Chapter 8 Notes

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8Chemistry: A Molecular Approach by Nivaldo J. TroChemistry: A Molecular Approach by Nivaldo J. TroCHEM 1411 General ChemistryCHEM 1411 General Chemistry

Mr. Kevin A. BoudreauxAngelo State Universitywww.angelo.edu/faculty/kboudrea

Mr. Kevin A. BoudreauxAngelo State Universitywww.angelo.edu/faculty/kboudrea

Periodic Properties Periodic Properties of the Elementsof the Elements

Chapter Objectives:• Learn how to write electron configurations for

neutral atoms and ions.• Learn how to predict trends in atomic radius,

effective nuclear charge, ionic radius, ionization energy, electron affinity, and metallic character from the positions of the elements on the periodic table.

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Periodic PropertiesPeriodic Properties• A number of element and ion properties vary in a

periodic, predictable way across the periodic table. These include:– atomic radius– effective nuclear charge– ionic radius– ionization energy– electron affinity– metallic character– electronegativity (Ch. 9)

Chapter 8 Notes

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The DevelopmentThe Developmentof theof the

Periodic TablePeriodic Table

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Periodic PropertiesPeriodic Properties• It has long been known that many of the elements

have similar chemical properties.– Lithium, sodium, and potassium all perform the

same reaction with water, 2M(s) + 2HOH(l) → 2MOH(aq) + H2(g)

the only difference being the masses of the metals themselves and the vigor and speed of the reaction.

Lithiumslow

Sodiumfast

Potassiumwarp speed

Chapter 8 Notes

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The Invention of the Periodic TableThe Invention of the Periodic Table• In 1869 Dimitri Mendeleev published a table in

which the elements that were known at the time were arranged by increasing atomic mass, and grouped into columns according to their chemical properties. The properties of the elements varied (more or less) in a periodic way in this arrangement.

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MendeleevMendeleev’’s Periodic Tables Periodic Table Atomic Sodium Weight Chlorides Salts H 1 HCl Li 7 LiCl Be 9.4 BeCl2 B 11 BCl3 C 12 CCl4 N 14 Na3N O 16 Na2O F 19 NaF Na 23 NaCl Mg 24 MgCl2 Al 27.3 AlCl3 Si 28 SiCl4 P 31 Na3P S 32 Na2S Cl 35.5 NaCl K 39 KCl Ca 40 CaCl2 As 75 Na3As Se 78 Na2Se Br 80 NaBr

• Mendeleev noticed that when he grouped the elements by their properties, there were some “holes”which he guessed corresponded to as-yet-unknown elements.

• Mendeleev predicted some of the properties for two of these, eka-aluminum (?=68), and eka-silicon (?=72), which corresponded well to gallium (Ga, discovered in 1875) and germanium (Ge, 1886)

Chapter 8 Notes

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The Periodic Table by Atomic NumberThe Periodic Table by Atomic Number• To make the properties of the elements “line up”

properly, it was sometimes necessary to exchange the order of the elements.– For instance, potassium (39.0983 g/mol) is slightly lighter than

argon (39.948 g/mol), so by increasing atomic weight, potassium should be in Group 8A, and argon in Group 1A, but that clearly doesn’t fit their observed properties.

• After the discovery of the nucleus and the proton, and with the development of X-ray spectroscopy, it was discovered that the periodic table could be written in order of increasing atomic number, with no need to “play around” with the order of the elements. It was also possible to count protons, and see exactly how many “missing” elements there were.

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The Modern Periodic Table of the ElementsThe Modern Periodic Table of the Elements

Chapter 8 Notes

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Electron Electron ConfigurationsConfigurations

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Energy Levels of OneEnergy Levels of One--Electron AtomsElectron Atoms• In one-electron atoms (H, He+, etc.) the energy of

the orbital depends only on n. The 2s and 2porbitals have the same energy, the 3s, 3p, and 3dorbitals all have the same energy, etc. (Orbitals having the same energy are said to be degenerate.)

