Chemistry 11 Final Examination Test Notes
Matter= Anything that has mass and takes up volume
Element= A pure substance, which cannot be separated into simpler substances as a result of any chemical process
Silver (Ag), Copper (Cu), Hydrogen (H)
Atom= Smallest possible unit of an element, which retains the fundamental properties
Molecule= Cluster of 2 or more atoms held together strongly by electrical forces
Ion= An atom or molecule which possesses an electrical charge
Particle= A general time to describe a small bit/amount of matter, such as an atom, molecule or ion.
Heterogeneous or Mechanical Mixture= Substance consisting of more than one phase or material
Homogeneous or Solution= Substance of one phase or materials of the same kind (gas, in gas, liquid in gas, liquid in liquid etc)
Solvent= Substance in greater quality
Solute= Substance in smaller quantity
Anion= Negative ion
Cation= Positive ion
Naming and Formula Writing for Ionic Compounds
Polyatomic Ions – Formula to Name
1.
2.
3.
Look for the left most element on the periodic table, record it as it is named on the periodic table.
Look for the right part of the formula on the polyatomic ions sheet
Write them out together (Cation positive ion then Anion negative ion)
Eg. K2SO4 = Potassium Sulphate
*Look for a metal or NH4 (Ammonium), this means the rest of the equation is a negative polyatomic ion.
NaCH3COO = Sodium Acetate
Fe(NO3)2 = Iron (ii) Nitrate
Name to Formula
*Use parentheses to denote more than one polyatomic ion
Eg. Ammonium Oxide = NH4 + O-2  (NH4)2O
Potassium Chloride= K+ ClO-3  KClO3
Copper (ii) Dichromate= Cu2 + Cr2 + O4-2 = CuCr2O4
Ionic Compounds – Element-to-Element – Formula to Name
1. List/Name the metal as it is on the periodic table
2. Change the non-metal ending (right side) to –ide
3. Metal always comes first in this case, then non-metal
Eg. MgF2  Magnesium Fluoride
Al2O3  Aluminum Oxide
Rb2S  Rubidium Sulphide
Name to Formula
1. Metal always comes first, then the non-metal
2. Charges of the metal and non-metal must cancel out (Use Crossover Technique)
Eg. Calcium Chloride  Ca2+ Cl-1  CaCl2
Strontium Chloride  Sr2+ N-3  Sr3N2
Calcium Oxide  Ca2+ O-2  CaO NOT Ca2O2
Covalent Compounds
Note: No charges to consider
No crossover technique
No metals!
Prefix Numbers List
Mono =1
Di
=2
Tri
=3
Tetra =4
Penta =5
Hexa
Hepta
Octa
Nona
Deca
=6
=7
=8
=9
=10
Formula to Name
1. Name the left most element in the formula
2. Name the second element, ends in –ide
3. Add the prefix to each element’s name in order to indicate the number of atoms of each element in the compound.
Eg. N2O2  Dinitrogen Dioxide
*If the first element doesn’t have more than one atom, do not add a prefix
Eg. NI3  Nitrogen Triiodide
Name to Formula
1. Name the left most element of the formula by using the atomic symbol that corresponds to it on the periodic table
2. Add a prefix if there is more than one atom of this element.
3. Name the element on the right side using the atomic symbol for it on the periodic table
4. Add a prefix to denote that there is more than one atom
Eg. Carbon Monoxide  C + O  Carbon Monoxide
Diatomic Elements
These elements always exist in groups of 2 atoms and must be treated as such. This is shown as a subscript 2 beside the
diatomic element
H O F Br I N Cl
Hydrogen, Oxygen, Fluorine, Bromine, iodine, Nitrogen, Chlorine
Naming Hydrides
-When a crystal of an ionic compound is grown from evaporation from an aqueous solution, it is frequently found that the
crystal structure will include water molecules.
