Chemistry Final Review

Unit 3: Solutions Introduction − A solution is a homogeneous mixture of 2 or more things − Homogeneous mixture – components cannot be “seen”, equally mixed − Saturated solution – contains the maximum amount of solute it can have at a given temperature − Unsaturated solution – contains less than the maximum amount of solute − Supersaturated solution – contains more solute than it can dissolve − Solute – the substance that dissolves − Solvent – the dissolving medium −

Types of Solutions Types of Solutions

Solid Solute

Liquid Solute

Solid Solution (solid is the solvent) Gaseous Solution (gas is the solvent) Liquid Solution (liquid is the solvent)

Alloys (brass)

Amalgams (mercury in Charcoal filters silver) Water vapour in air Air (O2 in N2)

Microscopic particles in air Salt water, Kool-aid

Vinegar, rum + coke

Gaseous Solute

Carbonated drink

Properties of solutions 1. As temperature changes, so does the concentration of solute you can dissolve in a liquid solution. − As temperature rises, amount of solid solute rises − As temperature rises amount of gas solute lowers As pressure rises, amount of gas solute rises − Pressure has little effect on the amount of solid solute that can dissolve − If a solid dissolves in liquid, it’s said to be soluble; if it doesn’t dissolve, it’s said to be insoluble − If 2 liquids mix, they are miscible; if they don’t mix, they are immiscible 2. Freezing Point Depression (FPD) − A solution will freeze/melt at a temperature lower than the solvent would on its own 3. Boiling Point Elevation (BPE) − A solution will boil at a temperature higher than that of the solvent alone 4. Dissociation − Most ionic solutions will dissociate into ions when they dissolve (ex: NaCl(aq) → Na+ + Cl-) − Some covalent compounds ionize when they dissolve (CH3COOH(aq) → H+(aq) + CH3COO-(aq)) − These solutions will form electrolytes and conduct electricity − Not everything will dissociate nor form an electrolyte (ex: C12H22O11(aq) →C12H22O11(aq)) 5. The amount of solute in a solution can be quantitatively represented by calculating the solution’s concentration

− There are many acceptable units for concentration a. Molarity (M) = mol solute units: mol = M “molar” L solution L b. Molality (m) = mol solute units: mol = m “molal” kg solvent kg Examples: 1. Calculate the concentration (Molarity), of a solution that contains 74.6g of sodium carbonate in a 500.0mL solution. 500.0mL = 0.5000L 74.6g Na2CO3 x 1 mol Na2CO3 = 0.70377mol Na2CO3 106.0g Na2CO3 M = 0.70377mol Na2CO3 = 1.41mol/L = 1.41M 0.5000 L 2. How many grams of silver nitrate will remain if 750mL of 2.6M AgNO3 solution is evaporated to dryness? 750mL = 0.750L 2.6M = mol AgNO3 0.750L mol AgNO3 = 1.95 1.95 mol AgNO3 x 169.8g AgNO3 = 330g AgNO3 1 mol AgNO3 3. What volume of a 7.2M barium acetate solution would contain 19.3g of solute? 19.3g Ba(C2H3O2)2 x 1 mol Ba(C2H3O2)2 = 0.0755884542 mol Ba… 255.33g Ba(C2H3O2)2 7.2M = 0.07559 mol L L = 0.07559 mol 7.2M L = 0.010

Characteristics of Solutions − A solution may exist as gas, liquid, or solid depending on the state of its solvent − Most solutions are liquids − Reaction can take place in aqueous solutions – solutions in which reactants and products are mixed in water − Water – most common solvent among liquid solutions − A substance that dissolves in a solvent is said to be soluble in that solvent (ex: sugar is soluble in water) − A substance that does not dissolve in water is said to be insoluble in that solvent (ex: sand in water) − Oil is insoluble in vinegar; and thus, oil and vinegar are said to be immiscible − Two liquids that are soluble in each other are said to be miscible (ex: antifreeze)

Solvation in Aqueous Solutions − To form a solution, solute particles must separate from one another and the solute and solvent particles must mix

