AP Chemistry 2018-2019 Summer Assignment

Dear AP Chemistry Students,

The purpose of this summer packet is to provide you with a complete review of first year Chemistry and to begin new material that will be covered in the course. Note, to complete this assignment, you DO NOT have to have your textbook. Instead, you will watch a number of different videos (QR codes accompany each section) while completing notes. You will then complete practice problems on your own to ensure that you understand the concepts.

This packet is lengthy as last year’s AP Chemistry students firmly believed that a complete review of first year Chemistry and introduction to AP material over the summer would have made the transition much easier. Do not wait to complete this assignment until the end of the summer (as you will note there are days during the summer for you to come in to review), and please make sure to email me if you have any questions or concerns (Rebecca.Poliner@providencehigh.org). We will review the material during the first class period and you will have your first test the second class period.

These are the STRONGLY suggested review meetings (will last as long as you need). Additional practice problems will be

assigned during these meetings. If you cannot make a meeting, you need to email me pictures of each completed

section before the meeting date, print out the practice problems, and have them ready to turn in on the first day of

school.

-June 25 @ 9AMà First Year Chemistry Review.

-August 6 @ 9AM à New AP Chemistry Material Review.

If you fail the summer review test, you will be asked to transfer out of the course. This is not meant as punishment, rather to

ensure that every AP Chemistry student succeeds in this course.

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Before you begin your work for AP Chemistry, please take a moment to read this letter of advice from a previous Chemistry student:

Dear Incoming AP Chemistry Student,

First of all, I would like to say welcome to AP Chemistry. This course may be one of the most challenging classes here at Providence. It will test your ability to manage your time effectively and step out of your comfort zone. There will be times when you get a problem easily, but there may also be times when you feel like you understand absolutely nothing. When you are in this situation, please don’t hesitate to ask Ms. Poliner questions. Even if you think it will make you look “stupid” or it’s something you should already know, ASK HER! She is a valuable resource and is eager to work to help you learn the material, pass the AP exam, and pass the class. This course is not like the many others where you can get a good grade with minimal studying. In this class, you need to get extra help on the concepts you are not strong in to be successful.

I personally struggled a lot with asking for help. I am generally a very shy person and for all my life, school has come very easy for me. I never felt the need to ask for help because usually I was able to teach myself or get by without understanding a lesson. However, when I took this class, I was unable to fully comprehend the challenging concepts. I was busy with sports and everyone else seemed to get it, so I never bothered to ask a question. In addition, I was afraid that I would be disturbing Ms. Poliner or she would realize how behind I was in the class. As a result, I failed every test second semester, regularly getting 30%. This was a big drop, since I’ve always received straight A’s and I even got an A+ in Chemistry Honors. Ms. Poliner tried to reach out and get me to come to tutoring, but there’s only so much a teacher can do. To be honest, I was scared to come in, because I felt vulnerable. At the time, I viewed seeking help as a weakness or a failure. I thought that if I came to tutoring, I wasn’t good or smart enough to understand the lesson. After having a D in the class and being at risk of getting my college admissions revoked and not graduating, I finally realized I needed come to tutoring and ask questions. When I came in, Ms. Poliner was really nice. She sat with me, reviewed the lesson, and went over problems with me. There was no judgment at all. Even when I asked fairly basic questions, she never once ridiculed me. After a few weeks of seeking help, my grade and my confidence improved. I now understand the importance of talking to your teachers. It still is uncomfortable to ask for help sometimes, but I know it is worth it.

It took me pretty much all my school life and the risk of not graduating to realize the value of using your teachers as a resource. Don’t let it take you that long. Get help from the start. Even if you are too afraid to ask questions in class, go after school and ask, or send your question through email. Ultimately, it’s your education and you need to take initiative and ask yourself: “What can I do now to get me where I want to be in the future?” and “How can I get the most out my education with the resources available to me?” Please remember that asking for help is not a failure, it’s a strength. It takes a strong person to acknowledge that they don’t understand something and find a way to fix it. The only thing that is a failure is not being able to admit your faults. As long as you ask questions, the class will be an enjoyable and rewarding experience. With that said, I hope you guys have a great time in AP Chemistry.

