Chemistry 3U Exam Review

Changes in MatterChemical= difficult to reverse, heat or light is produced, gas is formed, new colour, precipitatePhysical= change in form or state

Definitions Heterogeneous mixture- see different components (pizza) Homogeneous mixture- everything is mixed so they appear as one substance (kool-aid) Element- pure substance that cannot be broken down; 1 type of atom Compound- a pure substance that is composed or 2 or more atoms chemically combined in a fixed

proportion Atom- the basic unit of an element Molecule- the simplest structural unit of an element or compound

Pure Substances Solids-most a crystalline, are regular shapes with flat sides (faces), all faces are 90° to one another Gases- don’t have a definite shape or volume; they always fill their containers Liquids- don’t have a regular or fixed shape; always take the shape of their container Kinetic Theory- solids are made up of particles that are close together and not moving about. When

a solid is heated the particles vibrate faster. At the melting point they begin to move around, but remain close to keep a definite volume. As the liquid is heated, the particles increase their speed and take up more room. At the boiling point the particles are moving fast enough to escape from one another and the liquid becomes a gas.

Making a Pure Substance Magnetic Separation- magnets separate iron and steel from other non-magnetic metals. Flotation- Metal ores usually have a different density from the soil in which they are found, and the

ore may be floated off using oil, sometimes water. Filtration- the substances which dissolve may be separated from those that don’t. Sedimentation- Insoluble solids can usually be separated from liquids by allowing the solid to settle. Evaporation- this method is used to obtain any dissolved solid from a solution. Distillation- the whole process of evaporating a liquid and then condensing it again Chromatography- as the substance become spaced out they are separated from one another.

Line Spectra A wave has maximum and minimum values called crests and troughs, respectively. The distance

between successive crests and successive troughs is known as the wavelength. Wavelengths of visible light is usually measured in nanometers (nm)

Electromagnetic energy- is commonly known as light energy. Frequency- of a light wave is the number of cycles that pass a point in a second Line Spectrum- consists of distinct colored lines rather than a rainbow.

The Bohr Model of the Atom The energy of electrons is quantized An electron that occupies a higher energy level is said to be in an excited state If the electron is found in the lowest possible energy level it’s in the ground state

Atoms and their Composition Atomic #- # of protons/ electrons in a stable atom Mass #- total # of protons and neutrons in a stable atom (Mass#-Atomic #)=# of neutrons Isotopes- are atoms of an element that have the same number of protons but different number of

neutrons Radioisotope- unstable isotope which undergoes radioactive decay

Atoms-Inside Story Protons- positive, in nucleus, mass of 1 Neutrons- neutral, in nucleus, mass of 1 Electrons- negative, move in the space around nucleus, mass of 1/2000 Standard Atomic Notation- write the chemical symbol of the atom and place the atomic number to

the lower left and the mass number to the upper left Bohr-Rutherford Diagrams- circle drawn in the centre to represent the nucleus of the atoms, the

numbers of protons and neutrons are written in it. Electrons are shown in the circular orbits around the nucleus.

Valence Electrons 1st max=2 2nd max=8 3rd max= 18 4th max= 32 5th max= 50

Atoms, Elements and the Periodic Table Elements are arranged in seven number periods (horizontal rows) and 18 numbered groups (vertical

columns) Groups are numbered according to 2 different systems. The current system numbers the groups

from 1-18. An older system numbered the groups from I-VIII, and separates them into two categories labelled A and B.

Elements in the eight A groups are the main-group elements. Also called the representative elements. The elements in the ten B groups are known as the transition elements.

Within the B group transition elements are two horizontal series of elements called inner transition elements. They usually appear below the main periodic table. Notice, however, that they fit between the elements in Group 3 (IIB) and group 4 (IVB)

A bold staircase line runs from the top of group 13 (IIA) to the bottom of group 16 (VIA). This line separates the elements into three broad classes: metals, metalloids (semi-metals) and non-metals.

Group 1 (IA) elements are known as the alkali metals. They react with water to form alkaline or basic solutions.

Group 2 (IIA) elements are known as alkaline earth metals. They react with oxygen to form alkaline solutions.

