chemistry exam review package
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Chemistry Exam Review
By Sagar Arenja
Unit 1:............................................................................................................................................ 2
Periodic Table.............................................................................................................................2
Model of the Atom.....................................................................................................................3
Atomic Mass...............................................................................................................................4
Isotopes and Radioisotopes....................................................................................................4
Trends........................................................................................................................................ 4
Types of Bonds...........................................................................................................................5
Lewis Structures.........................................................................................................................6
Naming and Formulas................................................................................................................9
Unit 2........................................................................................................................................... 11
Chemical Reactions..................................................................................................................11
Types of Reactions................................................................................................................... 12
Unit 3....................................................................................................................................... 14
Moles....................................................................................................................................... 14
Percent Composition................................................................................................................14
Empirical vs. Molecular Formula..............................................................................................14
Stoichiometry...........................................................................................................................16
Limiting Reagent...................................................................................................................... 17
Percent Yield............................................................................................................................ 17
Empirical Formula Problem..................................................................................................17
Molecular Formula Problem.................................................................................................18
Limiting and Excess Reactant Problem.................................................................................18
Formulas.................................................................................................................................. 18
Unit 4........................................................................................................................................... 19
Solutions and Solubility............................................................................................................19
Solutions...............................................................................................................................19
Solubility...............................................................................................................................19
Concentration.......................................................................................................................21
Preparing Solutions.............................................................................................................. 21
Dilutions............................................................................................................................... 22
Ionic Equations.....................................................................................................................22
Water Quality and Purification.............................................................................................23
Stoichiometry in Solutions....................................................................................................23
Acids and Bases....................................................................................................................23
pH and pOH..........................................................................................................................24
Acid-Base Titration............................................................................................................... 25
Unit 5........................................................................................................................................... 26
Gases.................................................................................................................................... 26
Gas Laws...............................................................................................................................26
Molar Volume/ Ideal Gas Law..............................................................................................27
Calculations Involving Reactions producing Gases...............................................................28
Unit 1:
Periodic Table Elements are named according to IUPAC – International Union of Pure and Applied Chemistry Johann Dobereiner (1780-1849) noted the similarity among physical and chemical properties of
several groups of three elements in 1829. He called them triads (Lithium, Potassium, and Sodium are one triad). This is referred to as the law of triads even though it isn’t an actual law.
John Alexander Newlands (1837 – 1898) arranged elements in order of increasing atomic mass and noticed repeating patterns of physical/chemical properties every 8th element. It was called the law of octaves but it was not accepted.
Julius Lothar Meyer (1830 – 1895) arranged elements in order of atomic mass and found a repeating pattern in the relative volumes of individual atoms and he observed a change in length of that repeating pattern. By 1868, he had developed a tale of the elements that closely resembles the modern Period Table.
Dmitry Mendeleev published the first periodic law. It stated that the elements arranged in order of increasing atomic mass show a periodic recurrence of properties at regular intervals. He listed the elements in order of atomic mass in vertical columns which showed that when a column ended, the chemical properties would repeat themselves, at which point a new column was started.
He revised his P.T in a horizontal method that elements with similar properties were in the same column/group. This made it clear there were periodic trends. There were blank spaces that he left for future discovered elements.
Periodic Law (modern) – When elements are arranged in order of increasing atomic number, their properties show a periodic recurrence and gradual change
Group – elements with similar chemical properties in a vertical column in the main part of the table ; Period – elements arranged in a horizontal row, whose properties change from metallic on the left
to nonmetallic on the right
Model of the Atom Empirical knowledge – knowledge coming directly from observations Theoretical Knowledge – knowledge based on ideas created to explain observations – models,
analogies, words, symbols. Democritus all matter could be divided into smaller pieces until a single indivisible particle was
reached (atom) which was in constant motion, has different sizes, and regular geometric shapes with space in between.
Aristotle said continuous motion was illogical. Supported the four element theory, that all matter was made up of fire, air, water, and earth.
Dalton’s Atomic Theory o all matter is composed of tiny, indivisible particles called atomso All atoms of an element have similar propertieso Atoms of different elements have different properties
o Atoms of two/more elements can combine to form new substanceso When atoms join or separate, they are not destroyed.
Law of Conversation of mass – matter is neither created nor destroyed Law of constant composition – compounds always have the same percentage composition by mass J.J. Thomson used a modified cathode ray tube to measure the mass of the particle and its electric
charge and discovered atoms. created a model saying negatively charged electrons are distributed inside the atom
Ernest Rutherford shot alpha particles (particles produced by radioactive decay) through very thin pieces of gold foil which led to his hypothesis of the nucleus, a positively charged core, and that it was surrounded by electrons. He gave the positive charged particles in the nucleus the name proton.
