UNIT 2: Atomic Theory and Structure Part D: The Periodic Table of the Elements

Big Picture Idea: An element’s electronic structure determines its chemical and physical properties and therefore its placement on the periodic table.

Big Picture Question: How does the periodic table use physical and chemical properties to arrange the elements? Suggested Resources…Homework AssignmentsClasswork Assignments Laboratory ActivitiesFormative AssessmentsTextbook pages: Chapter 6Websites: www.webelements.com

Key Terms:1. Mendeleev2. Moseley3. Periodic Law4. period5. group/family6. Alkali Metals7. Alkaline Earth Metals8. Halogen9. Nobel Gases

10. metals

11. nonmetals12. semimetals/metalloid 13. valence electrons14. s,p,d,f blocks

15. atomic radius16. ionic radius17. ionization energy (+

ion)18. electronegativity19. transition metals24. Lanthanide series25. Actinide series26. diatomic elements

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Directions: Use this information as a general reference tool to guide you through this unit. Don’t hesitate to ask your teacher for help!

By the conclusion of this unit, you should know the following:

1. Contributions of Moseley and Mendeleev in the development of the current periodic table.

2. Arrangement of periodic table based on electron configuration and names of groups/periods.

3. Group/Period trends for atomic radius, ionization energy, electronegativity, and metallic character based on atomic structure.

4. Molar mass can be used to covert between grams, moles and molecules of a diatomic element.

By the conclusion of this unit, you should be able to do the following:

1. Given a group of atoms, arrange them based on periodic trends.

2. State the periodic law3. Know scientists and their contributions to the development

of the Periodic Table4. Explain why elements in a group have similar properties5. Identify the four blocks on the periodic table6. Identify four important periodic trends and explain how

each reflects the electron configurations of the elements7. Use dimensional analysis to convert between molecules,

moles and grams of a diatomic element.

Practice Problems:

1. What did Mendeleev contribute and how does this relate to Moseley?

2. What does the Periodic Law state?

3. Look at a periodic table- identify periods, groups, families, s block, p block, d block, f block, alkali metals, alkaline earth metals, halogens, noble gases, metals, semimetals, nonmetals, metalloids .Draw arrows for ionization energy, atomic radius and electronegativity. Which element is highest for each?

4. Looking at electron configurations, what stays the same and what changes in a family (do it again for a period)?

5. What happens to ion size when ions gain electrons/ when they lose electrons?

6. Rank in order of increasing electronegativity a) Sr, K, Ca b) S, Se, Cl c) B, C, N d) Na, Li, K

7. Rank in order of atomic radius and explain why these trends happen a) Li, Be, B b) S, P, Cl c) Rb, K, Cs d) Na, K, Ca

8. Rank the ions in order of ionic radius, why do these trends happen a) Mg2+, Na+, Al3+ b) Ba2+, Ca2+, Sr2+ c) Se2-, O2-, S2- d) Mg2+, Mg, Ca e) S, S2-, O

9.Rank the atoms from lowest to highest ionization energy a) Cl, Br, Ar b) K, Rb, Na c) C, N, O d) Rb, Ba, Sr

10. Write the short hand electron configuration for element 19, element 53, element 88. Which families do they belong to and give electron dot structure.

11. Specify properties of m, nm, sm. Given an unknown element with some general properties can you identify it as a metal, non-metal, or semimetal? 12. List the diatomic elements.

13. a. Find the mass of 2.5 moles of nitrogen gas. B. How many molecules are in 50.0g of chlorine?

PERIODIC TABLE ELECTRON CONFIGURATION ACTIVITY1. Complete the chart below.2. Find the square corresponding to each element on the blank periodic table. Copy the last part of the

electron configuration into the periodic table square. 3. Look for patterns, and fill in as many of the squares as you can.

Element Electron Configuration Dot Diagram1. Hydrogen

2. Helium

3. Lithium

4. Potassium

5. Francium

6. Beryllium

7. Magnesium

8. Barium

9. Vanadium

10. Chromium

11. Manganese

12. Zinc

13. Boron

14. Aluminum

15. Carbon

16. Germanium

17. Nitrogen

18. Flourine

19. Neon

20. Xenon

Reading Electron Configuration Directly from the Periodic Table

Element Electron ConfigurationCesium [Xe] 6s1

Strontium [Kr]5s2

Silicon [Ne] 3s23p2

Argon [Ne] 3s23p6

Chromium [Ar] 4s23d4

Chlorine [Ne] 3s23p5

Tin [Kr] 5s24d105p2

Radium [Rn] 7s2

Nickel [Ar] 4s23d8

Palladium (Pd) [Kr] 5s24d8

Antimony (Sb) [Kr] 5s24d105p3

Francium (Fr) [Rn] 7s1

Electron Configuration Ending Element[Ne] 3s2 Mg

[Ar] 4s23d104p5 Br[Kr] 5s24d3 Nb

[Rn] 7s25f146d107p4 116 (Uuh)1s22s22p6 Ne[Kr] 5s1

Rb

[Ar] 4s23d10

Zn

[Kr] 5s2

Sr

[Kr] 5s24d105p3

Sb

[Xe] 6s24f145d4

W

History:

Mendeleev 1st periodic table (arranged elements by mass)

Moseley revised periodic table (arranged by atomic number)

Periodic Law the elements physical and chemical properties follow periodic (repeating) trends (based on their location of the periodic table) ***Patters repeat***

Parts of the Periodic Table:

Group/Family column – same valence electrons or outer configuration

Period row, same energy level

Metals good conductor, solid @ room temp., shiny, ductile, malleable, form + ions

Non-Metals poor conductor, solid/liquid/gas @ room temp., brittle, form - ions

Metalloids elements that touch the staircase (except Al), have intermediate properties

Families

Reactivity Ion Formation Misc.

Alkali Metals violently +1

Not found free in nature, e-

configuration ends in s1

Alkaline Earth Metals

very +2 e- configuration ends in s2

Transition Metals More stable Form ions with more than one charge

Form colored solutions, d block

Inner Transition Metals

Many are radioactive Varying charges F block

Halogens Most reactive nonmetals -1Diatomic as elements, e- configuration ends

in p5

Noble Gases stable Do not form ionsFull valence electron

shell, e- configuration ends in p6

Blocks

s block – 1st 2 columns on left side of the table

p block – last 6 columns on right side of the table

d block – middle 10 columns of the table

f block – lower (removed) 2 rows of the table

CLASSIFYING THE ELEMENTS

Part A CompletionUse this completion exercise to check your understanding of the concepts and terms that are introduced in this section. Each blank can be completed with a term, short phrase, or number.

1

The periodic table displays the symbols and _______ the

elements along with information about the structures of their

_______. The Group 1A elements are called _______, and the

Group 2A elements are called _______. The elements in Groups

1A through 7A are called the _______. The nonmetals of Group

7A are _______, and the _______ make up Group 8A. Between

Groups 2A and 3A, there are _______in periods 4 through 7 and

_______ in periods 6 and 7.

The atoms of the noble gas elements have their highest

occupied s and _______ sublevels filled. The highest occupied s

and p sublevels of the representative elements are _______.

1. names

2. atoms

3. alkali metals

4. alkaline earth metals

5. representative elements

6. halogens

7. noble gases

8. transition metals

9. inner transition metals

10. p

11. not filled

2 3

4

5

6 78

9

10

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Part B True-FalseClassify each of these statements as always true, AT; sometimes true, ST; or never true, NT.

ST_____ 12. Group A elements are representative elements.

NT 13. Chlorine has the electron configuration 1s22s22p63s23p7.

NT 14. The element in Group 4A, period 3, is gallium.

AT 15. There is a relationship between the electron configurations of elements and their chemical and physical properties.

Part C MatchingColumn A

f 16. alkali metals

e 17. inner transition metal

g 18. representative element

d 19. transition metal

b 20. noble gas

c 21. alkaline earth metals

a 22. halogens

Column B

a. nonmetals of Group 7A

b. an element in which the highest occupied s and p sublevels are filled

c. Group 2A elements

d. an element whose highest occupied s sublevel and a nearby d sublevel contain electrons

e. an element whose highest occupied s sublevel and a nearby f sublevel generally contain electrons

f. Group 1A elements

g. an element whose highest occupied s or p sublevels are partially filled

Worksheet: Periodic Table

For each of the following, circle the correct element.

Li Si S metal

K Ca Sc largest atomic mass

S Cl Ar member of the halogen family

V Nb Ta largest atomic number

Te I Xe member of noble gases

Si Cs B member of alkali metals

As Se Br 6 valence electrons

H Li Na nonmetal

Hg Tl Pb member of transition metals

Na Mg Al electron distribution ending s2p1

Pb Bi Po metalloid

B C N gas at room temperature

Ca Sc Ti electron distribution ending in s2d2

K S Ba has an electron dot like: X:

U Zn Kr member of inner transition metals

Ca S Br forms a +2 ion

F Na Mg member of the alkaline earth metals

Al Cr P 3 valence electrons

Rb I Al forms a -1 ion

H He O a gas that is not reactive

NOTES – PERIODIC TRENDS

Li F

K

Definition Down a Group Across a PeriodAtomic Radius – size of an atom (1/2 the distance between 2 nuclei of different atoms)

increases down a group because of more energy levels and more shielding

elements gets smaller (stronger nuclear pull)