Group 17 (VIIA) elements are known as the halogens. They combine with other elements to form compounds called salts.

Group 18 (VIIIA) elements are known as the noble gases. Noble gases don’t combine naturally with any other element.

Periodic Trends are patterns that are evident when elements are organized by their atomic numbers Energy Levels are fixed, three dimensional volume in which electrons travel around the nucleus. Lewis Structures a symbolic representation of the arrangement of the valence electrons of an

element. Stable Octet an arrangement of 8 electrons in the valence shell of an atom. Periodic Law states that the chemical and physical properties of the elements repeat in a regular,

periodic patter when they are arranged occurring to their atomic numbers.

Arranging Elements in the Periodic Table Metals are shiny, good conductors of heat and electricity, ductile, and malleable. Non-Metals are dull in colour, good insulators and brittle. Metalloids are used to make semiconductors John Newlands (1884) introduced the concept of octaves since he noticed that the properties of

elements seemed to reoccur after every eighth one. Dmitri Mendeleev (1869) organized all the known elements by relative atomic mass and noticed

that the properties of the elements were periodic function of that mass.

Trends within Groups and Rows As one moves down a column or group in the periodic table one notices the following tendencies. Metals get softer the further down one goes 2. The reactivity of metals increases 3. The reactivity of

non-metals decreases 4. The boiling point of non-metals increases As one moves across a row or period from the left to right one notices the following tendencies The metallic nature of the elements decreases until they become non-metallic 2. The reactivity

declines until the middle of the row and then increases until the noble gases are reached where it rabidly declines.

The Atomic Model and the Periodic Table An ion is an electrically charged particle formed when an atom gains or losses electrons. Photo-ionization uses light energy to remove electrons from an atom while electron bombardment

uses fast moving electrons to form ions. The energy released when a neutral atom attracts an extra electron is called its electron affinity. All

these observations can be summed up as follows: Metals react by losing electrons. The lower the first ionization energy, the more reactive is the

metal. Non-Metals (except the noble gases) react by gaining electrons. The higher the first ionization

energy and the higher the electron affinity the more these non-metals attract extra electrons and hence they are more reactive.

Noble gases are very unreactive because they have stable electron arrangements and do not easily lose or gain electrons. Helium has two electrons in its full outer energy level and others have eight electrons.

As you go down each group in the periodic table, the size of an atom increases. As you go across a period, the size of an atom decreases. Ionization energy tends to decrease down a group Ionization energy tends to increase across a period.

Ionic and Covalent Bonds Ionic Bonding- electron transfer between metal atoms and non-metal atoms IDE- binary compound (2 element compounds)

Bonding by sharing “e” does not involve electrostatic charges

Comparing Ionic and Covalent CompoundsProperty Ionic Compound Covalent CompoundState at room temp. Crystalline solid Liquid, gas or solidMelting point High LowElectrical Conductivity in liquid Yes No Solubility in water Most have high solubility Most have low solubilityConducts electricity when dissolved in water

Yes Not usually

Electron Sea Bonding The force that holds metal atoms together is called metallic bonding Electrons are free to move, the metal ions are not rigidly held in a lattice formations, therefore,

when a hammer pounds a metal, the atoms can easily slide past one another. Polar Bonds- the shared electrons are strongly pulled to the more electronegative atoms Delta –δ Non-Polar Bonds- both atoms exert a similar pull on the shared electrons

Predicting Bond TypeNumber of Valence Electrons Bond TypeOne of the bonding atoms has less than half the max number of valence electrons (1,2,3) except for hydrogen; i.e., metals are involved

IONIC

Both of the bonding atoms have half or more of their max number of valence electrons; i.e., both are non-metals

COVALENT

Why and How Atoms Combine

Octet Rule- When atoms form ions or combine in compounds they obtain electron configurations of the nearest noble gas (usually this means that there will be 8 outer electrons)

Covalent bonding- 2 non-metals; electrons are shared instead of transferred Ionic bonding- non-metal and a metal

Ionic bonding involves the formation of ions: metals lose electrons to become positive ions, non-metals gain electrons to become negative ions

