Units 1-2

Unit 1-2 Review

Vocabulary

Matter - Anything that occupies space and has mass; the physical material of the universe. Atoms- the smallest representative particle of an element. Molecule-a chemical combination of two or more atoms. Pure substances- matter that has a fixed composition and distinct properties. Law of constant composition (law of definite proportions)- chemical formulas are in constant ratio/proportions to each other (Joseph Proust) A given compound always contains exactly the same proportions of the elements. Ex. Water always has 2 hydrogen and 1 oxygen. Mixtures - A combination of two or more substances in which each substance retains its own chemical identity. Solutions - homogenous mixtures Metric system- A system of measurement used in science and in most countries. The meter and the gram are examples. SI units- The preferred metric units for use in science. Mass- A measure of the amount of material in an object. It measures the resistance of an object to being moved. In SI units, mass is measured in kilograms. Precision- The closeness of agreement among several measurements of the same quantity; the reproducibility of a measurement. Accuracy-A measure of how closely individual measurements agree with the correct value. Subatomic particles- Particles such as protons, neutrons, and electrons that are smaller than an atom. Cathode rays- Streams of electrons that are produced when a high voltage is applied to electrodes in an evacuated tube. Radioactivity- The spontaneous disintegration of an unstable atomic nucleus with accompanying emission of radiation. Nucleus- The very small, very dense, positively charged portion of an atom; it’s composed of protons and neutrons. Protons- A positively charged subatomic particle found in the nucleus of an atom. Electrons- A negatively charged subatomic particle found outside the atomic nucleus; it’s a part of all atoms. An electron has a mass 1/1836 times that of a proton. Atomic mass units- (amu) A unit based on the value of exactly 12 amu for the mass of the isotope of carbon that has six protons and six neutrons in the nucleus. Mass number- The sum of the number of protons and neutrons in the nucleus of a particular atom. Isotopes- Atoms of the same element containing different numbers of neutrons and therefore having different masses. Metallic elements- Elements that are usually solids at room temperature, exhibit high electrical and heat conductivity, and appear lustrous. Most of the elements in the periodic table are metals. Nonmetallic elements- Elements in the upper right corner of the periodic table; nonmetals differ from metals in their physical and chemical properties. Diatomic molecule- A molecule composed of only two atoms. each element in one molecule of a substance.

Ionic compounds- A compound composed of cations and anions. Polyatomic ions- An electrically charged group of two or more atoms. Alkanes- Compounds of carbon and hydrogen containing only carbon-carbon single bonds. Alcohols-An organic compound contained by substituting a hydroxyl group (-OH) for a hydrogen on a hydrocarbon. Hydrocarbons- Compounds composed of only carbon and hydrogen. Law of conservation of mass/matter- Antoine Lavoisier- Made very precise measurements and developed law. Mass/matter is not created nor destroyed through a chemical reaction. Chemical reaction atoms rearrange. Simple Alcohols- contain carbon, hydrogen, and oxygen bonded with single covalent bonds. The number of carbons determines the root name of the alcohol.

Formulas/Conversions Mass number = # of protons + # of neutrons Density = Mass

Volume

Average Atomic mass = (massA) (%) + (mass B) (%) + …

Cn H2n+2 (formula to find alkanes)

(Simple Alcohol formula) Cn H2n+1 OH Kelvin = 273.15 + C° 1 Liter= (10 cm3) = 1000 cm3 = 1 dm3 1 cm3 = 1 mL

Big Concepts Uncertainty in Measurements 1. Record certain digits, and one uncertain digit 2. Measurements always have uncertainty. Degree of uncertainty is a function of the precision

of measure device. 3. Uncertainty of last digit is always assumed to be (+|-)1 unless otherwise reported. Rules for Significant Figures 1. Non zeroes ALWAYS important 2. Leading 0’s that precede non 0 digits are NEVER significant. Ex. 0.042 has 2 significant

figures 3. Captive 0’s between nonzero digits are ALWAYS important. 4. Trailing 0’s at the end of a number are only significant if the numbers is written with decimal

point. Ex. 10. (2 digits) 5. Exact numbers have unlimited number of significant figures. Exact numbers result from

counting (Ex. Apples) or by definition (12 = 1 dozen) 6. In scientific notation, the 10^X part of the number is never counted as significant. 7. Multiply/Divide- limit answer to fewest sig. fig in original numbers 8. Add/Subtract- limit answer to same number of decimal places that appear in the original data

with the fewest number of decimal places. Separating Mixtures Mixtures can be separated into pure substances by physical means 1. Filtration- separation of a solid from a liquid using filter paper, funnel, gravity. 2. Distillation- Using the volatility (boiling point) of different substances for separation. 3. Chromatography- using the relative affinity of particle for the medium that they must pass

through. Ancient History � Democritus (400BC) - first proposed that matter was made of some “smallest” particle- the

atom. � Alchemists - primarily focused on trying metals into gold. Learned to prepare the mineral

acids. � Robert Boyle - measured relationships between pressure and volume of air. � Joseph Priestly - discovered oxygen, disproved “phlogiston” theory. Dalton’s Atomic Theory 1. Each element is made up of tiny particles called atoms. (Democritus’ idea.) 2. Atoms of a given element are identical/ atoms of different elements are different in some

fundamental way or ways. (NOT TRUE- ISOTOPES!!) 3. Chemical compounds form when atoms combine with each other. A given compound always

has the same relative number and types of atoms. (Proust’s idea) 4. Chemical reactions involve reorganizing the atoms. The atoms do not change in a chemical

reaction. (Lavoiser-Law of conservation of mass)

Avogadro’s Principle � At same temp and pressure, equal volumes of different gases contain the same number of

particles. � This only made sense if the distance between particles was very great compared to the size of

the particles. Otherwise, the different sizes of particles would impact the volume of the gas. JJ Thompson- The Electron � Use cathode ray tube- partially evacuated. When high voltage was applied, a ray was

produced a the negative end. This ray was repelled by negatives. “Like repels like; opposites attract”

Plum Pudding Model 1. Thompson reasoned that since all metals he experimented with produced electrons, all atoms

must contain electrons. 2. If an atom is neutral, there must be something positive in the atom as well. 3. Therefore, Thompson proposed the plum pudding model (chocolate chip cookie) Ernest Rutherford - Finding the Nucleus (Gold Foil Experiment) � Observation - most particles passed straight through. Conclusion: most of the atom is empty

space. � Observed : few particles bounced back (like repels like) � Conclusion: Particles hit a very tiny but dense center (nucleus) Since He+ is positive, center

must be (+) � Observed: Few particles veered slightly off path � Conclusion: Since particles were partially charge and like charges repel, the nucleus must

contain positive charge. Naming Acids � Rules for naming acids depend on whether the anion contains oxygen or not. � If there are no oxygens: The acid is named with the prefix hydro- and the suffix -ic. � If oxygens are present: If the anions ends in -ate, use the suffix -ic. If the anion ends in -ite,

use suffix -ous.