-Molecules, which include water molecules in their crystal structures, are called Hydrates
Prefix List denoting the number of water molecules
Monohydride
=1
Hexahydride
=6
Dihydride
=2
Heptahydride
=7
Trihydride
=3
Octahydride
=8
Tetrahydride
=4
Nonahydride
=9
Pentahydride
=5
Decahydride
=10
Eg. CuSO4  5H2O
Copper Sulphate Pentahydride
Types of Reactions
Synthesis Reaction
A reaction that involves the combination of two or more substances
A + B = AB
Decomposition Reaction
A reaction that involves breaking down a molecule into simpler substances
AB  A + B
Single Replacement
A reaction that involves replacing one atom in a compound with another
A + BX  AX + B
Double Replacement
A reaction that involves an exchange of atoms or groups between two different compounds
AB + XY  AY + XB
Special Type of Double Replacement : Neutralization
Acid + Base = Salt + Water
Combustion (Anything with H and C in it or C5H12)
Reaction of the hydrocarbon with oxygen to produce Carbon Dioxide and Water
Hydrocarbon + O2  CO2 + H2O
Balancing Chemical Equations
1. Quickly scan the equation to identify atoms which occur only in one species on each side of the equation
You must pick one of these atoms to be your starting point
2. Metal atoms are often the leader of the reaction, meaning that they should be balanced first
If there are no metal atoms, look for something other than H and O
3.
4.
5.
6.
Hydrogen and Oxygen are often found several times throughout an equation and cannot be balanced until
most of the others are
When a coefficient is put in front of a particular species to balance a particular atom, at least one other atom in the
species now has the same coefficient.
Look on the other side of the equation to see if this atom appears more than once without a coefficient in front of it
and balance it. Continue until a coefficient has been put in front of every molecule, even ones.
Try to balance entire groups such as: SO4, PO4, NO3 etc.
If an atoms is part of a diatomic element, such as O in O 2
Notice that:
2 O-atoms = 1 O2
1 O-atom = ½ O2
5 O-atoms = 5/2 O2
If a fraction occurs (Eg. ½) during the balancing, multiply the equation by the denominator (Eg. 2), which eliminates the
fraction.
Exothermic and Endothermic Reactions
Exothermic
-Gives off heat to its surroundings
-Heat is released from the reactants
Less energy to break bonds than is given off to form bonds
A + B  C + Heat
Total Heat = Negative Value
Endothermic
-Absorbs heat from its surroundings
-Heat enters the reaction instead
Takes in more energy to break bonds than is given off to produce bonds
A + B + Heat  C
Total Heat = Positive Value
H = Change in enthalpy during a reaction
Total heat contained in a system
H = HProducts + HReactants
Bohr Diagrams
First shell = Maximum of 2 atoms
Second shell = Maximum of 8 atoms
Third shell = Maximum of 8 atoms
Continue with 8 atoms for more shells
The outermost shell of electrons are called valence electrons
Example:
Neon Atom
Sodium Atom
Potassium Atom
When making Bohr diagrams for ions, we put a square bracket around the entire diagram and put the charge in the top right
corner
Example
Oxygen Ion
Potassium Ion
Lewis diagrams only show valence electrons compared to Bohr diagrams
Example
Boron
Aluminum
Use of Significant Figures
Rules for Significant Figures in General:
1. Any non-zero number ALWAYS counts
456 = 3 sig figs
71567 = 5 sig figs
2.
Zeroes to the left of a decimal SOMETIMES count
690. = 3 sig figs
690 = 2 sig figs
3.
Any zeroes IN BETWEEN count
67.0001 = 6 sig figs
7001.1 = 5 sig figs
4.
Leading zeroes are NEVER significant
0.00005 = 1 sig fig
0.00302 = 3 sig figs
Multiplying with sig figs:
1. Never round off until all you calculations are finished
2. Round off to the smallest number of significant digits
Addition and Subtraction with sig figs:
1. Use all known digits in the calculation
2. Round off to the last decimal place that both numbers have in common
The Mole
Single Step Mole Conversions
Mole to grams = Molar mass
1 mol
Mole to particles/molecules =
Mole to liters =
22.4 L
1 mol
Mass of a substance
Atomic mass of the atom
6.02 x 1023
1 mol
Used when particles of a
substance needs to be found
This is or gases only and
When the volume needs to be found
*BE CAREFUL OF DIATOMIC MOLECULES (H, O, F, Br, I, N, Cl) They are 2 times worth.
