AP Chemistry Unit 2 Homework PacketFill in the Chart below:

Element Atomic #

# of Protons

# of Electrons

# of Neutrons

Nuclear Notation

# of Valence

Electrons

Oxidation Number

Sodium - 23

Calcium - 40

Aluminum-26

Chlorine - 35

Chlorine - 35

Fluorine - 19

Uranium-235

Iodine - 131

Oxygen - 16

Oxygen - 17

Phosphorus- 30

Phosphorus-32

AVERAGE ATOMIC MASS1. Calculate the average atomic mass of the element iron(Fe) using the following data:

Iron-54 6% Iron-56 92% Iron-57 2%

2. Calculate the average atomic mass of the element nitrogen(N) using the following data: Nitrogen-14 is 95% abundant, Nitrogen-15 is 3% abundant, and Nitrogen-16 is 2% abundant.

AP Chemistry Unit 2 Homework PacketPERIODIC TABLE REVIEW

1. Explain what property each element in a specific group has in common with each other.

2. Explain what elements in the same period have in common with each other.

3. Put a check in each box that correctly describes the element given.

Students should be able to demonstrate an understanding of the following essential knowledge: 1.B.1 The atom is composed of negatively charged electrons, which can leave the atom, and a positively charged

nucleus that is made of protons and neutrons. The attraction of the electrons to the nucleus is the basis of the structure of the atom.

1.C.1 Many properties of atoms exhibit periodic trends that are reflective of the periodicity of electronic structure.

Metal Metalloid Nonmetal AlkalineMetal

Alkaline EarthMetal

Transition metal

Halogen Noble gas

SbSrRnPPtCsS

FeBrArHSiBF

HeSeZnRa

AP Chemistry Unit 2 Homework Packet4. Write in the space, “metals”, “metalloids”, or “nonmetals” to indicate which type of element.

a.Located on the left side of the P.T.

b. Located on the right side of the P.T.c. Solids are brittled. Majority of the elementse. Gain electrons to form negative ionsf. Located along the “staircase”g. Have lusterh. Malleablei. Lose electrons to form positive ionsj. Ductilek. Excellent conductors of heat & electricityl. Poor electrical & heat conductorsm. Low electronegativity valuesn. Low ionization energy o. High ionization energyp. High electronegativity valuesq. Ions are larger than their atomsr. Ions are smaller than their atoms

5. Check all the boxes which describe the element.

Name

Physical Properties Chemical Properties

State at STP(s, l, or g) Brittle Malleable

/ductile

Conductor Ionizationenergy

Electro-negativity Electrons

Good Poor Low High Low High Lose Gain

CAgMg

IS

AuFeBrArH

Hg

ATOMIC RADIUS

Students should be able to demonstrate an understanding of the following essential knowledge: 1.B.1 The atom is composed of negatively charged electrons, which can leave the atom, and a positively charged

nucleus that is made of protons and neutrons. The attraction of the electrons to the nucleus is the basis of the structure of the atom.

1.C.1 Many properties of atoms exhibit periodic trends that are reflective of the periodicity of electronic structure.

AP Chemistry Unit 2 Homework Packet

Trends:

Across a period atomic radius _________________________

due to ____________________________________________________

____________________________________________________________

Down a group atomic radius _________________________

due to ____________________________________________________

____________________________________________________________

1. Identify and explain the trend in atomic size for the following transitions in the periodic table.

(a) Moving vertically from Ar to He

(b) Moving horizontally from Na to Ar

2. In each of the following pairs, pick the larger species. Explain you answer in each case.(a) Cu and Cu2+

(b) F and F-

(c) Na and K

3. Only one of the following statements is correct. Which one?(a) All cations are larger than their corresponding atoms(b) All anions are smaller than their corresponding atoms(c) Atomic size increases on transitioning from left to right across period

2 of the periodic table(d) The most common ion of chlorine is smaller than a chlorine atom(e) The most common ion of strontium is larger than a strontium atom (f) The most common potassium ion is larger than the most common sodium ion(g) The ions most commonly formed by group 16 elements are smaller than their corresponding atoms

4. Consider the plot below that shows atomic and ionic radii of the most commonly formed ion (in units of pm) for selected elements, plotted against atomic number.

AP Chemistry Unit 2 Homework Packet

(a) What do the elements that have smaller ionic radii than their corresponding atomic radii have in common?

(b) Suggest a reason for the absence of comparative atomic and ionic radii data for elements with atomic numbers of 2, 10 and 18.

(c) Identify the element with atomic number 19, identify the formula of the ion that it commonly forms, and convert the radii of both the atom and the ion to units of cm.

(d) What common feature can be identified for all of the non-metals on the plot?

(e) What accounts for the sharp increase in height of the lines that occurs at elements with atomic numbers 3, 11 and 19 respectively?

IONIZATION ENERGY (and more radii)

1. Consider the table:

(a) In which group does this element appear on the periodic table?

(b) What is the minimum number of electrons that this element must have?

Students should be able to demonstrate an understanding of the following essential knowledge: 1.B.1 The atom is composed of negatively charged electrons, which can leave the atom, and a positively charged

nucleus that is made of protons and neutrons. The attraction of the electrons to the nucleus is the basis of the structure of the atom.