From,

A Former AP Chemistry Student

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First Year Chemistry Review Density:

Ex. A sample of tetrachloroethylene, a liquid used in dry cleaning that is being phased out because of its potential to cause cancer, has a mass of 40.55 g and a volume of 25.0 mL at 25°C. What is the density at this temperature? Will tetrachlorethylene float in water? Explain.

Practice:

1. If an unknown solid weighs 84.0 grams and occupies 30.0 cm3 of space, what is its density? (Ans: 2.80 g/cm3)

2. An irregularly shaped stone was lowered into a graduated cylinder holding a volume of water equal to 2 mL. The height of the water rose to 7 mL. If the mass of the stone was 25 g, what was its density? (Ans: 5 g/mL)

Ex. What is the mass, in grams, of a liquid having a density of 1.50 g/mL and a volume of 3.5 liters?

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Practice:

1. What is the volume (in L) of a 200. gram sample of gold if its density is known to be 20.5 g/cm3? (Ans: 0.00976 L)

2. A 10.0 cm3 sample of copper has a mass of 89.6 g. What is the density of copper? (Ans: 8.96 g/cm3)

3. Silver has a density of 10.5 grams/cm3 and gold has a density of 19.3 g/cm3. Which would have the greater mass, 5cm3 of silver or 5cm3 of gold?

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Significant Figures:

Section 1.5- Uncertainty in Measurement

• What are significant figures?

2.2g versus 2.2405g

• Determining the number of Significant Figures:

Practice: Determine the number of sig figs in each of the following:

a. .034 ______________ b. 340______________ c. .340 ______________

1. I want to complete a lab in which I need to weigh .540 g of a substance. I use a balance that can measure to the nearest 0.01 g. Will I be able to determine the mass of this substance to three significant figures? If not, how many will I be able to measure it to?

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Mole Conversions:

Section 3.4- Avogadro’s Number and the Mole

Write the conversion factors on each line:

Mass (g) Mole Particles (atoms, molecules, formula units)

Volume (L)

Ex. 0.5 moles of C6H12O6 to grams

1. 35 g of CuSO4 • 5 H2O to moles Hint: The 5 molecules of water should be considered part of this molecule (Ans: 0.14 mol)

Ex. 45.6 g of Na2SO4 to formula units.

1. 4.3 x 1024 molecules of carbon dioxide to liter at STP (Ans: 160 L)

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Ex. What mass of iron can be recovered from 25.0 g of Fe2O3?

1. What mass of silver can be produced from 125 g of Ag2S? (Ans: 109 g)

Precipitation Reactions:

Section 4.2- Precipitation Reactions

• Precipitate:

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Double Replacement Reactions

Ex. Copper (II) Chloride + Sodium Sulfideà

Molecular Equation:

Complete Ionic Equation:

Net Ionic Equation:

1. Barium chloride + Sodium sulfate à

Molecular Equation:

Complete Ionic Equation:

Net Ionic Equation:

2. Lead (II) Nitrate + Potassium Iodide à

Molecular Equation:

Complete Ionic Equation:

Net Ionic Equation:

Stoichiometry and Limiting Reactants:

Section 3.6- Quantitative Information from Balanced Equations

Fill in the conversion factors on the lines.

mass A mass B

Mole A Mole B

volume A volume B

Particles A Particles B

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Ex: Given the formula: N2 + 3 H2 à 2 NH3, how many grams of NH3 will be produced if 67.2 L of N2 reacts with excess H2 at standard temperature and pressure (STP)?