Multi Step Mole Conversions (The Mole Bridge) or Stoichiometry
a) Grams, b) Molecules, c) Particles to Moles  Moles to Moles  Moles to a) Grams, b) Molecules, c) Particles
What you’re looking for is on the top of the equation (Numerator)
After you cancel out, what you’re left with is the unit you should be looking for
Example: The combustion of Propane (C3H8) produces CO2 and H2O
C3H8 + 5 O2  3 CO2 + 4 H2O
What is the mass of CO2 produced by reacting 2.00 mol of O2?
Central Calculation:
Moles of O2 to moles of CO2
5 Mols O2 = 3 Mols CO2
ROAD MAP
First convert moles of O2 to moles of CO2
Then convert moles of CO2 to mass of CO2
Isotopes
The isotope number comes from the number of protons and neutrons
Adding a proton to an element, creates a new element
Adding a neutron to an element creates a heavier version of the element
If you want to calculate the molar mass of an isotope,
Multiply the isotope number with the percent given.
If you want to calculate the molar mass of two isotopes,
Multiply both the isotopes by their percentages respectively and add the totals together for the average mass.
Example Gallium 60 and Gallium 71
60
Ga (Gallium)
Ga (Gallium)
71
Isotope
60
71
% (Percent)
60/100
40/100
Total
36
28.4
TOTAL: 64.4 grams
Atomic Weight (Molar Mass)
*REMEMBER MOLE BRIDGE
CONVERT MOLES TO GRAMS ETC
PAY ATTENTION TO THE NUMBER OF MOLES GIVEN IN THE PROBLEM
Percentage Yield
% Yield = Mass Obtained x 100 %
Mass Expected


Usually a yield of 100% cannot be expected
Helps you find out the actual amount of the product obtained compared to the expected amount.
Find the mass expected from the general equation after the mole bridge, conversion etc
Then put the obtained mass from the equation over the expected mass. Multiply that by 100% to get a percentage
Limiting and Excess Reagents
Limiting reagents set a limit on the amount that can be formed
Excess reagents are when there’s some substance left over
Coordination Compounds
-Transition metals have many electrons available to form bonds; therefore they are able to form coordination compounds
-Transition metals have 2 Valencies
Primary Valency (Combining Capacity, Oxidation State)
How many bonds are formed when the Transition metal forms with oppositely charged ions
Secondary Valency (Coordination Number)
The number of bonds between a transition metal and any surrounding ion, molecule or atom
General Structure
1. Metal ion
Transition metal ions have many possible combining capacities
2. Ligands
Ions, atoms or molecules bonded to the central metal atom or ion
Around 2-8 ligands in different shapes and ways
3. Counter Ion
The ion of an opposite charge that makes up the compound to produce a charge of zero
Rules for Naming Coordination Compounds
Eg. [CuCl2(en)2]Br)
Cu Cl
en Br
Oxidation state= X + (-1)2 + 0 + (-1) = 0
X-3=0
X=3
1. When naming any ionic compound, the Cation (Positive Metal) is named before the anion (Negative Metal)
Br- = Charge of -1 which is why it is in the back
2. When naming a complex ion, ligands are named before the metal ion in alphabetical order
Chloro ethylene diammine (Cl + en)
3. In naming ligands
Anionic Ligands= an “O” is added to the name (Eg. Fluoro, Cyano, Hyroxo, Chloro)
Neutral ligands= The name of the ligand is used
4. Make sure to use prefixes of di, tri, tetra etc. for ligands to denote the number of ligands
Mono = 1 Di = 2 Tri = 3 Tetra = 4 Penta = 5 Hexa = 6 Hepta = 7 Octa = 8 Nona = 9 Deca = 10
If the ligand already contains a prefix, or it is a polydentate, use Bis-, Tris-, Tetrakis- etc.
5. The oxidation state of the central metal ion is designated by a roman numeral in brackets
Dichloro bisethylenediammine Copper (iii)
6. When more than one type of ligand is present, they are named alphabetically (Prefixes don’t affect the order)
7. If the complex ion is “-ve” or an anion, the suffix “-ate” is added to the name of the metal, or if the metal has a latin
name
8. Name counter ion
Dichloro bisethylenediammine Copper (iii) Bromide
*Be sure to add ion to the back of the entire name if there is a charge in the original unnamed formula
Eg. [Ru(NH3)5Cl]2+ = Pentaammine ChloroRuthenium (iii) ion
Eg. Fe(CN)6-4 = HexacyanoFerrate (ii) ion
Eg. [Co(NH3)5Cl]Cl2 = Pentaammine ChloroCobaltChlorate (iii)
Generalized rules for naming coordination compounds
1.
