Chemistry 2: Periodic Table Name __________________________A. Quantum Mechanical Model (6.5 to 6.6)

1. electron’s exact position or velocity is unknowable a. Heisenberg uncertainty principleb. orbital = 90 % probable location of an electron

2. each electron has four quantum numbers: n, l, ml, ms a. principal energy levels (n)—defines orbital radiusb. sublevels (l)—defines orbital shape

1. l = 0, 1, 2, •••, (n - 1)2. 0 = s, 1 = p, 2 = d, 3 = f

c. orbital (ml)—defines spatial orientation1. ml = – l, ••• -1, 0, +1, •••, + l 2. number of orbitals: s (1), p (3), d (5), f (7)

d. spin (ms)—defines magnetic field +½ (), -½ ()Pauli exclusion principle—no two electrons can have the same spin in the same orbital

3. relationship among values of n, l, ml through n = 4n Possible l Sublevel

DesignationPossible ml

1 0 1s 0

2 0 2s 01 2p -1, 0, 1

30 3s 01 3p -1, 0, 12 3d -2, -1, 0, 1, 2

4

0 4s 01 4p -1, 0, 12 4d -2, -1, 0, 1, 23 4f -3, -2, -1, 0, 1, 2, 3

B. Electron Arrangements in Atoms and Ions (6.7 to 6.9)1. electrons fill from low to high energy (same for all atoms)

a. n: 1 < 2 < 3 < 4 < 5 < 6 < 7 b. l: s < p < next energy level < d < fc. (n, ml) equal energy = degenerate

2. outer (valence) electrons—interact with other atomsa. highest occupied principle energy levelb. s and p sublevels only (maximum 8 electrons)

Ene

rgy

5s4p

3d4s

3p

3s2p

2s

1s3. inner (core) electrons + nucleus = core charge (+1 to +8)4. organization of the periodic table

a. row (period)—same valence energy level (n)b. column (group)—same # of valence electrons

1. main groups (1: alkali metals, 2: alkaline earth metals, 17: halogens, 18: noble gases)

2. similar chemical properties1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

1 1s 1s2 2s 2p3 3s 3p4 4s 3d 4p5 5s 4d 5p6 6s 5d 6p7 7s 6d 7p

lanthanide

actinide

5. types of diagramsa. electron configuration

1. n (#), l (letter), # of electrons (superscript)Al: 1s22s22p63s23p1

2. abbreviated: replace inner (core) electrons with noble gas symbol—Al: [Ne]3s23p1

b. orbital diagrams1. electrons (arrows) fill specific orbital2. Hund's rule: maximum # of electrons with

the same electron spin (maximum number of half-filled degenerate orbitals)

Element # e- Orbital Diagram Electron Configuration

1s 2s 2p 3sLi 3 1s2 2s1

Be 4 1s2 2s2

B 5 1s2 2s2 2p1

C 6 1s2 2s2 2p2

N 7 1s2 2s2 2p3

O 8 1s2 2s2 2p4

F 9 1s2 2s2 2p5

Ne 10 1s2 2s2 2p6

Na 11 1s2 2s2 2p6 3s1

c. quantum numbers (arranged by periodic table position)

n l = 0(s) ms (½, -½)

l = 1(p)

1 ½ -½2 ½ -½ l = 2 (n = row # – 1)

(d)½ ½ ½ -½ -½ -½

3 ½ -½ ½ ½ ½ -½ -½ -½4 ½ -½ ½ ½ ½ ½ ½ -½ -½ -½ -½ -½ ½ ½ ½ -½ -½ -½

0 0 -2 -1 0 1 2 -2 -1 0 1 2 -1 0 1 -1 0 1ml

6. column 6 and 11: an s electron moves to the d sublevel to maximize half and/or full orbitals (lower energy state)

4s 3d24Cr [Ar]