− Attractive forces exist between the pure solute particles, between the pure solvent particles and between the solute and solvent particles − When a solid solute is placed in a solvent, the solvent particles completely surround the surface of the solid solute − If the attractive forces between the solvent and solute particles are greater than the one holding the solute particles together, the solvent particles pull the solute particles apart and surround them − These surrounded particles then move away from the solid solute , out into the solution − The process of surrounding solute particles with solvent particles to form a solution is called solvation − Hydration – solvation in water − To determine whether a solvent and solute are alike, you must examine the bonding and the polarity of the particles and the intermolecular forces between the particles

Aqueous Solutions of Ionic Compounds − When a crystal of an ionic compound (ex: NaCl) is placed in a beaker of water, the water molecules collide with the surface of the crystal − The charged ends of the water molecules attract the positive sodium ions and the negative chloride ions − This attraction between the dipoles and the ions is greater than the attraction among the ions in the crystal, and so the ions break away from the surface − The water molecules surround the ions and the solvated ions turn into a solution − This exposes more ions on the surface of the crystal and salvation continues until the entire crystal has dissolved and all ions are distributed throughout the solvent − Gypsum is insoluble in water because the attractive forces among the ions in CaSO4 are so strong that they cannot be over come by attractive forces exerted by the water molecules, and thus solvation does not occur

Heat of Solution – one of the factors affecting rate of solvation − During the solvation, the solute must separate into particles − Solvent particles move apart in order to allow solute particles to come between them − Energy is required to overcome the attractive forces within the solute and within the solvent, so both steps are endothermic − When solute and solvent particles mix, the particles attract each other and energy is released – exothermic − The overall energy that occurs during the solution formation process is called heat of solution − Some solutions release energy as they form, whereas others absorb energy during formation − Ex: NH4NO3 dissolves in water – container feels cool CaCl2 dissolves in water – container feels warm

Making a Solution of Required Concentration Method 1: Having a solid solute and H2O Step 1:Calculate the mass of solid that you need. − Ex: To make 1.5L of 1.0M HCl solution 1.0M = mol

1.5L Step 2:Measure out that amount… with a scale! Step 3:Mix the solute with a small amount of H2O − Acid → add acid to water − Anything else → add water Step 4:Swirl until the solution dissolves; add more H2O if required. Step 5:Fill container to desired volume Method 2: Dilute a stock solution Step 1:Calculate the amount of stock solution that you need HOW? Dilute is not changing the number of moles of solute, but changing the volume of the solvent

Ex: 1.5L of 1.0M HCl from a 6M stock solution V1 = (1.5L)(1.0M) = 0.25L 6M

Electronegativity − EN increases on the periodic table as we travel from left to right, therefore the most negative group is the halogens − EN decreases as you travel down the periodic table − Group 18 (noble gases) has no electronegativity − Different atoms have differing abilities to attract electrons in a bond (electronegativity) − Electronegativity is the attraction an atom has for the shared pair of electrons in a covalent bond − When atoms with different ENs form a bond, the atom with the higher EN will draw electrons to itself in that bond − Electrons will not be equally shared − 1 atom will be more negatively charged than the other

Polar bonds − Bonds formed between atoms of differing EN − Electrons are shared unevenly − One end will have a partial positive charge and the other will have a partial negative charge − Ex:

Non-Polar Covalent Bonds − The only true covalent bonds are bonds in which electrons are shared equally − Ex: When 2 of the same atoms form a molecule (diatomic molecules)

Water – A Polar Molecule

− − − − −

Water is called the universal solvent because it can dissolve a large number of substances Unique structure Made up of 2 H atoms covalently bound to a single O atom H atoms and O atom each share a pair of electrons unevenly Oxygen is more EN than H → electrons lie more towards the oxygen

− The bent shape of the water molecule combined with 2 polar bonds makes water a polar molecule and a strong solvent

Non-Polar Molecules − Some molecules can contain polar bonds, but because of the shape are considered to be nonpolar molecules − Ex: Carbon tetrafluoride − F atoms are equally distanced − Each C→F bond is polar, but the polar bonds “balance out” or cancel out each other because they are symmetrical

