being presented.
2. Write the correct formulas for the product(s).
3. Balance the equation.
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Redox reaction – a reaction in which electrons are transferred from one atom to another; charge is conserved
Oxidation – loss of electrons from atoms of a substance; oxidation # increases; substance that is oxidized acts as the REDUCING
AGENT (RA)
Ex –
Na  Na+ + e-
Sodium is oxidized
Reduction – gain of electrons by atoms of a substance; oxidation # decreases; substance that is reduced acs as the OXIDIZING
AGENT (OA)
Ex – Cl2 + 2e-  2ClChlorine is reduced
*Memory technique* LEO GER (Loss of Electrons is Oxidation, Gain of Electrons is Reduction)
oxidized
Ex – 2K (s) + Br2 (g)  2KBr (s)
Rules for assigning Oxidation Numbers
ALWAYS
GROUP I +1
GROUP II
Al
Zn
Ag
F
GROUP 17
GROUP 16
N and P
+2
+3
+2
+1
-1
-1 with metals
-2 with metals
-3 with metals
ALMOST ALWAYS
H +1
(except with active metals)
O -2
(except in H2O2 and w/ F)
**The sum of the oxidation numbers in a compound must equal zero.
**The sum of the oxidation number in a PAI is equal to the charge of the ion.
**All uncombined elements and diatomics have an oxidation number of zero.
RATE
Reaction rate – the change in concentration (molarity) of a reactant or product per unit of time
Factors That Affect Rate
1. Concentration
2. Surface area
3. Temperature
4. Catalysts
Concentration
The greater the concentration, the more often collisions can occur, the greater the reaction rate.
Surface Area
Increasing the surface area of reactants provides more opportunity for collisions with other reactants.
Temperature
Increasing temperature increases the average kinetic energy of the particles that make up a substance
Particles collide more frequently at higher temperatures than at lower temperatures.
Therefore, increasing the temperature at which a reaction occurs increases the rate of reaction.
Catalysts / Inhibitors
Catalyst – increases the rate of a chemical reaction without it self being consumed in the reaction.
It is not part of the product … … it does not yield more product … … it is not included in the written equation
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A catalyst decreases the activation energy by providing an alternate method for reactants to react.
Inhibitor – a substance that slows down the rate of reaction.
An inhibitor increases that activation energy.
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Enthalpy (H) – a measure of heat content of a system
∆H = heat absorbed or released from a system; these values will be given to you
∆Hrxn = Hproducts - Hreactants
ENTHALPY
Exothermic reactions
*chemicals react and give off heat
*∆H is negative
*products are more stable
Endothermic reactions
*chemicals need to absorb energy in order for the reaction to take place
*∆H is positive
*reactants are more stable
UNIT 12 STOICHIOMETRY/LIMITING REACTANTS
Using Ratios to Solve Stoichiometry Problems
Ratios can be used to solve almost any problem that could be solved using dimensional analysis. To use this approach, you must
recall “Mole Facts” for each substance.
Mole Facts for a particular substance:
1 mole = 6.02 x 1023 particles (molecules, formula units)
1 mole = molar mass (Periodic Table)
1 mole = 22.4 L @ STP (for gases)
Problem Solving Strategy:
1st: balance the equation
2nd: write the info from the word problem above the equation
3rd: use mole island diagram to map out your steps
4th: use dimensional analysis and mole ratios to solve
Example:
2KClO3  2KCl + 3O2
How many grams of KClO3 are required to produce 9.00 L of O2 at STP?
# grams KClO3 = 9.0 L O2 X 1 mole O2 X 2 moles KClO3 X 122.5 g KClO3 = 32.8 g
22.4 L O2
3 moles O2
1 mole KClO3
LIMITING/EXCESS REACTANTS: Used up first in a reaction -Limit how much product is made
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1st:
2nd:
3rd:
4th:
balance equation
pick a product (it is wise to choose the simplest product)
using the starting amounts of each reactant, use stoichiometry to calculate the amount (in grams) of chosen product
analyze your results; the limiting reactant is the reactant which gives you less product
If there are 100.0 grams of each reactant available, determine which one is the limiting reactant.
3Ca +
2AlBr3  3CaBr2 + 2Al
#g Al = 100 g Ca X 1mole Ca X 2 mole Al X 27g Al
= 45 g Al
40 g Ca
3 mole Ca
1 mole Al
# g Al = 100 g AlBr3 X 1 mole AlBr3 X 2 mole Al
X 27 g Al = 10.1 g Al
267 g AlBr3
2 mole AlBr3
1 mole Al
AlBr3 is the limiting reactant.
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Theoretical Yield: calculated amount of products
Actual Yield: amount of product formed in laboratory experiment
% yield = actual yield (from experiment)
x 100
theoretical yield (from calculations)
% yield = 14 g H2O x 100
= 82.4% yield
17 g H2O
UNIT 13 ACIDS AND BASES
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Acids
2. Give up H+ when dissolved in water
3. Turn litmus red
4. Clear in Phenolphthalein
5. Sour taste
6. Has pH < 7
7. Reacts with metal to form hydrogen
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Bases
1. Gives up (OH)- when dissolved in water
2. Turns litmus blue
3. Pink in Phenolphthalein
4. Bitter taste
5. Feels slippery
6. has pH > 7
*pH – power of hydrogen ino concentration
*pH = -log [H+]
[H+] = 2nd log –pH
 Neutralization
The reaction of an acid and a base that produces a salt and water.
KOH + HCl  HOH + KCl
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Titration – the process which involves finding the concentration of an unknown solution by
using a certain volume of a known solution.
End point – the point where neutralization is achieved and the indicator has changed color
Naming Acids
Anion
Ending
ACIDS
start with 'H'
Acid Name
-ide
hydro-(stem)-ic acid
-ate
(stem)-ic acid
-ite
(stem)-ous acid
2 elements
3 or more elements (with oxygen)
hydro- prefix
-ic ending
no hydro- prefix
-ate ending
becomes
-ic ending
-ite ending
becomes
-ous ending
Naming Bases
Just like all ionic compounds, write the symbol for each ion, its oxidation number, and then criss-cross to get the subscripts.
Example: potassium hydroxide
K+1 (OH)-1 KOH
Bronsted-Lowry Acid and Base
Acid – Donates H+ Base – Accepts H+
Arrhenius Acid and Base
Acid – gives an H+ in solution
Base- gives an OH- in solution
FORMULAS pH = - log[H+]
MaVa (#H+) = MbVb(#OH-)
M1V1 =M2V2