1.C.1 Many properties of atoms exhibit periodic trends that are reflective of the periodicity of electronic structure.

IE 1st 2nd 3rd 4th

578 1817 2745 11580

AP Chemistry Unit 2 Homework Packet2. Arrange the following species in order of increasing size. Rb+, Y3+, Br-, Kr, Sr2+ and Se2-.

3. Is it possible for two different atoms to be isoelectronic? If so give examples.

4. Is it possible for two different anions to be isoelectronic? If so give examples.

5. Consider the table below:

IE 1st 2nd 3rd 4th 5th 6th

737 1450 7732 10540 13360

17995

(a) In which group will X be found? Explain.

6. Explain carefully why rubidium tends only to form a +1 ion?

7. Explain carefully why elements in the same group react in similar ways?

8. Identify any (and all) isoelectronic species in the following list; Fe2+, Sc3+, Ca2+, F-, Co2+, Co3+,Sr2+, Cu+, Zn2+ and Al3+.

9. Arrange the following atoms into order of increasing first ionization energy. Sr, Cs, S, F and As.

10. Explain each of the following observations.

(a) Sodium has a lower first-ionization energy than lithium.

(b) Oxygen has a lower first-ionization energy than nitrogen.

(c) There is a general increase in the first ionization energy from sodium to argon.

(d) Boron has a lower first ionization energy than beryllium. 11. Consider the ionization energies of elements X and Y shown below in kJmol-1. X and Y are in the same period of

the periodic table and are adjacent to one another in the table.

IE 1st 2nd 3rd 4th 5th 6th 7th 8th 9th

AP Chemistry Unit 2 Homework Packetx 1680 3375 6050 8409 1102

215165 17868 92038 106440

y 2080 3950 6122 9370 12180

15239 20000 23068 115375

(a) In which group would one find element X? Explain.

(a) Does element X lie to the right or the left of element Y in the periodic table? Explain.

(b) Which is the first period on the periodic table that these elements could be in? Explain.

(c) Why are the second ionization energies of both elements larger than their respective first ionization energies?

(d) It is found that Y has the largest first ionization energy in the period that it is found. What does this tell us about Y?

(e) It is found that element Q, which is in the same period as X and Y but lies to the left of element X in the periodic table, only has values for its first four ionization energies. Suggest a reason for this observation.

12. (a) Define first ionization.

(b) Write an equation to show the second ionization energy of calcium.

13. Why does N have a higher first IE than O? Explain using orbital notations.

14. Why does Be have a higher IE than B? Explain using orbital notations.

AP Chemistry Unit 2 Homework PacketELECTRONEGATIVITY AND POLARITY

Across a period electronegativity _________________________ due to ________________________________________________ _________________________________________________________________________________________________________________________

Down a group electronegativity _________________________ due to __________________________________________________ _________________________________________________________________________________________________________________________

1. Which element has the highest electronegativity? Why?

2. Explain the trend in EN from P to S to Cl.

3. Explain the trend in electronegativity from Cl to Br to I.

Students should be able to demonstrate an understanding of the following essential knowledge: 1.B.1 The atom is composed of negatively charged electrons, which can leave the atom, and a positively charged

nucleus that is made of protons and neutrons. The attraction of the electrons to the nucleus is the basis of the structure of the atom.

1.C.1 Many properties of atoms exhibit periodic trends that are reflective of the periodicity of electronic structure. 2.C.1 In covalent bonding, electrons are shared between the nuclei of two atoms to form a molecule or polyatomic

ion. Electronegativity differences between the two atoms account for the distribution of the shared electrons and the polarity of the bond.

AP Chemistry Unit 2 Homework Packet

Half Life Problems1) Fluorine-21 has a half life of approximately 5 seconds. What fraction of the original

nuclei would remain after 1 minute?

2) Iodine-131 has a half life of 8 days. What fraction of the original sample would remain at the end of 32 days?

3) The half-life of chromium-51 is 28 days. If the sample contained 510 grams, how much chromium would remain after 56 days? How much would remain after 1 year? How much was present 168 days ago?

4) If 20.0 g of a radioactive isotope are present at 1:00 PM and 5.0 g remain at 2:00 PM, what is the half life of the isotope?

5) The half life of Uranium-238 is 4.5 billion years and the age of earth is 4.5 X 109 years. What fraction of Uranium-238 that was present when Earth was formed still remains?

6) Chromium-48 decays. After 6 half-lives, what fraction of the original nuclei would remain?

7) The half life of iodine-125 is 60 days. What fraction of iodine-125 nuclides would be left after 360 days?

8) Titanium-51 decays with a half life of 6 minutes. What fraction of titanium would remain after one hour?

9) A medical institution requests 1 g of bismuth-214, which has a half life of 20 min. How many grams of bismuth-214 must be prepared if the shipping time is 2 h?

10)The half life of radium 226 is 1602 years. If you have 500 grams of radium today how many grams would have been present 9612 years ago?

AP Chemistry Unit 2 Homework Packet

Nuclear Decay Equations

AP Chemistry Unit 2 Homework Packet