Chapter 8 Notes

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Energy Levels of MultiEnergy Levels of Multi--Electron AtomsElectron Atoms• In multi-electron atoms the orbital energy depends

on n and l (s<p<d<f), so there are differences in energy between the subshells.

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Energy Levels of MultiEnergy Levels of Multi--Electron AtomsElectron Atoms• There is some crossover of energies from one shell

to another. (A 3d orbital in some atoms is higher in energy than a 4s orbital.)

• These energy differences result from a balance of attractive forces between electrons and the nucleus, and repulsive forces between electrons.

• Adding more electrons increases electron-electron repulsions.

• Having more protons in the nucleus lowers the energy of close-lying electrons.

• Orbitals which have a greater electron density near the nucleus experience more of the nuclear charge, and are lowered in energy because of their greater penetration of the inner atomic core.

Chapter 8 Notes

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Energy Levels of MultiEnergy Levels of Multi--Electron AtomsElectron Atoms• Repulsion of the outer-shell electrons by ones in

lower-lying shells cause the outer-shell electrons to be pushed father away from the nucleus, and to be held less tightly. Part of their attraction for the nucleus is canceled, and we say that these electrons are shielded from the nucleus by the inner electrons.

Figure 8.4

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Putting Everything TogetherPutting Everything Together• Knowing the relative energies of the

various orbitals allows us to predict the electron configuration for any atom or ion — a list of how the electrons in that atom are distributed among its orbitals.

• Each successive electron added to an atom occupies the lowest-energy orbital available, resulting in the ground-state configuration of the atom.

• A set of rules called the aufbauprinciple (German, building-up) guides the order in which orbitals are filled.

Chapter 8 Notes

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The The AufbauAufbau PrinciplePrinciple• Lower-energy orbitals fill before higher-energy

orbitals. The order in which the orbitals are filled is:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f

• An orbital can hold only two electrons, which must have opposite spins (Pauli exclusion principle).

• Hund’s Rule: If two or more degenerate orbitals (orbital at the same energy level) are available, one electron goes into each orbital (spin up) until all are half-full; only then does a second electron fill one of the orbitals. (This is a consequence of the mutual repulsion between like charged-electrons.)

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Writing Electron ConfigurationsWriting Electron Configurations• To write an element’s electron configuration, write

in the orbitals that are occupied by electrons, followed by a superscript to indicate how many electrons are in the set of orbitals (e.g., H 1s1)

• Another way to show the placement of electrons is an orbital diagram, in which each orbital is represented by a circle (or a line, or a square), and the electrons as arrows pointing up (↑) or down (↓) (indicating the electron spin).

Chapter 8 Notes

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Examples: Electron ConfigurationsExamples: Electron Configurations1. Using the diagram on the next page, write the

electron configuration and orbital diagram for the following elements, and state whether they are diamagnetic or paramagnetic.

Li 1s22s1

Be 1s22s2

B 1s22s22p1

CNOFNe

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Chapter 8 Notes

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Abbreviated Electron ConfigurationsAbbreviated Electron Configurations• For atoms following neon, the full electron

configuration can be extremely cumbersome to write:Ba: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 6s2

• Since everything up to the 5p6 is exactly the same electron configuration as the noble gas xenon (Xe), this configuration can be abbreviated as:

Ba: [Xe] 6s2

• Abbreviated electron configurations are always based on the nearest preceding noble gas.

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Electron Configurations and the Periodic TableElectron Configurations and the Periodic Table• It is not necessary to memorize the orbital order in

the aufbau scheme, because this sequence can be read from the periodic table in the following fashion:

Chapter 8 Notes

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Electron Configurations and the Periodic TableElectron Configurations and the Periodic Table

(n)s (n)p(n-1)d

(n-2)f

n

1s1s

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Examples: Electron ConfigurationsExamples: Electron Configurations2. For the following elements, provide the abbreviated

electron configuration, orbital diagram, and state whether the element is diamagnetic or paramagnetic.