James Chadwick demonstrated the existence of neutrons in the nucleus Bohr’s Theory electrons within an atom possess discrete energies called energy levels and that
electrons travel along a pathway called the shell/orbito He theorized that if additional energy is supplied to an electron in a given energy level, then
the electron can jump to a higher, unfilled energy level farther away from the nucleus (transition). The quantity of energy required for the jump is equivalent to the difference in energy between the energy levels. When an electron jumps back, it releases the equivalent amount of energy.
o Lowest energy ground stateo Line spectrum a pattern of distinct lines, each of which corresponds to light of a single
wavelength.o Continuous spectrum pattern of colours observed when a narrow beam of white light is
passed through a prism or spectroscope.o Energy Levels:
6th down to 2nd violet 5th – 2nd indigo 4th – 2nd green 3rd – 2nd red
Quantum Mechanics: a theory of the atom in which electrons are described in terms of their energies and probability patterns
Due to Quantum Mechanics, you cannot know the precise position and motion of an electron at the same time.
Electron cloud: the region of an atom in which electrons are most probably located.
Atomic Mass Atomic number (z) refers to the # of protons in the nucleus and determines the identity of an
element. H.G.J. Moseley discovered this Mass number (A) – the sum of the number of protons and neutrons Number of neutrons = mass number – atomic number N = A – Z
Isotopes and Radioisotopes Isotopes: Variations of the number of neutrons from atom to atom within the same element Radioisotopes: Isotopes that decay and become radioactive
o Radiate alpha, beta and gamma particles Half-Life : Al = Ao x (1/2) t/T Carbon-14 dating: a technique that uses radioactive carvon-14 to identify the date of death of
once-living material
Trends Shielding : Electrons that are in inner shells “shield” the valence electrons from the force of
attraction exerted by the positive charge in the nucleus. Metals
o Reactivity decreases left to right its easier to lose less electronso Reactivity increases top to bottom less attraction between nucleus and valence
electrons Non-Metals
o Reactivity increases bottom to top there is more attraction as shells decrease allowing for easier attraction of electrons (fewer shells, the closer electrons are to the nucleus [positive charge], and so attraction is stronger.)
o Reactivity increases left to right more protons, more attraction Atomic Radius
o ½ the distance between nuclei of two identical adjacent atomso Atomic Radius decreases left to right more attraction as the protons attract
electrons closer.o Atomic Radius increases top to bottom more shells
Ionic Radiuso When ions form, nuclear (positive) charge always stays the sameo An atomo An atom’s valence electrons repel each other increasing atomic radiuso Add an electron (anion), increase repulsion, increase radiuso Remove an electron (cation), decrease repulsion, decrease radius
Ionization Energyo Energy required to remove an electron from an atomo Increases left to right there are more unshielded protons to attract and hold the
electrons Electron Affinity
o Energy released when an atom gains an electron (how much an atom is willing to pay to get an electron)
o Increases left to righto Increases bottom to top
Electronegativity
o Strength of attraction an element has for electrons within an intramolecular covalent bond
o Increases left to righto Increases bottom to top
Types of Bonds Ionic Bonds
o Occur between a metal and a non-metalo Involves transfer of electrons from less electronegativity (EN) atom to more EN atomo Occurs when different of EN is >1.7o Charged ions are attracted to each other due to being oppositely charged; Electrostatic
attractiono Ions form when an atom gains or loses an electrono Examples: NaCl, MgCl2
Properties of Ionic Bondso High Intermolecular forces (forces between molecules) which causes them to be solid at
room temperature and have high melting pointso Low intramolecular forces (forces within molecules), resulting in compounds breaking
apart in water to form electrolyteso Solid’s particles form crystal lattice structure (repeating pattern of particles)
Covalent Bondso Occurs between two non-metalso Involves sharing of 1,2, or 3 pairs of electrons between two atomso Occur when EN < 1.7o Examples: H2O, NH3
o Multiple Covalent Bonds: sharing of 2 or more pairs of electrons (CO2)o Polar and Non Polar Covalent Bonds
Polar: When EN difference is >0 and <1.7 results in inequal sharing and partial charges on atoms one end of molecule is slightly positive and one end is slightly negative
Non-Polar: When EN difference is 0 equal sharing of electronso Coordinate Covalent Bonds
Bond in which both of the shared electrons come from the same atom Many polyatomic ions contain this You can’t tell where electrons originated from once bond is formed NH4
+ and SO42-
Polar Molecules: a molecule that is positively charged at one end and negatively charged at the other end due to unequal sharing of electrons in a covalent bond
o Must contain polar covalent bondso However, if molecule is symmetrical in shape, charges cancel out and molecule is non-
polar
Properties of Ionic/Covalento Ionic
Electrostatic attraction ionic bond Forms crystal lattice Melting Point: in order to melt an ionic compound, you need to break ionic
bonds high intermolecular forces high melting point Volatility: particles in a volatile compound must be held together by weaker
forces so that some can break away and travel through the air to our noses (waft it we get whole molecules in our noses) high intermolecular forces low volatility
Solubility Ionics tend to be soluble in water because water is a polar compound that can exert enough force to overcome ionic bonds and cause the ions to go into solution. low intramolecular forces soluble (they form electrolytes)
Conductivity Ionics have conductivity because ionic compounds as well as bases form electrolytes in aqueous solutions
o Covalent Covalent compounds exist in the form of molecules which are held together in
clusters by weak forces (intermolecular forces) Melting Point: Low Intermolecular forces low melting point Volatility: Low intermolecular forces, high volatility Solubility: High Intramolecular forces low solubility Conductivity: Molecular compounds no conductivity
(exceptions include: acids which form electrolytes in aqueous solutions conductivity)
Lewis Structures Valence electrons are found in the outermost orbit Lewis structures show atoms with their valence electrons Rules:
Symbol of the element is in the center (represents P, N, and inner E) Electrons in outermost energy level are placed as dots on 4 sides of symbol Dots are placed singly, then paired up Octet Rule: 8 valence electrons Ions:
Cation (positive charge lost electron) don’t draw electrons around the symbol. Put the symbol inside [ ] and write the charge outside in the top right.
