1 unit 6: chapters 11-12. pages 295-366 atomic electron configurations and periodicity

Unit 6: Chapters 11-12. Pages 295-366 Unit 6: Chapters 11-12. Pages 295-366 ATOMIC ELECTRON ATOMIC ELECTRON

CONFIGURATIONS AND PERIODICITYCONFIGURATIONS AND PERIODICITY

• First model of the electron behavior• Vital to understanding the atom• Does not work for atoms with

more than 1 electron

• First model of the electron behavior• Vital to understanding the atom• Does not work for atoms with

more than 1 electron

Bohr Model

Collision of Ideas

Dalton

Thompson

Rutherford

Newton

Maxwell

Einstein

De Broglie

Matter

The Photoelectric EffectDuality of Light

• Wave behavior

• Particle behavior

de Broglie’s Novel NotionLight was “known” (thought) to be a wave, but

Einstein showed that it also acts particle-like

Electrons were particles with known mass & charge

What if ……

electrons behaved as waves also

Evidence for de Broglie’s Notion

Diffraction pattern obtained with firing a beam of electrons through a crystal.

This can only be explained if the electron behaves as a wave!

Nobel Prize for de Broglie in 1929

• Extremely small mass

• Located outside the nucleus

• Moving at very high speeds

• Have specific energy levels

• Standing wave behavior

• Extremely small mass

• Located outside the nucleus

• Moving at very high speeds

• Have specific energy levels

• Standing wave behavior

Electron Characteristics

A baseball behaves as a particle and follows a predictable path.

An electron behaves as a wave, and its path cannot be predicted.

All we can do is to calculate the probability of the electron following a specific path.

Baseball vs Electron

What if a baseball behaved like an electron?

Characteristic wavelength• baseball 10-34 m• electron 0.1 nm

All we can predict is…..

Werner Heisenberg(1901-1976)

• Proposed that the dual nature of the electron places limitation on how precisely we can know both the exact location and speed of the electron

• Instead, we can only describe electron behavior in terms of probability.

The Uncertainty Principle

speedspeed

positionposition

Erwin Schrodinger (1887-1961)

• In 1926, Austrian physicist, proposed an equation that incorporates both the wave and particle behavior of the electron

• When applied to hydrogen’s 1 electron atom, solutions provide the most probable location of finding the electron in the first energy level

• Can be applied to more complex atoms too!

Wave Equation & Wave Mechanics

Solutions to Schrodinger’s Wave Equation

Gives the most probable location of electron in 3-D space around nucleus (probability map)

- most probable location called an

orbital

- orbitals can hold a maximum of 2 e-

“Most Successful Theory of 20th Century”

Dalton

Thompson

Rutherford

Newton

Maxwell

Einstein

De Broglie

Matter

Schrödinger

Heisenberg

WaveMechanics

Quantum

Mechanics

Quantum Mechanics ModelDescribes the arrangement and space

occupied by electrons in atoms

Quantum

Mechanics

Electron’s energy is quantized

Mathematics of waves to define orbitals(wave mechanics)

Bohr Model v. Quantum Mechanics

Energy

Electron

Position/Path

Bohr Q. Mech.

Dartboard Analogy

Suppose the size of the probability distribution is defined

as where there is a % chance of all hits being confined.

The electron's movement cannot be known precisely.

We can only map the probability of finding the electron at various locations outside the nucleus.

The probability map is called an orbital.

The orbital is calculated to confine 90% of electron’s range.

Quantum Mechanics Model

Arrangement of Electrons in Atoms

Electrons in atoms are arranged asElectrons in atoms are arranged as

SHELLS (n) = distance from nucleus SHELLS (n) = distance from nucleus

1, 2, 3, …1, 2, 3, …

SUBSHELLS (l) = shape of region of probabilitySUBSHELLS (l) = shape of region of probability

s, p, d, fs, p, d, f

ORBITALS (mORBITALS (mll) = orientation in space) = orientation in space

• There is a relationship between the quantum number (n) and its the number of subshells.

Principal quantum number (n) = number of subshells

Arrangement of Electrons in Atoms

Representing s Orbitals

The 2s orbital is similar to the

1s orbital, but larger in size.

