o] = 1 2 •[ne] = 1 2 - simeon ca

Copyright © by Holt, Rinehart and Winston. All rights reserved.

ResourcesChapter menu

Objectives

• Relate the electron configuration of an atom to its chemical reactivity.

• Determine an atom’s number of valence electrons, and use the octet rule to predict what stable ions the atom is likely to form.

• Explain why the properties of ions differ from those of their parent atoms.

Section 1 Simple IonsChapter 5



Chemical Reactivity

• How much an element reacts depends on the electron configuration of its atoms.

• Example, oxygen reacts with magnesium. In the electron configuration for oxygen, the 2p orbitals, which can hold six electrons, have only four:

• [O] = 1s22s22p4

• Neon has no reactivity. Its 2p orbitals are full:

• [Ne] = 1s22s22p6




Chemical Reactivity, cont. NOTE: Noble Gases Are the Least Reactive Elements• Neon is a noble gas and the noble gases, Group 18 in

periodic table, show almost no chemical reactivity.

• The noble gases have filled outer energy levels.

• This electron configuration can be written as ns2np6

where n is the period # of the outer energy level.




Noble Gases Are the Least Reactive Elements, cont.• The 8 electrons in outer energy level fill the s

and p orbitals, making these noble gases stable.

• In most chemical reactions, atoms try to match the s and p electron configurations of the noble gases.

• Seeking either empty outer energy levels or full outer energy levels of 8 electrons is called the octet rule.




Chemical Reactivity, continuedAlkali Metals and Halogens Are the Most Reactive Elements• Atoms whose outer s and p orbitals do not match the

electron configurations of a noble gas will react to lose or gain electrons to fill outer orbitals.

• In water, potassium atom (an alkali metal) gives up one electron in its outer energy level.

• Then, it has the s and p configuration of a noble gas.

1s22s22p63s23p64s1 1s22s22p63s23p6




Chemical Reactivity, continuedAlkali Metals and Halogens Are the Most Reactive Elements, continued• Chlorine, a halogen, is also very reactive.

• An atom of Chlorine has seven electrons in its outer energy level.

• By gaining just one electron, it will have the s and pconfiguration of a noble gas.

1s22s22p63s23p5 1s22s22p63s23p6




Valence Electrons

• Potassium K after it loses one electron has the same electron configuration as chlorine Cl after it gains one.

• Both are the same as that of the noble gas argon.

[Ar] = 1s22s22p63s23p6

• The atoms of many elements become stable by achieving the electron configuration of a noble gas.

• The electrons in the outer energy level are known as valence electrons.




Valence Electrons, continuedPeriodic Table Reveals an Atom’s Number of Valence Electrons• To find out how many valence electrons an atom has,

check the periodic table.

• For example, the element magnesium, Mg, has the following electron configuration:

[Mg] = [Ne]3s2

• This configuration shows that a magnesium atom has two valence electrons in the 3s2 orbital.




Valence Electrons, THIS TABLE IS IN NOTEBOOKPeriodic Table Reveals an Atom’s Number of Valence Electrons, continued

• The electron configuration of phosphorus, P, is [Ne]3s23p3. • Each P atom has five valence electrons: two in the 3s orbital

and three in the 3p orbital.




Valence Electrons, continuedAtoms Gain Or Lose Electrons to Form Stable Ions• All atoms are uncharged because they have equal

numbers of protons and electrons.

• For example, a potassium atom has 19 protons and 19 electrons.

• After giving up one electron, potassium still has 19 protons but only 18 electrons.

• Because there are more protons than electrons, there is a net electrical charge (+1).




Valence Electrons, continuedAtoms Gain Or Lose Electrons to Form Stable Ions, continued• An ion is an atom, radical, or molecule that has

gained or lost one or more electrons and has a negative or positive charge.

• The following equation shows how a potassium atom forms an ion with a 1+ charge. Positive ions are cations.