NaMgAlClArFeSnPbI

Chapter 8 Notes

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Valence ElectronsValence Electrons• Valence electrons are the electrons in the outermost

shell (with the highest value of n). All the elements in the same group on the periodic table have similar electron configurations for their valence shells, and therefore have similar chemical properties.

• Core electrons are those in complete principal energy levels, and do not participate in bonding.

Fig.8.6

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Periodic Table of Valence Shell ConfigurationsPeriodic Table of Valence Shell Configurations

Figure 8.7

Chapter 8 Notes

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Some Anomalous Electron ConfigurationsSome Anomalous Electron Configurations• Half-filled and filled subshells have special stability,

leading to some unexpected electron configurations:– Cr [Ar]3d44s2 actual: [Ar]3d54s1

– Cu [Ar]3d94s2 actual: [Ar]3d104s1

– Ag [Kr]4d95s2 actual: [Kr]4d105s1

– Au [Xe]4f145d96s2 actual: [Xe]4f145d106s1

• Most of the anomalous electron configurations occur in elements with atomic numbers greater than Z = 40, where the energy differences between subshells are very small. In all of these cases, the transfer of an electron from one subshell to another lowers the total energy of the atom because of a decrease in electron-electron repulsion.

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The Explanatory Power of the QM ModelThe Explanatory Power of the QM Model• The chemical properties of elements are largely

determined by the number of valence electrons they contain.– The Group 8A elements are inert because their

valence shells are already full.– The Group 1A elements (ns1) can attain a noble

gas configuration by losing their single valence electrons, forming 1+ charges.

– The Group 2A elements (ns2) lose their two valence electrons to form 2+ charges.

– The Group 7A elements (ns2np5) gain one electron to complete their valence shells.

Chapter 8 Notes

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Atomic RadiusAtomic Radiusandand

Effective Nuclear Effective Nuclear Charge Charge

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Atomic RadiusAtomic Radius• The atomic radius of an atom can be measured

either by taking one-half of the distance between atoms in an atomic solid (the nonbonding atomic radius or van der Waals radius) or by taking one half of the distance between two identical bonded nuclei or between atoms in a metallic crystal (the bonding atomic radius or covalent radius).

Chapter 8 Notes

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Trends in Atomic RadiiTrends in Atomic Radii• An average atomic radius for a particular element

can be determined from the measurement of a large number of elements and compounds.

• When the atomic radius is graphed against the atomic number (Figures 8.9 and 8.10), there is a periodic pattern to the radii across a group:– the radius is largest for the Group 1A elements,

and becomes smaller as we move towards the Group 8A elements.

– for elements in a particular group, the size increases from top to bottom within the group.

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Trends in Atomic RadiiTrends in Atomic Radii

Figure 8.9

Chapter 8 Notes

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Atomic Radius: Top to BottomAtomic Radius: Top to Bottom• As we add larger valence shells (larger values of n),

the size of the atom increases. Therefore, atomic size increases as we move from top to bottom in a group.

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Effective Nuclear ChargeEffective Nuclear Charge• The net nuclear charge actually felt by an outer-shell

electron, called the effective nuclear charge (Zeff), is often substantially lower than the actual nuclear charge Z, because the core electrons “shield” the outer electrons from the full effect of the nuclear charge:

Zeff = Zactual - Electron shielding

Figure 8.11

Chapter 8 Notes

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Atomic Radius: Left to RightAtomic Radius: Left to Right• As we move to the right in a period, there are more

protons in the nucleus, but no increase in shielding (the number of inner electrons isn’t changing, and valence electrons don’t shield each other). Zeffincreases from left to right across a period.