”Larger” means that the

highest probability for

finding the electron lies

farther out from the nucleus.

Each can hold a maximum of

electrons.

Comparison of 1s and 2s Orbitals

Probability Maps of the Three 2p Orbitals

The 2p orbital is in the n = energy level.

There are 2p orbitals oriented in three directions.

Each orbital can hold a maximum of electrons.

The maximum number of electrons in the 2p sublevel is .

Adding all 2p orbitals would result in a sphere.

The five 3d orbitals are generally oriented in different directions.

Adding all five orbitals, would result in a sphere.

The five orbitals, taken together, make up the d subshell of the n = 3 shell.

Each orbital can hold a maximum of two electrons.

This sublevel has a maximum of electrons.

Probability Maps of the Five 3d Orbitals

Probability Maps of 7 f Orbitals Probability Maps of 7 f Orbitals

Each orbital can be assigned no Each orbital can be assigned no

more than 2 electrons! And more than 2 electrons! And

each electron spins in opposite each electron spins in opposite

directions. directions.

Arrangement of Electrons in AtomsArrangement of Electrons in AtomsElectron Spin Quantum Number- mElectron Spin Quantum Number- mss

Electron Spin Quantum Electron Spin Quantum NumberNumber

DiamagneticDiamagnetic: NOT attracted to a magnetic : NOT attracted to a magnetic fieldfieldParamagneticParamagnetic: substance is attracted to a : substance is attracted to a magnetic field. Substance has magnetic field. Substance has unpaired unpaired electronselectrons..

n ---> shelln ---> shell 1, 2, 3, 4, ...1, 2, 3, 4, ...

l ---> sublevell ---> sublevel s, p, d, fs, p, d, f

mmll ---> orbital ---> orbital -l ... 0 ... +l-l ... 0 ... +l

mmss ---> electron spin ---> electron spin +1/2 +1/2

and -1/2and -1/2

n ---> shelln ---> shell 1, 2, 3, 4, ...1, 2, 3, 4, ...

l ---> sublevell ---> sublevel s, p, d, fs, p, d, f

mmll ---> orbital ---> orbital -l ... 0 ... +l-l ... 0 ... +l

mmss ---> electron spin ---> electron spin +1/2 +1/2

and -1/2and -1/2

4 QUANTUM4 QUANTUMNUMBERSNUMBERS

Summary:

Pauli Exclusion Principle- Pauli Exclusion Principle- No No two electrons in the same atom can two electrons in the same atom can

have the same set of 4 quantum have the same set of 4 quantum numbers.numbers.

Pauli Exclusion Principle- Pauli Exclusion Principle- No No two electrons in the same atom can two electrons in the same atom can

have the same set of 4 quantum have the same set of 4 quantum numbers.numbers.

Determine the quantum numbers for Determine the quantum numbers for the outer two valence electrons in the outer two valence electrons in the lithium atom.the lithium atom.

Aufbau Principle-Electrons fill open Aufbau Principle-Electrons fill open lower energy levels sequentiallylower energy levels sequentially

lower energy to higher energy lower energy to higher energy

Writing Electron Writing Electron ConfigurationsConfigurations

value of nvalue of l

no. ofelectrons

spdf notationfor H, atomic number = 1

Two ways of Two ways of writing configs. writing configs. One is called One is called thethe spdf spdf notation.notation.

Broad Periodic Table Classifications

• Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O)

• Transition Elements: filling d orbitals (Fe, Co, Ni)• Lanthanide and Actinide Series (inner transition elements):

filling 4f and 5f orbitals (Eu, Am, Es)

Writing Orbital NotationsWriting Orbital NotationsWriting Orbital NotationsWriting Orbital Notations

Two ways of Two ways of writing writing configs. Other configs. Other is called theis called the orbital box orbital box notation.notation.