K → K+ + e−




Valence Electrons, continuedAtoms Gain Or Lose Electrons to Form Stable Ions, continued• In the case of chlorine, far less energy is required for

an atom to gain one electron rather than give up its seven valence electrons to be more stable.

• The following equation shows how a chlorine atom forms an ion with a 1− charge. Negative ions are anions.

Cl + e− → Cl−




Valence Electrons, continuedCharacteristics of Stable Ions• Both an atom and its ion have the same number of

protons and neutrons, so the nuclei are the same.

• Because chemical properties of an atom depend on the number and configuration of its electrons, an atom and its ion have different chemical properties.




Valence Electrons, continuedMany Stable Ions Have Noble-Gas Configurations• Many atoms can form stable ions with a full octet. For

example, Ca, forms a stable ion.

• The electron configuration of a calcium atom is:

[Ca] = 1s22s22p63s23p64s2

• By giving up its two valence electrons in the 4sorbital, it forms a stable cation with a 2+ charge:

[Ca2+] = 1s22s22p63s23p6

• This electron configuration is like that of argon.




Some Ions with Noble-Gas Configurations

Section 1 Simple Ions IN NOTEBOOK



Valence Electrons, continuedSome Stable Ions Do Not Have Noble-Gas Configurations• Not all stable ions have an electron configuration like

those of noble gases. Transition metals often form ions without complete octets.

• With the lone exception of rhenium, Re, the stable transition metal ions are all cations.

• Also, some elements, mostly transition metals, form stable ions with more than one charge.




Stable Ions Formed by the Transition Elements and Some Other Metals

Section 1 Simple Ions in notebook



Atoms and Ions

• Both sodium and chlorine are very reactive.• When they are mixed, a violent reaction takes place,

producing a white solid—table salt (sodium chloride).

• It is made from sodium cations and chloride anions.

Ions and Their Parent Atoms Have Different Properties

• Having identical electron configurations does not mean that a sodium cation is a neon atom.

• They still have different numbers of protons and neutrons.




Atoms and Ions, continuedAtoms of Metals and Nonmetal Elements Form Ions Differently• Nearly all metals form cations. For example,

magnesium metal, Mg, has the electron configuration:

[Mg] = 1s22s22p63s2

• To have a noble-gas configuration, the atom must either gain six electrons or lose two.

• Less energy to lose two electrons than gaining six.




Atoms and Ions, continuedAtoms of Metals and Nonmetal Elements Form Ions Differently, continued• Atoms of all nonmetal elements form anions. For

example, oxygen, O, has the electron configuration:

[O] = 1s22s22p4

• For noble-gas configuration, an oxygen atom must either gain two electrons or lose six.

• Acquiring two electrons requires less energy than losing six.




Objectives HIGHLIGHT THE DETAILS

• Describe the process of forming an ionic bond.

• Explain how the properties of ionic compounds depend on the nature of ionic bonds.

• Describe the structure of salt crystals.

Section 2 Ionic Bonding and SaltsChapter 5



Ionic Bonding

• Pyrite is a mineral that is shiny like gold, but it is made of iron cations and sulfur anions.

• Because opposite charges attract, cations and anions attract one another and an ionic bond is formed.

• The iron cations and sulfur anions of pyrite attract one another to form an ionic compound.




Ionic Bonding, continuedIonic Bonds Form Between Ions of Opposite Charge• When sodium and chlorine react to form sodium

chloride, sodium forms a stable Na+ cation and chlorine forms a stable Cl− anion.

• The force of attraction between the 1+ charge on the sodium cation and the 1− charge on the chloride anion creates the ionic bond in sodium chloride.

• Sodium chloride is a salt, the scientific name given to many different ionic compounds.




Ionic Bonding, continuedIonic Bonds Form Between Ions of Opposite Charge, continued• Salts are electrically neutral ionic compounds made

up of cations and anions held together by ionic bonds in a simple, whole-number ratio. i.e. 1:1 2:1 3:1

• Attractions between ions in a salt do not stop with a single cation and a single anion.