Ionization Energy - energy needed to remove an electron from an atom

Decreases down a group because the electron is far from the nucleus and shielding

increases because nucleus is stronger, better protects electron

Electronegativity - ability of an atom to remove and electron from another atom

Decreases down a group because the electron is far from the nucleus and shielding from nucleus

increases because nucleus is stronger, better able to take electrons (EXCEPT noble gases)

NUCLEAR CHARGEamount of positive charge that reaches valence electrons

SHEILDINGinner energy levels block positive charge in nucleus from reaching valence electrons

PROBLEMS:

1. Circle the member of each pair that has the greatest radius:a. oxygen (O) or carbon (C) b. calcium (Ca) or barium (Ba)

c. magnesium (Mg) or phosphorus (P) d. strontium (Sr) or silver (Ag)

2. Arrange the elements below in order of increasing radius:sodium (Na), potassium (K), nickel (Ni), bromine (Br)

Na <Br < Ni < K

3. Circle the member of each pair that has the greatest ionization energy:a. oxygen (O) or carbon (C) b. calcium (Ca) or barium (Ba)

c. magnesium (Mg) or phosphorus (P) d. strontium (Sr) or silver (Ag)

4. Arrange the elements below in order of increasing ionization energy:sodium (Na), potassium (K), nickel (Ni), bromine (Br)

K < Ni < Na < Br

5. Select the member of each pair that has a greater electronegativity:a. oxygen or carbon b. calcium or barium

c. magnesium or nitrogen d. sulfur (S) or argon (Ar)

1. Circle the member of each pair that has the greatest radius:a. nitrogen or arsenic (As) b. calcium (Ca) or zinc (Zn)

c. manganese (Mn) or technetium (Tc) d. iodine (I) or rubidium (Rb)

e. boron (B) or neon f. krypton (Kr) or neon (Ne)

2. Arrange the elements below in order of increasing radius:Cesium (Cs), Potassium, Bromine, Br<K<Cs

3. Circle the member of each pair that has the greatest ionization energy:a. nitrogen or arsenic (As) b. calcium (Ca) or zinc (Zn)

c. manganese (Mn) or technetium (Tc) d. iodine (I) or rubidium (Rb)

e. boron (B) or neon f. krypton (Kr) or neon (Ne)

4. Arrange the elements below in order of increasing ionization energy:Cesium (Cs), Potassium, Bromine, Cs<K<Br

5. Select the member of each pair that has a greater electronegativity:a. nitrogen or arsenic (As) b. calcium (Ca) or zinc (Zn)

c. manganese (Mn) or technetium (Tc) d. iodine (I) or rubidium (Rb)

e. boron (B) or neon f. krypton (Kr) or neon (Ne) neitherFormation of Ions

PERIODIC TRENDS WORKSHEET

Octet Rule: elements try to get 8 valence electrons when they bond

Cations: example: Na vs. Na+ the ion is smaller than the atom it came from (less energy levels)!

Na Na+

Bohr Diagram

Electron Configuration 1s22s22p63s1 1s22s22p63s1 [ 1s22s22p6 ]+1

Anions: example: Cl vs. Cl- The ion is bigger than the atom it came from (nucleus has a harder time holding onto the electron if there are more electrons in the ion)!

Cl Cl-

Bohr Diagram

Electron Configuration 1s22s22p63s23p5 1s22s22p63s23p6

PROBLEMS: Arrange the following from smallest to largest in size (atomic and ionic radius).

a. N < N-3 b. Mg2+ < Mg

c. Be+2 < Li+1 <Li

d. F < F-1 < O-2

MOLECULE/MOLE/MASS CONVERSIONS FOR DIATOMIC ELEMENTS

Atomic mass: mass of one atom of an element

Molecular mass: total mass of all atoms in a molecule

Diatomic Elements: always fond as two atoms of the same element together H2 O2 N2 Cl2 Br2 I2 F2

Examples: oxygen: O2 1 mol = 32 g

bromine: Br2 1 mol = 160g

1. Find the mass of 1.20 moles of bromine

1.2 moles Br2 159.8 g = 191.76g1 mole Br2

2. How many molecules are in 2.4 moles of nitrogen

2.4 moles N2 6.02 x 1023 molecules = 1.4 x 1024 molecules 1 mole N2

3. Find the mass of 4.5 x 1025 molecules of iodine

4.5 x 1025 molecules I2 1 mole I2 253.8 g = 18971.8g 6.02 x 1023 molecules 1 mole

4. How many molecules are in 15.0g of oxygen?

15.0 g O2 1 mole O2 6.02 x 1023 molecules = 2.8 x 1023 molecules 32.0 g

5. *** Find the mass of 0.25moles of radon

0.25 moles 222 g = 55.5 g Rn1 mole

6. ***How many atoms are in 4.0g of oxygen?