Arrowsdepictelectronspin

ORBITAL BOX NOTATIONfor He, atomic number = 2

Arrowsdepictelectronspin

ORBITAL BOX NOTATIONfor He, atomic number = 2

One electron has n = 1, l = 0, mOne electron has n = 1, l = 0, m ll = 0, m = 0, mss = + 1/2 = + 1/2

Other electron has n = 1, l = 0, mOther electron has n = 1, l = 0, m ll = 0, m = 0, mss = - 1/2 = - 1/2

Different subshells within the same principal shell have different energies.

The more complex the subshell, the higher its energy. This explains why the 3d subshell is higher in energy than the 4s subshell.

Energy ordering of orbitals for multi-electron atoms

Rules for Filling Orbitals

Bottom-up (Aufbau’s principle)

Fill orbitals singly before doubling up (Hund’s Rule)

Paired electrons have opposite spin (Pauli exclusion principle)

CobaltSymbol

Atomic Number

Full Configuration

Valence Configuration

Shorthand Configuration

Orbital diagram and electron configuration for a ground

state lithium atom

state carbon atom

Hund’s Rule- electrons in the same sublevel will spread out into their own orbital before doubling up.

Silicon's valence electronsSilicon's valence electrons

Selenium's valence electronsSelenium's valence electrons

Core electrons and valence electrons in germanium

Outer electron configuration for the elements

The periodic table gives the electron configuration for AsThe periodic table gives the

electron configuration for As

Valence Electrons by GroupValence Electrons by Group

Ion charges by groupIon charges by group

Periodic LawPeriodic Law

All the elements in a group have the same electron configuration in their outermost shells

Example: Group 2Be 2, 2

Mg 2, 8, 2Ca 2, 2, 8, 2

QuestionQuestion

Specify if each pair has chemical properties that are similar (1) or not similar (2):

A. Cl and Br

B. P and S

C. O and S

General Periodic General Periodic TrendsTrends

1. Atomic and ionic size1. Atomic and ionic size 2. Electron affinity2. Electron affinity

3. Ionization energy 3. Ionization energy 4. Metallic Character 4. Metallic Character

Higher effective nuclear chargeElectrons held more tightly

Larger orbitals.Electrons held lesstightly.

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

• Z* is the nuclear charge experienced by the Z* is the nuclear charge experienced by the outermost electrons.outermost electrons. Screen 8.6. Screen 8.6.

• Explains why E(2s) < E(2p)Explains why E(2s) < E(2p)• Z* increases across a period owing to Z* increases across a period owing to

incomplete shielding by inner electrons.incomplete shielding by inner electrons.• Estimate Z* by --> [ Estimate Z* by --> [ Z - (no. inner electrons) Z - (no. inner electrons) ]]• Z = number of electronsZ = number of electrons• Charge felt by 2s e- in Li Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Z* = 3 - 2 = 1• Be Be Z* = 4 - 2 = 2Z* = 4 - 2 = 2• B B Z* = 5 - 2 = 3Z* = 5 - 2 = 3 and so on!and so on!

Effective Effective Nuclear Nuclear ChargeCharge

Electron cloud for 1s electrons

Figure 8.6

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

• Atom Z* Experienced by Electrons in Valence Orbitals

• 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

Lithium Beryllium

Sodium

Atomic Atomic SizeSize

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

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

•Size goes DOWNSize goes DOWN on going on going across a period.across a period.

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

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

•Size goes DOWNSize goes DOWN on going on going across a period.across a period.

Atomic RadiiAtomic RadiiAtomic RadiiAtomic Radii Figure 8.9Figure 8.9

Trends in Atomic SizeTrends in Atomic SizeSee Figures 8.9 & 8.10See Figures 8.9 & 8.10

0 5 10 15 20 25 30 35 40

2nd period

3rd period 1st transitionseries

Radius (pm)

Atomic Number

0 5 10 15 20 25 30 35 40

2nd period

3rd period 1st transitionseries

Radius (pm)

Atomic Number

Ion SizesIon SizesIon SizesIon Sizes

Li,152 pm3e and 3p

Li+, 60 pm2e and 3 p

+Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?

Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?

• CATIONSCATIONS are are SMALLERSMALLER than the than the atoms from which they come.atoms from which they come.

