Atomic and Molecular Structure Michael Abosch Brian Pflaum 2 nd Period.

Atomic and Molecular Atomic and Molecular StructureStructure

Michael AboschMichael Abosch

Brian Pflaum Brian Pflaum

22ndnd Period Period

Unit OutlineUnit Outline

Atomic and Electronic Structure, Atomic and Electronic Structure, and Quantum Mechanicsand Quantum Mechanics

Periodic TrendsPeriodic Trends Molecular StructureMolecular Structure Bonding TheoryBonding Theory

The Wave Nature of The Wave Nature of LightLight

Electromagnetic Radiation- all visible light, Electromagnetic Radiation- all visible light, radio waves, infrared, X-rays etc.radio waves, infrared, X-rays etc.

Electromagnetic Spectrum- shows Electromagnetic Spectrum- shows radiation arranged in order of increasing radiation arranged in order of increasing wavelengthwavelength

Visible light is only a small portion of Visible light is only a small portion of spectrum.spectrum.

http://steve.files.wordpress.com/2006/03/Elcetromagnetic%20spectrum.jpg

The Wave Nature of The Wave Nature of LightLight

ffλλ= c= c (frequency)(wavelength)= Speed of light (frequency)(wavelength)= Speed of light

(2.9979x10(2.9979x1088 m/s) m/s) Frequency measured in sFrequency measured in s-1 -1 (often Hz)(often Hz) Wavelength measured in meters (often Wavelength measured in meters (often

nm,nm,μμm)m)

The Quantization of The Quantization of EnergyEnergy

Quantum=The smallest quantity of Quantum=The smallest quantity of energy that can be emitted or energy that can be emitted or absorbed as electromagnetic absorbed as electromagnetic radiation.radiation.

Energy, E, of a single quantum equals Energy, E, of a single quantum equals a constant times the frequency of a constant times the frequency of radiation.radiation.

E=hfE=hf h=planck’s constant=6.626X10h=planck’s constant=6.626X10--

3434Joule-seconds.Joule-seconds.

Photoelectric EffectPhotoelectric Effect

When photons of sufficiently high When photons of sufficiently high energy (greater than the individual energy (greater than the individual metal’s threshold energy) strike a metal’s threshold energy) strike a metal surface, electrons are emitted metal surface, electrons are emitted from the metalsfrom the metals

Energy of Photon, E=hf (planck’s Energy of Photon, E=hf (planck’s constant)(frequency)constant)(frequency)

Kinetic Energy of ejected electrons: Kinetic Energy of ejected electrons: KEKEee=E=Ephoton-photon-EEthreshold of metalthreshold of metal

Wave Behavior of MatterWave Behavior of Matter

Dual nature of radiant energy: both Dual nature of radiant energy: both particle and wave-like propertiesparticle and wave-like properties

DeBroglie wavelength: DeBroglie wavelength: wavelength=(Planck’s wavelength=(Planck’s constant)/(momentum)=(h)/(mv)constant)/(momentum)=(h)/(mv)

Mass in Kg, Velocity in m/sMass in Kg, Velocity in m/s

OrbitalsOrbitals

An allowed energy state of an An allowed energy state of an electron in the quantum mechanical electron in the quantum mechanical model of the atom; describes the model of the atom; describes the spatial distribution of the electron. spatial distribution of the electron. The orbital is defined by the values The orbital is defined by the values of quantum numers n, l, and mof quantum numers n, l, and mll

The Principal Quantum The Principal Quantum NumberNumber

The principal quantum number, n, The principal quantum number, n, can have positive integral values of can have positive integral values of 1,2,3 etc…As n increases, the orbital 1,2,3 etc…As n increases, the orbital becomes larger, and the electron becomes larger, and the electron spends more time farther from spends more time farther from nucleusnucleus

The Azimuthal Quantum The Azimuthal Quantum NumberNumber

The azimuthal quantum number, l, can The azimuthal quantum number, l, can have integral values from 0 to n-1 for have integral values from 0 to n-1 for each value of n. This quantum number each value of n. This quantum number defines the shape of an orbital. The defines the shape of an orbital. The value of l is generally designated by the value of l is generally designated by the letters s,p,d, and f.letters s,p,d, and f.