Polarizations Electronegativity is the tendency for an atom to attract the shared electrons inn a covalent bond

more strongly to itself. Polar Covalent Bonds- atoms have significantly different electronegativities Electron pairs that are not involved in bonding are called lone pairs Electron pairs that are involved in bonding are called bonding pairs

Polar and Non-Polar Bonding Dipole-Dipole δ+ is attracted to δ- Hydrogen Bonding- H with N, O or F Dispersion Force- the natural force between molecules (like methane)Water- boiling point 100C, melting point 0C, # of electrons 10e-

Methane- boiling point-161C, melting point -183C, # of electrons 10e-

Writing Chemical Formulas and Naming Chemical Compounds Chemical Nomenclature is a system of naming chemical species including elements, ions and

compounds

some elements were named after their appearance or source some elements were named after famous people or places (Einsteinium, Californium) some elements were named after their properties (gold (Aurium) means shining dawn) some are named creatively (Mercury, Cobalt) Losing an electron forms a cation Gaining an electrons forms an anion Metatomic ions have only one element; metatomic anions end in IDE

Formulas and Naming Compounds Binary Compounds- 2 elements; the name of the first element and the periodic table name of the

second element ending in IDE. (Lithium fluoride) Prefix System (covalent bonding)- 1-mono; 2-di; 3-tri; 4-tetra; 5-penta; 6-hexa; 70hepta; 8-octa; 9-

nona; Ex Cl2O = Dichlorine (mono)oxide Suffix System (ionic bonding)- When multivalent metal atoms have only 2 valence the high charged

is assigned the ending –IC and the lower is assigned the ending –OUS. Ex FeCl3= Ferric Chloride; Hg2O= Mercurous Oxide

Stock System (ionic bonding)- this system simple puts the charge that is being used by the first element in brackets and Roman numerals after the element name Ex. Iron (III) Chloride

Acids-consist of one or more hydrogen ions in front of radicals consisting of a non-metallic element, in some cases, oxygen. These normal oxyacids all end in -ATE- the word hydrogen followed by the radical name ending in ate. They can also be referred to by the nonmetallic element in the compound ending in -IC, followed by the word acid. No oxygen= hydro-IC ;2oxygen atoms less than normal= hypo-ous ;1 oxygen atom less than normal=-ous; normal=-IC ;1 oxygen atom more than normal= per-IC

Radicals- these are a combination of different atoms that have a charge equal to the number of hydrogen atoms they normally combine with. (look at chart in notes)

Bases- most cases made by a metal followed by the hydroxide radical. Bases are named by taking the name of the metal and having it followed by the word hydroxide. When the metal is of a valence greater than 1+ the hydroxide must be put on brackets and have the criss-crossed valence number placed outside the bracket.

Salts- are named by naming the metal or ammonium radical first than by naming the radical that is involved. When an acid and a base react he products are salt and water.

Polyprotic Acids- these are acids that can give off more than one hydrogen atom. When only one or two hydrogens are given off the metal or ammonium ion that replaces it creates and ACID SALT. The single hydrogen tern can also be replaced with the prefix bi as in sodium bicarbonate instead of sodium hydrogen carbonate. The prefix system is used in naming these products.

Acids Reacting with Bases Monoprotic Acids- contain only one hydrogen ion that can dissociate. Diprotic Acids- contains 2 hydrogen ions that can dissociate. Give rise to two possible salts Triprotic Acids- contain 3 hydrogen ions; give rise to 3 possible salts.

Chemical Equations Steps to correctly putting a chemical equation together 1. The word equation 2. Constructing the skeletal/ skeleton equation - convert words into the correct chemical formulae

3. Balancing the skeletal equation to observe the "law of conservation of matter" we use coefficients to do this.

Types of Reactions 1. Synthesis - X+Y>XY 2. Decomposition- PQ> P+Q 3. Single Replacement- X + YZ>Y +XZ 4. Double Replacement- PQ+RS> PS+RQ