29Cu [Ar] 7. electron arrangements in monatomic ions

a. ions with noble gas structure1. elements that are ± 3 of the noble gas lose or

gain electrons to reach noble gas electron configuration

3- 2- 1- 0 1+ 2+ 3+N3- O2- F- Ne Na+ Mg2+ Al3+

2. ions with the same # of e: isoelectronicb. transition metal ions

1. transition metal lose s electrons first2. may lose d electrons if it eliminates sublevel

or reduces doubling up4s 3d

29Cu+ [Ar]

26Fe3+ [Ar] 8. magnetic properties

a. element is magnetic (paramagnetic) if it contains unpaired electrons, whose magnetic fields are reinforcing respond to external magnetic field

4s 3d26Fe [Ar]

b. elements with all paired electrons (diamagnetic) are unaffected by magnetic fields (columns 2, 12, 18)

C. Periodic Properties—Main Groups (7.1 to 7.6)1. effective nuclear charge (Zeff) or shielding effect

a. atom holds valence electrons because of attraction between atom core and valence shell

b. atomic core = nucleus + core electrons 1. core charge = Zeff (# p – # core e)2. other valence electrons reduce Zeff because of

electron-electron repulsion2. atomic size (radius)

a. minimum distance between two gas atoms or distance between nuclei in diatomic molecule

b. group: increase as energy level increasesc. period: decrease as Zeff increases from +1 to + 8d. transition metals

1. group: increase with energy level2. period: no change (Zeff constant)

3. ionic size (radius) compared to parent atoma. smaller cations (lose energy level)b. larger anions (more electron-electron repulsion)c. isoelectronic series (± 3 from noble gas) largest

(fewest protons) to smallest (most protons)N3- > O2- > F- > Ne > Na+ > Mg2+ > Al3+

4. ionization energya. E to remove electron from a gaseous atom

1. X(g) X+(g) + 1e-

2. all +E (greater value = harder to ionize)b. inversely proportional to atomic radius (E 1/r)

1. weaker hold on distant electron less energy to remove electron

2. anomalies: ionized electron comes from a relatively less negative energy level less energy to reach zero energy (ionization) a. 13: ionized electron comes from a higher

energy sublevel p vs. sb. 16: ionized electron comes from a full

orbital (higher energy than ½-filled)c. successive ionization energies

Successive Ionization Energies in kJ/molElement I1 I2 I3 I4

Na 496 4562*** (core electrons)Mg 738 1451* 7733***Al 578 1817** 2745* 11,577**

*1. *small increases within a sublevel2. **greater increase between sublevels3. ***greatest increase between energy levels

5. electron affinitya. E to add electron to gaseous atom

1. X(g) + 1 e- X-(g)2. E = EX X- + Ee-

a. EX X- > 0 (added electron increases atom's overall energy level)

b. Ee- < 0 (added electron's energy decreases from zero)

c. –E when added electron enters relatively low-energy orbital (stable anion)

d. +E when added electron enters relatively high-energy orbital (unstable anion)

b. electron affinity becomes more negative from left to right because receiving orbital energy decreases

c. anomalies: receiving orbital energy is relatively high 2: p orbital energy > s orbital energy 15: full orbital energy > ½-filled orbital energy

(electron-electron repulsion raises level) 18: next higher energy level > valence level

d. group: little change for next energy level because greater orbital volume reduces electron-electron repulsion, which counterbalances reduced core attraction

6. metals, nonmetals and metalloidsa. metals—left side of stair step

1. shiny, conduct heat and electricity, malleable and ductile, mostly solids (except Hg)

2. form ionic compounds with nonmetals3. small positive ionization energy4. positive or small negative electron affinity5. lose electrons during reactions