2.
3.
4.
5.
6.
7.
Identify the central metal ion
Identify the charge on the central metal ion (Parentheses)
Identify the ligands
Calculate the total charge on the ligands
Calculate the total charge on the complex ion
Write the formula with central metal ions first, then ligands
If more than one ligand is present, anion ligands are given before neutral ligands
Isomers
General
Formula
CM + 2L
Coordination
Number
2
CM + 3L
Geometry
Examples
Notes
Name
-Not Common
-M+L are on the same plane
-Ligands are 180 degrees
Linear
3
-Not Common
-M+L are on the same plane
-Ligands are 120 degrees
Trigonal planar
CM + 4L
4
Tetrahedral
CM + 4L
4
-Very Common
Complex anions + halides
-Ligands are 109 degrees
-No chance of geometric
isomers (Cis/Trans)
M+2L are on the same plane
-Common for Nickel (ii) and
Platinum (ii)
-M+4L’s on the same plane
-Ligands are 90 degrees
-Geometric Isomers are
possible
Square Planar
CM + 5L
5
-Rare
-M+3L’s are on the same
plane
Trigonal
Bipyramidal
CM + 6L
6
-Most Common
-Ligand bonds are 90
degrees
Octahedral
There are 2 types of General Isomers
Structural Isomers (Different bonds)  Coordination isomers and linkage isomers
and Stereoisomers (Same bonds, different arrangement)  Geometric isomers and Optical isomers
Coordination isomers
The composition of the complex ion varies
Linkage Isomers
The complex ion is the same in structure, but the point of attachment of 1 or more ligand(s) differs
Geometric Isomers
Cis/Trans isomers
Occur as square planar and octahedral complexes
Cis = side by side
Trans = across each other
Octahedral
Optical Isomers
Mirror images of each other
“Left-Right Gloves” also called “Enatiomers”
Organic Chemistry
Alkanes
Single bonded Carbon and Hydrogen atoms
Have the general formula of: Cn + H2n+2 (Eg. C2H6)
All bonds follow the tetrahedral arrangement
109 degrees
C1 to C4 = Gases
C5 to C15 = Liquids
C16 and above are solids (Waxes or Parrafins)
Structure and formula – 3 ways to represent a hydrocarbon
Eg. Propane (C3H8)
Full Structure
Condensed Structure
Molecular Formula
Since each carbon is bonded to the maximum number of Hydrogens available, they are known as Saturated Hydrocarbons.
Naming typical compounds – Alkanes
1. Select the longest continuous chain of carbon atoms for the basic name
2. Number the carbons in the chain from the end nearest the first branch (to give the lowest set of numbers)
3. Identify substituents. If there are more than on substituents of the same kind present, use the prefixes “di”, “tri”,
“tetra” etc.
4. Locate the substituents by the number of the carbon to which they are attached
5. Put substituents in alphabetical order
6. Separate numbers from letters by a dash and write the whole name as one word with the basic name at the end
Naming Unbranched Alkanes
1. Count the number of carbon atoms in the carbon backbone. Use a prefix to denote the number counted
2. Add the ending “ane” to the end Eg. CH4 = Methane
C5H12 = Pentane
1= Meth
2= Eth
3= Prop
4= But
5= Pent
6= Hex
7= Hept
8= Oct
9= Non
10= Dec
Naming Side Chains
Some hydrocarbons have side chains = Alkyl Groups
Alkyl Groups= An alkane that has lost one hydrogen atom. To name an alkyl, use the prefix of the number of carbon atoms and
add the ending “yl”
Eg. CH4  CH3 = One Carbon  Methyl
Functional Groups
Specific groups can be found on organic molecules, known as functional groups
Allows for:
Acids, bases, both
Pleasant or yucky smell
Saturation of molecules helps to identify possible shapes
Naming Branched Alkanes
1. Identify the longest continuous chain of carbon atoms (Parent Chain)
2. Identify all the side chains that branch off the parent chain
3. Number the carbons on the parent chain consecutively that gives us the lowest possible number to the side chain
4.