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Value of lValue of l 00 11 22 33

Letter Letter UsedUsed

ss pp dd ff

The Magnetic Quantum The Magnetic Quantum NumberNumber

The magnetic quantum number, mThe magnetic quantum number, mll, , can have integral values between -l can have integral values between -l and l. Describes orientation of and l. Describes orientation of orbital in space.orbital in space.

Relationship amongst Relationship amongst Quantum NumbersQuantum Numbers

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Spin Magnetic Quantum Spin Magnetic Quantum Number and Pauli Number and Pauli Exclusion PrincipleExclusion Principle

The Spin Magnetic Quantum The Spin Magnetic Quantum Number, mNumber, mss, has two possible , has two possible values: +1/2, -1/2.values: +1/2, -1/2.

No two electrons in an atom can No two electrons in an atom can have the same set of four quantum have the same set of four quantum numbers n, l, mnumbers n, l, mll, and m, and mss

Thus, an orbital can hold a maximum Thus, an orbital can hold a maximum of two electrons, and they must have of two electrons, and they must have opposite spins.opposite spins.

Electron ConfigurationsElectron Configurations

Electron Configuration=A particular Electron Configuration=A particular arrangement of electrons in the arrangement of electrons in the orbitals of an atom.orbitals of an atom.

The orbitals are filled in order of The orbitals are filled in order of increasing energy, with no more increasing energy, with no more than two electrons per orbital.than two electrons per orbital.

Orbital DiagramsOrbital Diagrams Each orbital is denoted by Each orbital is denoted by

a box, and each electron a box, and each electron by a half arrow (which by a half arrow (which represents spin-up or spin-represents spin-up or spin-down)down)

Electrons having opposite Electrons having opposite spins are said to be paired spins are said to be paired when they are in the same when they are in the same orbitalorbital

An unpaired electron is An unpaired electron is one not accompanied by a one not accompanied by a partner of opposite spin.partner of opposite spin.


Hund’s RuleHund’s Rule Hund’s Rule=For orbitals of the same energy level, Hund’s Rule=For orbitals of the same energy level,

the lowest energy is attained when the number of the lowest energy is attained when the number of electrons with the same spin is maximized.electrons with the same spin is maximized.

Note how in the diagram below, all three p orbitals Note how in the diagram below, all three p orbitals are filled singularly before an electron is pairedare filled singularly before an electron is paired


The Periodic Table and The Periodic Table and Electron Filling OrderElectron Filling Order


Condensed Electron Condensed Electron ConfigurationsConfigurations

The Electron configuration of the most recent nobel gas is represented by its chemical symbol in brackets. From there, Just proceed in the normal filling order until you reach the element.

In Potassium, the previous noble gas is argon, and its remaining Electron occupies just one of the s orbitals, hence why it is denotedAs 4s1


IonsIons

Start by writing the electron Start by writing the electron configuration for the normal elementconfiguration for the normal element

Then remove (or add) electrons as Then remove (or add) electrons as necessary, always taking (or adding) necessary, always taking (or adding) from the highest principle quantum from the highest principle quantum number first (ignoring the filling number first (ignoring the filling order).order).