Practice Problems

How many significant digits are shown: 1. 0.459 2. 0.45900 3. 459,000 4. Average Atomic Mass problem: What’s the (aam) of silver if one isotope has a mass of 106.90509 amu and is 51.86% abundant and the other isotope has a mass of 108.90470 and is 48.14% abundant? 5. Lithium with 2 isotopes, Lithium-6 and Lithium-7 has an aam of 6.124 amu. Lithium-6 has a mass of 6.047 amu and Llithium-7 has a mass of 7.056 amu. 6. Write the formula for Pentanol

Solution Section

1. 3 digits 2. 5 digits 3. 3 digits 4. (106.90509)(.5186) + (108.90470)(.4814) = 107.87 aam 5. 6.124 amu = (6.047)(x) + (7.056)(1-x) = 92.37% Li - 6 ; 7.630% Li -7 6. C5H11OH

Unit 3

Unit 3 Key Terms Stoichiometry- The relationships among the quantities of reactants and products involved in chemical reactions. Law of Conservation of Mass- A scientific law stating that the total mass of the products of a chemical reaction is the same as the total mass of the reactants, so that mass remains constant during the reaction. Chemical Equations- Chemical reactions are represented in a concise way. Reactants- The chemical formulas on the left of the arrow represent the starting substances. Products- The chemical formulas on the right of the arrow represent substances produced in the reaction. Combination Reaction- A reaction in which two or more substances react to form one product. Decomposition Reaction- A reaction in which one substance undergoes a reaction to produce two or more other substances. Combustion Reaction- A rapid reaction that produces a flame. Formula Weight- The mass of the collection of atoms represented by a chemical formula. Molecular Weight- The mass of the collection of atoms represented by the chemical formula for a molecule. Mole- A collection of Avogadro’s number (6.02·1023) of objects. Avogadro’s Number- The number of 12C atoms in exactly 12 g of 12C; it equals 6.02·1023. Molar Mass- The mass of one mole of a substance in grams; it is numerically equal to the formula weight in atomic mass units. Limiting Reactant (or Limiting Reagent)- The reactant present in the smallest stoichiometric quantity in a mixture of reactants; the amount of product that can form is limited by the complete consumption of the limiting reactant. Theoretical Yield- The quantity of product that is calculated to form when all of the limiting reactant reacts. Percent Yield- The ratio of the actual yield of a product to its theoretical yield, multiplied by 100.

Not on AP Info Sheet Percentage Composition- %Element= (Number of atoms of that element)(Atomic weight of element)/(Formula weight of compound) · 100% Percent Yield= (Actual Yield/Theoretical Yield) ·100% Vodcast 1: The Mole A mole of an atom has 6.02214 x 1023 atoms per mole. Molar mass is the mass of 1 mole of any substance. The percent mass of is 100 times the mass of the element in 1 mole divided by the mass of 1 mole of the molecule. Vodcast 2: Empirical and Molecular Formulas Difference between an empirical formula and a molecular formula: Empirical Formula: The formula of a substance in which the atoms are at lowest whole number ratio. Molecular Formula: The formula of a substance based on what the molecule looks like. Steps to finding the empirical formula: 1 – Change the percent composition to grams. 2 – Convert the grams from #1 to moles. 3 – Make a mole ratio of whole numbers from the moles in #2. 4 – The ratio of the elements to each other are the subscripts. To determine the molecular formula from the empirical formula: 1 – Determine the formula mass of the empirical formula. 2 – Divide the number by the molar mass of the molecular compound. 3 – Multiply the formula by the number from #2. Vodcast 3: Chemical Reactions: *In a chemical reaction, atoms are neither created nor destroyed. Chemical bonds are broken and atoms are rearranged. States of matter should be included in a chemical reaction [Solid (s); liquid (l); gas (g); dissolved in water –or aqueous- (aq)] When balancing a reaction, the formulas must never change. The most complicated substance should always go first in the equation. Types of reactions: Addition Reactions (Synthesis): Two or more substances or elements chemically react to form a single product. A + B à AB Decomposition Reactions :A substance breaks down into the elements or compounds that make up that substance. AB à A + B Single Replacement: One element replaces another element in a compound A + BC à AC + B Double Replacement: Two substances react with each other to form two new compounds. All double replacement reactions must have a source of energy to remove a pair of ions from the solution. AB + CD à CB + AD Combustion Reaction: Elements or compounds combine with oxygen to make a new substance.

Example: XH + O2 à XO2 + H2O Oxidation Reduction Reaction (Redox reactions): Electrons are transferred during the reaction. Therefore the oxidation numbers of at least two elements of the equation must change. Redox reactions can also be single replacement, synthesis, and decomposition reactions. Acid-Base Neutralization Reactions: Acids react with bases to produce a salt and water. One mole of H+ will react with one mole of OH- to produce one mole of water (H2O). Vodcast 4 Stoichiometry: Reaction Stoichiometry: Calculating amounts of reactants and products of a reaction. Steps: 1 – Balance the reaction 2 – Convert known substance to moles 3 – Used balanced reaction to set up appropriate mole ratios 4 – Convert from moles of new substance to appropriate units Limiting Reactant – The reactant that is consumed first that determines the amount of product formed Theoretical Yield – the amount of product formed in the lab Percent Yield – the ration of actual yield to theoretical yield Unit 3 Practice Problems

1. Isopentyl acetate, the compound responsible for the scent in bananas, can be produced commercially. Interestingly enough, bees release about 1 µg of this compound when they sting. The resulting scent attracts other bees to the attack. How many molecules of isopentyl acetate are released in a typical sting? How many atoms of carbon are present?

2. Determine the percent composition for ethanol. 3. What is the empirical formula when mercury forms a compound with chlorine that is

73.9% Hg and 26.1% Cl? 4. Suppose that a compound is composed of carbon, hydrogen, and oxygen. The compound

is 0.4092 g C, 0.0458 g H, and 0.5450 g O. What is the empirical formula for this compound?

5. Determine the molecular formula for the compound with the following data: 24.27% C 4.07% H

71.65% Cl Molar Mass = 98.96 g/mol

6. Adipic acid (H2C6H8O4) is used to produce nylon. The acid is made commercially by a controlled reaction between cyclohexane and oxygen. If 25.0 g of C6H12 and 25.0 g O2 are used, how many grams of H2C6H8O4 can be produced? 2C6H12 (l) + 5O2 (g) � 2H2C6H8O4 (l) + 2H2O (l)

7. Methanol is used as fuel in race cars. It can be manufactured by combining gaseous carbon monoxide and hydrogen. Suppose 8.60 kg H2 is reacted with excess CO. Calculate the theoretical yield of methanol. If 3.57 • 104 g CH3OH is actually produced, what is the percentage yield?

CO (g) + 2H2 (g) � CH3OH (l)

Unit 3 Solutions 1.

2.

C = (12.01)(2)

• 100 = 52.13% (46.08)

H = (1.01)(5)

• 100 = 13.15% (46.08)

O = (16.00)(1)

• 100 = 34.72% (46.08)

3.