5.
Starting with the first side chain in alphabetical order, name the carbon number the alkyl group is attached to followed
by a dash and then the name of the alkyl group (Eg. 2-methyl)
If you have more than one of the same alkyl group, list each carbon number where the repeated group is attached,
separated by commas, with a dash, and ad a prefix to the alkyl group to show how many identical groups are attached
(Eg. 2,3 – dimethyl)
Naming Alcohols
1. Select the longest continuous chain of carbon atoms which contains the carbon bearing the hydroxyl group
2. Change the ending to “-ol” from the “-ane” of the corresponding alkane
3. Umber the carbons in the chain from the end nearest the carbon bearing the –OH
4. Locate the position of the –OH by the number of the carbon to which it is attached
5. Name substituents as with alkanes
6. Hydroxyl (-OH) groups have priority over double bonds in terms of numbering
Naming Carboxylic Acids
1. Select the longest continuous chain of carbon atoms which contains the carbon bearing the carboxylic grouo.
2. Change the ending to “-oic” from the “-ane” of the corresponding alkane
3. The carboxylic carbon atom is assigned number 1
4. Locate and name substituents as with alkanes.
Naming Alkynes
1. Select the longest continuous chain of carbon atoms which contains both carbons of the triple bond for the basic name.
2. Change the ending to “yne” from the “-ane” of the corresponding alkane
3. Number the carbons in the chain from the end nearest the first carbon of the triple bond
4. Locate the position of the triple bond by the number of the first carbon involves in the triple bond
5. Name substituents as with alkanes.
Naming Alkenes
1. Select the longest continuous chain of carbon atoms which contains both carbons of the double bond for the basic
name.
2. Change the ending to “-ene” from the “-ane” of the corresponding alkane
3. Number the carbon atoms in the chain from the end nearest the first carbon of the double bond
4. Locate the position of the double bond by the number of the first carbon involved in the double bond
5. Name substituents as with alkanes
6. In cyclic system number the carbons of the double bond as 1 and 2
Esters
End in “-oate”
Made of alcohol and carboxylic acid
Name
Alkene
Alkyne
Halide
(F, Cl, Br, I)
Alcohol
“ol”
Aldehyde
“-al”
Ketone
“-one”
Functional group
Name
Ether
“-oxy”
Amine
“-amino”
Amide
“-amide”
Carboxylic acid
“-oic acid”
Ester
“-oate”
Aromatic Ring
Alkyl Halides
These compounds follow the same rules as branched alkanes except:
The halogen groups F, Cl, Br and I are named as:
Fluoro, Chloro, Bromo, Iodo
Functional group
Structural Isomers
Both branched and unbranched hydrocarbons exist when given a molecular formula. Both of these have the same structure, but
different arrangement of atoms. They are called structural isomers.
Cycloalkanes
Hydrocarbons can exist in Cyclic (ring) forms
These are known as Cycloalkanes (Cyclic Hydrocarbons)
Cycloalkanes have the general formula of CnH2n
Naming Substituted Cycloalkanes
A single substituent does not require a number to indicate the position of attachment. All carbons are considered identical.
If more than one substituent is present, the first substituent is assumed to be Carbon “1” and the remaining substituents are
numbered either clockwise or anticlockwise to have the lowest set of overall number values.
