Chapter 8Chapter 8Periodic Periodic Properties Properties of the of the ElementsElements

Electron Spin ExperimentElectron Spin Experiment

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Electron SpinElectron Spinexperiments by Stern and Gerlach showed a

beam of silver atoms is split in two by a magnetic field

the experiment reveals that the electrons spin on their axis

as they spin, they generate a magnetic field◦ spinning charged particles generate a

magnetic field if there is an even number of electrons, about

half the atoms will have a net magnetic field pointing “North” and the other half will have a net magnetic field pointing “South”

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Spin Quantum Number, Spin Quantum Number, mmssspin quantum number describes how the

electron spins on its axis◦ clockwise or counterclockwise◦ spin up or spin down

spins must cancel in an orbital◦paired

ms can have values of ±½

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Pauli Exclusion PrinciplePauli Exclusion Principle no two electrons in an atom may have the same set of 4

quantum numbers therefore no orbital may have more than 2 electrons,

and they must have with opposite spins knowing the number orbitals in a sublevel allows us to

determine the maximum number of electrons in the sublevels sublevel has 1 orbital, therefore it can hold 2

electronsp sublevel has 3 orbitals, therefore it can hold 6

electronsd sublevel has 5 orbitals, therefore it can hold 10

electronsf sublevel has 7 orbitals, therefore it can hold 14

electrons

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Quantum Numbers of Quantum Numbers of Helium’s ElectronsHelium’s Electronshelium has two electrons both electrons are in the first energy levelboth electrons are in the s orbital of the first

energy levelsince they are in the same orbital, they must

have opposite spins

n l ml ms

first

electron1 0 0 +½

second

electron1 0 0 -½

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Electron Configurations the ground state of the electron is the lowest

energy orbital it can occupy the distribution of electrons into the various orbitals

in an atom in its ground state is called its electron configuration

the number designates the principal energy level the letter designates the sublevel and type of orbital the superscript designates the number of electrons

in that sublevel He = 1s2

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Orbital DiagramsOrbital Diagrams

we often represent an orbital as a square and the electrons in that orbital as arrows◦the direction of the arrow represents the

spin of the electron

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orbital with1 electron

unoccupiedorbital

orbital with2 electrons

Sublevel Splitting in Sublevel Splitting in Multielectron AtomsMultielectron Atoms the sublevels in each

principal energy level of Hydrogen all have the same energy – we call orbitals with the same energy degenerate◦ or other single

electron systems for multielectron atoms,

the energies of the sublevels are split◦ caused by electron-

electron repulsion the lower the value of

the l quantum number, the less energy the sublevel has◦ s (l = 0) < p (l = 1) <

d (l = 2) < f (l = 3)

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Penetrating and ShieldingPenetrating and Shielding the radial distribution function

shows that the 2s orbital penetrates more deeply into the 1s orbital than does the 2p

the weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force, they are more shielded from the attractive force of the nucleus

the deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively

the result is that the electrons in the 2s sublevel are lower in energy than the electrons in the 2p

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En

erg

y

1s

7s

2s

2p

3s

3p3d

6s6p

6d

4s

4p4d

4f

5s

5p

5d5f

Notice the following:1. because of penetration, sublevels

within an energy level are not degenerate

2. penetration of the 4th and higher energy levels is so strong that their s sublevel is lower in energy than the d sublevel of the previous energy level

3. the energy difference between levels becomes smaller for higher energy levels

Filling the Orbitals with Filling the Orbitals with ElectronsElectronsenergy shells fill from lowest energy to highsubshells fill from lowest energy to high◦ s → p → d → f◦Aufbau Principle

orbitals that are in the same subshell have the same energy

no more than 2 electrons per orbital◦Pauli Exclusion Principle

when filling orbitals that have the same energy, place one electron in each before completing pairs◦Hund’s Rule

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Electron Configurations of Electron Configurations of Multielectron AtomsMultielectron Atoms

n = 1

s orbital (l = 0)

1 electronH: 1s1

1s2

n = 1

s orbital (l = 0)

2 electronsHe:

n = 2

s orbital (l = 0)

1 electrons1s2 2s1Li:

Lowest energy to highest energy

Valence ElectronsValence Electrons the electrons in all the

subshells with the highest principal energy shell are called the valence electrons

electrons in lower energy shells are called core electrons

chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons

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ExamplesExamplesFor the following atom, write:◦ the Ground State Electron Configuration◦Use short hand notation to write orbital