73.9 g Hg 1 mol Hg = 0.368 mol Hg

.368 = 1

200.59 Hg .368

26.1 g Cl 1 mol Cl 0.736 mol Cl

0.736 = 2

35.45 g Cl .368

Mercury (II) Chloride, HgCl 2

4. 0.4092 g C 1 mol C 3.407 mol C

=1 • 3 =3 12.01 g C 3.406

0.0458 g H 1 mol H 4.53 mol H =1.33 • 3 = 6

1.01 g H 3.406

0.5450 g O 1 mol O 3.406 mol O = 1 • 3 = 3

16.00 g O 3.406 C3H6O3

5. 24.27 g C 1 mol C = 2.021 mol C

=1 12.01 g C 2.021

1 µg C7H12O2 1 g 1 mol 6.02 • 1023 = 5 • 1015 molecules C7H12O2 106 µg 128.19 g 1 mol

5 • 1015 molecules C7H12O2 7 carbon atoms = 3 • 1016 atoms carbon 1 molecule

7.65 g Cl 1 mol = 2.021 Cl =1

35.45 g 2.021 CH2Cl 12.01 + 2(1.01) + 35.45 = molar mass of 49.48 g

98.96 g/mol = 2

49.48 g/mol (CH2Cl) • 2 = C2H4Cl2

6.

25.0 g C6H12 1 mol C6H12 2 mol A. Acid 146.16 g A. Acid = 43.4 g A. Acid

84.18 g C6H12 2 mol C6H12 1 mol A. Acid

25.0 g O2 1 mol O2 2 mol A. Acid 146.1 g A. Acid = 45.7 g A. Acid

52.00 g O2 5 mol O2 1 mol A. Acid C6H12 is the limiting reactant.

7. 8.60 kg H2 1000 g 1 mol H2 1 mol CH3OH 32.05 g CH3OH

= 68,200 g CH3OH 1 kg 2.02 g H2 2 mol H2 1 mol CH3OH

3.57 • 104 g

• 100 = 52.3 % 68,200 g

4.07 g H 1 mol H = 4.03 mol H = 2

1.01 g H 2.021

Unit 4

Unit 4-Reactions in water Vocab: Dissolving (Hydration): When the positive ends of water molecules attracted to the

anion in a salt. Or when the negative end of water molecules attracted to the cation in a salt.

Solubility: The quantity of a particular substance that can dissolve in a particular solvent.

(Solubility constant is Ksp pg. 1045)

Electrolytes: Solutions in a homogeneous mixture.

Solute: Substance being dissolved.

Solvent: Substance doing the dissolving.

Bronsted-Lowry Acid: Donates a Hydrogen.

Bronsted-Lowry Base: Accepts a Hydrogen.

Amphoteric-Can act as an acid or an base.

Polyprotic acid: Acids that can give more than one Hydrogen.

Re-dox reaction: Reaction when electrons are transfered between elements.

Titration: Process of reacting a solution of unknown concentration with one of a known

concentration.

Formulas: M1V1=M2V2

MaVaCb=MbVbCa

M=mol/L

Big Concept: Several types of reactions occur in water. When water is the solvent in the solution then it is an aqueous solution. Three other reactions are precipitation, acid base, and oxidation-reduction(redox). Strong electrolytes are strong acids and strong bases and dissolves completely in water Weak electrolytes are weak acids and weak bases and dissolves partially in water

Molecular compounds tend to be nonelectrolytes When a compound in a reaction is insoluble then a precipitate is formed The molecular equation lists the reactions and products in their molecular form To write net ionics cross out anything that remains unchanged from left side to right and then write the equation with what is left In an acid-base reaction the acid donates a hydrogen ion to the base An oxidation occurs when an atom or an ion loses an electron A reduction occurs when an atom or an ion gains an electron LEO goes GER Lose Electrons Oxidize, Gain Electrons Reduce Two solutions could have the same compounds but be different due to the molarity Molarity measures the concentration of a solution Practice Problems:

1. Why is salt water a stronger electrolyte than sugar water?

2. What kind of solutions are strong electrolytes?

3. If a solute partially dissolves in water, what kind of electrolyte is the resulting solution?

What would the arrow for the reaction of this compound dissolving look like?

4. What has to be in the products for a double replacement reaction to happen in solution?

5. Why does a precipitate form in double replacement reactions?

6. What kinds of compounds are ionized in chemical reactions?

7. What is a Bronsted-Lowry acid and a Bronsted-Lowry base?

8. What are the strong acids and bases? What do they make when they react with eachother?

9. How can you tell if a redox reaction happens in acidic, basic, or neutral solution?

10. What is the oxidation number of Cr in CrO42-?

11. What would the chromate ion become in a redox reaction and what would its new

oxidation number be?

12. Which has a higher molarity of H+ ions, 1 liter of 0.3M H3PO4 or 1 liter of 0.6M HCl?

13. Describe how to do a dilution.

14. What Volume of 0.115M HClO4 solution is needed to neutralize 50.00 mL of 0.0875M

NaOH?

Solutions:

1. Salt is an ionic solute which has ions dissolved in water that conduct electrons better than

sugar molecules, which don’t have any ions.

2. Strong acids and bases and ionic compounds in solution.

3. A weak electrolyte. The arrow would point in both directions.

4. A precipitate, gas, or water.

5. Because pairs of oppositely charged ions attract each other so strongly that they make an

insoluble compound.

6. Strong acids, strong bases, and soluble ionic compounds, all in solution.

7. Bronsted-Lowry acids are substances that ionize a H+ ion in solution. Bronsted-Lowry

bases are substances that accept or react with a H+ in solution.

8. Strong acids: hydrochloric acid, hydrobromic acid, hydroiodic acid, nitric acid, sufuric

acid, and perchloric acid.

Strong bases: group 1 metal hydroxides and heavy group 2 metal (Ca and down)

hydroxide.

They react with eachother to make a salt and water.

9. By looking for H+ ions or OH- ions for acidic or basic, respectively. If neither ions are

present, the solution is neutral.

10. +6

11. Cr3+ and 3+.

12. 1 liter of 0.3M H3PO4.

13. First place the solution in a volumetric flask, then slowly add water and swirl until the

bottom of the meniscus reaches the mark on the volumetric flask.

14. 38.0 mL of 0.115M HClO4

Unit 5

Unit 5 - Gases

Big Concepts:

1. Gas Variables

• If amount of gas doubles, from 1 mol to 2 mol, then the volume

increases.

• If the Kelvin temperature is doubled, then the volume increases.

• If one mol of gas pressure is doubled, then the volume decreases.

2. How do gases resemble each other?

• Gas particles are very spread out, this allows them to be compressible.

• Particles are in constant random motion.

3. When using PV=nRT, you must understand:

• Temperature must be expressed in absolute units… AKA, Kelvin!

(K=°C+273.15)

• Volume must be measured in Liters.

• Amount of gas must be expressed in moles � NOT grams.

• Pressure must be measured in atmospheres! (P= Force/Area)

• The “R” stands for the ideal gas constant- (L · atm / mol K)

4. Gas Laws:

• Boyle’s Law: P1V1 = P2V2 (Ex: Breathing, popping a balloon, balloon in bell

jar)

• Charles’ Law: V1 / T1 = V2 / T2 (Ex: Pool ring, hot air balloons, balloons in the

cold)

• Gay-Lussac’s Law: P1 / T1 = P2 / T2 (Ex: Cold tire = low pressure, Metal

canister warnings, egg in the bottle)

• Combined Gas Law: P1V1 / T1 = P2V2 / T2

*When using any of these laws, you do not have to be in liters, moles, or

atmospheres, but you must always measure in Kelvin!