(alkali metals are most reactive)b. nonmetals—right side of stair step and H

1. opposite properties to metals2. form molecules in addition to ionic compounds3. large positive ionization energy4. large negative electron affinity except

columns 15 and 185. gain or share electrons during reactions except

noble gases (halogens are most reactive)c. metalloids—touch stair step except Al

1. intermediate properties depending on physical and chemical conditions

2. Po and At classification unresolved

Experiments1. Halogen Lab (Wear Goggles)—Mix halogen molecules

with halide ions to determine if an electron is transferred and in doing so, rank Cl, Br and I from most negative electron affinity to least negative.Add 20 drops of bromine water to three small test tubes. Add 20 drops of hexanes (HEX) to each test tube. Stopper the test tubes and shake until the bromine color is mostly in the HEX layer. (AVOID BREATHING OR TOUCHING HALOGENS). Record the color of the top (HEX) layer. Add 20 drops of 0.1 M NaCl to the first tube, 20 drops of 0.1 M NaBr to the second tube, and 20 drops of 0.1 M NaI to the third. Stopper and shake each tube. Record the color of the HEX layer. Repeat with chlorine water and iodine water. a. How do you know if the halogen molecule took an

electron from the halide ion?

b. Record the color of the hex layer after adding the halogen, and after adding the halide ions. Highlight the box where a reaction occurred.

Halogen

water

HEXColor

HEX layer Color

With Br- With Cl- With I-

Br2

Cl2I2

c. Highlight which reaction (1 or 2) occurred for each pair.Pair Reaction 1 Reaction 2

Br2 & I- Br2 + 2 I- 2 Br- + I2 I2 + 2 Br- 2 I- + Br2

Cl2 & Br- Cl2 + 2 Br- 2 Cl- + Br2

Br2 + 2 Cl- 2 Br- + Cl2

Cl2 & I- Cl2 + 2 I- 2 Cl- + I2 I2 + 2 Cl- 2 I- + Cl2d. The halogen with the more negative electron affinity will

take an electron from the halogen with the less negative electron affinity. Rank the halides from most negative electron affinity to least negative electron affinity.

e. How do the results compare to the electron affinity data?Cl Br I

-349 kJ/mol -325 kJ/mol -295 kJ/mol

f. What is the correlation between electron affinity and atomic radius? Give a reason for this correlation.

Practice ProblemsA. Quantum Mechanical Model

1. Complete the number of electrons in each sublevel and total for the following energy levels. Level

nMaximum in sublevels, 2(2l +1) Total

2n2s (l = 0) p (l = 1) d (l = 2) f (l = 3)1234

2. Fill in the quantum numbers for the first three principle energy levels, and then answer the questions below.n 1 2 3l

ml

ms

a. Give the four quantum numbers for the two electrons in the 3d sublevel where ml = 0.

b. How many electrons have the quantum numbers?2,1, __, __ 3, __, __, +½B. Electron Arrangements in Atoms and Ions

3. Write the order that electrons fill sublevels from 1s to 7p.

4. Complete the chart for each element.

Symbol Period #

Group # Group Name Metal or

Nonmetal5 Halogen4 1

Be3 noble gas

5. In the electron configuration (1s2) what does each indicate?1 s 2

6. Write the electron configuration formagnesium silicon Iron

7. What does the symbol [Ar] represent in [Ar]4s23d8?

8. Write the two ways to show the abbreviated configuration for iodine.

emphasize filling order emphasize valence shell

9. Write the orbital diagram for the following.1s 2s 2p 3s 3p 4s 3d

S

Co10. What is the abbreviated electron configuration for each?

Cr Cu Mo Ag

11. Write the quantum numbers for the boxed electrons.1s 2s 2p 3s 3p 4s 3d

12. Write the set of quantum numbers for the electron position (a-d) on the periodic table below.12 a34 b d5 c

a b c d

C. Periodic Properties—Main Groups13. List the ions with Ar noble gas structure.

cations anions

14. Write the abbreviated electron configurations for each ion.Sc3+ Ag+ Ru3+ Zn2+

15. Label the following as paramagnetic or diamagnetic.Cu Mg Cu+

16. Use the relative sizes of atoms and their common ion for columns 1, 2, 3, 16 and 17 to answer the questions.

a. Indicate whether the trend is increase or decrease for.

Relative Size Increase

Decrease

Atoms from top to bottomCations from top to bottomAnions from top to bottomAtoms from left to rightCations from left to rightAnions from left to rightIsoelectric ions from low Z to high ZSize of cation compared to atomSize of anion compared to atomb. Explain why cations are smaller than their atoms.

c. Explain why anions are larger than their atom.