Diagram◦Determine the core electrons and valence

electrons Carbon Sulfur Potassium

Electron configuration of Electron configuration of transition metal and atoms in transition metal and atoms in higher energy statehigher energy state

For the following atom, write:◦ the Ground State Electron Configuration◦Use short hand notation to write orbital

Diagram◦Determine the core electrons and valence

electrons Cr Br Bi

Trend in Atomic Radius – Main Trend in Atomic Radius – Main GroupGroup

Different methods for measuring the radius of an atom, and they give slightly different trends◦ van der Waals radius = nonbonding◦ covalent radius = bonding radius◦ atomic radius is an average radius of an

atom based on measuring large numbers of elements and compounds

Atomic Radius Increases down group◦ valence shell farther from nucleus◦ effective nuclear charge fairly close

Atomic Radius Decreases across period (left to right)◦ adding electrons to same valence shell◦ effective nuclear charge increases◦ valence shell held closer

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Effective Nuclear ChargeEffective Nuclear Charge in a multi-electron system, electrons are

simultaneously attracted to the nucleus and repelled by each other

outer electrons are shielded from full strength of nucleus◦ screening effect

effective nuclear charge is net positive charge that is attracting a particular electron

Z is nuclear charge, S is electrons in lower energy levels◦ electrons in same energy level contribute to

screening, but very little◦ effective nuclear charge on sublevels trend, s > p >

d > f

Zeffective = Z - S

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Screening & Effective Screening & Effective Nuclear ChargeNuclear Charge

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Trends in Atomic Radius Transition Trends in Atomic Radius Transition MetalsMetals increase in size down the Group atomic radii of transition metals roughly the same size

across the d block◦ must less difference than across main group elements◦ valence shell ns2, not the d electrons◦ effective nuclear charge on the ns2 electrons

approximately the same

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Example – Choose the Larger Example – Choose the Larger Atom in Each Pair Atom in Each Pair

1) N or F, 1) N or F

2) C or Ge

3) N or Al

4) Al or Ge?

1) N or F

2) C or Ge,

1) N or F

2) C or Ge

3) N or Al,

Ionization Energyminimum energy needed to remove an

electron from an atom ◦gas state◦ endothermic process◦ valence electron easiest to remove

◦M(g) + IE1 M1+(g) + 1 e-

◦M+1(g) + IE2 M2+(g) + 1 e-

first ionization energy = energy to remove electron from neutral atom; 2nd IE = energy to remove from +1 ion; etc.

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General Trends in 1General Trends in 1stst Ionization Ionization EnergyEnergy

larger the effective nuclear charge on the electron, the more energy it takes to remove it

the farther the most probable distance the electron is from the nucleus, the less energy it takes to remove it

1st IE decreases down the group◦ valence electron farther from nucleus

1st IE generally increases across the period◦ effective nuclear charge increases

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Trends in Ionic RadiusTrends in Ionic Radius Ions in same group have same charge Ion size increases down the group

◦ higher valence shell, larger Cations smaller than neutral atom; Anions bigger than

neutral atom Cations smaller than anions

◦ except Rb+1 & Cs+1 bigger or same size as F-1 and O-2 Larger positive charge = smaller cation

◦ for isoelectronic species◦ isoelectronic = same electron configuration

Larger negative charge = larger anion ◦ for isoelectronic series

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Electron Configuration of Electron Configuration of Cations in their Ground Cations in their Ground StateState cations form when the atom loses electrons from the

valence shell for transition metals electrons, may be removed from

the sublevel closest to the valence shell

Al atom = 1s22s22p63s23p1

Al+3 ion = 1s22s22p6

Fe atom = 1s22s22p63s23p64s23d6

Fe+2 ion = 1s22s22p63s23p63d6

Fe+3 ion = 1s22s22p63s23p63d5

Cu atom = 1s22s22p63s23p64s13d10

Cu+1 ion = 1s22s22p63s23p63d10

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Example 8.8 – Choose the Atom in Each Example 8.8 – Choose the Atom in Each Pair with the Higher First Ionization Pair with the Higher First Ionization Energy Energy

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1) Al or S, 1) Al or S

2) As or Sb,

1) Al or S

2) As or Sb

3) N or Si,

1) Al or S

2) As or Sb

3) N or Si

4) O or Cl?

Irregularities in the TrendIrregularities in the Trend Ionization Energy generally increases from

left to right across a Periodexcept from 2A to 3A, 5A to 6A

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Be

1s 2s 2p

B

1s 2s 2p

N

1s 2s 2p

O

1s 2s 2p

Which is easier to remove an electron from B or Be? Why?Which is easier to remove an electron from N or O? Why?