Naming Compounds -A Summary Write the name of the element with the lowest electonegativity first. If this element has more than

one valence, indicate the valence in the compound using roman numerals

Guidelines for Balancing Equations 1.Balance the element, other than hydrogen and oxygen, that has the greatest number of atoms in

any reactant or product 2. Balance the other elements, other than hydrogen and oxygen 3. Balance oxygen and hydrogen, whichever one is present in the combination state. Leave until last

whichever one is present in the uncombination state. 4. Check that the equation is balanced by counting the number of atoms of each element on each

side of the equation Physical states must be shown ! L, S, Aq, G

Simple Nuclear Reactions Rules for Balancing Nuclear Equations 1. The sum of the mass numbers (written as subscripts) on each side of the equation must balance. 2. The sum of the atomic numbers (written as subscripts) on each side of the equation must balance. Alpha Decay(α)- involves the loss of one alpha particle, which is a helium nucleus, 4

2He, composed of 2 protons and 2 neutrons

22688Ra>222

86Rn + 42He

Beta Decay (β) - occurs when an isotope emits an electron, called a beta particle. Represented as 0-1e

31H> 3

2He +0-1e 1

0n> 11H + 0

-1e Gamma Radiation (γ)- is high energy electromagnetic radiation. Often accompanies either alpha

or beta particle emission. 13735Cs> 137

56Ba +0-1e + 0

0γMode Change In……………………………………

Mass Numbers Atomic Numbers Number of Neutrons(α) -4 -2 -2(β) 0 +1 -1(γ) 0 0 0 Nuclear Fission- occurs when a highly unstable isotope splits into smaller particles. Usually has to be

induced in a particle accelerator. Here, an atom can absorb a stream of high energy particles like neutrons, 1

0n. This will cause the atom to split into smaller fragments. 23592U+1

0n>8735Br+146

57La+310n

Nuclear Fusion- occurs when a target nucleus absorbs an accelerated particle. The reaction that takes place in a hydrogen bomb is a fusion reaction, as are the reactions that take place within the Sun. Fusion reactions require very high temp. to proceed but produce enormous amounts of energy. The fusion reaction that takes place in a hydrogen bomb is > 6

3Li + 10n > 3

1H +42He

How to Calculate Average Atomic MassAverage Atomic Mass = Relative mass of Isotope A + Relative mass of Isotope B + Relative mass of

Isotope C + ...... etc.Relative Mass of Isotope A= percentage abundance of isotope A X mass of isotope ARelative Mass of Isotope B= percentage abundance of isotope B X mass of isotope BETC………………………………Percentage abundance should be shown as a fraction… 72% > 72/100EXAMPLE : B-10; 19.78% abundance , atomic mass of 10.01 u B-11; 80.22% abundance, atomic mass of 11.01uStep 1. Determine the relative mass of each isotope

Relative mass of B-10= 19.78/100 X 10.01= 1.9799uRelative mass of B-11= 80.33/100 X 11.01= 8.8322u

Step 2. Put the answers from step 1 into the following equationAverage atomic mass = relative mass of B-10 + relative mass of B-11So the average atomic mass of Boron= 1.9799u + 8.8322u

Therefore the average atomic mass of boron is 10.81u

How to Count Atoms Symbol of an element represents 1 atom of that element (Na= 1 atom) Subscript (lower right, behind symbol) indicates how many atoms there are (Na3= 3 Na atoms) Subscript outside a bracket multiples all the elements inside the bracket (Mg3(PO4)2= 3 Mgs, 2 PO4s) Coefficient in front of a chemical formula indicates the number of molecules of that compound

2H2O= 4 H and 2 O . 2 lots of H2O

Avagadro’s Number6.02X1023

Converting Moles to Number of Particles and Converting Number of Particles to Moles

N N= number of particles n=number of moles NA= Avagadro’s Number

n NA

Converting from Mass to moles and Converting from Moles of mass

m m= actual or given mass of a substance n=number of moles M=molar mass of the substance

n M

Converting Between Moles, Mass and Number of ParticlesN=mNA/M m=NM/NA M=mNA/N

The Law of Definite ProportionsThe law of definite proportions states that the elements in a chemical compound are always present in the same proportions. Another way of stating it is to say that a specific compound has the same composition anywhere in the universe.