17. Which period 3 element has the following successive ionization energies?

First Second Third Fourth Fifth786 1577 3232 4356 16,091

18. Graph the data on the grids below.a. Atomic Radius (x 10-10 m)

Group # 1 2 13 14 15 16 17 18Period

2 1.34 0.90 0.82 0.77 0.75 0.73 0.71 0.69

Period 3 1.54 1.30 1.18 1.11 1.06 1.02 0.99 0.97

1.5

1.3

1.1

0.9

0.7 b. Ionization Energy (kJ/mol)Group # 1 2 13 14 15 16 17 18Period

2 520 899 801 1086 1402 1314 1681 2081

Period 3 496 738 578 786 1012 1000 1251 1521

2100

1900

1700

1500

1300

1100

900

700

500

c. Electron Affinity (kJ/mol)Group # 1 2 13 14 15 16 17 18Period

2 -60 > 0 -27 -122 > 0 -141 -328 > 0

Period 3 -53 > 0 -43 -134 -72 -200 -349 > 0

> 00

-100

-200

-300

19. Use the graphed data above to complete the table.Atomic Radius

Period Pattern left to rightAnomaly group numbers

Ionization EnergyPeriod Pattern left to rightAnomaly group numbers

Electron AffinityPeriod Pattern left to rightAnomaly group numbers

20. Explain the following observations.a. Atomic radii of Li = 1.34 Å and Na = 1.54 Å.

b. Atomic radii of Al = 1.18 Å and Si = 1.11 Å.

c. First ionization energies for B = 801 kJ/mol and Al = 578 kJ/mol.

d. First ionization energies for Si = 786 kJ/mol and P = 1012 kJ/mol.

e. First ionization energies for Mg = 738 kJ/mol and Al = 578 kJ/mol.

f. First ionization energies for P = 1012 kJ/mol and S = 1000 kJ/mol.

g. For Na, the first ionization energy I1 = 495 kJ/mol and second ionization energy I2 = 4562 kJ/mol.

h. The gap between the first and second ionization energies is greater for Al (I1 = 578 kJ/mol and I2 = 1817 kJ/mol) than Si (I1 = 786 kJ/mol and I2 = 1577 kJ/mol).

i. The common ion for magnesium is Mg2+, where I1 = 738 kJ/mol, I2 = 1451 kJ/mol and I3 = 7733 kJ/mol.

j. The electron affinities for Mg > 0 and Na = -53 kJ/mol.

k. The electron affinities for Si = -134 kJ/mol and P = -72 kJ/mol.

l. The electron affinities for S = -200 kJ/mol and Cl = -349 kJ/mol.

m. The electron affinities for Cl = -349 kJ/mol and Ar > 0.

Practice Multiple ChoiceBriefly explain why the answer is correct in the space provided.Questions 1-3

(A) Heisenberg uncertainty principle (B) Pauli exclusion principle (C) Hund's rule(D) Shielding effect

1. Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic.

2. States that an orbital can hold no more than two electrons.

3. Predicts that it is impossible to determine simultaneously the exact position and the exact velocity of an electron.

4. Which set of quantum numbers describes the highest energy valence electron in a ground-state gallium atom (Z = 31)?(A) 4,0,0,½ (B) 4,0,1,½ (C) 4,1,1,½ (D) 4,1,2,½

5. Which has the outer electronic configuration, s2p3? (A) Si (B) Cl (C) Se (D) As

Questions 6-9 refer to atoms with the atomic orbitals shown. (A) 1s 2s (B) [He] 2s 2p (C) [He] 2s 2p (D) [Ar] 4s3d