• One cation attracts several anions, and one anion attracts several cations.

• They are all pulled together into a tightly packed crystal structure.




Characteristics of Ion Bonding in a Crystal Lattice

Visual Concepts in notebookChapter 5



Ionic Bonding, continuedTransferring Electrons Involves Energy Changes• Ionization energy is the energy that it takes to remove

the outermost electron from an atom.

• The equation below shows this process for sodium.

Na + energy → Na+ + e−

• some elements, such as chlorine, release energy when an electron is added.

Cl + e− → Cl− + energy




Ionic Bonding, continuedTransferring Electrons Involves Energy Changes, continued• The energy released when chlorine accepts an

electron is less than energy required to remove an electron from a sodium atom.

• Adding/removing electrons is only part of forming an ionic bond.

• The rest of the process of forming a salt supplies enough energy to make up the difference so that the overall process releases energy.




Ionic Bonding, continuedSalt Formation Involves Endothermic Steps• The process of forming the salt sodium chloride can

be broken down into five steps.

1. Energy is needed to make solid sodium a gas.

Na(solid) + energy → Na(gas)

2. Energy is also required to remove an electron from a gaseous sodium atom.

Na(gas) + energy → Na+(gas) + e−




Ionic Bonding, continuedSalt Formation Involves Endothermic Steps, continued

3. Chlorine exists as a molecule containing two chlorine atoms. Energy must be supplied to separate the chlorine atoms so that they can react with sodium.

Cl–Cl(gas) + energy → Cl(gas) + Cl(gas)

• To this point, the first three steps have all been endothermic. These steps have produced sodium cations and chlorine atoms.




Ionic Bonding, continuedSalt Formation Also Involves Exothermic Steps

4. An electron is added to a chlorine atom to form an anion. This step releases energy.

Cl(gas) + e−→ Cl−(gas) + energy

5. When a cation and anion form an ionic bond, it is an exothermic process. Energy is released.

Na+(gas) + Cl−(gas) → NaCl(solid) + energy

• The last step is the driving force for salt formation.




Ionic Bonding, continuedSalt Formation Also Involves Exothermic Steps, continued• The energy released when ionic bonds are formed is

called the lattice energy.

• This energy is released when the crystal structure of a salt is formed as the separated ions bond.

• Without this energy, there would not be enough energy to make the overall process spontaneous.




Ionic Bonding, continuedSalt Formation Also Involves Exothermic Steps, continued• If energy is released when ionic bonds are formed,

then energy must be supplied to break these bonds.

• As sodium chloride dissolves in water, water supplies energy for the Na+ and Cl− ions to separate.

• Because of its much higher lattice energy, magnesium oxide does not dissolve well in water.

• There is not enough energy to separate the Mg2+

and O2− ions from one another.




Ionic Compounds Do Not Consist of Molecules

• The ratio of cations to anions is always such that an ionic compound has no overall charge.

Ionic Compounds

• Water is a molecular compound, so individual water molecules are each made of two hydrogen atoms and one oxygen atom.

• Sodium chloride is an ionic compound, so it is made up of many Na+ and Cl− ions all bonded together to form a crystal. There are no NaCl molecules.




Ionic Compounds, continuedIonic Compounds Do Not Consist of Molecules, continued• Metals and nonmetals tend to form ionic compounds

and not molecular compounds.

• The formula CaO likely indicates an ionic compound because Ca is a metal and O is a nonmetal.

• In contrast, the formula ICl likely indicates a molecular compound because both I and Cl are nonmetals.

• Lab tests are used to confirm such indications.




Ionic Compounds, continuedIonic Bonds Are Strong• Repulsive forces in a salt crystal include those

between like-charged ions.

• Each Na+ ion repels the other Na+ ions. Each Cl− ion repels the other Cl− ions.

• Another repulsive force exists between the electrons of ions that are close together.

• Attractive forces include those between the positive nucleus of one ion and electrons of other ions.