Anion (negative charge gained electron) place electrons around the symbol. Put the symbol in a [ ] and write the charge outside in the top right
A line between two elements indicates a covalent bond Lewis structures for polyatomic:
Most polyatomics have co-ordinate covalent bonds
1. Add up the total number of valence electrons 2. Draw the skelton structure (least EN element is in the center) 3. Distribute valence electrons around atom until each atom has a complete
valence shell (8 electrons except Hydrogen which has 2) 4. Ensure all electrons have been used up and all atoms have 8 electrons (except
H) 5. Determining charge:
Divide covalent bonds, distribute electrons. Compare total electrons around each individual atom from Lewis
Structure to that of neutral atom. Determine net formal charge of polyatomic ion.
Some atoms require double bonds to meet the octet rule (CO2) If you have to many electrons in your diagram, eliminate unpaired
electrons and add double bonds. Check central atom and see if it has an octet. If not, move lone pairs of
electrons from surrounding atom to form double/triple covalent bonds
Naming and Formulas Ionic Compounds
o Formulas Write the symbols of elements with the metal first, then the non-metal Write the ionic charge (valence) of each element on the top right, for each
element Criss-cross the #’s down to become subscripts on the opposite element Cancel to the lowest ratio by dividing by the largest common factor
o Names Metal first, then the non-metal First element’s name doesn’t change, second element ends in –ide Multivalent Metals:
IUPAC system: Write the name of the metal followed by its charge in Roman numeral in brackets. Ex. Cu (II) copper 2+
Classic System:
Polyatomic (oxyanions)
Name Formula Charge/Valence
Carbonate CO32- 2-
Nitrate NO3- 1-
Phosphate PO43- 3-
Sulphate SO42- 2-
Hydroxide OH- 1-
Chlorate ClO3- 1-
o Different forms of Oxyanions:
o Formulas
Write the symbols of the elements with the metal first, then the radical(polyatomic)
Write the ionic charge (valence) for each Criss-cross the #’s down to become subscripts and place brackets around the
polyatomic ion if there is more than one and therefore it must have a subscript. Cancel to the lowest ratio
o Names The metal is first, then the polyatomic ion.
Hydrated Salts (salts with water molecules associated with each formula unit)o Ex: Copper (II) sulphate * pentahydrate Cu(SO4) * 5H2O
Acid Salts (salts whose anions contain one or more covalently bonded hydrogen atoms)
o
Acidso Binary Acids: acids that contain hydrogen and one other element (HBr, HCl)
Start the name with “hydro” and connect the name with the name of the second element changing the element to ic
HBr Hydrobromico Oxyacids:
If oxyanion ending is ate, ending is changed to ic If oxyanion ending is ite, ending is changed to ous
o
Molecular Compounds (formed when electrons are shared between non-metals making the compound)
o Formulas Write the symbols first – Put the Non-metal on the left side of P.T first Put combining capacities as superscripts Criss-cross superscripts to get subscripts for the formula Reduce subscripts to lowest form and exclude 1 in formula.
o Naming Change the name of the second non-metal to ide If the first non-metal has more than one atom, indicate it with a prefix Indicate the number of atoms of the second non-metal with a prefix.