5. PV=nRT is used when gases are not at STP. At STP, 1 mol of gas takes up 22.4 L

of volume.

6. Kinetic Theory of Gases:

• Gases consist of particles in continuous, random motion. *There are frequent

collisions with one another & the container.

• Collisions between gas particles are complete elastic; one particle may slow

down but the other will speed up. *No kinetic energy converted to heat.

• Average energy of motion of the gas particle is directly proportional to

temperature.

7. Kinetic energy:

• ≠ speed of molecules

• = Temperature

Formulas

Kinetic Energy (per molecule)** (KE= Kinetic Energy, m= mass, v= velocity, aka speed)

KE = ½ mv2

Temperature in Kelvin** K= ºC + 273.15

Pressure (F= force, A= area) P = F / A

Dalton’s Gas Law** Ptotal = Pa + Pb

Combined Gas Law** P1V1 / T1 = P2V2 / T2

Ideal Gas Law** PV = nRT

**Included on the AP Resource Sheet

Vocabulary:

• Thermodynamics- the study of energy and its transformations

• Thermochemistry- the transformations of energy, especially heat, during

chemical reactions

• Kinetic energy- the energy due to motion of the object

• Potential energy- the energy that an object possesses by virtue of its position

relative to other objects

• Joule- SI unit of energy 1 J = 1 kg-m2/s2

• Calorie- another SI unit of energy, Originally defined as the quantity of energy

necessary to increase the temperature of 1 g of water by 1 degree C: 1 cal = 4.184

Joule

• System- a specific amount of matter

• Surroundings- everything outside the system

• Work- the energy expended to move an object against a force

• Heat- the energy that is transferred from a hotter object to a colder one

• Energy- the capacity to do work or to transfer heat

• Internal energy- the sum of all the kinetic and potential energies of its

component parts

• Endothermic- the system absorbs heat from the surroundings

• Exothermic- the system releases heat to the surroundings

• State of function- the internal energy (E)

• Enthalpy- thermodynamic function that accounts for heat flow in chemical

changes occurring at constant pressure when no forms of work are performed

other than P-V work

• Enthalpy of reaction- the enthalpy of the products minus the enthalpy of the

reactants:

• ∆Hrxn = H (products) - H (reactants)

• Calorimetry- the amount of heat transferred between the system and the

surroundings that is measured experimentally

• Calorimeter- measures the temperature change accompanying a process

• Heat Capacity- the amount of heat required to raise its temperature by 1 K

• Molar Heat Capacity- the heat capacity for 1 mol of a pure substance

• Bomb Calorimeter- constant-volume calorimetry that is carried out in a vessel

of a fixed volume

• Hess’s Law- if a reaction is carried out in a series of steps, ∆H for the reaction

will be equal to the sum of enthalpy changes for the steps

• Enthalpy of Formation- the enthalpy change for the reaction in which the

substance is formed from its constituent elements ( ∆Hf )

• Standard Enthalpy- the enthalpy change when all reactants and products are a

1 atm pressure and a specific temperature, usually 298 K ( ∆H○ )

• Standard Enthalpy of Formation- the change in enthalpy for the reaction

that forms 1 mol of the substance from its elements with all reactants and

products at 1 atm pressure and usually 298 K ( ∆H○

f )

• Fuel Value- the heat released when 1 g of the substance is combusted

Practice Problems:

1. The temperature of a sample of an ideal gas confined in a 2.0 L container was raised from 27˚C to 77˚C. If the initial pressure of the gas was 1200 mmHg, what was the final pressure of the gas?

(A) 300 mmHg (B) 600 mmHg (C) 1400 mmHg (D) 2400 mmHg (E) 3600mmHg

2. Nitrogen gas was collected over water at 25˚C. If the vapor pressure of water at

25˚C is 23 mmHg, and the total pressure in the container is measured at 781 mmHg, what is the partial pressure of the nitrogen gas?

(A) 23 mmHg (B) 46 mmHg (C) 551 mmHg (D) 735 mmHg (E) 758 mmHg

3. If the atmospheric pressure is 0.975 atm, what is the pressure of the enclosed gas

in each of the three cases depicted in the drawing?

(i) (ii) (iii)

4. Chlorine is widely used to purify municipal water supplies and to treat swimming pool waters. Suppose that the volume of a particular sample of Cl2 gas is 9.22 L at 1124 torr and 24˚C. How many grams of Cl2 are in the sample?

5. Which gas is the most dense at 1.00 atm and 298 K? CO2, N2O, or Cl2?

6. How does a gas differ from a liquid with respect to density and compressibility?

Solutions:

1. C

1200 mmHg

= P2

300 K 350 K P2 = 1400 mmHg

2. E 781 mmHg – 23 mmHg = 758 mmHg

3.

(i) 52 cm 10 mm = 520 mm Pgas = 0.975 - 520 mmHg = 0.291 atm 1 cm 760 mm Hg

(ii) Pgas = 0.975 + 67 mmHg = 1.063 atm 760 mm Hg

(iii) 10.3 cm 10 mm = 103 mm Pgas = 0.975 + 103 mmHg = 0.136 atm 1 cm 760 mm Hg

4. 1124 torr 9.22 L = n 0.0821 Latm/mol K 297 K

760 torr n = 0.5592 mol Cl2

0.5595 mol Cl2 70.906 g Cl2 = 39.7 g Cl2 1 mol Cl2

5. Cl2 because it has the greatest molar mass. 6. Gases have a lower density and more compressibility than liquids.

Unit 6

Unit Six: Electrons Vocabulary: Electronic structure: Describes the energies and arrangement of the electrons around an atom. Electromagnetic radiation: A form of energy that has wave characteristics that propagates through a vacuum at the characteristic speed of 3.00x108 m/s

Wavelength: The distance between identical points on successive waves. Quantum number: The smallest increment of radiant energy that may be absorbed or emitted; the magnitude of radiant energy is hv. Photons: Smallest increment of radiant energy; a photon of light with frequency v as in energy = hv. Spectrum: Distribution among various wavelengths of the radiant energy emitted or absorbed by an object. Line spectrum: A spectrum that contains radiation at only certain specific wavelengths. Ground state: The lowest energy, or most stable state. Excited state: Higher energy state then the ground state. Momentum: The product of mass and velocity of a particle. Wave Functions: A mathematical description of an allowed energy state for an electron in the quantum mechanical model of the atom. Orbital: An allowed energy state of an electron in the quantum mechanical model of the atom; the term orbital is used to describe the spatial distribution of the electron. An orbital is defined by the values of the three quantum numbers. Electron shell: A collection of orbital’s that have the same value of n. Subshell: One or more orbitals with the same set of quantum numbers. Electron spin: A property of the electron that makes it behave as though it were a tiny magnet. The electron behaves as if it were spinning on its axis. Pauli exclusion principle: A rule stating that no two electrons in any atom may have the same four quantum numbers. Electron configuration: A particular arrangement of electrons in the orbital’s of an atom.