• An increasing Zeff means there is a stronger attraction between the nucleus and the valence electrons. Therefore, atomic radius decreases from left to right in a period.

nucleusinner electronsvalence electrons

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Trends in Atomic RadiiTrends in Atomic Radii

Figure 8.10

Chapter 8 Notes

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Effective Nuclear Charge and Atomic RadiusEffective Nuclear Charge and Atomic Radius

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Examples: Atomic Radii and the Size of AtomsExamples: Atomic Radii and the Size of Atoms3. Which atom is larger, C or N?

4. Which atom is larger, Si or Ge?

5. Which atom is larger, Mg or Ba?

6. Which atom is larger, Al or Ge?

7. Arrange the following elements in order of increasing size: Cs, Rb, Sr, Ca.

Chapter 8 Notes

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Properties of IonsProperties of Ions

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The Valence ShellThe Valence Shell• The electrons in the outermost or valence shell (the

one that has the highest value of n) have the highest energies, and are on average the farthest away from the nucleus, and so are the ones which are most exposed to other atoms.– For main-group metals, the electrons lost in

forming cations are taken from the valence shell.– For main-group nonmetals, the electrons gained

in forming anions are added to the valence shell.• The inner, core electrons (inner shell) do not usually

play a role in chemical bonding.

Chapter 8 Notes

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The Formation of Sodium ChlorideThe Formation of Sodium Chloride• When sodium reacts with chlorine, the 3s1 electron

in sodium is lost, producing an sodium ion :Na 1s22s22p63s1 → Na+ 1s22s22p6 + e-

– The next shell down (n=2) is now the outermost shell, which is full; there is little tendency to lose more electrons.

– The Na+ ion has the same configuration as the noble gas neon, with 8 e-’s in the valence shell.

• The chlorine atom accepts an electron into its 3psubshell, filling the n=3 shell:

Cl 1s22s22p63s23p5 + e- → Cl- 1s22s22p63s23p6

– The Cl- ion the same electron configuration as the noble gas argon, with 8 e-’s in the valence shell.

MOV: Formation of Sodium Chloride

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The Octet RuleThe Octet Rule• The Na+ and Cl- ion’s electron configurations are the

same as that of the nearest noble gas (the ions are said to be isoelectronic with the nearest noble gas). Atoms “prefer” to have a filled outermost shell because this is more electronically stable.

• This can be generalized into the octet rule: Main-group elements tend to undergo reactions that leave them with eight outer-shell electrons.– That is, main-group elements react so that they

attain a noble gas e- configuration with filled sand p sublevels in their valence electron shell.

• There are many exceptions to the octet rule, but it is useful for making predictions about some chemical bonds.

Chapter 8 Notes

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Group 1AGroup 1A--3A Cations3A Cations• Elements with similar outer shell configurations

form similar ions.• For instance, the alkali metals form ions with a +1

charge; the valence s1 electron is the one that’s lost:1A Li 1s22s1 Li+ 1s2

1A Na 1s22s22p63s1 Na+ 1s22s22p6

1A K 1s22s22p63s23p64s1 K+ 1s22s22p63s23p6

• The Group 2A and 3A metals also tend to lose all of their valence electrons to form cations.2A Be 1s22s2 Be2+ 1s2

2A Mg 1s22s22p63s2 Mg2+ 1s22s22p6

3A Al 1s22s22p63s23p1 Al3+ 1s22s22p6

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Group 4A and 5A CationsGroup 4A and 5A Cations• The Group IV and V metals can lose either the

electrons from the p subshell, or from both the s and p subshells, thus attaining a pseudo-noble gas configuration.

4A Sn [Kr]4d105s25p2

Sn2+ [Kr]4d105s2

Sn4+ [Kr]4d10

4A Pb [Xe]4f145d106s26p2

Pb2+ [Xe]4f145d106s2

Pb4+ [Xe]4f145d10

5A Bi [Xe]4f145d106s26p3

Bi3+ [Xe]4f145d106s2

Bi5+ [Xe]4f145d10

Chapter 8 Notes

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Transition Metal CationsTransition Metal Cations• Transition metals lose their ns electrons before the

(n-1)d electrons, usually forming 2+ charges. Some can also lose electrons from their highest d levels.