Unit 2
Chemical Reactionso For a chemical reaction to occur, a chemical change must occur
o Chemical reactions are accompanied by changes in energy Reactions that release energy exothermic Reactions that absorb energy endothermic Reactions that occur immediately when two substances are mixed together
spontaneous reactionso For some reactions you need to input activation energy
o Ways to tell a chemical reaction occurred: Colour change Gas formation Precipitate formed Energy (heat/light) change New products are formed ** Need a chemical test to show a new compound has formed**
o Writing Chemical Formulaso Reactants Products (law of conversation of matter atoms cannot be created nor
destroyed)o 5g of A + 5g of B 10 g Producto Equation must be balanced and have states in subscripts: solid (s), aqueous (aq), liquid
(l), and gas (g)o Word equation just use chemical names (no balancing)
Water + carbon dioxide glucose + oxygeno Chemical Equation
Skeletal equation (unbalanced) Ex. H2O (l) + CO2 (g) C6H12O6 (aq) + O2 (g)
Balanced Equation Ex. 6H2O (l) + 6CO2 (g) C6H12O6 (aq) + 6O2 (g)
How to Balanceo 1. Start with Word Equationo 2. Write the Skeleton Equationo 3. Balance elements in compounds one at a time (leave H and
O). Start with an element that is only in one reactant and one product
o 4. Balance H and O last –trickyo If polyatomics are the same on both sides of the equation count
them together
Types of Reactionso Synthesis: When two substances combine to form a more complex substance
A + B AB Non Metal Oxides (also called Acidic Oxides) combine with water to form acids:
CO2 + H20 H2CO3
Metal Oxides (also called Basic Oxides) react with water to form bases: Li2O + H2O 2 LiOH
Acidic oxides and basic oxides react to form a salt containing an oxyanion: CaO + CO2 CaCO3
o Decomposition: Reverse of synthesis, when a more complex compound is broken down into two or more simpler substances
AB A + B Thermal Decomposition (heating solid substances) Ex. 2 HgO 2Hg + O2
Acids don’t usually decompose but carbonic acid does (Ex. H2CO3 CO2 + H2O Most bases decompose into a basic oxide and water (Ex. Ca(OH)2 CaO + H2O) Salts with oxyanions can decompose to form acidic and basic oxides
(Ex. CaCO3 CaO + CO2) Sometimes hard to predict (NH4NO3 N2O + 2H2O Catalysts can be used (placed above the )
o Single Displacement: one element replaces another element in a compound A + BC AC + B Determined by the activity series
Metals higher on the activity series can displace metals lower on the activity series
o Double Displacement: Reaction between two soluble ionic compounds where cation from species changes places with cation from other species.
AB + CD AD + BC Produces a precipitate Uses Solubility Chart Reactions producing Gas
o Combustion: The reaction of a substance with oxygen producing oxides, heat and light Hydrocarbons: substance composed only of carbon and hydrogen
Can undergo complete and incomplete combustion dependent on amount of oxygen present
Lots of oxygen complete combustion CxHy + O2 CO2 + H2O Colour of the flame indicates whether it is complete/incomplete
o Yellow more oxygen neededo Blue complete combustion
Incomplete Combustion Too little oxygen Yellow flame Produces carbon monoxide (toxic) along with carbon dioxide (CO2) and
elemental carbon called soot (fine particles consisting mostly of carbon).
Carbon Monoxide: Colourless, odourless, and tasteless gaso Shaped like an oxygen molecule allowing it to bind to blood
reducing the amount of oxygen getting to the body
o Initial symptoms of CO – headache, nausea, dizzinesso High exposure can result in vomiting, collapse, loss of
consciousness and eventual suffocationo Neutralization: Special type of double displacement reaction in which an acid and a base
combine to form water and an ionic compound (salt) The Hydrogen ions from the acid join with the hydroxide ions of the base to
form water.
Unit 3
Moleso 1 mole = 6.02 x 1023
o N = n x NA (N = number of entities, n = moles, NA = Avogadro’s constant)o Avogadro’s Constant (NA) = 6.02 x 1023 it is usefule becomes scientists have
experimentally determined that there are 6.02 x 1023 carbon atoms in 12 of 12C Therefore 1 mol of C atoms = 12 g 1 C atom = 1.m.u
o Atomic Mass vs. Molar Masso Atomic Mass: the mass of an atom expressed in atomic mass unitso Molecular Mass: The sum of the atomic masses of the atoms in a moleculeo Molar Mass: The mass in grams of 1 mole of atomso Average Atomic Mass = Molar Mass
Na (Sodium) 22.99 a.m.u 22 g/molo n = m/M (n = moles, m = mass, M = Molar Mass)o Two step calculations:
Given the # of particles, the mass can be determined Particles moles Mass
Given the mass, the # of particles can be determined Mass moles particles
Percent Compositiono Instead of stating actual masses of each element in a compound, it is often more useful
to express the composition in terms of percentageo Percent composition of a compound is always the same regardless of the mass of a
sample 1. Find the molar mass of the substance first 2. Now find the percentage of each substance using the formula % composition = melement/ mmolecule
Empirical vs. Molecular Formulao Empirical Formula: represents the simplest whole number ratio of ions or atoms in a
compound
o Molecular Formula: refers to the exact number of atoms covalently bonded to form a single molecule.
Two different compounds can have the same empirical formula Molecular formula = n x empirical formula (where n = anything [1,2,3, any whole
#]) When given percent composition of a compound, assume 100g
What is the empirical formula of a compound that is 17.6 % H and 82.4 % N? Assume you have 100 g so 17.6 g is H and 82.4 g is N.