Hund’s rule: A rule stating that electrons occupy degenerate orbital’s in such a way as to maximize the number of electrons with the same spin. Formulas On the AP chemistry resource sheet:

E = hU C = λ*U λ = h / mU p = mU En = -2.178 x 10-18 / n2 joules Not on the AP chemistry resource sheet: En = Rh (1/n2) E f-I = Rh (1/nf

2 – 1/ni2)

D = s*t S = d/t

Big Concept = -Understanding how light interacts with matter. This provides insight into the behavior of electrons in atoms. -The fact that atoms give off characteristic colors of light (light spectra) provides clues about how electrons are arranged in atoms, leading to two important ideas: Electrons exist only in certain energy levels around nuclei, and energy is involved in moving an electron from one level to another. - How electrons are arranged in atoms is described by quantum mechanics in terms of orbital’s. - Knowing the energies of orbital’s as well as some fundamental characteristics of electrons allows us to determine the ways in which electrons are distributed among various orbital’s in an atom (electron configuration). - The electron configuration of an atom is related to the location of the element on the periodic table. Practice problems =

1) Which of the following sets of quantum numbers is unacceptable? A) (4, 3, -2, +1/2) B) (3, 0, +1, -1/2) C) (3, 0, 0, -1/2) D) (3, 1, `, +1/2) E) (2, 0, 0, -1/2)

2) In a given atom , no two electrons can have the same set of four quantum numbers. This statement is known as the

A) Pauli exclusion principle B) Hund’s Rule C) Einstein principle D) Heisenberg uncertainty principle E) Bohr law

3) The element with the ground state electron configuration of [Ar] 3d74s2 is A) Mg B) K C) Ar D) Co E) Ni

4) Atomic radii decrease from left to right across a period because of A) an increase in ENC B) an increase in gross energy level C) an increase in subenergy level D) an increase in shielding E) a decrease in effective nuclear charge

5) The correct ordering of atoms in progressively decreasing ionization energy is A) F > O > C > Li > Na B) Na > Li > C > O > F C) F > O > C > Na > Li D) C > O > F > Li > Na E) O > F > C > Na > Li

Answers; 1) B 2) A 3) D 4) A 5) A

Unit 7

Unit 7—Periodicity

Vocabulary: • Effective Nuclear Charge: The electric field is equivalent to one generated by a charge

located at the nucleus.

• Bonding Atomic Radius: The radius of an atom, defined by the distances separating it from the other atoms to which it is chemically bonded.

• Isoelectronic Series: A series of atoms, ions, or molecules that has the same number of

electrons.

• Ionization Energy: The energy required to remove an electron from a gaseous atom when the atom is in its ground state.

• Electron Affinity: The energy change that occurs when an electron is added to a gaseous

atom or ion.

• Metallic Character: When an element exhibits the physical and chemical properties characteristics of metals, like luster, malleability, ductility, and electrical conductivity.

NO FORMULAS NEEDED FOR THIS UNIT.

Big Concept: Effective Nuclear Charge: Going Left � Right on the periodic table increases because of added protons to the nucleus. Going Top � Bottom on the periodic table decreases because of the shielding effect. Atomic Radius: Going Left � Right on the periodic table size decreases because of the increased effective nuclear charge. Going Top � Bottom on the periodic table size increases because of the increased shielding effect and there are more energy levels. Ionization Energy: Going Left � Right on the periodic table, energy increases. Going Top � Bottom on the periodic table, energy decreases. Electronegativity: Going Left � Right on the periodic table, electronegativity increases. Going Top � Bottom on the periodic table, electronegativity decreases.

Practice Problems 1.) Account for each of the following observation in terms of atomic theory and/or quantum

theory a.) Atomic size decreases from Na to Cl in the periodic table. b.) The first ionization energy of K is less than that of Na.

2.) Using principles of atomic fluorine, oxygen, and xenon, as well as some of their

compounds. a.) Account for the fact that the first ionization energy of atomic fluorine is greater than

that of atomic oxygen. ( You must discuss both atoms in your response) b.) Predict whether the first ionization energy of atomic xenon is greater than, less than,

or equal to the first ionization energy of atomic fluorine. Justify your prediction. 3.)a.)What is meant by the term effective nuclear charge?

b)How does the effective nuclear charge experienced by the valence electrons of an atom vary going from left to right across the period of the periodic table.

4.)a.) What is meant by the terms acidic oxide and basic oxide?

b.) How can we predict whether an oxide will be acidic or basic based on its composition?

5.)Write a balanced equations for the reaction that occurs in each of the following cases: a) Potassium metal burns in an atmosphere of chlorine gas. b) Strontium oxide is added to water. c) A fresh surface of lithium metal is exposed to oxygen gas. d) Sodium metal is reacted with molten sulfur. 6) Which will experience the greater effective nuclear charge, the electrons in the n = 3 shell in Ar or the n = 3 shell in Kr? Which will be closer to the nucleus? Explain.

7a) Why are monoatomic cations smaller than their corresponding neutral atoms? b) Why are monoatomic anions larger than their corresponding neutral atoms? c) Why does the size of ions increase as on proceeds down a column in the periodic table? 8a) What is the general relationship between the size of an atom and its first ionization energy? b) Which element in the periodic table has the largest ionization energy? Which has the smallest? 9) Based on their positions in the periodic table, predict which atom of the following pars will have the larger first ionization energy: a) O, Ne b) Mg, Sr c) K, Cr d) Br, Sb

e) Ga, Ge 10) Identify the element whose ions have the following electron configurations: a) a 3+ ion with [Ar]3d³ b) a 2+ ion with [Kr]4d¹º5s². How many unpaired electrons does each ion contain?

11) The atoms and ions Na, Mg⁺, Al²⁺, and Si³⁺ are isoelectronic. a) For which of these will the effective nuclear charge acting on the outermost electron be smallest? b)For which will it be the greatest?

Solutions 1a) Across the periodic table from Na to Cl, the number of electrons in the s- and p- orbitals of the valence shell increases, as does the number of protons in the nucleus. The added electrons only partially shield the added protons, resulting in an increased effective nuclear charge. This results in a greater attraction for the electrons, drawing them closer to the nucleus, making the atom smaller. b) Both Na and K have an s1 valence-shell electron configuration (Na: [Ne] 3s1; K: [Ar] 4s1). The K atom valence has a higher n quantum number, placing it farther from the nucleus than the Na atom valence electron. The greater distance results in less attraction to the nucleus. Because its valence electron is less attracted to its nucleus, the K atom has the lower ionization energy. 2a)In both cases the electron removed is from the same energy level (2p), but fluorine has a greater effective nuclear charge due to one more proton in its nucleus (the electrons are held more tightly and thus take more energy to remove). b)The first ionization energy of Xe should be less than the fires ionization energy of F. To ionize the F atom, an electron is removed from a 2p orbital. To ionize the Xe atom, an electron must be removed from a 5p orbital. The 5p is a higher energy level and is farther from the nucleus than 2p, hence it takes less energy to remove an electron from Xe. 3a) Effective nuclear charge is a representation of the average electrical field experienced by a single electron. It is the average environment created by the nucleus and the other electrons in the molecule, expressed as a net positive charge at the nucleus. b) Going from left to right across a period, effective nuclear charge increases. 4a) An acidic oxide is a compound when placed in water creates an acidic solution. A basic oxide is a compound when placed in water creates a basic solution

b) Nonmetal oxides are acidic oxide. Metal oxides are basic oxides.