Fe=[Ar]4sFe=[Ar]4s223d3d66

Fe(II)=[Ar]3dFe(II)=[Ar]3d66

Anomalous Electron Anomalous Electron ConfigurationsConfigurations

Electron configurations of certain Electron configurations of certain elements appear to violate the “rules”elements appear to violate the “rules”

Frequently occurs when there are enough Frequently occurs when there are enough electrons to lead to precisely half-filled electrons to lead to precisely half-filled sets of degenerate (same energy-level) sets of degenerate (same energy-level) orbitals, or to completely fill an orbital. orbitals, or to completely fill an orbital. This conserves EnergyThis conserves Energy

No universal pattern or predictabilityNo universal pattern or predictability Ex: Chromium is [Ar]4sEx: Chromium is [Ar]4s113d3d55 instead of instead of

[Ar]4s[Ar]4s223d3d44

Practice Practice What’s the What’s the

electron electron configuration for configuration for Lead?Lead?

Answer: Answer: [Xe]6s[Xe]6s224f4f14145d5d10106p6p22

Assign Quantum Assign Quantum numbers to it’s last numbers to it’s last filled electron.filled electron.

Answer: n=6, l=1, Answer: n=6, l=1, mmll=0, m=0, mss=+1/2=+1/2

http://www.elementsdatabase.com/

Periodic TrendsPeriodic Trends

Atomic SizeAtomic Size Ionic SizeIonic Size Ionization EnergiesIonization Energies ElectronegativityElectronegativity

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Atomic SizeAtomic Size

Within each group, size increases from top to Within each group, size increases from top to bottom, results primarily from the increase in bottom, results primarily from the increase in principle quantum number of electronsprinciple quantum number of electrons

In each period, atomic radius tends to In each period, atomic radius tends to decrease from left to right. Increase in the decrease from left to right. Increase in the effective nuclear charge as we move across a effective nuclear charge as we move across a row steadily draws valence electrons closer to row steadily draws valence electrons closer to nucleusnucleus

Exceptions: The addition of a paired electron Exceptions: The addition of a paired electron produces increased repulsion that sometimes produces increased repulsion that sometimes leads to an increase in size (Like from a pleads to an increase in size (Like from a p33 to a to a pp44 element.) element.)

Atomic SizeAtomic Size

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Ionic vs. Atomic SizeIonic vs. Atomic Size

Cations: Compared to its neutral atom, Cations: Compared to its neutral atom, cations are smaller because electrons cations are smaller because electrons have vacated the biggest orbitalhave vacated the biggest orbital

Anions: Compared to its neutral atom, Anions: Compared to its neutral atom, anions are larger because adding anions are larger because adding electrons increases repulsions, which electrons increases repulsions, which leads to more space.leads to more space.

Ionic vs. Atomic SizeIonic vs. Atomic Size


Isoelectronic SeriesIsoelectronic Series

Isoelectronic Series=A group all Isoelectronic Series=A group all containing the same number of containing the same number of electrons. As the atomic number electrons. As the atomic number increases, the radius decreases.increases, the radius decreases.

Ex: ClEx: Cl--, Ar, K, Ar, K++

Size: ClSize: Cl-->Ar>K>Ar>K++

Ionization EnergyIonization Energy

Ionization Energy=The minum Ionization Energy=The minum energy required to remove an energy required to remove an electron from the ground state of the electron from the ground state of the isolated gaseuous atom or ionisolated gaseuous atom or ion

The Greater the ionization energy, The Greater the ionization energy, the more difficult it is to remove an the more difficult it is to remove an electron.electron.

Variations in Successive Variations in Successive Ionization EnergiesIonization Energies

II11>I>I22>I>I33 etc… etc… It’s more difficult to It’s more difficult to

pull away an electron pull away an electron from an increasingly from an increasingly more-positive ionmore-positive ion

There is a sharp There is a sharp increase in ionization increase in ionization energy to remove a energy to remove a core electron, as they core electron, as they are closer to the are closer to the nucleus.nucleus.


Periodic Trends in First Periodic Trends in First Ionization EnergyIonization Energy

Within each period, ionization energy generally Within each period, ionization energy generally increases with increasing atomic number.increases with increasing atomic number.(Smaller atomic radius)(Smaller atomic radius)

Within each group, Ionization generally Within each group, Ionization generally decreases from top to bottom (Larger atomic decreases from top to bottom (Larger atomic radius).radius).