Fe [Ar] 3d64s2

Fe2+ [Ar] 3d6

Fe3+ [Ar] 3d5

——————————————————————Zn [Ar] 3d10 4s2

Zn2+ [Ar] 3d10

——————————————————————Ag [Kr] 4d10 5s1

Ag+ [Kr] 4d10

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Nonmetal Anions and Noble GasesNonmetal Anions and Noble Gases• The Group 6A – 8A non-metals gain electrons until

their valence shells are full (8 electrons).4A C 1s22s22p2 C4- 1s22s22p6

5A N 1s22s22p3 N3- 1s22s22p6

6A O 1s22s22p4 O2- 1s22s22p6

7A F 1s22s22p5 F- 1s22s22p6

• The Group 8A noble gases already possess a full outer shell, so they have no tendency to form ions.

8A Ne 1s22s22p6

8A Ar 1s22s22p63s23p6

Chapter 8 Notes

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Examples: Predicting Ion ConfigurationsExamples: Predicting Ion Configurations8. Predict the ground-state electron configuration for

each of the following ions, and state whether they are diamagnetic or paramagnetic. (sim. to Ex. 8.6)a. Ra2+

b. La2+ , La3+

c. Mn2+, Mn3+

d. Ni2+

e. Ti4+

f. N3-

g. Te2-

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Examples: Predicting Ion ConfigurationsExamples: Predicting Ion Configurations9. What 2+ ion has the ground-state electron

configuration 1s22s22p63s23p63d10 ?

Chapter 8 Notes

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Ionic RadiiIonic Radii• Cations are smaller than their parent atoms, since

electrons are being removed from the valence shell.• Anions are larger than their parent atoms, since the

number of electron-electron repulsions increases when electrons are added.

• The greater the ionic charge, the smaller the ionic radius (e.g., Fe3+ < Fe2+).

• Trends in ion size are the same as for neutral atoms:– Ionic size increases down a group.– Ionic size decreases for cations across a period.– Ionic size decreases for anions across a period.

MOV: Gain and Loss of Electrons

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←Figure8.12

Figure→8.13

Chapter 8 Notes

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Examples: Predicting Periodic TrendsExamples: Predicting Periodic Trends10. Arrange the following ions in order of increasing

size: O2-, P3-, S2-, As3-.

11. Which species is larger, Na or Na+?

12. Which species is larger, Br or Br-?

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Examples: Predicting Periodic TrendsExamples: Predicting Periodic Trends13. Which species is larger, Sn2+ or Sn4+?

14. Which species is larger, Rb+ or Br-?

15. Arrange the following in order of increasing size: Ca2+, S2-, Ar, K+, Cl-.

Chapter 8 Notes

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Periodic Properties: Periodic Properties: Ionization Energy,Ionization Energy,Electron Affinity, Electron Affinity,

andandMetallic CharacterMetallic Character

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Ionization EnergyIonization Energy• Ionization Energy (IE) — the amount of energy

needed to remove an electron from a mole of neutral atoms in the gas phase:M(g) → M+(g) + e-(g); ΔH° = IE (always > 0)

• Ionization energy is a positive energy value (an endothermic process); energy must be provided in order to remove an electron from an atom.

• The energy required to remove the first electron is the first ionization energy (IE1).

MOV: Ionization Energies

Chapter 8 Notes

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Trends in First Ionization EnergyTrends in First Ionization Energy• As we move down a group, atomic size increases, so

the outermost electrons become easier to remove, since they are farther from the nucleus. Therefore, ionization energy decreases as we move down a group.

• As we move left-to-right across a period, Zeffincreases and atomic size decreases. As a result, the outer electrons are held more tightly, and are harder to remove. Therefore, ionization energy increasesas we move from left to right across a period.