Steps to doing empirical formula 1. Assume 100g of the sample 2. Convert each mass to moles 3. Determine the lowest whole # ratio by dividing both molar amounts
by the lowest molar amount 4. If you end up in a decimal that’s really close to a whole #, round up to
a whole #o However, if the decimal is not close use this chart
0.125
1/8 Multiply by 8
0.25 1/4 Multiply by 4
0.33 1/3 Multiply by 3
0.375
3/8 Multiply by 8
0.5 1/2 Multiply by 2
0.625
5/8 Multiply by 8
0.66 2/3 Multiply by 3
0.75 3/4 Multiply by 4
0.875
7/8 Multiply by 8
o You can check your solution by working backward (IE. Find Percent Composition to see if it matches)
o Molecular Formula Molecular formula = n x empirical formula Molar mass of a compound = n x molar mass of empirical formula
Molar Mass of the compound must be given because it is determined experimentally
N = molar mass of compound/ molar mass of empirical formulao Finding Empirical Formulas from experimenting
Analysis of metal oxides
Heating some metal oxides releases O2 and leaves the pure metal Take known amount of metal oxide and heat with Bunsen burner.
Determine the mass of the pure metal. Subtract this mass from the original to determine the amount of oxygen
o 4.608 g of silver oxide is heated; 4.306 g of silver metal is left. What is the empirical formula?
o M of silver = 4.306, m of oxygen = 4.608 – 4.306o Then divide by lowest mole to find empirical
Analysis of Hydrated Salts Hydrated compounds have water associated with them Need to know how much water they contain
o You heat a known mass to release the water. You then determine the mass of the anhydrous salt (no water) at which point you can determine the mass of water released.
5.742 g of hydrated magnesium sulfate is heated. 2.801 g of anhydrous magnesium sulfate is left.
5.742 – 2.801 = m of water2.801 = m of MgSO4
Divide by lowest moles to find formula Carbon-Hydrogen Analysis
Known mass is burned and water and CO2 produced are collected From the mass of each, you can calculate the mass of H and C in the
original compoundo Unknown compound with C and H is burned, 3.94 g of water
and 9.62 g of CO2 is produced. Molar mass is 84.0 g/mol. What is the molecular formula?
o Find moles of CO2 and H2O Use that to find moles of C, H, and O Divide C and H by lowest moles to find empirical
Stoichiometry Mole Ratios
Coefficients in a chemical reaction N2 + 3H2 2NH3
o N2 to H2 = 1:3 ratio (N2/H2 = 1/3) Steps to Stoichiometry
1. Write out proper balanced equation and put all givens underneath the equation.
2. If given mass, determine number of moles of reactants: n = m/M 3. Set up mole ratio of desired compounds 4. Calculate moles of required substance based on mole ratio 5. Determine mass using: n = m/M
6. Write your therefore statement
Limiting Reagento In typical reactions, reactants are usually NOT stoichiometrically equivalent. If they give you
BOTH reactants, you need to do limiting reagento There is one limiting, and one excess
o Steps to solving problems involving limiting reagent: 1. Start with a balanced chemical equation 2. Determine the amount of moles in BOTH reactants 3. Choose ONE of the reactants and set up a mole ratio 4. Compare the mole ratio results with what you actually have 5. From this information determine the limiting reagent 6.Using the limiting reagent determine the amount of product that can be made Don’t forget the therefore statement
Percent Yieldo Theoretical Yield of a reaction refers to the amount of product you predict should be
present (found through stoichiometry)o Actual Yield refers to the amount of product you actually get (experimental)o % Yield = [Actual Yield / Theoretical Yield] x 100%
Empirical Formula ProblemA compound consists of 72.2% Mg (Magnesium), 27.8% N (Nitrogen). What is the empirical formula? The Molar Masses:
1. n (Mg) = 24.3050
gmoles
2. n (N) =14.0067
gmoles
Elements % Mass (g)n= m
MrDivide by small Multiply till whole
Mg 72.2 72.2 2.97 2. 971. 98
=1 . 5 1 .5 × 2=3
N 27.8 27.8 1.98 1. 981. 98
=1 1× 2=2
Answer: Mg3N2
Molecular Formula Problem
If the molar mass of the compound is 100.9
gmoles , what is the molecular formula?
Empirical formula = Mg3N2
Empirical mass
= (3× 24 .3050 )+(2×14 . 0067 )=100 . 99 moles
Molecular formula = (Empirical Formula) ¿ X
X
=Molecular massEmpirical Mass
=100 .99 moles ¿100 .99 moles ¿=1 ¿
¿
Limiting and Excess Reactant ProblemSilver nitrate and sodium phosphate have reacted in equal amounts of 200 grams each. What is the limiting reactant?