Different Compounds from the Same ElementsThe law of multiple proportions states that when two elements combine to form more than one compound, the masses of one element will combine with a fixed mass of the other element in small, whole number ratios

Percentage Composition from a Chemical FormulaThe law of definite proportions allows for practical calculations when the proportions of elements in a compound are known. Proportions are often stated in the form of percentage composition, which is the percent by mass of each element in a compound. The percent composition can be determined by experiment or calculated from the formula of a known compound.

Ex. What is the percentage composition of mercury in mercuric oxide (HgO) and mercuric sulphide (HgS)HgO: molar mass (M)= 200.59+16.00=216.59g/molSo in HgO %Hg=200.59/216.59X100=92.6%HgS: molar mass (M)= 200.59+32.07=232.66g/molSo in HgS %Hg=200.59/232.66X100=86.2%

From %Composition to Simplest FormulaThe simplest formula of a substance has all component elements present in it reduced to their lowest numerical terms. This formula is referred to as the EMPERICAL FORMUAL. C2H2>CH; N2O4>NO2

Empirical Formula of a Compound1. Divide the element’s mass per 100 gram sample (% composition) by its own molar mass2. Divide each element’s number of moles by the SMALLEST of all calculated values3. ONLY IF NECESSARY multiply the mole ratio by the SAME FACTOR to obtain whole number answers

Element Mass/100g (m)

Molar Mass (M)g/mol

Number of Moles (n)

Mole Ratio Revised whole number ratio

C 63.10 12.01 63.10/12.01= 5.254

5.254/1.975= 2.66

2.66x3= 8

H 5.31 1.01 5.31/1.01=5.257

5.257/1.975= 2.66

2.66x3= 8

O 31.60 16.00 31.60/16.00= 1.975

1.975/1.975= 1.00

1.00x3= 3

Molecular (Actual) FormulaWe sometimes have to determine the molecular formula from the substance’s empirical formula. To do this we follow all the same rules as before to find the empirical formula and the providing we know the mass of the formula unit of the molecule we can easily determine the molecular formula.Element Mass in 100g Sample

(%value)# of moles (m/M) Mole ration (Divide all

by smallest # moles)Whole #ratio (X all by same factor, if necessary)

C 39.95g 39.95/12=3.33 3.33/3.33=1H 6.71g 9.71/1=6.71 6.71/3.33=2.01O 100-(39.93+6.71)= 53.34/16=3.33 3.33/3.33=1

53.34gThe empirical formula of this compound in CH2O and its mass is 30g The molar mass is 180g so the number of empirical units in the real formula is 180/30=6The molecular formula will be (CH2O)6 or C6H12O6

The Carbon-Hydrogen Combustion AnalyzerUsed to determine the composition for organic compounds containing carbon and hydrogen. Works according to the following principals: Complete combustion of a hydrocarbon produces carbon dioxide and water

CaCl2(S)+2H2O(G)>CaCl2 . 2H2O(S)

If the mass of the trap is take before the reaction and again following the combustion of the compound, the amount of water can be determined.

The carbon dioxide produced by the combustion reaction is collected in a similar trap, this time containing NaOHNaOH(S)+CO2(G)>NaHCO3(S)

If the mass of the rap is determined before and after the combustion reaction, the amount of carbon dioxide produced can be determined.

Since all of the hydrogen in the compound must be converted into water, the mass of hydrogen in the sample can be calculated using the following formula: mass of water collected X (mass of hydrogen in water/molar mass of water)= mass of hydrogen

Similarly, the mass of carbon can be found using the mass of carbon dioxide produced:Mass of CO2(G) collected X (mass of carbon in CO2/ molar mass of CO2)= mass of carbon

If there is any oxygen in the sample, it can be determined using the law of conservation of mass

Hydrated Ionic Compounds Many ionic compounds crystallize from a water solution with water molecules incorporated into

their crystal structure, forming a HYDRATE Hydrates have a specific number of water molecules chemically bonded in each formula unit. Compounds that have no water molecules incorporated into them are called ANHYDROUS to

distinguish them from their hydrated forms.