Irregularities: Added “p” orbital sometimes Irregularities: Added “p” orbital sometimes leads to decrease in ionization energy because leads to decrease in ionization energy because the “p”the “p” orbitals have more space than the “s” orbitals. orbitals have more space than the “s” orbitals. Adding a paired electron can also lead to a Adding a paired electron can also lead to a decrease in ionization energy, as there is decrease in ionization energy, as there is increased electron-electron repulsion.increased electron-electron repulsion.

Periodic Trends in First Periodic Trends in First Ionization EnergyIonization Energy


ElectronegativityElectronegativity Electronegativity=An Electronegativity=An

order of an atom’s order of an atom’s overall ability to overall ability to attract electrons. It attract electrons. It combines atomic size combines atomic size and ionization energy and ionization energy into a single summary into a single summary number.number.

http://www.green-planet-solar-energy.com/images/PT-small-electroneg.gif

Covalent BondingCovalent Bonding Created when two Created when two

atoms share electronsatoms share electrons Strive to fulfill the Strive to fulfill the

Octet rule- “atoms Octet rule- “atoms tend to gain, lose, or tend to gain, lose, or share electrons until share electrons until they are surrounded they are surrounded by eight valence by eight valence electrons”electrons”

Many covalent bonds Many covalent bonds are exceptions to the are exceptions to the octet ruleoctet rule

www.ider.herts.ac.uk/.../covalent_bonding.gif

http://academic.brooklyn.cuny.edu/biology/bio4fv/page/covalent-hydrogen.jpeg

Lewis SymbolsLewis Symbols

Consists of the Atom’s chemical symbol, Consists of the Atom’s chemical symbol, plus one dot for every valence electron it plus one dot for every valence electron it hashas

Anions have extra dots, cations fewer Anions have extra dots, cations fewer dotsdots

Examples:Examples:

. . . .. . . . HH•• : : Ar Ar :: ::FF:: •• C C ••

. . . .. . . .

Drawing Lewis Drawing Lewis StructuresStructures

Write the Chemical symbols for every atom in the Write the Chemical symbols for every atom in the moleculemolecule

The atom that makes the most bonds is generally the The atom that makes the most bonds is generally the central atomcentral atom

Determine the total amount of Valence Electrons in Determine the total amount of Valence Electrons in the moleculethe molecule

Place single bonds between all atoms in the molecule Place single bonds between all atoms in the molecule that bondthat bond

With remaining electrons, fill up octets on all the With remaining electrons, fill up octets on all the atomsatoms

If extra electrons exist, place them on the central If extra electrons exist, place them on the central atomatom

If too few electrons exist, create double, or triple If too few electrons exist, create double, or triple bonds, keeping the octet rule in mind. bonds, keeping the octet rule in mind.

Drawing Lewis Drawing Lewis StructuresStructures

Example- COExample- CO22

O C O

Carbon makes more bonds (4) than oxygen (2)

O+O+C = 6+6+4= 16

- -

Write all Chemical symbols

Place single bondsFill all Octets

Not Enough! Must make double bonds

= =This Creates 16 electrons, while satisfying the octet rule

Formal ChargeFormal Charge

Formal charge= the charge the atom Formal charge= the charge the atom would have if each bonding electron pair would have if each bonding electron pair were shared evenly between its two atomswere shared evenly between its two atoms

To determine formal charge draw Lewis To determine formal charge draw Lewis structure, andstructure, and Count all unshared electrons per atomCount all unshared electrons per atom Add half of the single, double, or triple bonds Add half of the single, double, or triple bonds

electrons to the total (either 1,2, or 3 electrons)electrons to the total (either 1,2, or 3 electrons) Subtract this number from that atom’s usual Subtract this number from that atom’s usual

amount of valence electronsamount of valence electrons

Formal ChargeFormal Charge

Example- CNExample- CN--

[:C≡N:]-

Count all unshared electrons

2 2

Add half of bond total

+(6/2) +(6/2)

5 5

Subtract from Atom’s usual amount of valence electrons

4- = 5- =

-1 0

Electron DomainsElectron Domains Any Bond (only single bonds) plus Any Bond (only single bonds) plus

electron pairs (or last unpaired electron) electron pairs (or last unpaired electron) counts as an electron domain.counts as an electron domain.