Ionic Compounds, continuedIonic Bonds Are Strong, continued• Attractive forces exist between oppositely charged

ions and involve more than a single cation and anion.

• Six Na+ ions surround each Cl− ion and vice versa.

• As a result, the attractive force between oppositely charged ions is significantly greater in a crystal than it would be if the ions existed only in pairs.

• Overall, the attractive forces are much stronger than the repulsive ones, so ionic bonds are strong.




Ionic Compounds, continuedIonic Compounds Have Distinctive Properties• Most ionic compounds have high melting and boiling

points because of the strong attraction between ions.

• To melt, ions cannot be in fixed locations.

• Because the bonds between ions are strong, a lot of energy is needed to free them.

• Still more energy is needed to move ions out of the liquid state and cause boiling, so ionic compounds are rarely gaseous at room temperature.




Ionic Compounds, continuedIonic Compounds Have Distinctive Properties, continued




Ionic Compounds, continuedLiquid and Dissolved Salts Conduct Electric Current• To conduct an electric current, a substance must satisfy

two conditions: • it must contain charged particles• those particles must be free to move

• Ionic solids, such as salts, generally are not conductors because the ions cannot move.

• When a salt melts or dissolves, the ions can move about and are excellent electrical conductors like saltwater.




Ionic Compounds, continuedSalts Are Hard and Brittle• Like NaCl, most ionic compounds are hard and brittle.

• Hard means that the crystal is able to resist a large force applied to it.

• Brittle means that when the applied force becomes too strong to resist, the crystal develops a widespread fracture rather than a small dent.

• Both properties are due to the patterns in which the cations and anions are arranged in all salt crystals.




Ionic Compounds, continuedSalts Are Hard and Brittle, continued• The ions in a crystal are arranged in a repeating

pattern, forming layers.

• Each layer is positioned so that a cation is next to an anion in the next layer. The attractive forces between opposite charges resist motion.

• As a result, the ionic compound will be hard.

• Also, it will take a lot of energy to break all the bonds between layers of ions.




Ionic Compounds, continuedSalts Are Hard and Brittle, continued• If a force causes one layer to move, ions of the same

charge will be positioned next to each other.

• The cations in one layer are now lined up with other cations in a nearby layer. The anions are also.

• Because like charges are next to each other, they will repel each other and the layers will split apart.

• This is why all salts shatter along a line extending through the crystal known as a cleavage plane.




Ionic Compounds, continuedHow to Identify a Compound as Ionic• All ionic compounds are solid at room temperature. • Tap the substance.

• Ionic compounds do not break apart easily and they fracture into tiny crystals.

• Heat the substance. • Ionic compounds generally have high melting and boiling

points.

• Use a conductivity device to find if the dissolved or melted substance conducts electricity.

• Dissolved and molten ionic compounds conduct electricity.




Salt Crystals

• Despite their differences, the crystals of all salts are made of simple repeating units.

• These repeating units are arranged in a salt to form a crystal lattice, the regular pattern in which a crystal is arranged.

• These repeating patterns within a salt are the reason for the crystal shape that can be seen in most salts.




Salt Crystals, continuedCrystal Structure Depends on the Sizes and Ratios of Ions• Formulas indicate ratios of ions.

• For example, the formula for NaCl indicates there is a 1:1 ratio of sodium cations and chlorine anions.

• Within a NaCl crystal, each Na+ ion is surrounded by six Cl− ions, and each Cl− ion by six Na+ ions.

• Because the edges of the crystal do not have this arrangement, they are locations of weak points.




Salt Crystals, continuedCrystal Structure Depends on the Sizes and Ratios of Ions, continued• The arrangement of cations and anions to form a

crystal lattice depends on the size of the ions and the ratio of cations to anions.

• For example, the salt calcium fluoride has one Ca2+ ion for every two F− ions.

• The cations and anions in calcium fluoride also have a greater difference in size than those in NaCl.