1. 2KClO3 à 2KCl + 3O2

How many grams of potassium chloride are produced if 25.0 g of potassium chlorate decompose? (Ans: 15.3 g)

Section 3.7- Limiting Reactants

• Limiting Reactant:

• Excess Reactant:

Ex.

Al + Cl2 à AlCl3

A mixture of 1.5 mol Al and 3 mol Cl2 are allowed to react. How many grams of product are formed? Which one is the limiting reactant?

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1. A 2.00 g strip of zinc is placed in a solution containing 2.50 g silver nitrate causing this reaction to occur:

Zn + AgNO3 à Ag + Zn(NO3)2

How many grams of silver will form? (Ans: 1.58 g)

• Theoretical Yield-

• Actual Yield-

• Percent Yield-

Ex.

Fe2O3 + 3 CO à 2 Fe + 3 CO2

A. You determined the limiting reactant of this experiment to be Fe2O3. If you start with 150 g of Fe2O3, what is the theoretical yield of Fe?

B. In the lab, you produce 87.9 grams of iron. Calculate the percent yield.

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1. A student performs the following experiment: CaCO3(s) à CaO(s) + CO2(g). The student uses 270 g of CaCO3 and obtains 120 g of CaO. Calculate the percent yield. (Ans: 79%)

Concentration Calculations:

Section 13.4- Expressing Solution Concentration

• Qualitatively-

• Quantitatively:

Ex. You are given 250 ml of a 0.347 M glucose solution. Now add to this solution 2.39 g more of solid glucose all of which will dissolve. (Assume with no change in volume). What is the new molarity?

1. What is the molar concentration of 0.125 g Na2CO3 dissolved in 50.0 mL solution? (Ans: 0.0236 M)

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2. Calculate the molarity of a substance that has 15 g of MgCl2 dissolved to produce 450 mL of solution.

Concentration Stoichiometry:

Ex. How many grams for silver chromate will precipitate when 100.0 mL of 0.400 M silver nitrate solution is added to potassium chromate? 1. What volume (in mL) of a 0.455 M sulfuric acid (H2SO4) solution is required to neutralize 25.00 mL of a 0.800 M sodium hydroxide (NaOH) solution? (Ans: 22.0 mL)

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2. What volume (in mL) of 0.455 M silver nitrate (AgNO3) solution is required to precipitate all of the bromide ions in 40.00 mL of a 0.120 M strontium bromide (SrBr2) solution? (Ans: 21.1 mL) 3. What volume (in mL) of 2.50 M hydrobromic acid (HBr) solution will be consumed in reacting with 30.00 g of aluminum metal? (Ans: 1334 mL)

Gases:

Section 10.1- Characteristics of Gases

• Characteristics of Gases:

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Section 10.2- Pressure

• Pressure-

• Atmospheric Pressure and the Barometer:

Ex. 1277 mmHg to atm

1. 11.5 kPa to mmHg (Ans: 86.3 mmHg)

Section 10.4- The Ideal-Gas Equation

• Ideal Gas Law:

Ex. A flashbulb of volume 2.6 cm3 contains O2 gas at a pressure of 2.3 atm and a temperature of 26 °C. How many moles of O2 does the flashbulb contain?

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b. How many grams of oxygen is this?

c. What is the density of oxygen in the container?

1. Many gases are shipped in high-pressure containers. Consider a steel tank whose volume is 42.0 L and which contains O2 gas at a pressure of 18,000 kPa at 23°C. What mass of O2 does it contain? (Ans: 9800 g)

2. How many moles of air are there in the lungs of an average adult with a total lung capacity of 3.8 L? Assume that the person is at 1.0 atm pressure and has a normal body temperature of 37 °C. (Ans: 0.15 mol)

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Oxidation Numbers:

Assigning Oxidation Numbers

1. A way of keeping track of electrons gained and electrons lost.

2. Complete the following rules for assigning oxidation numbers:

a. Elements always have an oxidation number of _________. b. The sum of the oxidation numbers of compounds is ________. c. The sum of the oxidation numbers for polyatomic ions is

______________________________________________. d. Oxygen is normally _________ except when it is a peroxide (O22-) in which case the oxidation number is

___________. e. Hydrogen is ________ when bound to nonmetals and ______________ when bonded to metals. f. Fluorine is always _______. Other halogens typically have an oxidation number of ______ but when

combined with oxygen in molecules, they have ________________________ oxidation states. g. Group 1 Metals are always _______. Group 2 Metals are always_______. h. Positive Charge generally comes first.

