Atomic Electron Configurations and Chemical Periodicity Goals: 1.Understand the role magnetism plays...

Atomic Electron Atomic Electron Configurations and Configurations and

Chemical PeriodicityChemical Periodicity

Goals:

1. Understand the role magnetism plays in determining and revealing atomic structure.

2. Understand effective nuclear charge and its role in determining atomic properties.

3. Write the electron configuration for elements and monoatomic ions.

4. Understand the fundamental physical properties of the elements and their periodic trends.

Arrangement of ElectronsArrangement of Electronsin Atomsin Atoms

Electrons in atoms are arranged as:Electrons in atoms are arranged as:• Shells (n)

• Subshells (l)• Orbitals (ml)

Electrons have _____.• ms, __________________ quantum

number, = +1/2 and -1/2Complete description of electrons

requires _______ quantum numbers.

Electron Spin Magnetic Electron Spin Magnetic Quantum NumberQuantum Number

Electron Spin and Electron Spin and MagnetismMagnetism

• ____________: NOT attracted to a magnetic field

• ___________: substance is attracted to a magnetic field.

• Substances with unpaired electronsunpaired electrons are ______________.

Electron Spin and Electron Spin and MagnetismMagnetism

• H atoms, each has a single electron, they are paramagnetic – when an external magnetic field is applied, the electron magnets align with the field.

• He atoms, with two electrons, are diamagnetic.– We assumed opposite spin

orientations – spins are __________.

The Pauli Exclusion PrincipleThe Pauli Exclusion Principle

No two electrons in an atom can have the same set of four quantum numbers.

Therefore,

Each orbital can be assigned no more than ____ electrons!

Orbital Box DiagramsOrbital Box Diagrams

When n = 1, then l = 0• this shell has a single orbital (1s) to which

2e- can be assigned.

H (1e) n=1, l =0, ml =0ms = +1/2

He (2e) n=1, l=0, ml = 0, ms = +1/2

n=1, l=0, ml = 0, ms = -1/2

1s

1s

Electrons in AtomsElectrons in Atoms

A 2p electron can be designated by A 2p electron can be designated by which set of quantum numbers?which set of quantum numbers?

n l ml ms

a. 1 0 0 +1/2b. 2 1 0 +1/2c. 2 2 +1 -1/2d. 3 1 +2 +1/2e. 3 2 +1 +1/2

Students should be familiar with the values and meaning of quantum

numbers.

Electrons in AtomsElectrons in Atoms

• Electrons generally assigned to Electrons generally assigned to

orbitals of successively higher energy.orbitals of successively higher energy.

• For For H atomsH atoms, E = - C(1/n, E = - C(1/n22). E depends ). E depends

only on _____.only on _____.

• For For many-electron atomsmany-electron atoms, energy , energy

depends on both ____ and depends on both ____ and __________..

Assigning Electrons to Assigning Electrons to SubshellsSubshells

1 e- atom

• In many-electron atom:In many-electron atom:a) subshells increase in a) subshells increase in

energy as value of energy as value of _______increases._______increases.

b) for subshells of same _____, b) for subshells of same _____, subshell with lower n is subshell with lower n is lower in energy.lower in energy.

Effective Nuclear Charge, Z*Effective Nuclear Charge, Z*

• Z* - the nuclear charge experienced by a particular electron in a multielectron atom, as modified by the presence of the other electrons.

• Li has 3 p (+) and 3 e (-)2 e in 1 s orbital ; 1 e in 2 s orbitale- in 2s should “see” a +1 charge,but it sees 1.28• C has 6 p (+) and 6 e (-)2 e- in 1s ; 2 e- in 2s ; 2 e- in 2pe- in 2s should “see” +3, but see

3.22e- in 2p should “see” a +2 charge,

but see 3.14

Electron cloud for 1s electrons


• Z* is _______ for s electrons than for p electrons.– s electrons always have a

lower energy than p electrons in the same quantum shell.

• The Z* _________ across a period.

• The 2s electron PENETRATES the region occupied by the 1s electron. – 2s electron experiences a

_________ positive charge than expected.

Atomic Electron ConfigurationsAtomic Electron Configurations

• The arrangements of __________ in the elements in the ground state.

• In general, electrons are assigned to orbitals in order of increasing ________.

• Electron configuration can be given with the orbital box diagram, or with the spdf notation.

11s

value of n

label of l

no. ofelectrons

spdf notation

for H, atomic number = 1for H, atomic number = 1

Orbital Box notation

1s

Electron ConfigurationsElectron Configurations

• The outermost electrons of an element are assigned to the indicated orbitals.

LithiumLithiumGroup 1AGroup 1A

Atomic number = 3Atomic number = 3

________ ---> 3 total electrons________ ---> 3 total electrons

1s

2s

3s3p

2p

BoronBoronGroup 3AGroup 3A


___________ ---> ___________ --->

5 total electrons5 total electrons

1s

2s

3s3p

2p

CarbonCarbon

Here we see for the first Here we see for the first time time ___________ RULE:___________ RULE: When placing electrons in When placing electrons in a set of orbitals having a set of orbitals having the same energy, we the same energy, we place them singly as long place them singly as long as possible.as possible.