Irregularities in the Irregularities in the First Ionization Energy First Ionization Energy TrendsTrends

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Be

1s 2s 2p

B

1s 2s 2p

Be+

1s 2s 2p

To ionize Be you must break up a full sublevel, cost extra energy

B+

1s 2s 2p

When you ionize B you get a full sublevel, costs less energy

Irregularities in the Irregularities in the First Ionization Energy First Ionization Energy TrendsTrends

Tro, Chemistry: A Molecular Approach 31

To ionize N you must break up a half-full sublevel, cost extra energy

N+

1s 2s 2p

O

1s 2s 2p

N

1s 2s 2p

O+

1s 2s 2p

When you ionize O you get a half-full sublevel, costs less energy

Trends in Successive Ionization Trends in Successive Ionization EnergiesEnergies removal of each successive

electron costs more energy◦ shrinkage in size due to

having more protons than electrons

◦ outer electrons closer to the nucleus, therefore harder to remove

regular increase in energy for each successive valence electron

large increase in energy when start removing core electrons

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Trends in Electron AffinityTrends in Electron Affinity energy released when an neutral atom gains an

electron◦ gas state◦ M(g) + 1e- M-1(g) + EA

defined as exothermic (-), but may actually be endothermic (+)◦ alkali earth metals & noble gases endothermic, WHY?

more energy released (more -); the larger the EA generally increases across period

◦ becomes more negative from left to right◦ not absolute◦ lowest EA in period = alkali earth metal or noble gas◦ highest EA in period = halogen

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Metallic CharacterMetallic Character Metals

◦ malleable & ductile◦ shiny, lusterous, reflect light◦ conduct heat and electricity◦ most oxides basic and ionic◦ form cations in solution◦ lose electrons in reactions – oxidized

Nonmetals◦ brittle in solid state◦ dull◦ electrical and thermal insulators◦ most oxides are acidic and molecular◦ form anions and polyatomic anions◦ gain electrons in reactions – reduced

metallic character increases left metallic character increase down

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Example – Choose the Example – Choose the More Metallic Element in Each Pair More Metallic Element in Each Pair

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1) Sn or Te

2) P or Sb

3) Ge or In

4) S or Br?

Trends in the Alkali MetalsTrends in the Alkali Metalsatomic radius increases down the column ionization energy decreases down the columnvery low ionization energies◦good reducing agents, easy to oxidize◦ very reactive, not found uncombined in

nature◦ react with nonmetals to form salts◦ compounds generally soluble in water

found in seawaterelectron affinity decreases down the columnmelting point decreases down the column◦ all very low MP for metals

density increases down the column◦ except K◦ in general, the increase in mass is greater

than the increase in volume

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Trends in the HalogensTrends in the Halogensatomic radius increases down the column ionization energy decreases down the columnvery high electron affinities◦good oxidizing agents, easy to reduce◦ very reactive, not found uncombined in

nature◦ react with metals to form salts◦ compounds generally soluble in water

found in seawaterreactivity increases down the columnreact with hydrogen to form HX, acidsmelting point and boiling point increases

down the columndensity increases down the column◦ in general, the increase in mass is greater

than the increase in volume38

Trends in the Noble GasesTrends in the Noble Gasesatomic radius increases down the column ionization energy decreases down the column◦ very high IE

very unreactive◦ only found uncombined in nature◦used as “inert” atmosphere when reactions

with other gases would be undersirablemelting point and boiling point increases

down the column◦ all gases at room temperature◦ very low boiling points

density increases down the column◦ in general, the increase in mass is greater

than the increase in volume

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Example– Write a balanced Example– Write a balanced chemical reaction for the following.chemical reaction for the following.

reaction between potassium metal and bromine gas

K(s) + Br2(g)

(ionic compounds are all solids at room temperature)

reaction between rubidium metal and liquid water

Rb(s) + H2O(l)

reaction between chlorine gas and solid iodineCl2(g) + I2(s)

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Magnetic properties Magnetic properties electron configurations that result in unpaired

electrons mean that the atom or ion will have a net magnetic field – this is called paramagnetism◦ will be attracted to a magnetic field

electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field – this is called diamagnetism◦ slightly repelled by a magnetic field

ExamplesExamplesWrite the Electron Configuration and

Determine whether the following atoms or their ions are Paramagnetic or Diamagnetic

Al and Al3+

O and O-2

Ag and Ag+