Dissolving Process 1.Solvent particles move apart from other solvent particles - Requires energy to overcome forces of attraction → endothermic 2.Solute particles separate from the other solute particles - Requires energy to overcome forces of attraction → endothermic - The energy in this step is known as lattice energy (amount of energy required to separate the molecules or ions from each other in a solid crystal) 3. Solute particles move between the solvent particles - Attraction between solute and solvent releases energy → exothermic - This is known as the heat or energy of hydration - When water is the solvent = molecules/ions are said to be hydrated - When water is not the solvent the processes of the solvent particles surrounding the solute particles is called solvation

Energy Changes During Dissolving

− The total heat change in the dissolving process is called the heat of the solution − Heat of solution = the sum of the heat changes in the dissolving process − Sum of heat absorbed > heat released → endothermic − Container will feel cooler as the substance dissolves − Ex: Cold packs contain a larger bag of solid ammonium nitrate with a smaller inner bag of water − Sum of heat absorbed < heat released → exothermic − Ex: Hot packs contain a large bag of solid calcium chloride with a smaller inner bag of water

Effect of Solute on Vapour Pressure − Recall: evaporation is vapourization at the surface − In a closed container: − Evaporation and condensation are at equilibrium − The pressure created by the vapour is called vapour pressure *Pvap − When a solute is added to a solvent: [figure 2] − Solute is evenly distributed throughout solvent as it dissolves − Some solute particles will occupy the surface of the liquid − Less solvent particles at the surface − Decreases # of solvent particles that can evaporate − Rate of condensation is unchanged

Factors that Affect Solubility − Pressure affects the solubility of gaseous solute and gaseous solutions − Temperature affects the solubility of all substances

Pressure and Solubility − The solubility of a gas in any solvent increases as its external pressure (the pressure above the solution) increases − Soda pops contain CO2 dissolved in an aqueous solution − In bottling the carbonated beverages, CO2 is dissolved in the solution at a pressure higher than the atmospheric pressure − When the container is opened, the pressure in the neck of the bottle decreases → bubbles of CO2 gas form in the solution, rise to the top and escape − This will continue until the solution loses almost all of its CO2 gas and goes flat

Temperature and Solubility − Many substances are more soluable at high temperatures than low temperatures − Ex: CaCl2 has a solubility of about 64g CaCl2 per 100g water at 10°C → increasing the temperature to 27°C increases the solubility to 100g CaCl2 per 100g water − All gaseous solutes in liquid solvents are less soluble at a higher temperature than at a lower temperature − Kinetic energy of gas particles allows them to escape from a solution more readily at a higher temperature − As a solution’s temperature increases, the solubility of a gaseous solute decreases

− A supersaturated solution contains more dissolved solute than a saturated solution at the same temperature − A saturated solution is formed at a high temperature and then cooled slowly and this allows the excess solute to remain dissolved in solution at the lower temperature − Unstable → if a seed crystal (tiny solute) is added, the excess solute precipitates quickly − Crystallization can occur if the inside of the container is scratched or if it undergoes a physical shock, like stirring or tapping the container − Ex: rock candy, mineral deposits at the edges of mineral springs, cloud seeding

Henry’s Law − States that at a given temperature, the solubility (S) of a gas in a liquid is directly proportional to the pressure (P) of the gas above the liquid − Formula: S1 = S2 P1 P2 − S1 → solubility of a gas at a pressure P1 − S2 → solubility of the gas at the new pressure P2 − To solve for S2: S1 = S2 Cross-multiplying yields→ S1P2=P1S2 P1 P2 S2 = (P2S1)/P1

Unit 4: Gases and Atmosphere Properties of Gases (Ideal or Perfect Gases) 1. Least dense of any of the phases - mostly empty space 2. Takes the shape of its container - no definite volume 3. Easily compressed - if compressed too much they will liquefy (ex: liquid nitrogen) 4. Particles are in constant motion - in an ideal gas, Brownian Motion – motion in a straight line with perfectly elastic collisions 5. Gases exert pressure on the walls of their container - units: force/unit area = N/m2 = Pa 6. Gases are fluid - they can be poured 7. Expand and contrast significantly when the temperature is changing

Four Main Variables − − 1. 2.