Salt Crystals, continuedCrystal Structure Depends on the Sizes and Ratios of Ions, continued• Because of the size differences of its ions and their

ratio in the salt, the crystal lattice structure of calcium fluoride is different from that of sodium chloride.

• Each calcium ion is surrounded by eight fluoride ions.

• At the same time, each fluoride ion is surrounded by four calcium ions.




Salt Crystals, continuedSalts Have Ordered Packing Arrangements• All salts are made of repeating units. The smallest

repeating unit in a crystal lattice is called a unit cell.

• The ways in which a salt’s unit cells are arranged are determined by X-ray diffraction crystallography.

• X-rays that strike ions in a crystal are deflected, while X-rays that pass through the crystal form a pattern.

• By analyzing this pattern, scientists can calculate the positions that the ions in the salt must have.




Converting Between Mass, Amount, and Number of Particles

Section 1 Avogadro’s Number and Molar ConversionsChapter 7



Understanding Formulas for Polyatomic Ionic Compounds

Section 2 Relative Atomic Mass and Chemical FormulasChapter 7



• Remember that an answer must never be given to more significant figures than is appropriate.

• Round molar masses from the periodic table to two significant figures to the right of the decimal point.

Section 1 Avogadro’s Number and Molar Conversions

Molar Mass Relates Moles to Grams, continuedRemember to Round Consistently

Chapter 7



Calculating Molar Mass for Ionic Compounds

Section 2 Relative Atomic Mass and Chemical FormulasChapter 7

1

Evidence of a Chemical Reaction (Notes on odd pages)

Section 1 Describing Chemical ReactionsChapter 8

Equations and Reaction Information

Section 1 Describing Chemical ReactionsChapter 8

2

Section 3 Classifying Chemical Reactions

Displacement Reactions, continued• The activity series ranks the reactivity of elements

Chapter 8

Visual ConceptsChapter 8Double-Displacement Reaction



Visual Concepts

Reading a Chemical Equation

Chapter 8



Writing a Net Ionic Equation

Chapter 8



Fuel-Oxygen Ratio

Section 3 Stoichiometry and CarsChapter 9



The illustration below shows the parts of an airbag system on an automobile. Use it to answer questions 10–13.

Standardized Test Preparation

Interpreting Graphics

Chapter 9



10.What is the purpose of the igniter in this system?

F. pump air into the air bag

G.prevent any reaction until there is a crash

H.provide energy to start a very fast reaction that produces a gas.

I. provide energy to expand air that is stored in the bag, inflating it like a hot air balloon



Chapter 9



10.What is the purpose of the igniter in this system?

F. pump air into the air bag

G.prevent any reaction until there is a crash

H.provide energy to start a very fast reaction that produces a gas.

I. provide energy to expand air that is stored in the bag, inflating it like a hot air balloon


Chapter 9 Standardized Test Preparation



11.Why does the designer of the air bag need to understand the stoichiometry of the reaction that produces the gas?





11.Why does the designer of the air bag need to understand the stoichiometry of the reaction that produces the gas?

Answer: The stoichiometry of the reaction is needed in order to calculate the amount of gas produced. If the wrong amount of reactant is used, the air bag may over-inflate or under-inflate, making it ineffective.



Chapter 9



12. Which of the following is a reason why most automobile manufacturers have replaced NaN3 with other compounds as the reactants for filling air bags?

A. The sodium produced by the reaction is dangerous.

B. The nitrogen produced by the reaction can be harmful.

C. The materials used to make sodium azide are rare and expensive.

D. The decomposition of sodium azide is too fast so it fills the air bags too quickly.





12. Which of the following is a reason why most automobile manufacturers have replaced NaN3 with other compounds as the reactants for filling air bags?