StoichiometryStoichiometry describes the chemical equations using mathematical relationships.EX) how much chalk (CaCo3) can be made when 2.8 g CaO is reacted with excess (XS) CO2?Therefore 1g +(44/56g)= 100-59 (divide by 56)So 2.8g+ (44/56g)X2.8g=100.56X2.8 (X by 2.8) Route 1 MethodTherefore 2.8g>>>>>>>>>>5 g chalk (EX)Given 2.8 g +XS >>>?Determine mole ration CaO:CaCO3 its is 1:1 from bal. equation. Find # mole CaO used ie/ n=m/M = 2.8/56= 0.05 moles. Calc # mole CaCO3 that will result by using mole ratio in bal equation.Therefore 0.05 moles of CaO will produce 0.05X1=0.05 moles of CaO3. Convert moles to mass M=nM. Therefore m CaCO3=0.05X100=5 g MOLE METHOD !

Percentage Yield The actual yield is the quantity of product that is actually produced in a chemical reaction. The theoretical yield is the quantity of a product calculated from the balanced equation.

The percent yield is the ratio expressed as a percentage, of the actual or experimental quantity or product obtained (actual yield) to the maximum possible quantity of product (theoretical yield) derived from a Stoichiometric calculation

Percentage yield = actual yield/theoretical yield X100%

Determining the Limiting and Excess Reactant Limiting reactant is the reactant that is completely used up in a chemical reaction Excess reactant is the reactant that remains after a reaction is over. CaF2+H2SO4(L)>2HF(G)+CaSO4(S)

10g 15.5g78g + 66g1g + 66/78gSo 10 g+ 66/78gX10

8.46g therefore CaF2 is the limiting reactant

Solutions If a solute (solid) dissolves in a solvent we say that it is soluble If it does not dissolve or dissolves only very slightly, we say that it is insoluble If a liquid dissolves in a solvent it is regarded as being miscible If one liquid does not dissolve in a solvent we say that it is immiscible Whether or not a solute will dissolve in a given solvent depends upon their INTERMOLECULAR

FORCES, these are:o Dispersiono Dipole-Dipole forces (occur between polarized molecules which attract each other)o Hydrogen Bonds (special dipole-dipole attractions between hydrogen and oxygen, nitrogen

or fluorine in molecules) Concentration Of Solutions

n c= conc. mol L-1 n=# or moles of solute v= volume of solution in L

C V

Dilution of SolutionsIf C1=n/V1 and C2=n/V2 > we get n=C1V1 and n=C2V2 Therefore C1V1=C2V2

Concentration of Ions in SolutionK2SO4(S) > 2 K+

(AQ) + SO42-

(Aq)

Solutions The process of an ion forming intermolecular attractions with solvent particles is called solution 1. The solute breaks apart > in ionic compounds, attractions among ions break apart > in molecular

substances, attractions among individual molecules break apart2. Attractions among solvent particles are overcome.3. Attractions form among solute and solvent particles.

3 Factors that affect how quickly a solute dissolves are: temperature (dissolves faster at higher temperature) stirring (increases speed and moves the solvent particles towards the solute, creates intermolecular attractions faster) surface area (provides a larger surface for which the solvent can act)

Expressing the Concentrations of Solutionsg/100mL Is a measure of mass of solute per 100mL of solution

(v/v) % Is a measure of volume of solute per volume of solution

(m/v) % Is a measure of mass of solute per volume of solution

(m/m) % Is a measure of mass of solute per mass of solution

Ppm Is a measure of mass of solute per million times more mass of solution 1X106

Ppb Is a measure of mass of solute per billion times more mass of solution 1X109

Ppt Is a measure of mass of solute per trillion times more mass of solution 1X1012

mol/L Is a measure of amount in moles of solute per litre of solution

UNITS MUST BE THE SAME!!!!!

Solubility RulesSoluble> more than 10grams per litre/ more than 0.1 moles per liter Slightly or sparingly soluble> 1 gram to 10 grams per liter/ 0.01 to 0.1 moles per literInsoluble> less than 1 gram per liter/ less than 0.01 moles per liter

Strategies for Solving Stoichiometric Problems1. Write out the balance chemical equation2. Set-up a table that contains the information provided in the balanced chemical equation as well as

the data given from the problem (mole ration, molar mass, and given amount)3. Convert the given mass to moles4. Use the mole-to-mole ration to find the required number of moles of the second substance5. Convert the number of moles of the second substance into the desired quantity (i.e., mass,

molecules)6. Write a concluding statement.