4 main variables to predict behaviour of gases A change in any variable results in a change in another variable moles pressure - Standard Pressure – 101.325 kPa, 760 mmHg, 1 atm 3. volume

4. temperature - in Kelvin (0K = -273.15°C = absolute zero) - Standard Temperature → 273.15K = 0°C

Measuring Gas Pressure a) Measuring Atmospheric Pressure Device used: Barometer Torricelli Barometer – created by Evangelista Torricelli in 1600s − he was attempting to disprove that “nature abhors a vacuum” b) Measuring pressure for a sample of gas Device used: Manometer (can be either closed or opened) Two Types Three outcomes for open-ended manometer: - Pgas = Patm - Pgas > Patm → Pgas = Patm + x - Pgas < Patm → Pgas = Patm - x

Gas Laws Kinetic Theory − Gas particles do not attract or repel each other – Gases are free to move within their containers without interference from other particles − Gas particles are much smaller than the distances between them – Gases have no volume (mostly empty space) – Gases can be compressed by moving gas particles closer together because of this low density of particles − Gas particles are in constant random motion – They spread out and mix with each other because of this motion – They move in straight lines until they collide with each other or with the walls of their container − No kinetic energy is lost when gas particles collide with each other or with the walls of their container – As long as the temperature stays the same − All gases have the same average kinetic energy at a given temperature – As temperature rises, the total energy of the gas system rises – As temperature lowers, the total energy of the gas system lowers

Units of Measurement for Pressure − Millibar – meteorological unit of atmospheric pressure (1 mb = 1 atm) − Atmosphere – derived from standard atmospheric pressure at sea level (1 atm = 101.325 kPa = 760 mmHg) − mmHg is not commonly used, except in science labs – use ratio for conversions: 101.325 kPa = x kPa___ 760 mmHg x mmHg

Factors Affecting Gas Pressure

1. Volume and Pressure − “If moles and pressure are held constant, then the pressure of a gas is inversely proportional to the volume of the container that it occupies.” - Robert Boyle − Boyle’s Law: P ∝ 1/V; P1V1=P2V2 − Ex: What pressure would 3.0 mol of CO2 exert on a 5.0L container if it exerts 720mmHg in a 15.0L container? − (720mmHg)(15.0L) = P2(5.0L) 5.0L 5.0L – P2 = 2160mmHg – P2 = 2200mmHg

2. Volume and Temperature − “If moles and pressure are held constant, then volume and temperature exhibit a directly proportional relationship.” - Jacques Charles − Charles’s Law: V ∝ T; V1/T1 = V2T2 − Must be in Kelvin − Ex: A gas at 40.0°C occupies a volume of 2.32L. If the temperature is raised to 75.0°C, what will be the new volume of the gas? − (348K)2.32L = V2(348K) 313K 348K – V2 = 2.58L

3. Pressure and Moles − “If volume and temperature are held constant, than the pressure of gas and number of moles are directly proportional.” - Amadeo Avogadro − Avogadro’s Principle: P ∝ n; P1n1 = P2n2 − Ex: What pressure would 5.0 moles of H2 gas exert in a 36L container at 12°C if, under the same conditions, 3.0 moles of H2 gas exerted 1.9 atm of pressure? − 1.9 atm = _P2_ 3.0 mol 5.0 mol – P2 = 3.2 atm

4. Pressure and Temperature − “If moles and volume are held constant, then gas pressure (P), is directly proportional to the gas’ temperature (in K)” - Joseph Louis Gay-Lussac − Gay-Lussac’s Law: P ∝ T; P1/T1 = P2/T2

5. Combined Gas Law − P1V1 = P2V2 n1T1 n2T2 − Combined law of the four formulas − Use STP, 22.4L, and 1 mol as one side as default values − Temperature is in Kelvin, use pressure units in question − Ex: 1.0g of an unknown gas occupies a volume of 4.2L at 180°C and 95.0 kPa. Calculate its molecular mass. − (95.0kPa)(4.2L) = (101.325kPa)(22.4L) n1(453.15K) (1 mol)(273.15K) – n1 = 0.106…mol – MM = (1.09)/(0.106…) = 9.4μ