StepsSodium nitrate Sodium phosphate
Chemical 3AgNo3 Na3Po4
Mass (g) 200g 200g
Molar Mass
gmoles
169.87 163.94
Ratio moles 200169 . 87
=1.18 mol
200163 . 94
=1.22 moles
Divide by coefficients 1. 18 3
=0.392
1. 221
=1.22
Answer: Limiting reactant is Sodium nitrate
Formulas
n = m/MM Percentage yield=Actual yield
Theoretical yield×100
Answer: Mg3N2
Molecular Formula=(Empirical Mass )׿ ¿=(Mg3 N2)×1
¿Mg3 N2
X
Unit 4
Solutions and Solubility
Solutionso Solution: homogenous mixture of two or more substanceso Composed of a solute (liquid in which substances are dissolved) and a solvent (the substances
being dissolved)o When water is the solvent, the solutions are given the specific name of acqueous
o Solution Combinations:
Solute\ Solvent Solid Liquid Gas
Solid Copper in zinc
(a brass alloy)
Salt in water (brine) Naphthalene in air (a moth repellent)
Liquid Mercury in gold or silver (a dental amalgam)
Ethylene glycol in water (automotive antifreeze mixture)
Water in air (humidity)
Gas Hydrogen in platinum Carbon dioxide in water (carbonated water)
Oxygen in nitrogen (the main components of air)
Solubilityo Solubility: Physical property of a substance that describes how soluble or insoluble it is
o When a substance dissolves better than another in a solvent, we say it’s more solubleo Defined as the amount of a solute that will dissolve in a given volume or mass of a
solvento When the max amount of a solid or gaseous solute has been dissolved in a given volume
of solvent at a given temperature, a saturated solution is obtained.o If more solute is added to a saturated solution, the excess will remain undissolved and
the solution is called supersaturated.o If a solution can still dissolve more solute, then the solution is unsaturated.o Dissolving a solute into a solvent is a Physical Change
o Factors affecting rate of dissolvingo Temperature
Solids: increase temperature increase rate of dissolving At a high temperature, solvent molecules have greater kinetic energy.
They collide with undissolved solid molecules more frequently breaks apart faster increased rate of dissolving
Liquids
Temperature does not greatly affect solubility of liquids because attractive forces within the liquid solute are not strong to begin with
Gases: Temperature increases Solubility decreases Gas particles have more kinetic energy than liquid particles. When
dissolved in liquid, they lose some of this energy. At higher temperatures gain back energy gas particles come out of solution solution becomes less soluble
o Agitation/stirring/shaking: Increase agitation Increased rate of dissolving It increases rate of collisions It brings fresh solvent molecules into contact with undissolved solid molecules
o Size of Particles: Decrease size increase rate of dissolving Smaller particles have greater surface area than larger particles and so more
solid is exposed to solvento Molecule Size: Smaller molecules are more soluble than larger molecules
Intermolecular forces (London forces) are greater in larger molecules than smaller molecules so it’s harder to break them apart.
o Pressure Changes in pressure does not affect solubility of liquids/solids Solubility of gases is directly proportional to the pressure of the gas above the
liquid (increase pressure above the liquid increase solubility)o When a solute will dissolve in a solvent?
Forces of attraction between solute and solvent determine whether a solute will dissolve in a solvent
Forces of Attraction:o Solvent-Solvento Solvent-Soluteo Solute-Solute
When the forces of attraction between solute and solvent are greater than solute-solute and solvent-solvent, the solution will form
Like Dissolves Like Polar Compounds dissolve in Polar solvents Non-Polar Compounds will dissolve in Non-Polar Solvents
Ionic Compounds in Water Water is a polar covalent molecule and since it is bent in shape, it
doesn’t cancel charges out Partially Negative Oxygen atoms other positive charges Partially Positive Hydrogen atoms other negative charges Partial positive Hydrogen surrounds negative ions while partial negative
oxygen surrounds positive ions
Intermolecular attractive forces between water and the individual ions are greater than the attraction between the ions in the compound (electrostatic attraction)
Covalent Compounds in water Polar dissolves in water Non Polar doesn’t dissolve in water (Oil in Water)
Concentrationo Concentration: the amount of solute present in a specific amount of solution
o C = amount of solute/ amount of solutiono Expressed as the mass of solute per 100 mL of solvent (Ex. Solubility of NaCl is 36g/100
mL means that 36g of NaCl can dissolve in 100 mL of water for a saturated solution)o Concentration Expressions
o Mass/Volume %: mass/volume % = mass of solute (g) / volume of solution (mL) x 100 %o Mass/mass %: mass of solute/ mass of solution x 100 %
Used when you have mass of solute dissolved in mass of solid solution (ex. Alloys)
When it says something dissolve in solution, don’t add mass of solute to that of solvent
When it says dissolved in x g of a solvent, add the masses of solute and solvent for the denomination
o Volume/Volume %: volume of solute (mL)/volume of solution (mL) x 100 % Used when two liquids are mixed to form a solution
o Parts per Million (1 part in 106): ppm = mass of solute/mass of solution x 106
o Parts per billion (1 part in 109): ppb = mass of solute/ mass of solution x 109
o Ion Concentrations: one formula unit of an ionic compound can release multiple ions of a specific element
So Moles Per Litre (mol/L): c = n/v or concentration = amount of solute (moles)/volume of
solution (L) Most commonly used Concentration in terms of moles/L is often referred to as molarity It is the # of moles of solute that can be dissolved in 1L in solution