Group 4AGroup 4A


___________ ---> ___________ --->


1s

2s

3s3p

2p

Hund’s RuleHund’s Rule

• The most stable arrangement of electrons is that with the __________ _____________________, all with the same ________ direction.

• This arrangement makes the total energy of an atom as low as possible.

Nitrogen and OxygenNitrogen and OxygenGroup 5AGroup 5A


_________---> _________--->


1s

2s

3s3p

2p

Group 6AGroup 6A

1s

2s

3s3p

2p

1s

2s

3s3p

2p

Fluorine and NeonFluorine and Neon

• Note that we have Note that we have reached the end of the reached the end of the 2nd period, and the 2nd period, and the 2nd shell is ________!2nd shell is ________!

Group 7AGroup 7A


____________---> ____________--->


Group 8AGroup 8A


_____________ ---> _____________ --->


1s

2s

3s3p

2p

Sodium and Potassium –Sodium and Potassium –Noble gas notationNoble gas notation

Na - Group 1ANa - Group 1AAtomic number = 11Atomic number = 111s1s2 2 2s2s2 2 2p2p6 6 3s3s11 or “neon core” + 3sor “neon core” + 3s11

[Ne] 3s[Ne] 3s1 1 (uses rare gas notation)(uses rare gas notation)Note that we have begun a new period (3Note that we have begun a new period (3rdrd ) )

All Group 1A elements haveAll Group 1A elements have[core]ns[core]ns11 configurations, n=period configurations, n=period

numbernumber

K – Atomic number = 19, 4K – Atomic number = 19, 4thth period period______________________

PhosphorusPhosphorus

All Group 5A elements have [core ] ns2 np3

configurations where n is the period number.

Group 5AGroup 5A


spdf: ______________spdf: ______________

short: __________short: __________

1s

2s

3s3p

2p

Transition MetalsTransition Metals

3d orbitals used for Sc-Zn (Table 8.4)3d orbitals used for Sc-Zn (Table 8.4)

and so are d-block elements.and so are d-block elements.

Transition MetalsTransition MetalsAll 4th period elements have the configuration

[argon] nsx (n - 1)dy and so are d-block elements.

Copper

Iron

Chromium

26e-

Lantanides and ActinidesLantanides and Actinides

4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 8.2)4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 8.2)

and so are f-block elements.and so are f-block elements.

Lantanides and ActinidesLantanides and ActinidesAll these elements have the configuration

[core] nsx (n - 1)dy (n - 2)fz and so are f-block elements.

CeriumCerium[Xe] 6s[Xe] 6s22 5d 5d11 4f 4f11

UraniumUranium[Rn] 7s[Rn] 7s22 6d 6d11 5f 5f33

Ion ConfigurationsIon Configurations

To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)].

P ---> P3+ [Ne] 3s2 3p3 - 3e- [Ne] 3s2 3p0

1s

2s

3s3p

2p

1s

2s

3s3p

2p

Ion ConfigurationsIon Configurations

For transition metals, remove ns electrons and then (n - 1) electrons.

Fe [Ar] 4s2 3d6 loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6

To form cations, always remove electrons of highest n value first!

3d4s

Fe3+

4s 3d 3d4s

Fe Fe2+

PracticePractice• Which substance will be paramagnetic?V+5 or Fe+3

Students should be familiar with writing electronic configurations and identifying diamagnetic vs.

paramagnetic materials.

Ion Size Makes a BIG Ion Size Makes a BIG DifferenceDifference

• About 20% of the CO2 binds to hemoglobin and is released in the lungs. About 70% is converted by Carbonic Anhydrase into HCO3- ion, which remains in the blood plasma until the reverse reaction releases CO2 into the lungs.

• Carbonic Anhydrase catalyzes the reversible hydration of CO2 to form bicarbonate anion and a proton:CO2 + H2O <==> HCO3- + H+

• Toxic metals like Cd2+ replace Zn2+

inactivating the enzyme.

PERIODIC PERIODIC TRENDSTRENDS

PERIODIC PERIODIC TRENDSTRENDS

Periodic TrendsPeriodic Trends

• Atomic and ionic sizeAtomic and ionic size• Ionization energyIonization energy• Electron affinityElectron affinity

Higher effective nuclear chargeElectrons held ______ tightly

Larger orbitals.Electrons held ____tightly.

Atomic sizeAtomic size

• Size goes ____ on going down a group. See Figure 8.9.

• Because electrons are added further from the nucleus, there is ______ attraction.

• Size goes _______ on going across a period.


• Atom Z* Experienced by Electrons in Valence Orbitals

(Outermost) • Li +1.28• Be -------• B +2.58• C +3.22• N +3.85• O +4.49• F +5.13

Increase in Increase in Z* across a Z* across a periodperiod

[Values calculated using Slater’s Rules][Values calculated using Slater’s Rules]

Atomic size- Transition Atomic size- Transition MetalsMetals

• 3d subshell is inside the 4s subshell.• 4s electrons feel a more or less constant Z*.• Sizes stay about the same and chemistries are

similar!