Trends and properties of the Groups in the Periodic Table
Group 1: Alkali Metals
-1 electron on the outermost shell
-More reactive than Alkali Earth Metals
-Colorless ions formed
-Ions have a 1+ charge
-Lower melting, boiling point
-Fewer properties than Alkali Earth Metals
-Soft enough to cut with a knife
Group 2: Alkali Earth Metals
-2 electrons on the outermost shell
-Generally insoluble
-Less reactive than Alkali Metals
-Lose 2 electrons per atom
-Ions have a 2+ charge
-Not soft enough to be cut with a knife
-Generally harder
Groups 3-12: Transition Metals
-Good conductors of heat + electricity
-Higher boiling point and melting point than group 1 and 2 elements
-Form brightly colored compounds
-Transition metals have multiple oxidation states or valencies
-Transition metals form complex ions and coordination compounds
Group 13: Boron Family
-Naturally found in their elemental form
-Atomic Radius increases
-Ionic Radius increases
-Ionization energy decreases
-Boron is derived from Borax
-The elements in this family are generally scarce, except for Aluminum, which is abundant
Atoms have 3 electrons on the outermost shell
Group 14: Organic Chemistry (Carbon Family)
-Anything that contains Carbon in it is considered Organic
-Carbon is a non-metal, Silicon and Germanium are metalloids, and Tin and Lead are metals
-They have 4 electrons in their outermost shells
-Tend to be unreactive and form covalent compounds
-Atomic Radius increases
-Ionization energy decreases
-Melting point decreases
-Boiling point decreases
Group 15 Nitrogen Family
-Atomic radius increases
-Ionic radius increases
-Ionization energy decreases
-Electronegativity decreases
-Boiling point increases
-Melting point jumps from low to high
-Nitrogen is most commonly found as atmospheric gas
-Range from very abundant to very rare
-Consist of 5 electrons on their outermost shell
-Oxidation numbers of +3 and +5
Group 16 Oxygen Family
-Atomic radius increases
-Ionic radius increases
-Ionization energy decreases
-Electronegativity decreases
-Boiling Point increases
-Melting Point increases
-Density increases
-Occur naturally as elements or in a combined form
-Oxidation number of -2
-Exist in allotropic forms
-Consist of 6 electrons in their outermost shell
-Oxygen is found as a clear, colorless, odorless gas at room temperature, very reactive with most other elements
Group 17 Halogen Family
-Are all non metals and found in a combined form
-Atomic radius increases
-Ionic radius increases
-Ionization energy decreases
-Electronegativity decreases
-Found in the Earth’s crust and in sea water
-Abundant to rare
-Consist of 7 electrons on their outermost shell
-Reactive
-Tend to gain one electron but also share electrons and have positive oxidation states (We’ll use -1)
-Boiling point increases
-Melting point increases
-Density increases
Lewis Structures
Simple Ionic Compounds
-Determine the charge expected for each atom
-Arrange the non-metal atom symmetrically around the atom
-Fill in the valence electrons for each atom
-Remove the electrons from the outer shell of the metal atom to form the ion
-Distribute the electrons equally to each non-metal atom to form the ion
Metal ion gives away electrons!
Covalent Compounds – Two or more gases, No metals
-Count the total number of valence electrons for each atom in the molecule
-Adjust the total number by subtracting an electron for every positive charge in the molecule/add an electron for every
negative charge
-Determine which atoms are bonded together and show this bond by drawing a line connecting 2 atoms.
This bond reflects two electrons
-Subtract the number of bonding electrons from the total # of valence electrons
-Place the remaining valence electrons on the terminal atoms first to complete the octet
-Place all left over electrons around the central atom if necessary
-If a central atom has less than an octet of electrons, the neighbor can share electrons
Nitrogen Cycle
The Nitrogen cycle is important because:
-Plants depend on the nitrogen cycle; plants keep the world and us alive
-Nitrogen also equals protein, no nitrogen equals no life
-Ammonia helps promote plant growth and is made directly from the air + Fertilizers replenish nutrients
≥ Made by the Haber Process from Hydrogen and Nitrogen at a high temperature and high pressure in the
presence of a catalyst
-Many countries depend on a limited source of Ammonia to keep food supplies up
-Ammonia synthesis = basis of Nitric Acid (HNO3) used in making explosives in WW1 and WW2
Nitrogen Fixation
The process of N2 becoming NH3 (Peas, Clovers, Alfalfa and Locust Trees)
Nitrogen fixing bacteria = NH3 to Nitrate (NO3) (Wheat, Rice, Corn and Potatoes)
Fertilizers are made from NH3 (Ammonia), obtained from Sodium Nitrate (NaNO3)
Ammonium Sulphate (NH4)2SO4
Ammonium Nitrate (NH4)NO3
Urea (NH2)2CO
Fertilizers with Nitrogen, Phosphorous and Potassium = Complete Fertilizers
Nitrogen stimulates overall plant growth
Phosphorous promotes root growth and flowering
Potassium regulates structures in leaves that allow CO2 to enter and H2O + O2 to leave
Regulates production of carbs and sugar
Proportion and percentage on package called N-P-K formula
More sugar= K
Grain = P
Lawns = N in spring and P in winter
Calculations for Groups 15 and 16
Percentage Composition
-If compounds have the same elements, you can use percentage composition to differentiate them
-Write out the compound given
-Separate it into its separate elements
-Find each element’s mass (Listed on the periodic table under their atomic symbol)
-Multiply the masses according to how many molecules are present in the compound
-Add together all the masses for a total mass to divide with
-Divide the masses of each element separately with the total mass to get the percentage
Molarity
EQUATION:
Molarity (Concentration) equals moles per liter
Example: What is the molarity of a solution containing 0.20 mol of NaCl in 120.0 ml of water?