Li,152 pm3e and 3p

Li +, 78 pm2e and 3 p

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

F,64 pm9e and 9p

F- , 136 pm10 e and 9 p

-Does the size go up or Does the size go up or down when gaining an down when gaining an electron to form an electron to form an anion?anion?

Does the size go up or Does the size go up or down when gaining an down when gaining an electron to form an electron to form an anion?anion?

• ANIONSANIONS are are LARGERLARGER than the than the atoms from which they come.atoms from which they come.

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

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

Trends in Ion SizesTrends in Ion SizesTrends in Ion SizesTrends in Ion Sizes

Figure 8.13Figure 8.13

Ionization EnergyIonization EnergySee Screen 8.12See Screen 8.12

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

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

Mg (g) + 735 kJ ---> MgMg (g) + 735 kJ ---> Mg++ (g) + e- (g) + e-

MgMg+ + (g) + 1451 kJ ---> Mg(g) + 1451 kJ ---> Mg2+2+ (g) + e- (g) + e-

MgMg2+2+ (g) + 7733 kJ ---> Mg (g) + 7733 kJ ---> Mg3+3+ (g) + e- (g) + e-

Energy cost is very high to dip into a Energy cost is very high to dip into a shell of lower n. shell of lower n. This is why ox. no. = Group no.This is why ox. no. = Group no.

Ionization EnergyIonization EnergySee Screen 8.12See Screen 8.12

Trends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization EnergyTrends in Ionization Energy

1 3 5 7 9 11 13 15 17 19 21 23 25 27 29 31 33 350

1st Ionization energy (kJ/mol)

Atomic NumberH Li Na K

Trends in Ionization Trends in Ionization EnergyEnergy

• IE increases across a IE increases across a period because Z* period because Z* increases.increases.

• Metals lose electrons more Metals lose electrons more easily than nonmetals.easily than nonmetals.

• Metals are good reducing Metals are good reducing agents.agents.

• Nonmetals lose electrons Nonmetals lose electrons with difficulty.with difficulty.

Trends in Ionization Trends in Ionization EnergyEnergy

• IE decreases down a group IE decreases down a group • Because size increases.Because size increases.

• Reducing ability generally Reducing ability generally increases down the increases down the periodic table. periodic table.

• See reactions of Li, Na, KSee reactions of Li, Na, K

Electronegativity• A measure of the ability of an atom

that is bonded to another atom to attract electrons to itself.

Electron AffinityElectron AffinityElectron AffinityElectron Affinity

A few elements A few elements GAINGAIN electrons electrons to form to form anionsanions..

Electron affinity is the energy Electron affinity is the energy involved when an atom gains involved when an atom gains an electron to form an anion.an electron to form an anion.

A(g) + e- ---> AA(g) + e- ---> A--(g) (g)

E.A. = ∆EE.A. = ∆E

Electron Affinity of OxygenElectron Affinity of OxygenElectron Affinity of OxygenElectron Affinity of Oxygen

∆∆E is E is EXOEXOthermic thermic because O has because O has an affinity for an an affinity for an e-.e-.

[He] O atom

EA = - 141 kJ

+ electron

O [He] - ion

• See Figure 8.12 and See Figure 8.12 and Appendix FAppendix F

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

• Affinity decreases down Affinity decreases down a group (EA becomes a group (EA becomes less negative).less negative).

Atom EAAtom EAFF -328 kJ-328 kJClCl -349 kJ-349 kJBrBr -325 kJ-325 kJII -295 kJ-295 kJ

Trends in Electron AffinityTrends in Electron AffinityTrends in Electron AffinityTrends in Electron Affinity

Metallic character trends in the periodic table

Metallic CharacterMetallic Character

The text links metallic character to the tendency to lose electrons in chemical reactions, and nonmetallic character to the tendency to gain electrons in chemical reactions. The metallic character trends therefore follow the ionization energy trends

The metallic character trends explain the location of metals,

metalloids, and nonmetals

The metallic character trends explain the location of metals,

metalloids, and nonmetals

Which is the more metallic element, Sn or Te?

Which is the more metallic element, Si or Sn?

1 unit 6: chapters 11-12. pages 295-366 atomic electron configurations and periodicity

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