Preparing Solutionso Concentration refers to the volume of solution, NOT solvento Steps to preparing a solution
o From the # of moles needed, determine the mass of solute required to prepare the solution (c=n/V, n = m/MM)
o Add solute to volumetric flasko Add enough solvent to dissolve solute (do not add required, just enough to
dissolve)
o Mix! (swirl)o Top up solvent to desired solution level
Dilutionso Most stock solutions are very concentrated because it is easier to store and transport but they
are more dangerouso When diluting, moles stays the sameo C1V1 = C2V2 or M1V1=M2V2
Ionic Equationso Spectator Ions: ions that remain as ions before and after a reactiono Total Ionic Equation: shows the dissociated ions of the soluble ionic compoundso Net Ionic Equation: an ionic equation that is written without the spectator ionso Guidelines to writing a net ionic equation
o 1. Include only ions and compounds that have reactedo 2. Write soluble ionic compounds as ionso 3. Write insoluble ionic compounds as formulaso 4. Covalent compounds, gases, and water are written as formulas (since they do not
produce ions)o 5. Write strong acids in their ionic forms since they dissociate in water
HCl H+ + Cl-
HNO3 H+ + NO3-
H2SO42H+ + SO42-
o 6. Make sure the net ionic equation is balanced for charges as well as atoms!o Identifying ions in aqueous solutions : Qualitative Analysis
o Colour
ion colour
Cr2+, Cu2+ blue
Cr3+, Cu+, Fe2+, Ni2+ green
Fe3+ yellow
Co2+, Mn2+ pink
CrO42- (chromate) yellow
Cr2O72- (dichromate) orange
MnO4- (permanganate) purple
o Flame Test
ion colour
Li+ red
Na+ yellow
K+, Cs+ violet
Ca2+, Sr2+ red
Ba2+ Yellowish-green
Cu2+ Bluish-green
Water Quality and Purificationo Hard water
Contains an appreciable concentration of calcium and magnesium ions
Stoichiometry in Solutionso Find moles through c = n/vo Use stoichiometry to find moles of other substances
Acids and Baseso General Properties
o Acids Sour Tasting Conduct Electricity Corrosive Turn blue litmus paper red
o Bases Slippery to touch Conduct Electricity Turn Red Litmus paper blue
o Arrhenius Theory An acid is a substance that produces hydrogen ions when dissolved in water
HCl H+ and Cl- A base is a substance that produces hydroxide ions when dissolved in water.
NaOH Na+ and OH- Problems with Arrhenius
Only substances with H+ form acids Only Substances with OH- ions form bases Meant that all salts must be neutral but certain salts dissolve in water to
be acids/bases (H2PO4 dissolves in water to be a base; NH3 dissolves in water to form a base)
o Bronsted Lowry Definition An acid is any molecule/ion that can give up a hydrogen ion A base is any molecule/ion that can combine with a hydrogen ion
Acids: donate hydrogen ions, Bases: accept hydrogen ions Any molecule with a hydrogen is a potential acid; any molecule with a
lone pair of electrons is a potential base Ionization: process that results in the production of ions in solution
Hydronium Ion (H3O+) A free H+ ion cannot exist in aqueous solution as it is simply a proton,
and it was a significant discovery when the hydronium ion was provent to exist
H+ and H3O+ are treated as the same species Conjugate Acid Base Pairs
Every Acid has a corresponding base and every base has a corresponding acid.
A conjugate acid-base pair differs by one proton (hydrogen) only. Strong and Weak Acids; Strong and Weak Bases
A strong acid is when protons completely ionize/dissociate when dissolved in water
A weak acid is when protons partially ionize/dissociate in water Bases that dissociate completely are strong bases Bases that dissociate partially are weak bases
pH and pOHo pH Scale
o range is from 0-14o 7 is neutralo <7 is an acido >7 is a base
o pH is a logarithmic relationship, meaning an change of 1 pH unit results in a concentration change of x10
o 1 sig fig = 1 decimal after the numbero Determining the pH
o You need to know the hydrogen ion/hydronium ion concentrationo pH = -log (H3O+)
o Determining pOH (pH of a basic solution)o pOH = -log(OH-)o pH + pOH = 14
o pH measurementso approximate value of pH can be obtained using an indicator (organic dye that changes
colour depending on pH)
Acid-Base Titrationo Acid-base titration is used to determine the concentration of an unknown acidic/basic solution
experimentallyo To do this, you need to know the concentration of the solution (called standard
solution)o You have to find the pH at which the number of moles of the standard is
stoichimetrically equal to the number of moles of the unknown (# of moles of H+ = # of moles of OH-) (equivalence point) which means neutralization of the acid/base has occurred.
o You need an indicator which tells you when neutralization has occurredo When titration ends, it is called the end point
o Steps for Titrationo Obtain a known volume of unknown in a flask or a beaker.o Add a couple of drops of the indicator to the unknown.o Wash and rinse out the buret. Make sure the stopcock is closed. Then add 30 – 50 mL
of the standard to the buret and set it up on the retort stand.o Start the titration by slowly opening the stopcock and letting the known solution flow
into the unknown. Swirl flask after each addition. When you add NaOH into the solution it should turn pink for a second where the drops land. Swirl the solution and the pink will disappear.
o Keep alternating between small amounts of NaOH, swirling and checking pH until it starts taking a while for the colour to disappear. Once this happens, start adding one drop at a time (don’t check the volume and pH at this point since it would be difficult).
o Colour change will occur when the pink no longer disappears. Record this volume and pH. You have reached the end point at this time (when the colour does not disappear with swirling).
o Do not go over the end point! Or else your calculations will be affected.