5a) 2K(s) + Cl₂(g) → 2KCl(s) b) SrO(s) + H₂O(l) →Sr(OH)₂(aq) c) 4Li(s) + O₂ →2Li₂O(s) d) 2Na(s) + S(l) → Na₂S(s) 6) The n = 3 electrons in Kr experience a greater effective nuclear charge and thus have a greater probability of being closer to the nucleus. 7a) Electrostatic repulsions are reduced by removing an electron from a neutral atom, effective nuclear charge increases, and the cation is smaller. b) The additional electrostatic repulsion produced by adding an electron to a neutral atom decreases the effective nuclear charge experienced by the valence electrons, and increases the size of the anion. c) Going down a column, valence electrons are further from the nucleus, and they experience greater shielding by core electrons. The greater radial extent of the valence electrons outweighs the increase in Z. 8a) The smaller the atom, the larger its first ionization energy (of the nonradioactive elements). b) He has the largest, and Cs the smallest first ionization energy. 9a) Ne b) Mg c) Cr d) Br e) Ge 10a) Cr³⁺, 3 unpaired electrons b) Sn²⁺, 0 unpaired electrons 11a) Na b) Si³⁺

Units 8-9

Chapter 8 Vocabulary Chemical Bond – Attractive forces that hold groups of atoms together and makes them function as a unit Ionic Bond – The electrostatic forces that form between metals and nonmetals Covalent Bond – Bonds formed by the sharing of electrons Metallic Bond – Bond between atoms of a metal Lewis Symbol – The chemical symbol for the element plus a dot for each valence electron Octet Rule – Atoms tend to gain, lose or share electrons until they have eight valance electrons Lattice Energy – The energy needed to separate an ionic crystal lattice into gaseous ions Nonpolar Covalent Bond – Bond in which the electrons are shared equally between two atoms Polar Covalent Bond – Bond in which one of the atoms exerts a greater attraction for the bonding electrons Electronegativity – The ability of an atom in a molecule to attract electrons to itself Dipole – When two electrical charges of equal magnitude but opposite sign are separated by a distance Dipole Movement – The measure of the magnitude of a dipole Formal Charge – The charge that the atom in a molecule would have if all atoms had the same Electronegativity Resonance Structure – A Lewis Structure where the placement of the electrons is different Bond Enthalpy – The enthalpy change for the breaking of a particular bond in a mole of gaseous substance Bond Length – Distance between the nuclei of the atoms involved in the bond

Chapter 9 Vocabulary

Bond Angles – The angles made by the lines joining the nuclei of the atoms in a molecule VESPER Model – Valence-Shell Electron-Pair Repulsion Model Bonding Pairs – In a Lewis structure, a pair of electrons that is shared by two atoms Nonbonding Pairs – Lone pairs of electrons in a Lewis structure not assigned completely to one atom Molecular Geometry – The arrangement of atoms Bond Dipoles – The dipole moment due to the two atoms of a covalent bond Hybridization – The mixing of different types of atomic orbitals to produce a set of equivalent hybrid orbitals Sigma Bond – Covalent bonds in which the electron density lies along the line connecting the atoms Pi Bond – Bonds that are formed from the side-to-side overlap Delocalized Electrons – Electrons that are spread over a number of atoms in a molecule rather then localized between a pair of atoms Molecular Orbital Diagram – A diagram that shows the energies of molecular orbitals relative to the atomic orbitals from which they are derived

Chapter 8 Formulas

Coulomb’s Law K•Q1•Q2 d

K – constant d – rcation + ranion

Q – charge on ion Formal charge

Cf = Ev – (Eu + ½ Eb) Cf – formal charge Ev – number of valence electrons Eu – number of unshared electrons Eb – number of bonded electrons

Practice Problems for Chapters 8 & 9

Multiple Choice: Use answers for 1-4:

A) Metallic Bonding B) Network Covalent Bonding C) Hydrogen Bonding D) Ionic Bonding E) London Dispersion Forces

1. Solids exhibiting this kind of bonding are excellent conductors of heat. 2. This kind of bonding is the reason that water is more dense than ice. 3. This kind of bonding exists between atoms with very different electronegativities. 4. The stability exhibited by diamonds is due to this kind of bonding. Use answers for 5-7:

A) BF3 B) CO2 C) H2O D) CF4 F) PH3

5. The central atom in this molecule forms sp2 hybrid orbital. 6. This molecule has a tetrahedral structure.

7. This molecule has a linear structure 8. Which of the following lists of species is in order of increasing boiling points?

a. H2, N2, NH3, b. N2, NH3, H2 c. NH3, H2, N2 d. NH3, N2, H2 e. H2, NH3, N2

Free Response: 1. Predict the chemical formula of the ionic compound formed between the following pairs of

elements: (a) Al and F; (b) K and S; (c) Y and O; (d) Mg and N

2. Which of the following bonds are polar:

(a) P—O; (b) S—F; (c) Br—Br; (d) O—Cl? Which is more electronegative atom in each polar bond?

3. Indicate the hybridization and bond angles associated with each of the following electron-domain geometries:

(a) linear; (b) tetrahedral; (c) trigonal planar; (d) octahedral; (e) trigonal bipyramidal. 4. Predict the molecular geometry of (a) H2Se; (b) PCl4

+; (c) NO2-; (d) BrF3; (e) I3

-.

Solutions Multiple Choice:

1. A- In metallic bonding, the freedom of electrons to move allows them to conduct heat and

electricity. 2. C- Molecules are farther apart n the solid (ice) than in the liquid making the solid (ice) less

dense than the liquid. 3. D- In an ionic bond, the difference in electronegativties is large enough, and then the more

electronegative atom will take an electron away from the other atom. These atoms are now bonded together by a strong ionic bond.

4. B- The carbon atoms in diamond are held together by a network of covalent bonds. The

tetrahedral structure of the carbon atoms is very stable and has no simple breaking points.

5. A- Boron forms three bonds with the fluorine atoms with no unbonded valence electrons, forming a sp2 hybrid orbital.

6. D- CF4 forms a tetrahedral structure, forming a sp3 hybrid orbital. 7. B- CO2 forms a linear structure with a sp hybridized orbital. 8. A- H2 only forms dispersion forces, N2 only forms dispersion forces as well but is larger than

H2 and has more electrons, and NH3 is polar and undergoes hydrogen bonding so it has the strongest intermolecular force and the highest boiling point.

Free Response: 1. (a) AlF3

(b) K2S (c) Y2O3

(d) Mg3N2

2. (a) Polar-O (b) Polar-F (c) Nonpolar (d) Nonpolar

3. (a) sp – 180º (b) sp3 – 109º (c) sp2 – 120º (d) sp3d2 – 90º and 180º (e) sp3d – 90º, 120º, and 180º

4. (a) bent (b) tetrahedral (c) bent (d) t-shaped (e) linear

Units 10

Thermochemistry Review

Equations:

• ∆Esystem = -∆Esurroundings

• ∆E = q + w

• ∆E = heat + work

• qreaction = -qcalorimeter

• qwater = mc∆T = -qreaction

• ∆Hoverall = ∆Hreaction 1 + ∆Hreaction 2

On AP Resouce Sheet: • ∆Hreaction = ∑∆Hf°product - ∑∆Hf°reactants

• q = mc∆T

Key:

• ∆E – change in energy (kJ)

• q – heat (kJ)

• w – work (kJ)

• m – mass (g)

• C – specific heat of element (J/g °C)

• ∆T – change in temperature (°C)

• ∆H – change in enthalpy

• ∑∆Hf° - sum of enthalpy change (kJ/mol)

Vocab • Thermodynamics: the study of energy and its transformations.