Ex. Find the oxidation states for the following compound:

a. KMnO4

b. SO42-

Practice: Assign oxidation numbers to the following substances.

1. KNO3 6. Al(NO3)3

2. NH3 7. S8

3. OH- 8. H2O2

4. H2SO3 9. CO3-2

5. KClO3 10. OF2

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Subatomic Particles:

Section 2.3- The Modern View of Atomic Structure

• Protons (positively charged) and neutrons (neutrally charged) in the center, electrons (negatively charged) around the outside.

• Because the size of an atom is so small, rather than using grams, we use amu (atomic mass unit). • Isotopes- same number of protons, different number of neutrons (have same chemical properties but different

physical properties) • Mass Number= # protons + # neutrons • Atomic Number = # protons • !"

#C à Mass Number-12; Atomic Number- 6.

Practice:

Ion Number of Protons

Number of Electrons

Number of Neutrons

Mass Number

N3-

14

19

18 39

Br-

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3

2 4

Average Atomic Mass Calculations:

Section 2.4- Atomic Weights

Ex. Naturally occurring chlorine is 75.78% 35Cl (atomic mass 34.969 amu) and 24.22% 37Cl (atomic mass 36.966 amu). Calculate the atomic weight of chlorine.

1. Three isotopes of silicon occur in nature: 28Si (92.23%), which has an atomic mass of 27.97693 amu; 29Si (4.68%), which has an atomic mass of 28.97649 amu; and 30Si (3.09%), which has an atomic mass of 29.97377 amu. Calculate the atomic weight of silicon. (Ans: 28.085 amu)

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Ex. Magnesium has two naturally occurring isotopes, Magnesium- 24 and Magnesium-26. If the average atomic mass is 24.31 amu, what is the percent abundance of each isotope?

1. Silver has two naturally occurring isotopes with the following isotopic masses:

107 Ag 109Ag

106.9051 amu 108.9047 amu

The average atomic mass of silver is 107.8682 amu. What is the percent abundance of the lighter of the two isotopes? (Ans: 51.84% and 48.16%)

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Periodic Table and Empirical/Molecular Formulas:

Section 2.5- The Periodic Table

Section 2.6- Molecules and Molecular Compounds

• Are created by nonmetal atoms. Covalent bonds are responsible for the formation of these compounds.

• Molecular Formulas:

• Empirical Formulas:

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Section 3.5- Empirical Formulas from Analyses

• Empirical Formula:

Exa. A compound with elements C, H, and O is found to have 9.1% hydrogen and 54.5% carbon. What is the empirical formula?

Exb. A compound with an empirical formula of C2H4O has a molecular mass of 176 g/mol. What is the molecular formula?

*Tips*

1. Carry decimal places out at least 4 digits.

2. You can only round to a whole # if it is within 1 tenth of that number (Ex. 1.03à 1)

3. When you are not within 1 tenth you will need to multiply all of your answers so they are whole numbers. (Ex. 1.5 x 2 = 3; 1.2 x 5 = 6)

1. A compound is found to be 64.9 % carbon, 13.5% hydrogen, and 21.6% oxygen. Its molecular mass is 148 g/mol. What is its molecular formula? (Ans: C8H20O2)

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2. Write the empirical formula for each of the following, and determine the percent composition for each:

a. anthracene, C14H10 94..34 5.66

b. mercury (II) sulfate 67.61

Wavelength, Frequency, and Energy:

Section 6.1: The Wave Nature of Light

• All types of electromagnetic radiation move through a vacuum at 3.00 x 108m/s (speed of light).