• Atoms with low ionization energies (metals) tend to form cations; those with high ionization energies (nonmetals) tend to form anions.

MOV: Periodic Trends Ionization Energies

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Ionization Energy vs. Atomic NumberIonization Energy vs. Atomic Number

Figure 8.14

Chapter 8 Notes

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Trends in Ionization EnergyTrends in Ionization Energy

Figure 8.15

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Exceptions to Trends in First Ionization EnergyExceptions to Trends in First Ionization Energy• There are some variations in the general trends; the

variations result from the ionization of atoms which have either one electron in a p subshell (B, Al), or are one electron away from being half-filled (O, S).

Chapter 8 Notes

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Higher Ionization EnergiesHigher Ionization Energies• Of course, it is possible to form more than just 1+

cations:M(g) → M+(g) + e-(g) first ionization energy (IE1)M+(g) → M2+(g) + e-(g) second ionization energy (IE2)M2+(g) → M3+(g) + e-(g) third ionization energy (IE3)

• More energy is required for each successive ionization step (IE3 > IE2 > IE1) because it is harder to remove an electron from a positively charged ion.

• Large jumps in ionization energy occur when we are trying to remove electrons from filled shells.

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Higher Ionization EnergiesHigher Ionization Energies• Why is the third ionization

energy of Be so much higher than the second and first?

Be 1s2 2s2

Be+ 1s2 2s1

Be2+ 1s2

Be3+ 1s1

IE1

IE2

IE3

Chapter 8 Notes

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Higher Ionization EnergiesHigher Ionization Energies

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Higher Ionization EnergiesHigher Ionization Energies

Chapter 8 Notes

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Electron AffinityElectron Affinity• Electron Affinity (Eea) — the energy change

accompanying the addition of electrons to 1 mole of atoms in the gas phase:M(g) + e-(g) → M-(g); ΔH° = Eea (usually < 0)

• Electron affinity is usually a negative energy value (an exothermic process); energy is usually released when an electron is added to a neutral atom.

• The more negative the electron affinity, the greater the tendency of the atom to accept an electron and the more stable the resulting anion will be.

• In general, electron affinity decreases as we move down a group, and increases from left to right, but the trend is not as smooth as for ionization energy and size.

MOV: Periodic Trends Electron Affinity

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Electron AffinityElectron Affinity

Figure 8.16

Chapter 8 Notes

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Trends in Metallic CharacterTrends in Metallic Character• Metals are good conductors of heat and electricity,

they are malleable and ductile, are often shiny, and easily lose electrons in chemical reactions.

• Nonmetals are poor conductors of heat and electricity, their physical states vary from solid to gas, and they tend to gain electrons in chemical reactions.

• Moving from left to right across a period, ionization energy increases and electron affinity becomes more negative. Therefore, metallic character decreases as we move from left to right across a period.

• As we move down a group, ionization energy decreases, making electrons more likely to be lost. Therefore, metallic character increases as we move down a group.

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Trends in Metallic CharacterTrends in Metallic Character

Figure 8.17

Chapter 8 Notes

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Trends in Metallic CharacterTrends in Metallic Character

Figure 8.18

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Summary of the Periodic TrendsSummary of the Periodic Trends

Chapter 8 Notes

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Examples: Predicting Periodic TrendsExamples: Predicting Periodic Trends16. Which atom has the larger IE: Al or S?

17. Which atom has the larger IE: As or Sb?

18. Which atom has the larger IE: N or Si

19. Which atom has the larger IE: O or Cl?

20. Which atom has the largest IE: Na, Sr, Be, Rb?

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Examples: Predicting Periodic TrendsExamples: Predicting Periodic Trends21. Arrange the following in order of increasing IE:

F, S, Cl.

22. Which has the larger electron affinity, O or F?

23. Which element is more metallic: Sn or Te?

24. Which element is more metallic: Ge or In?