Ion SizesIon Sizes

• CATIONS are __________ than the CATIONS are __________ than the atoms from which they come.atoms from which they come.

• The electron/proton attraction has The electron/proton attraction has gone ______ and so size ____________.gone ______ and so size ____________.

Li,152 pm3e and 3p

Li+, 78 pm2e and 3 p

+Forming Forming a cation.a cation.Forming Forming a cation.a cation.

Ion SizesIon Sizes

• ANIONS are __________ than the ANIONS are __________ than the atoms from which they come.atoms from which they come.

• The electron/proton attraction has The electron/proton attraction has gone ______and so size __________.gone ______and so size __________.

• Trends in ion sizes are the same as Trends in ion sizes are the same as atom sizes. atom sizes.

Forming Forming an anion.an anion.Forming Forming an anion.an anion.F, 71 pm

9e and 9pF-, 133 pm10 e and 9 p

-

Ion SizesIon Sizes

Redox ReactionsRedox Reactions

Why do metals lose Why do metals lose

electrons in their electrons in their

reactions? reactions?

Why does Mg form Why does Mg form

MgMg2+2+ ions and not ions and not

MgMg3+3+??

Why do nonmetals Why do nonmetals

take on electrons?take on electrons?

Why do metals lose Why do metals lose

electrons in their electrons in their

reactions? reactions?

Why does Mg form Why does Mg form

MgMg2+2+ ions and not ions and not

MgMg3+3+??

Why do nonmetals Why do nonmetals

take on electrons?take on electrons?

Ionization EnergyIonization Energy

IE = energy required to remove an electron from an atom in the gas phase.

Mg (g) + 738 kJ ---> Mg+ (g) + e-


Mg (g) + 738 kJ ---> MgMg (g) + 738 kJ ---> Mg++ (g) + e- (g) + e-

Mg+ (g) + 1451 kJ ---> Mg2+ (g) + e-

Mg+ has 12 protons and only 11 electrons. Therefore, IE for Mg+ > Mg.

Ionization EnergyIonization EnergyMg (g) + 735 kJ ---> Mg+ (g) + e-

Mg+ (g) + 1451 kJ ---> Mg2+ (g) + e-

Mg2+ (g) + ---> Mg3+ (g) + e-

Energy cost is very high to dip into a shell of lower n. This is why oxidation number = Group number.

7733 kJ

Effective nuclear chargeEffective nuclear charge

• Sodium

11+

10-

1-

A valance electron in an atom is attracted to the nucleus of the atom and it is repelled by the other electrons in the atom: inner e- shield or screen the outer electrons from attraction of the nucleus.Effect = 11-10 = +1

core

valance

Effective nuclear chargeEffective nuclear charge

Ra

dia

l ele

ctro

n d

ensi

ty [Ne] core

3s

For the 3s e- (valance e- of Na) there is a probability of being found close to the nucleus – there is a probability of experiencing a greater attraction than suggested. Zeff = +2.5Electrons in 3s orbitals has a higher Zeff than 3p orbitals: subshells energy trend is: ns < np < nd

Atomic sizeAtomic size

0

50

100

150

200

250

0 5 10 15 20 25 30 35 40

Li

Na

K

Kr

He

NeAr

2nd period

3rd period 1st transitionseries

Radius (pm)

Atomic Number

0

50

100

150

200

250

0 5 10 15 20 25 30 35 40

Li

Na

K

Kr

He

NeAr

2nd period

3rd period 1st transitionseries

Radius (pm)

Atomic Number

Size decreases across a period owing to increase in Z*. Each added electron feels a greater and greater + charge.

LargeLarge SmallSmall

Increase in Z*Increase in Z*

Ionization Ionization EnergyEnergy

IE ___________ across a period and ___________ down a group.


As Z* increases, orbital energies “drop” and IE increases.

Trends in Ionization EnergyTrends in Ionization Energy• IE increases across a period because Z* IE increases across a period because Z*

increases.increases.– Metals lose electrons more easily than

nonmetals.– Metals are good reducing agents.– Nonmetals lose electrons with difficulty.

• IE decreases down a group .IE decreases down a group .– Because size increases.– Reducing ability generally increases down the

periodic table (easier to give e-).– See reactions of Li, Na, K

Electron AffinityElectron Affinity

A few elements GAIN electrons to form anions.

Electron affinity is the energy change that occurs when _______________ to ___________________________.

A(g) + e- ---> A-(g) Electron Affinity = ∆E


Affinity for electron _________ across a period (EA becomes more negative).

Affinity __________ down a group (EA becomes less negative).

SummarySummary

PracticePractice

• Which of the following elements has the greater difference between its first and second ionization energies: C, Li, N, Be?

• Which should be smaller: the sulfide ion, S2-, or a sulfur atom, S?

Students should be familiar with periodic trends – IE, EA, Atomic size.

RememberRemember

• Go over all the contents of your textbook.

• Practice with examples and with problems at the end of the chapter.

• Practice with OWL tutor.• Work on your assignment for Chapter

8.• Practice with the quiz on CD of

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