Ex. A certain microwave has a wavelength of 0.032 meters. Calculate the frequency of this microwave.

1. A radio station broadcasts at a frequency of 590 KHz. What is the wavelength of the radio waves in nanometers?

(Ans: 5.1 x 1011nm)

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Section 6.2: Quantized Energy and Photons

• Planck’s constant: E=hν

Ex. A photon has a frequency (n) of 2.68 x 106 Hz. Calculate its energy.

1. The wavelength of green light from a traffic signal is centered at 5.20 x 10-5 cm. Calculate its energy in J. (Ans: 3.82 x 10-19 J)

2. An FM radio station has a frequency of 88.9 MHz. What is the wavelength of this radiation in meters? (Ans: 3.37 m)

3. Violet light has a wavelength of about 410 nm. What is its frequency? Calculate the energy of one photon of violet light. What is the energy of 1.0 mol of violet photons? (Hint: When calculating energy from frequency, the units are J/photon. There are 6.022 x 1023 photons per 1 mole). (Ans: 7.3 x 1014Hz; 4.85 x 10-19J/photon; 2.92 x 105 J/mol)

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4. The energy of a mole of photons of red light from a laser is 175 kJ/mol. Calculate the energy of one photon of red light in J/photons. What is the wavelength of red light in meters? In nm? (Ans: 2.91 x 10-19 J/photon; 6.84 x 10-7 m)

Electron Configurations and Orbital Diagrams:

Practice: For the following atoms, write the electron configuration AND the orbital diagram.

Ex. O:

1. N:

2. Al:

3. Fe:

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• Condensed Electron Configurations: Allows us to focus on valence electrons and represent the core electrons with a noble gas in brackets.

1. Find the noble gas (group 8A) from the row above your element. Put that symbol in brackets

2. Write the configuration from there.

Practice: Write the short hand electron configuration for the following atoms.

Ex. Ni- __________________________________ 3. P- __________________________________

1. Y- __________________________________ 4. V- __________________________________

2. K- __________________________________ 5. At- _________________________________

6. Write the longhand electron configuration for the following. For the starred problems, include the orbital diagram:

a. Ge:

b*. Zn:

*c. O2-:

d. Fe3+:

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Lewis Dot Diagrams:

Drawing Lewis Dot Diagrams:

Steps: Cà

Sà

Là

Cà

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Ex. H2O

Lewis Dot Diagram Without Model Kits

Lewis Dot Diagram With Model Kits

Ex. PBr3

1. CCl4

2. SCl2

3. F2

VSEPR Model: Model used in chemistry to predict the shape that molecules assume in order to minimize electron repulsions.

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Practice:

1. Draw the Lewis structure for the following. Make sure to use the VSEPR model to draw your molecules.

a. PCl3 b. CH3Br

c. F2O d. NH2Cl

Trickier Lewis Dot Diagrams:

1. O2

2. CO2

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3. N2

4. SO3

5. C2H4

6. CO32-

7. NO2+

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New AP Chemistry Material: Molecular Geometry:

Note: These molecular geometries must be memorized for the first day of school!

Section 9.2- The VSEPR Model

• Electron Domain Geometry:

o Bonding Pair:

o Nonbonding Pair:

• Molecular Geometry:

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• The effect of Nonbonding Electrons and Multiple Bonds on Bond Angle:

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• What happens when you have too few valence electrons?

• What happens when you have too many valence electrons?

Practice: For each of the molecules, identify the molecular geometry.

1. 2. 3.

4. 5. 6.

7. 8. 9.

10. 11. 12.

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13.

14. NF3 and PF5 are stable molecules. Write the electron-dot formulas for these molecules. On the basis of structural and bonding considerations, account for the fact that NF3 and PF5 are stable molecules but NF5 does not exist.