A. The sodium produced by the reaction is dangerous.

B. The nitrogen produced by the reaction can be harmful.

C. The materials used to make sodium azide are rare and expensive.

D. The decomposition of sodium azide is too fast so it fills the air bags too quickly.



Chapter 9



13. If the reaction that fills the air bag is the decomposition of sodium azide, represented by the equation, 2NaN3(s) 2Na(s) + 3N2(g), how many moles of products are produced by the decomposition of 3.0 moles of sodium azide?



Chapter 9



13. If the reaction that fills the air bag is the decomposition of sodium azide, represented by the equation, 2NaN3(s) 2Na(s) + 3N2(g), how many moles of products are produced by the decomposition of 3.0 moles of sodium azide?

Answer: 7.5 mol



Chapter 9

PASTE IN HANDOUTS FOR NOTEBOOK Name____________________________________

Page 1 of 3

Formation of a Covalent Bond


Page 2 of 3

Predicting Bond Character from Electronegativity Differences

• LEWIS ELECTRON DOT DIAGRAM: As you go from element to element across a period, you add a dot to each side of the element’s symbol.


Page 3 of 3

Lewis Electron Dot diagrams: You do not begin to pair dots until all four sides of the element’s symbol have a dot.

ELECTRON DOT DIAGRAM EXAMPLES

Change each bonding pair to a long dash. Place brackets around the ion and a 2− charge outside the bracket to show that the charge is spread out over the entire ion:

8A1A

2A

3B 4B 5B 6B 7B 8B 11B 12B

3A 4A 5A 6A 7APeriodic Table of the Elements

Los Alamos National Laboratory Chemistry Division

11

1

3 4

12

19 20 21 22 23 24 25 26 27 28 29 30

37 38 39 40 41 42 43 44 45 46 47 48

55 56 57

58 59 60

72 73 74 75 76 77 78 79 80

87 88 89

90 91 92 93 94 95 96

104 105 106 107 108 109 110 111 112

61 62 63 64 65 66 67

97 98 99

68 69 70 71

100 101 102 103

114 116 118

31

13 14 15 16 17 18

32 33 34 35 36

49 50 51 52 53 54

81 82 83 84 85 86

5 6 7 8 9 10

2H

Li

Na

K

Rb

Cs

Fr

Be

Mg

Ca

Sr

Ba

Ra

Sc Ti V Cr Mn Fe Co Ni Cu Zn

Y Zr Nb Mo Tc Ru Rh Pd Ag Cd

La* Hf Ta W Re Os Ir Pt Au Hg

Ac~ Rf Db Sg Bh Hs Mt Ds Uuu Uub Uuq Uuh Uuo

B C N O F

Al Si P S Cl

Ga Ge As Se Br

In Sn Sb Te I

Tl Pb Bi Po At

He

Ne

Ar

Kr

Xe

Rn

39.10

85.47

132.9

(223)

9.012

24.31

40.08

87.62

137.3

(226)

44.96

88.91

138.9

(227)

47.88

91.22

178.5

(257) (260) (263) (262) (265) (266) (271) (272) (277) (296) (298) (?)

50.94

92.91

180.9

52.00

95.94

183.9

54.94

(98)

186.2

55.85

101.1

190.2 190.2

102.9

58.93 58.69

106.4

195.1 197.0

107.9

63.55 65.39

112.4

200.5

10.81

26.98

12.01

28.09

14.01

69.72 72.58

114.8 118.7

204.4 207.2

30.97

74.92

121.8

208.9 (209) (210) (222)

16.00 19.00 20.18

4.003

32.07 35.45 39.95

78.96 79.90 83.80

127.6 126.9 131.3

140.1 140.9 144.2 (147) (150.4) 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0

232.0 (231) (238) (237) (242) (243) (247) (247) (249) (254) (253) (256) (254) (257)

Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

hydrogen

barium

francium radium

strontium

sodium

vanadium

berylliumlithium

magnesium

potassium calcium

rubidium

cesium

helium

boron carbon nitrogen oxygen fluorine neon

aluminum silicon phosphorus sulfur chlorine argon

scandium titanium chromium manganese iron cobalt nickel copper zinc gallium germanium arsenic selenium bromine krypton

yttrium zirconium niobium molybdenum technetium ruthenium rhodium palladium silver cadmium indium tin antimony tellurium iodine xenon

lanthanum hafnium

cerium praseodymium neodymium promethium samarium europium gadolinium terbium dysprosium holmium erbium thulium ytterbium lutetium

tantalum tungsten rhenium osmium iridium platinum gold mercury thallium lead bismuth polonium astatine radon

actinium

thorium protactinium uranium neptunium plutonium americium curium berkelium californium einsteinium fermium mendelevium nobelium lawrencium

rutherfordium dubnium seaborgium bohrium hassium meitnerium darmstadtium

1.008

6.941

22.99

Lanthanide Series*

Actinide Series~

1s1

[Ar]4s23d104p3[Ar]4s23d3[Ar]4s13d10

[Ne]3s23p6[Ne]3s23p4

[Ar]4s1[Ar]4s23d10

1s2

[He]2s1 [He]2s2

[Ar]4s23d7

[Ne]3s23p5

[He]2s22p1[He]2s22p2 [He]2s22p3

[Ar]4s23d5

[He]2s22p4 [He]2s22p5 [He]2s22p6

[Ar]4s23d104p5

[Ne]3s1 [Ne]3s23p1 [Ne]3s23p3[Ne]3s23p2

[Rn]7s25f146d2

[Ne]3s2

[Ar]4s2 [Ar]4s23d1 [Ar]4s23d2 [Ar]4s13d5[Ar]4s23d6

[Ar]4s23d8 [Ar]4s23d104p1 [Ar]4s23d104p2 [Ar]4s23d104p4 [Ar]4s23d104p6

[Kr]5s1 [Kr]5s2 [Kr]5s24d1 [Kr]5s24d2 [Kr]5s14d4 [Kr]5s14d5 [Kr]5s24d5 [Kr]5s14d7 [Kr]5s14d8 [Kr]4d10 [Kr]5s14d10 [Kr]5s24d10 [Kr]5s24d105p1 [Kr]5s24d105p2 [Kr]5s24d105p3 [Kr]5s24d105p4 [Kr]5s24d105p5 [Kr]5s24d105p6

[Xe]6s1 [Xe]6s2 [Xe]6s25d1

[Xe]6s24f15d1 [Xe]6s24f3 [Xe]6s24f4 [Xe]6s24f5 [Xe]6s24f6 [Xe]6s24f7 [Xe]6s24f75d1 [Xe]6s24f9 [Xe]6s24f10 [Xe]6s24f11 [Xe]6s24f12 [Xe]6s24f13 [Xe]6s24f14 [Xe]6s24f145d1

[Xe]6s24f145d2 [Xe]6s24f145d3 [Xe]6s24f145d4 [Xe]6s24f145d5 [Xe]6s24f145d6 [Xe]6s24f145d7 [Xe]6s14f145d9[Xe]6s14f145d10

[Xe]6s24f145d10 [Xe]6s24f145d106p1 [Xe]6s24f145d106p2 [Xe]6s24f145d106p3 [Xe]6s24f145d106p4 [Xe]6s24f145d106p5[Xe]6s24f145d106p6

[Rn]7s1 [Rn]7s2 [Rn]7s26d1

[Rn]7s26d2 [Rn]7s25f26d1 [Rn]7s25f36d1 [Rn]7s25f46d1 [Rn]7s25f6 [Rn]7s25f7 [Rn]7s25f76d1 [Rn]7s25f9 [Rn]7s25f10 [Rn]7s25f11 [Rn]7s25f12 [Rn]7s25f13 [Rn]7s25f14 [Rn]7s25f146d1

[Rn]7s25f146d3 [Rn]7s25f146d4 [Rn]7s25f146d5 [Rn]7s25f146d6 [Rn]7s25f146d7 [Rn}7s15f146d9

o] = 1 2 •[ne] = 1 2 - simeon ca

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