Water Treatment1. Collection- large particles and debris are removed by travelling screens as the water enters the

treatment plant2. Coagulation, Flocculation, and Sedimentation- chemicals known as coagulants are rapidly mixed

with the water to make the small particles in the water clump together. Flocculation is gentle mixing

to form a light, fluffy, precipitate called a floc. During sedimentation the floc settles very slowly, sinking and carrying suspended particles with it and thereby clearing the water.

3. Filtration- the reaming floc, other chemical and physical impurities, and most of the biological impurities (bacteria, etc.) are removed. The water flows by gravity through efficient filters made up of layers of sand and anthracite (carbon)

4. Disinfection- chlorine is added to kill microorganisms and to react with most organic molecules present. Alternative disinfectants include ozone or chlorine dioxide, ammonia, potassium permanganate, and even ultraviolet light.

5. Aeration- air, ozone or oxygen, “activated” charcoal, ammonia, chlorine dioxide, or potassium permanganate may be mixed with the water to further reduce taste and color problems.

6. Softening- hard water may be treated with sodium carbonate and calcium hydroxide or a phosphate to reduce water hardness by precipitating the calcium and magnesium ions.

7. Fluoridation- a small amount of fluoride is added to drinking water in some areas, as it makes the enamel layer of teeth more resistant to decay.

8. Post-chlorination- a final chlorine disinfection treatment kills any remaining microorganisms, and pH is adjusted to be slightly basic (since even slightly acidic water will corrode metallic pipes)

9. Ammoniation- ammonia is added to the end of the treatment process to stabilize the chlorine so that it remains dissolved in the treated water for longer periods of time.

Waste Water Treatment System1. Primary Treatment- involves screening, flotation, settling, and filtering out solid particles. It has no

effect on dissolved materials or microorganisms. This treatment removes about 40% of BOD2. Secondary Treatment- is a two-step process. The first step usually involves aerating the water to

support oxygen-using organisms, which react with dissolved organic substances to produce a sludge precipitate. The second step is chlorination, which further purifies the water. The resulting sludge may be used as landfill or fertilizer, although any heavy metal pollutants present are not removed. About 90% of BOD is removed from the water by the end of this stage. After this treatment, the water could be returned to the environment as it is suitable for most non-drinking purposes.

3. Tertiary Treatment- if used may involve a wide variety of systems and processes resulting in water clean enough for drinking. These processes may include reverse osmosis steam distillation, chemical precipitation –anything that will remove virtually all remaining organic chemicals and any harmful dissolved ionic compounds. This stage is the most expensive.

Predicting Precipitate Formation1. Identify type of reaction and possible products2. Look up solubility of both products3. Indicate states or reactants and products4. Write chemical equation for reaction5. Balance equation

Writing Net Ionic Equation6. Write total ionic equation7. Write net ionic equation

Acid-Base IndicatorsIndicator Colour in Acid Colour in Base

Litmus Red Blue

Bromothymol Blue Yellow Blue

Methyl Orange Red Orange

Phenophthalein Colourless Magenta

Acids and Bases Acids are proton donors Bases are proton acceptors Acids taste sour, turn blue litmus red, conduct electricity, no characteristic feel, produces hydrogen

gas with active metals, produces carbon dioxide gas with carbonate compounds. Bases taste bitter, turns red litmus blue, conduct electricity, feel slippery, don’t react with active

metals, don’t react with carbonate compounds

Arrhenius TheoryDefines an acid as a compound that can dissociate in water to yield hydrogen ions, H+, and a base as a compound that can dissociate in water to yield hydroxide ions, OH-.Problems >the solvent has no role play in the theory. This is wrong and the nature of the solvent plays a critical role in acid-base properties of substances; >all salts in the theory should produce solutions that are neither acidic nor basic. This is not the case; >the need for hydroxide as the base led Arrhenius to propose the formula NH4OH as the formula for ammonia in water. This led to the misconception that NH4OH is the actual base, not NH3.