6. Ideal Gas Law

− Use STP, 22.4L, and 1 mol as one side of combined gas law − P1V1 = R= 8.314 n1T1 − PV = nRT − Ex: 1.0g of an unknown gas occupies a volume of 4.2L at 180°C and 95.0 kPa. Calculate its molecular mass. − (95kPa)(4.2L) = n(8.314)(453.15K) – n = 0.106…mol – MM = (1.09)/(0.106…) = 9.4μ

The Scientists − 1564–1642: Galileo Galilei developed the suction pump. − 1643: Evangelista Torricelli developed the first barometer. − 1643–1645: Otto von Guericke made a pump that could create a vacuum so strong that a team of 16 horses could not pull two metal hemispheres apart. He reasoned that the hemispheres were held together by the mechanical force of the atmospheric pressure rather than the vacuum. − 1648: Blaise Pascal used Torricelli’s “barometer” and traveled up and down a mountain in southern France. He discovered that the pressure of the atmosphere increased as he moved down the mountain. − 1661: Christiaan Huygens developed the manometer to study the elastic forces in gases. − 1801: John Dalton stated that in a mixture of gases the total pressure (Patm) is equal to the sum of the pressure of each gas (P1 + P2 + P3…), as if it were in a container alone. The pressure exerted by each gas is called its partial pressure. − 1808: Joseph Louis Gay-Lussac observed the law of combining volumes. He noticed that, for example, two volumes of hydrogen combined with one volume of oxygen to form two volumes of water. (2H2 + O2 → 2H2O) − 1811: Amadeo Avogadro suggested, from Gay-Lussac’s experiments conducted three years earlier, that the pressure in a container is directly proportional to the number of particles in that container.

Applications of Gases − Airbags – has NaN3, KNO3, SiO2 – electrical impulse breaks NaN3 into Na and N2, causing the nitrogen gas to fill up the bag − Airships – H or He − Propane appliances – mostly C3H8 − SCUBA – O2, N2, and a small mix of others to make regular air

Unit 5: Organic Chemistry − The chemistry of carbon-containing compounds

Alkanes − − − −

Simplest of the hydrocarbons Only contain carbon-carbon single bonds General formula:CnH2n+2 Simplest: methane CH4 – molecular formula

− − − − − − −

- structural formula CH4 – condensed formula (CH3CH2CH3 – propane) Different prefixes for 1 – 4 (1 – meth; 2 – eth; 3 – prop; 4 – but) If the hydrogens are not visible in the structural formula, assume they are there Straight chained hydrocarbons – without branches (ex: butane) Branched hydrocarbons contain branches (ex: methylpropane) Structural isomer – same components but different structure (ex: methyl propane is a structural isomer of butane)

IUPAC Rules for Naming 1. Find the longest continuous chain of carbons 2. Number the carbons in the longest chain so that the branches are on the lowest numbered possible 3. Put the pieces of the name together Ex: 2,4-dimethylhexane NOTE: 1. If there is only one place to put the branch, omit the # 2. If you have two or more lengths of branches, list alphabetically 3. If you have multiple numbers of the same branch, you can use prefixes (mono, di, tri…) to indicate how many you have

Alkenes − − − − − −

CnH2n Contain a carbon-carbon double bond Are “unsaturated hydrocarbons” Simplest: Ethene C2H4 All carbons must have 4 bonds When naming, the double bond should be off of the lowest number (must pass through)

− Ex: 2-methyl-1-propene 2-methylpropene Methylpropene Naming: 1. If 2 or more bonds are possible, you must indicate where the double bond is 2. When numbering carbons in the longest chain: a. Make sure to indicate the lowest double bond b. The double bond gets the lowest numbered carbon possible

Cis-Trans (Geometric) Isomers − When you have a double bond, different arrangements can exist around the double bond − Geometric isomers exist if each carbon in the double bond is bonded to two different groups

− − Cis – has groups attached to the common side of the double bond − Trans – has groups attached to carbons on the opposite side of the double bond − For naming, add the prefix “cis-“ or “trans-“ before the number that indicates where the double bond is

Alkynes − − − −

General formula: CnH2n-2 Contain at least one triple bond Considered to be an unsaturated hydrocarbon The triple bond makes these compounds more reactive