4.0 g O2 1 mole 6.02 x 1023 molecules 2 atoms = 1.5 x 1023 at32.0 g 1 mole 1 molecule

MOLEPARTICLE

ATOMS&

MOLECULESMASS

(g)

1 mole = molar mass

(look it up on the PT!)

1 mole = 6.02 x 1023

atoms

PRACTICE MOLE CONVERSIONS:

1. List the diatomic elements: __H2__ __N2___ __O2___ __F2___ __Cl2___ __Br2___ __I2___

2. Find the mass of:a. 3.0 moles of oxygen __96 g___________

b. 1.15 moles of chlorine ___80.5 g__________

c. 0.35 moles of argon ___13.97 g___________

d. 1.14 x 1026 molecules of bromine ___30261.1 g___________

e. 2.55 x 1022 atoms of barium ___5.8 g___________

3. Convert to moles:a. 4.0g of hydrogen ___2 moles____________

b. 7.68 x 1024 molecules of iodine ___12.76 moles________

4. Find the number of molecules in:a. 16.0g of chlorine ___1.3 x 1023 molecules__

b. 2.5 moles of nitrogen __1.5 x 1024 molecules___

***5. How many atoms are in 4.0g of fluorine? __1.27 x 1023 atoms______

THE PERIODIC TABLE REVIEW

SECTION 6.1 ORGANIZING THE ELEMENTS1. Which element listed below should have chemical properties similar to fluorine (F)?

a. Li

b. Si

c. Br

d. Ne

2. Identify each element as a metal, metalloid, or nonmetal.

a. fluorine

b. germanium

c. zinc

d. phosphorus

e. lithium

3. Which of the following is not a transition metal?

a. magnesium

b. titanium

c. chromium

d. mercury

4. Name two elements that have properties similar to those of the element potassium.

5. Elements in the periodic table can be divided into three broad classes based on their general characteristics. What are these classes and how do they differ?

SECTION 6.2 CLASSIFYING THE ELEMENTS1. Use the periodic table to write the electron configuration for silicon. Explain your thinking.

2. Use the periodic table to write the electron configuration for iodine. Explain your thinking.

3. Which group of elements is characterized by an s2p3 configuration?

4. Name the element that matches the following description.

a. one that has 5 electrons in the third energy level

b. one with an electron configuration that ends in 4s24p5

c. the Group 6A element in period 4

5. Identify the elements that have electron configurations that end as follows.

a. 2s22p4

b. 4s2

c. 3d104s2

6. What is the common characteristic of the electron configurations of the elements Ne and Ar? In which group would you find them?

7. Why would you expect lithium (Li) and sulfur (S) to have different chemical and physical properties?

8. What characterizes the electron configurations of transition metals such as silver (Ag) and iron (Fe)?

SECTION 6.3 PERIODIC TRENDS1. Explain why a magnesium atom is smaller than atoms of both sodium and calcium.

2. Predict the size of the astatine (At) atom compared to that of tellurium (Te). Explain your prediction.

3. Would you expect a Cl– ion to be larger or smaller than an Mg2+ ion? Explain.

4. Which effect on atomic size is more significant, an increase in nuclear charge across a period or an increase in occupied energy levels within a group? Explain.

5. Explain why the sulfide ion (S2–) is larger than the chloride ion (Cl–).

6. Compare the first ionization energy of sodium to that of potassium.

7. Compare the first ionization energy lithium to that of beryllium.

8. Is the electronegativity of barium larger or smaller than that of strontium? Explain.

9. What is the most likely ion for magnesium to form? Explain.

10. Arrange oxygen, fluorine, and sulfur in order of increasing electro negativity.

Colored PeriodicTable Instructions:

1. First number each group from left to right (1-18). See the big periodic table at the front of the room.

2. Second, number each period from top to bottom (1-7) .

3. Draw in the stepladder separating metals and nonmetals. Label metals vs. nonmetals on the table. Describe the properties of each group.

4. Create a color coded key for each of the following families or series of elements, or include the information within the colored block: Include a description of the properties of each group or block of elements.

Alkali Metal Family (***don’t include hydrogen) +1Alkaline Earth Metal Family +2Halogens -1Nobel GasesTransition MetalsInner Transition MetalsMetalloids

5. Write the ionic charge of the family (shown above) at the top of the column. Copy additional charges from the board.

6. Indicate the trends down and across for each of the following properties using arrows: Atomic Radii Ionization Energy Electronegativity

7. Identify fluorine as the most reactive non-metal

8. Identify francium as being in the spot for the most reactive metal

9. Label the diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2, At2

10. Label Al3+, Zn2+, Ag+1