PCl5

ED = ____

ED Geometry ________________

Molecular Geometry _____________

XeF4

ED = ____

ED Geometry ________________

Molecular Geometry _____________

SO3

ED = ____

ED Geometry ________________

Molecular Geometry _____________

XeF2

ED = ____

ED Geometry ________________

Molecular Geometry _____________

AlH3

ED = ____

ED Geometry ________________

Molecular Geometry _____________

GeF2

ED = ____

ED Geometry ________________

Molecular Geometry _____________

SiH4

ED = ____

ED Geometry ________________

Molecular Geometry _____________

SF6

ED = ____

ED Geometry ________________

Molecular Geometry _____________

NO2-

ED = ____

ED Geometry ________________

Molecular Geometry _____________

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Bond Order and Length/Strength of Bonds:

Practice:

1. Create the Lewis Structure for each of the following and use them to justify the claims below.

(a) The bond length between the two carbon atoms is shorter in C2H4 than in C2H6. (b) The H-N-H bond angle is 107.5º, in NH3. (c) The bond lengths in SO3 are all identical and are shorter than a sulfur-oxygen single bond. (d) The I3- ion is linear.

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Polarity of Molecules:

Electronegativity:

Polarity of Molecules:

Nonpolar: Polar:

Practice: Go back to pg. 30-31 and for each of the molecules, note whether they are polar or nonpolar and EXPLAIN WHY (Dipoles cancel, or Dipoles do not cancel).

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1. Answer the following questions using principles of chemical bonding and molecular structure.

(a) Consider the carbon dioxide molecule, CO2, and the carbonate ion, CO32–. (i) Draw the complete Lewis electron-dot structure for each species. (ii) Account for the fact at the carbon-oxygen bond length in CO32– is greater than the carbon-oxygen bond length in

CO2.

(b) Consider the molecules CF4 and SF4.

(i) Draw the complete Lewis electron-dot structure for each molecule. (ii) In terms of molecular geometry, account for the fact that the CF4 molecule is nonpolar, whereas the SF4 molecule is

polar.

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Intermolecular Attractions:

Section 11.2- Intermolecular Forces

1. London Dispersion Forces:

• Strength of London Dispersion Forces: o Polarizability:

o Molecular Shape:

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2. Dipole-Dipole Forces:

3. Hydrogen Bonding:

Additional Notes:

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Practice:

1. This graph shows the BP’s of analogous compounds using elements from periods 2, 3, 4, and 5.

a. Explain why the BP of Xe > Kr > Ar > Ne:

b. Why is the BP of H2O > the others in its group?

2. Why is ΔHvap (energy required to vaporize a substance) much greater than ΔHfus (energy to melt a substance)? What does this reveal concerning changes in intermolecular forces in going from solid to liquid to vapor?

3. For which molecule in each of the following pairs would you expect the stronger intermolecular forces? Make sure to draw the Lewis Structure for each of these molecules before answering the question (a rough drawing, without counting valence electrons is OK!) and identify the type of intermolecular force EACH experiences.

a. CH3CH2CH2NH2 or H2NCH2CH2NH2

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b. CH3CH3 or H2CO

c. CH3OH or H2CO

d. HF or HBr

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Mass Spectrometry:

There is no video for this concept. Mass spectrometry simply provides a graph that separates isotopes based on their mass numbers. Therefore, the work required for calculating average atomic mass (see pg. 16 and 17) is the same…you just need to retrieve the data from the graph!

Ex. Calculate the atomic weight of copper.