• Thermochemistry: the aspect of thermodynamics examining the relationship

between chemical reactions and energy changes involving heat

• Potential Energy: a form of energy by a body by virtue of its position relative to

other objects

• Kinetic Energy: the energy of motion

• Force: any kind of push or pull exerted on an object

• Work: energy used to cause an object move against a force

• Heat: energy transferred from a hotter object to a colder object

• Energy: the capacity to do work or transfer heat

• First Law of Thermodynamics: energy is conserved, any energy lost by the system is

gained by the surroundings

• Internal Energy: Sum of all kinetic and potential energy of all components of the

system

• Endothermic: system absorbs heat

• Exothermic: system exerts heat

• Enthalpy: A thermodynamic quantity equivalent to the total heat content of a

system. It is equal to the internal energy of the system plus the product of pressure

and volume

• Calorimetry: The measurement of heat flow

• Calorimeter: Apparatus used to measure heat flow

• Hess’s Law: If a reaction is carried out in a series of steps, the change in H for the

reaction will equal the sum of the enthalpy changes for each step

• Enthalpy/heat of formation: enthalpy change associated with the process of the

formation of a compound from its original elements.

• Standard Enthalpy: The enthalpy change when all reactants and products are in

their standard states

• Calorie: amount of energy needed to raise 1 gram of water 1 °C

Big Concept Section

Basics to Thermochemisty and the First Law of Thermodynamics • Units to Use: Calorie vs. Joule, 1 cal = 4.184 J

• q = magnitude and direction of heat flow

• Endothermic v. Exothermic:

o q+ = endothermic q- = exothermic

• First Law of Thermodynamics

o ∆Esystem = -∆Esurroundings

• Energy and Work

o ∆E = heat + work (∆E = q + w)

o +W = work done to a gas, or compression

o –W = work done by a gas, or expansion

Enthalpy and Calorimetry 1. Enthalpy: heat of reaction

• +∆H = endothermic

• -∆H = exothermic

2. Calorimetry: process for measuring heat flow

1. Measure change in temperature of surrounding

2. Determine hear gained/lost by water

a. Change in H2O temperature (∆H)

b. Mass of H2O

c. Specific heat of H2O

Thermochemical Equations

1. Sign of ∆H indicates whether the reaction, if under pressure, is endothermic

or exothermic

2. Coefficients represent the number of moles needed for specific amount of

heat

3. Phases of matter of all species must be identified

4. Value of ∆H applies when products and reactants are at the same

temperature

5. Magnitude of ∆H is directly proportional to amount of reactant or product

• Solid � liquid = heat of fusion = liquid � solid

• Liquid � gas = heat of vaporization = gas � liquid

Hess’s Law

• The value of ∆H for a reaction ids the same whether it occurs in one step or a

series of steps.

o ∆Hoverall = ∆Hreaction1 + ∆Hreaction2

• Standard Heat of Formation: the enthalpy change when 1 mole of a

compound is formed at constant temperature and pressure from its elements

in their most stable form. (∆Hf°)(kJ/mol)

o ∆Hreaction = ∑∆Hf°product - ∑∆Hf°reactants

� For elements, 0 kJ/mol

• Enthalpies of Formation for Ions: ∆Hf° for H+ = 0, therefore, all other

calculations for ions can be found.

- - Questions from the AP Test

Substance Combustion Reaction Enthalpy of Combustion, ∆Ho

comb, at 298 K (kJ mol-1) H2(g) H2(g) + ½O2(g) � H2O(l) -290 C(s) C(s) + ½O2(g) � CO2(g) -390

CH3OH(l) -730 1. In the empty box in the table above, write a balanced chemical equation for the complete combustion of one mole of CH3OH(l). Assume products are in their standard states at 298K. Coefficients do not need to be in whole numbers. 2. On the basis of your answer to question q and the information in the table, determine the enthalpy change for the reaction C(s) + H2(g) + H2O(l) → CH3OH(l) 3. Write the balanced chemical equation that shows the reaction that is used to determine the enthalpy of formation for one mole of CH3OH(l).

4. The standard enthalpy of formation for nitrogen dioxide is the enthalpy change of the reaction… (A) ½ N2O4(g) → NO2(g) (B) ½ N2(g) + O2(g) → NO2(g) (C) N2(g) + 2O2(g) → 2NO2(g) (D) NO(g) + ½ O2(g) → NO2(g) 5. For endothermic reactions at constant pressure… (A) ∆H < 0. (B) ∆H > 0. (C) ∆G < 0. (D) ∆S < 0. (E) ∆S > 0. 6. Which of the following has a standard enthalpy of formation which is not zero? (A) Na(s) (B) Hg(l) (C) H2O(l) (D) N2(g) (E) C(s)

- Questions not from the AP Test 1. Given the data N2(g) + O2(g) → 2NO2(g) ∆H = +180.7 kJ 2NO(g) + O2(g) → 2NO2(g) ∆H = -113.1 kJ 2N2O(g) → 2N2(g) + O2(g) ∆H = -163.2 kJ use Hess’s law to calculate ∆H for the reaction N2O(g) + NO2(g) → 3NO(g) 2. Consider the following reaction: 2Mg(s) + O2(g) → 2MgO(s) ∆H = -1204 kJ

(a) Is this reaction endothermic or endothermic? (b) Calculate the amount of heat transferred when 2.4g of Mg(s) reacts at constant pressure.

(c) How many grams of MgO are produced during an enthalpy change of -96.0 kJ?

(d) How many kilojoules of heat are absorbed when 7.50g of MgO(s) are decomposed into

Mg(s) and O2(g) at constant pressure?

Solutions - - Questions from the AP Test

1.

2. Adding the following three equations, C(s) + O2(g) → CO2(g) -390 kJmol-1 H2(g) + ½O2(g) → H2O(l) -290 kJmol-1

CO2(g) + 2H2O(l) → CH3OH(l) 3/2O2(g) +730 kJmol-1

Yields this equation: C(s) + H2(g) + H2O → CH3OH(l) +50 kJmol-1 3. C(s) + 2H2(g) + ½O2(g) → CH3OH(l) 4. B 5. B 6. C

- - Questions not from the AP Test 1. ∆H= 155.7 kJ N2 + O2 → 2NO ∆H=180.7 kJ NO2 → NO + ½ O2 ∆H= 56.5 kJ N2O → N2 + ½ O2 ∆H=-81.6 kJ 2. a. exothermic reaction b. 2.4g Mg 1 mol Mg -1204 kJ = 59 kJ 24.31g Mg 2 mol Mg c. -96.0 kJ 2 mol MgO 40.31g MgO = 6.43g MgO -1204 kJ 1 mol MgO d. 7.50g MgO 1 mol MgO 1204 kJ = 112 kJ 40.31g MgO 2 mol MgO