How many moles of Zinc Chloride (ZnCl) are in 200.0 ml of 0.050 M solution
Dilution
50 ml of water is added to 100 ml of NaCl at 0.1M. What is the concentration of NaCl after dilution?
A student adds 25.0 ml of H2O to a 65.0 ml sample of 3.0 ml. What is the new [HCL]?
Double Dilutions
What is the [KOH] when you combine Solution A of 80.0 ml, O.15M KOH and Solution B of 50.0 ml, 0.25M KOH together?
Ions in Solution
a. L of 0.40M CaCl2 are mixed with 1.5 L of 0.2M NaCl. Find the [ ] of the ions.
Titrations
-Based on neutralization reaction
Reaction between an acid and a base to produce a salt and water
-Acids are compounds that begin with “H” – Hydro
-Bases are compounds that end with “OH” –Hydroxides
NaOH = Sodium Hydroxide
(NH4)OH = Ammonium Hydroxide
Ca(OH)2 = Calcium Hydroxide
Al(OH3) = Aluminum Hydroxide
-A salt is neither an acid nor a base
-Titration is the process of finding the unknown concentration of a chemical in a solution based upon the results of the
neutralization reaction
A solution with a known concentration is reacted with another of an unknown concentration until an equivalence point
or end point is reached
The end point = when all the moles of the acid have reacted with all the moles of the base
Titration Concentration Problem Example
Initial Volume
Final Volume
Volume Used
Average Volume
Rough
0
11.6
11.6
Test 1
12.0
23.2
11.2
Test 2
25.0
36.6
11.6
11.25
Test 3
30.0
41.3
11.3
You use 11.2 ml in Test 1 and 11.3 ml in Test 3 to calculate the average
volume because they are 0.2 ml apart.
Sodium Hydroxide [NaOH] was titrated with Sulphuric Acid (H2SO4) at 0.11M, 10.0 ml. Calculate the concentration of NaOH
Grams to Moles
If 2.5g of NaCl are dissolved in 50 ml of water, what is [NaCl]?
Electron Configuration
2 electrons in an s-type subshell
6 electrons in an p-type subshell
10 electrons in a d-type subshell
14 electrons in an f-type subshell
1.
2.
3.
4.
Write out the full configuration of energy levels
Find the noble gas closest, but lower than the atomic number given (Noble gas in the row before)
In the full configuration, draw a line at where the noble gas configuration ends
Write the noble gas atomic symbol (Ne, Ar, Kr, Xe, Rn) in brackets and write the remainder of the configuration
Example: Titanium 22
1.
2.
3.
4.
Electron Configuration for Ions – Positive Ions
1.
2.
3.
4.
Write out the starting atom outside the bracket on the left (Don’t write the charge)
Then write out the core notation of the atom inside the bracket
Write the charge in negative form in electrons in front of the whole bracket
Cross out all other shells in the bracket if any, leaving only one
Cross out in this order ( P  S  D)
Write out what’s left without the negative electron(s)
Example Tin 50
1.
2.
3.
4.
Negative Ions
1.
2.
Write out the notation like in step one ^ with core notation, add an electron depending on the charge to fill up the shell
Write the final answer with the electron number in the place of the charge on the atom and add it to the front orbital
Example Oxygen 8
1.
2.