Unit 5
Gaseso Properties
o Fill any space available to themo Low densityo Most are colourless and odourlesso Volumes change dramatically with changes in pressureo Can be mixed in any proportion to give a homogenous solution
o Atmosphere: 78.08 % Nitrogen, 20.95 % Oxygen, 0.93 % Argon, 0.03 % Carbon Dioxide, 0.002 % Neon, 0.008% Other
o Pressure (SI Unit = Pa = N/m2)o The more particles, higher the temperature (faster molecules)o Smaller the container, the more pressure there will beo 1 Standard Atmosphere (atm) = 101.3 kPa o 1 atm = 760 mm Hgo 1 atm = 14.7 lb * in-2
o Measuring Pressure Barometer: Measures the pressure of the atmosphere on its surroundings Manometer: Measures the pressure of a gas relative to the atmospheric
pressureo Pabsolute = Prelative + Patmospheric
o Kelvin Temperatureo Absolute Zero = 0 K or -273 degrees Co Conversion: Degrees C + 273
-30 degrees C -30 + 273 = 243 K 30 degrees C 30+273 = 303 K
o Standard Ambient Temperature and Pressure (SATP)o 25 degrees C, 100 kPa
o Standard Temperature and Pressure (STP)o 0 degrees C, 101.3 kPa
Gas Lawso Boyle’s law
o At a constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure (1/P)
o Pa 1/V P = k/V (k is an unknown constant slope (when you graph the relationship)o P 1V1 = P2V2: Lets you solve problems stemming from changing environmental conditions
and its effect on gaseso Charles’ Law
o At a constant pressure, the volume of a gas is directly proportional to its temperatureo V = kT V/T = ko V 1/T1=V2/T2
o Gay Lussac’s Law (LOL)o At constant volume, the pressure of a fixed mass of gas is directly proportional to its
Kelvin temperatureo P 1/T1=P2T2
o Combined Gas Lawso P1V1/T1 = P2V2/T2
o Only for changing environmental conditions, NO reaction is taking place
Molar Volume/ Ideal Gas Lawo Molar Volume: the volume occupied by one mole of gas under STP/SATP
o SATP: Volume of 1 mole is 24.8 Lo STP: Volume of 1 mole is 22.4 L
o Density = mass/volume = Molar Mass/ Molar Volumeo Ideal Gas Law
o PV=nRTo R = 8.31 (universal gas constant)
o Vapour Pressure/ Dalton’s Lawo When water is placed in a sealed container, some of the water evaporates to form
water vapouro It exerts a pressure known as water vapour pressureo At any given temperature and pressure, only a certain amount of water will evaporate
If the air above the water becomes saturated, it can no longer hold any more water
o Boiling Point: Every liquid evaporates until the air above it is saturated or liquid is vaporized
Every liquid exerts vapour pressure which increases with increased temperatures
When vapour pressure equals external pressure, the liquid begins to boil Boiling Point is defined as: the temperature at which its vapour pressure equals
the external pressureo Kinetic Molecular Theory
Vapour pressure can be explained using Kinetic Molecular Theory Average kinetic molecule implies that molecules have different kinetic
energies Molecules with more kinetic energy overcome attractive forces
between particles become gases
Substances have different boiling points because their particles have different levels of attraction
Liquids with strong attractive forces low vapour pressure Liquids with weak attractive forces high vapour pressure
o Dalton’s law of Partial Pressure To make sure gas molecules do not react with each other in a mixture, they
have different pressures P T = P1 + P2 + P3 …
o Ideal Gas Law and Dalton’s Law of Partial Pressures PTotalV = ntotalRT
PAV = nART/ PTV = nTRT to find out about gas A in a mixture
Calculations Involving Reactions producing Gaseso Gay Lussac’s Law of Combining Volumes (LOL)
o In any chemical reaction involving gases (when T and P is constant), the volumes of the gases are always in small whole number ratios (coefficients stoichiometry)
o Avogadro’s Hypothesis and the Law of Combining Volumeso The volume of a gas is directly proportional to the number of moles of gas when the
temperature and the pressure are constant. It also applies to two different gaseso N1/V1 = N2V2
o Gases collected through displacement of watero Gases can be bubbled through water water is displaced and gas is collected
When water is displaced, pressure inside vessel = atmospheric pressure Gas collected is not pure, it has water vapour Gas collected is a mixture (gas collected + water vapour)