The Brönsted Lowry TheoryWater can be considered an acid or a base since it can lose a proton to form a hydroxide ion, OH-, or accept a proton to form a hydronium ion, H3O+. When an acid loses a proton, the remaining species can be a proton acceptor and is called the conjugate base of the acid. Similarly when a base accepts a proton the resulting species can be a proton donor, this is called the conjugate acid of that base.

Conjugate Acid-Base PairsHX(aq) + H2O(l) <-> X- + H3O+

When the forward reaction occurs, HX donates a proton to water (acts like a base) to the form hydronium. When the reverse reaction occurs, the hydronium ion acts as the acid donating a proton to the X-. Conjugate acid-base pairs are compounds that differ by the presence of one proton, or H+.

HCN(aq) + H2O(l) > H3O+(aq) + CN-

Measuring pH[H3O+] pH 1 0 10-1 1

10-2 210-3 3………………………..10-14 14

Strong AcidsA strong acid is an acid that ionizes quantitatively (completely) in water to form hydrogen ions. The percent ionization of strong acids is greater than 99%. We will assume that it is 100% in calculations. (HCL(aq), HBr(aq), H2SO4(aq), HNO3(aq),H3PO4(aq))

Strong BasesIonic hydroxides have varying solubility in water, but all are strong bases that dissociate quantitatively (completely) when the dissolve in water. Any base that is capable of dissolving completely in water is also referred to as an alkali and an alkaline solutions is produced. Group 2 elements also form quite strong hydroxide solutions. When theses bases dissolve in water, two moles of hydroxide ions are formed for every mole of metal hydroxide that dissolves in solutions.

Weak Acids & BasesWeak acids and bases ionize incompletely when the react with water. We acids are acids that partially ionize in solution and exist primarily as complete molecules. They are defined by Bronsted and Lowry as proton donors. Weak bases are defined by Bronsted and Lowry as proton acceptors. So a weak base is any compound that dissociates very little to form an equilibrium that includes hydroxide ions, OH-(aq).

Hard and Soft Water→ HARD WATER is water with high concentrations of CA2+

(aq), Mg2+(aq),Fe2+

(aq),Fe3+(aq), and So4

2-(aq),

→ SOFT WATER is water with relatively low concentrations of these ions.→ The extent on the hardness depends on the types of rocks through which the water flows. It also

depends on the length of time that the water is in contact with the rocks.

Treating Water at Home→ These ions are not always removed at municipal treatment plants. If you wish, you can remove some

of these ions (mainly Ca2+, Mg2+ and Fe2+)yourself. For small volumes, you can add sodium carbonate decahydrat, Na2CO3·10H2O (washing soda). The sodium ions in the washing soda behave as spectator ions, leaving the water soft.

→ For large volumes, people install an ion exchange water softener. The hard water passes through a column packed with beads. The beads are made from an insoluble plastic material and are coated with sodium slats, often NaCl. (The salt coated beads are referred to as an ion exchange resin.) As the hard water passes through the column, the ions in the water displace the sodium calcium ions (and other hardness causing ions), the resin is regenerated. This is done by passing a very concentrated solution of sodium chloride (brine) through the column. The calcium ions are flushed out of the system along with excess sodium chloride solution.

Gases0°C=273Kelvin-273K=0°C

P= force/areaearthsP= weight/area

Avogadro’s Lawn α V or n=kV or n1/V1=n2/V2

Gay-Lussac’s LawP α T or P=kT or P1/T1=P2/T2

Boyle’s LawV α 1/P or PV=k or P1V1=P2V2

Charles’ LawV α T or V=kT or V1/T1=V2/T2

Combined Gas LawP1V1/T1=P2V2/T2

Dalton’s Law of Partial PressuresPtotal=P1+P2+P3+..............................

S.I. Standard Temperature and (Atmospheric) Pressure> (STP)Temperature 0°C equivalent to 273K (Kelvin used with gas equations)Pressure 101.3 kPa (kilopascals)Volume 22.4 Litres

The Ideal Gas LawV α nT/P V=RnT/P OR PV=NrtR (universal gas constant) = 8.31 kPa.L/mol.K

Standard Ambient Temperature and Pressure (SATP)Temperature 25°C/298KPressure 100kPa