6. Represents an atom that is chemically unreactive.

7. Represents an atom in an excited state.

8. Represents an atom that has four valence electrons.

9. Represents an atom of a transition metal.

Questions 10-12(A) 1s2 2s22p5 3s23p5

(B) 1s2 2s22p6 3s23p6

(C) 1s2 2s22p62d10 3s23p6

(D) 1s2 2s22p6 3s23p63d3 4s2

10. An impossible electronic configuration

11. The ground-state configuration for a halogen anion

12. The ground-state configuration for an alkaline earth cation

13. Which represents the ground state for the Mn3+ ion?(A) 1s2 2s22p6 3s23p63d4

(B) 1s2 2s22p6 3s23p63d5 4s2

(C) 1s2 2s22p6 3s23p63d2 4s2

(D) 1s2 2s22p6 3s23p63d8 4s2

14. In which group are the three species isoelectronic? (A) S2-, K+, Ca2+ (B) Sc, Ti, V2+ (C) O2-, S2-, CI- (D) Mg2+, Ca2+, Sr2+

15. The ionization energies for element X are listed in the table below. On the basis of the data, element X is most likely

Ionization Energies for element X (kJ mol-1)First Second Third Fourth Fifth580 1,815 2,740 11,600 14,800

(A) Na (B) Mg (C) Al (D) Si

16. In the periodic table, as the atomic number increases from 11 to 17, what happens to the atomic radius? (A) It remains constant. (B) It increases only. (C) It decreases only. (D) It increases, then decreases.

17. Which elements have most nearly the same atomic radius?(A) Be, B, C, N (B) Ne, Ar, Kr, Xe(C) Mg, Ca, Sr, Ba (D) Cr, Mn, Fe, Co

Questions 18-19 Use the following options.(A) O (B) Rb (C) N (D) Mg

18. What element has the most negative electron affinity?

19. Which of the elements above has the smallest ionic radius for its most commonly found ion?

20. Which gaseous atoms (Ca, V, Co, Zn, As) are paramagnetic?(A) Ca and As only (B) Zn and As only(C) Ca, V, and Co only (D) V, Co, and As only

21. Which property generally decreases across the periodic table from sodium to chlorine? (A) 1st ionization energy (B) Atomic mass (C) Ionic radius (D) Atomic radius

Questions 22-23 Consider a ground state atom of each element. (A) S (B) Ca (C) Ga (D) Sb

22. The atom that contains exactly two unpaired electrons

23. The atom that contains only one electron in the highest occupied energy sublevel

Practice Free Response1. For the following elements:

a. Write the abbreviated electron configurationb. Write the abbreviated orbital diagram c. Circle the electron in (b) that has the quantum numbers.d. Write the abbreviated electron configuration for the ion

Mg Cu P

a

b

c (3, 0, 0, -½) (3, 2, -2, +½) (3, 1, 0, ½)

d Mg2+ Cu+ P3-

2. The table shows the first three ionization energies in kJ/mol for third period elements, which are numbered randomly. Use the information to answer the questions.Elemen

t First Second Third

1 1,251 2,300 3,820

2 496 4,562 6,9103 738 1,451 7,7334 578 1817 2745

a. Which element is most metallic in character? Explain.

b. Identify element 3. Explain.

c. Complete the following.Electron configuration of element 3

Common ion charge of element 2

Chemical symbol of element 2Element with the smallest atomic radiusChemical symbol for element 4

3. Use the principles of atomic structure to explain each. a. The atomic radius of Li is larger than that of Be.

b. The electron affinity for K is less than Ca.

c. The first ionization energy of Se is less than As.

4. Consider the element strontium (Sr). Justify each answer.a. What is the outer electron configuration of Sr?

b. How does the atomic radius of Sr compare to Rb?

c. How does the atomic radius of Sr compare to Ca?

d. Compare Sr to Ca and Rb in first ionization energy.

e. How does the Sr2+ ion compare in size to the Sr atom?

f. How does the Sr2+ ion compare in size to the Br- ion?

g. As successive electrons are removed from the Sr atom, where does the largest jump in ionization energy occur?

h. Is strontium diamagnetic or paramagnetic?

i. Highlight the correct option when completing the following sentence: compared to Br, Sr is shinier/duller and a better/worst conductor.