Substance Combustion Reaction Enthalpy of Combustion, ∆Hocomb, at

298 K (kJ mol-1) H2(g) H2(g) + ½O2(g) → H2O(l) -290 C(s) C(s) + ½O2(g) → CO2(g) -390 CH3OH(l) CH3OH(l) + 3/2O2(g) → CO2(g) + 2H2O(l) -730

Units 11

Unit 11 Vocab: Vodcast 1

• Intramolecular Force- Within a molecule • Intermolecular Force- Between two molecules (ionic, covalent, metallic, covalent

network

Vodcast 2 • Dispersion Forces- Temporary dipole that’s created just because a compound has

electrons (every molecule has this) • Polarizability- The more electrons an atom has, the “stronger” the dispersion force,

which makes it have a higher melting and boiling point • Dipole-Dipole- Permanent region on a molecule that is partially positively charged and

partially negatively charged • Hydrogen Bonds-The strongest bonds, requires a super electronegative atom (N, O, F)

Vodcast 3 • Viscosity- Resistance to flow (greater viscosity = stronger IMF) • Volatility- Ease of evaporation (less IMF = easier to evaporate)

Vodcast 4 • Critical Point- Temperature where you can’t turn a gas into a liquid no matter what

pressure • Super Critical Fluid- The state between the gas and the liquid, past the critical point • Triple Point- The temperature and pressure where you could have a solid, liquid or gas • Deposition- Phase change from a gas to a solid • Sublimation- Phase change from a solid to a gas • Normal Boiling/Melting Point- at 1 atm

Vodcast 5 • Vapor Pressure- Pressure exerted by the vapor in the space above the liquid (depends on

IMF) • Equilibrium Vapor Pressure- Reached when the rate of evaporation equals the rate of

condensation

Formulas: No formulas in this unit

Concepts: Vodcast 1

• Ionic Bonding -Consist of cations and anions -Has a crystal lattice -High melting and boiling points -Good conductor -Soluble in water

• Metallic Bonding -It is described with the “Electron Sea” model -High conductivity, ductility and malleability -Insoluble in water

• Covalent Bonding -Non-conductors -Most are insoluble in water -Low melting/boiling point -Gases have weak IMF

• Covalent Network Solids -Held together by covalent bonds -Non-conductors -Very high melting/boiling point -Have crystal like structures

Vodcast 2 • Dispersion Forces (London Dispersion Forces)

1. Everything has a dispersion force 2. Bonds are weak and temporary

Factors Determining the Strength of Dispersion Forces 1. Polarizability- The more electrons, the stronger the force 2. Shape- Plays a little role in determining strength

• Dipole-Dipole Forces

1. Not everything has dipole-dipole 2. Solid- Fixed in place 3. Liquid- Molecules are less organized, not as many attractions 4. Gas- none

• Hydrogen Bonds 1. Strongest bonds (Super strong dipole-dipole) 2. Attached to a super electronegative atom (Nitrogen, Oxygen, Fluorine) 3. Causes high boiling/melting points

Vodcast 3

• Properties based on the IMF 1. Boiling/Melting Points

-Directly related to the strength of IMF (Stronger forces = higher points) -More electrons = higher melting/boiling points

2. Surface Tension -Liquids that from a skin (droplets or meniscus) -Only happens with strongest IMF

3. Viscosity

-Resistance to flow (thickness of a liquid) -Greater viscosity = slower it flows -Greater viscosity = stronger IMF

4. Volatility

-Ease of evaporation -Higher volatility = lower IMF

Vodcast 4

• Melting breaks some IMF bonds • Vaporizing breaks all IMF bonds • When altitude increase, boiling point decreases because there is less amounts

of molecules at higher altitudes • Pressure affects boiling point and melting point

-As pressure drops, so does the boiling point -As pressure increases, melting point decreases

• Phase diagram -A negative slope (between solid and liquid) means the liquid is denser -A positive slope (between solid and liquid) means the solid is denser

Vodcast 5

• Factors that affect equilibrium vapor pressure (EVP) 1. Surface area

-no affect 2. IMF -Stronger IMF = lower EVP -Weaker IMF = higher EVP

3. Temperature -As temperature increase, EVP increases

Practice AP Questions: 1.) Use principles of atomic structure, bonding and/or intermolecular forces to respond to each of the following. Your responses must include specific information about all substances referred to in each question. (a) At a pressure of 1 atm, the boiling point of NH3 (l) is 144 K. (i) Identify the intermolecular force(s) in each substance.

(ii) Account for the difference in the boiling points of the substances (b) The melting point of KCl(s) is 776 ˚C, whereas the melting point of NaCl(s) is 801˚C.

(i) Identify the type of bonding in each substance. (ii) Account for the difference in the melting points of the substances.

2.) ●

T(˚C)

50 100 150 200

1.5

1.0

0.5

P (atm)

If the pressure of the substance shown in the diagram is decreased from 1.0 atm to .5 atm At a constant temperature of 100˚C, which phase change will occur? (A). Freezing (B). Vaporization (C). Condensation (D). Sublimation (E). Deposition. 3). Under what conditions in the phase diagram above can all three phases of the substances shown in the diagram exists simultaneously in equilibrium. (A). Pressure = 1.0 atm, Temperature = 150˚C (B). Pressure = 1.0 atm, Temperature = 100˚C (C). Pressure = 1.0 atm, Temperature = 50˚C (D). Pressure = 0.5 atm, Temperature = 100˚C (E). Pressure = 0.5 atm, Temperature = 50˚C

4.) Account for each of the following observations about pairs of substances. In your answers, use appropriate principles of chemical bonding and/or intermolecular forces. In each part, your answer must include references to both substances. (a.) Even though NH3 and CH4 have similar molecular masses, NH3 has a much higher normal boiling point (-33˚C) than CH4 (-164˚C). (b.) At 25˚C and 1.0 atm, ethane (C2H6) is a gas and hexane (C6H14) is a liquid. (c.) Si melts at a much higher temperature (1,410˚C) than Cl2 (-101˚C). (d.) MgO melts at a much higher temperature (2,852˚C) than NaF (993˚C)

Solutions 1.) a.) (i) The intermolecular for NH3 are hydrogen bonding, dipole-dipole and dispersion forces, while NF3 only has dipole-dipole and dispersion forces. (ii) Because NH3 has hydrogen bonding, the bonds are harder to break which means it is harder to boil and makes the boiling point increase.

b.) (i) KCl only has dispersion forces for its bonding type and NaCl has dipole-dipole and

dispersion forces for its bonding types. (ii) NaCl has dipole-dipole bonds, making it harder to break apart and raising the melting point. 2). The answer is B. the liquid is more dense then the gas therefore it will vaporize at .5

atmospheres and at 100˚C 3). The answer is E. At this point, the triple point. All the phase change lines converge and all 3 phases are at equilibrium 4). a.) NH3 has a higher melting point because it has hydrogen bonding and it has a shape of trigonal pyramidal. CH4 doesn’t have hydrogen bonding and it has a shape of tetrahedral which makes NH3 harder to break b.) They are both nonpolar but, Hexane has more electrons which makes it have an increased dispersion force. c.) Si has a strong network and covalent bonds, while Cl2 has a weak dispersion force with weak bonds. d.) MgO melts at a higher temperature because it has a larger charge which creates a stronger attraction. (Mg2+O2 and Na+F-)