______1. Which statement is true regarding the relative abundances of the 6lithium or 7lithium isotopes?

A) The relative proportions change as neutrons move between the nuclei

B) 7Lithium is much more abundant

C) The relative ratio depends on the temperature of the element

D) 6Lithium is much more abundant

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______2. All isotopes of an element possess the same:

A) number of electrons, atomic number and mass, but have nothing else in common

B) atomic number and mass, but have nothing else in common

C) chemical properties and mass, but have nothing else in common

D) number of electrons, atomic number and chemical properties

______3. 63Cu is 69% of the naturally occurring isotope of Cu. If only one other isotope is present for natural copper, what is it?

A) 59Cu

B) 65Cu

C) 61Cu

D) 62Cu

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Determining Empirical Formula Using Combustion Reactions:

Ex. Answer the following questions about a pure compound that contains only carbon, hydrogen, and oxygen.

A 0.7579 g sample of the compound burns in O2(g) to produce 1.9061 g of CO2(g) and 0.3370 g of H2O.

a. Calculate the individual masses of C, H, and O in the 0.7579 g sample. b. Determine the empirical formula

1. 12.915 g of a biochemical substance containing only carbon, hydrogen, and oxygen was burned in an atmosphere of excess oxygen. Subsequent analysis of the gaseous result yielded 18.942 g carbon dioxide and 7.749 g of water. Determine the empirical formula of the substance. (Ans: CH2O)

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2. In the course of the combustion analysis of an unknown compound containing only carbon, hydrogen, and nitrogen, 12.923 g of carbon dioxide and 6.608 g of water were measured. Treatment of the nitrogen with H2 gas resulted in 2.501 g NH3. What the compound’s empirical formula? (Ans: C2H5N)

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Photoelectron Spectroscopy:

PES is a method used to identify the placement of electrons for a SINGLE atom. Data from PES experiments are

displayed as follows: (Note that these are ionization energies for different electrons in the SAME atom)

• Each peak is relative to the others. This indicates the relative number of electrons. If the peak is twice as big, there are twice as many electrons.

1. Note that the y axis shows the relative number of electrons in each peak. The peak heights therefore show how

many electrons would have a given ionization energy relative to another peak. Recall that the closer an electron is to

the nucleus, the more energy would be needed to remove that electron from an atom. In the diagram above, the

peak at 349 represents electrons CLOSEST to the nucleus, therefore, in the 1st shell. We know from electron

configurations that 2 electrons can fit into the 1s sublevel, so we can surmise that the peak at 349 represents two 1s

electrons (1s2).

2. The next peak at 37.1 MJ/mol would represent electrons in the ____________ subshell. Note the height of this peak is the same as the height of the first peak, so the peak should represent_____________ electrons. The electron configuration should be ____________.

3. The next peak at 29.1 MJ/mol should be the ____________ subshell. Note how large this peak is compared to the prior peak. The peak at 29.1 MJ/mol represents ____________ 2p electrons with an electron configuration of ____________.

4. The peak at 3.93 MJ/mol represents ____________ 3s electrons.

5. The peak at 2.38 MJ/mol represents ____________ 3p electrons (note the height of this peak is the same as the height for the six 2p electrons at 29.1 MJ/mol).

6. The last peak at 0.42 MJ/mol would be for electrons in the ____________ subshell. Based on the height of this peak, the number of electrons is ____________ because the height of the peak is ½ the height of the peaks at 3.93 MJ/mol, 37.1

MJ/mol and 349 MJ/mol.

What is the electron configuration of this element? ___________________________________________________

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1.

2.

Ex. It takes 208.4 kJ of energy to remove one mole of electrons from the atoms on the surface of rubidium metal. If rubidium metal is irradiated with 254-nm light, what is the maximum kinetic energy the released electrons can have in J/photon?

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1. 450 kJ/mol is required to remove an electron from cesium.

a. A ray of red light has a wavelength of 700. m. Will exposure to red light cause electrons to be emitted from cesium? (Ans: 1.71 x 102 kJ/mol; no)

b. What is the kinetic energy of the emitted electrons when cesium is exposed to UV rays of frequency 1.9×1015 Hz? (Ans: 308 kJ/mol).