Electron Domains are important in Electron Domains are important in understanding molecular shapeunderstanding molecular shape

Shapes are categorized by the amount of Shapes are categorized by the amount of total electron domains, then described total electron domains, then described further by the amount of bonding domainsfurther by the amount of bonding domains

If an atom has 5 electron domains, but If an atom has 5 electron domains, but only 3 are bonding domains, the other 2 only 3 are bonding domains, the other 2 are considered non bonding domains, and are considered non bonding domains, and are lone pairs.are lone pairs.

Molecular Molecular ShapesShapes

5 Basic Shapes5 Basic Shapes

All Pictures: chemlab.truman.edu/.../MM1Files/Linear3.gif

Linear

Trigonal Planar

Tetrahedral

Trigonal Bipyramidal

Octahedral

•Shape based on number of electron domains in the molecule

LinearLinear

•One or Two electron Domains

•1 or 2 bonding domains

•Bond angles = 180˚

www.renewacycle.com/2007_02_01_archive.html

•Example- CO2

Trigonal PlanarTrigonal Planar

Three Electron Three Electron DomainsDomains

Bond angle = 120Bond angle = 120˚̊ 3 bonding domains- 3 bonding domains-

trigonal planartrigonal planar Ex. BFEx. BF33

2 bonding domains- 2 bonding domains- bent moleculebent molecule

Ex. bent- NOEx. bent- NO22

Trigonal Planar

Bent

TetrahedralTetrahedral Four Electron DomainsFour Electron Domains Bond Angle109.5Bond Angle109.5˚̊ 4 bonding domains- 4 bonding domains-

TetrahedralTetrahedral ex. CHex. CH44

3 bonding domains- 3 bonding domains- trigonal pyramidaltrigonal pyramidal

ex. NHex. NH33

2 bonding domains- 2 bonding domains- bentbent

Ex. HEx. H22OO

Bent

Trigonal pyramidal

Tetrahedral

Trigonal BipyramidalTrigonal Bipyramidal Five Electron DomainsFive Electron Domains Bond Angles- Equatorial Bond Angles- Equatorial

120120˚ Polar 180˚˚ Polar 180˚ 5 bonding domains- 5 bonding domains-

trigonal bipyramidal- ex. trigonal bipyramidal- ex. PClPCl55

4 bonding domains- 4 bonding domains- Seesaw-ex. SFSeesaw-ex. SF44

Trigonal Bipyramidal

T-Shaped

See-Saw

Linear

•3 bonding domains- T-shaped- ex. ClF3

•2 bonding domains- Linear- ex. XeF2

OctahedralOctahedral 6 Electron Domains6 Electron Domains Bond Angles- Equatorial- 90Bond Angles- Equatorial- 90˚, ˚,

Polar 180˚Polar 180˚ 6 bonding domains- Octahedral6 bonding domains- Octahedral Ex. SFEx. SF66

5 bonding domains- Square 5 bonding domains- Square PyramidalPyramidal

Ex. BrFEx. BrF55

4 bonding domains- Square Planar4 bonding domains- Square Planar Ex. XeFEx. XeF44

Octahedral

Square Pyramidal

Square Planar

DodecahedralDodecahedral

Just KiddingJust Kidding

Bond Order & LengthBond Order & Length

Double bond= bond order of 2Double bond= bond order of 2 Triple bond = bond order of 3Triple bond = bond order of 3 As Bond order increases, bond length As Bond order increases, bond length

decreases decreases As Bond order increases, greater As Bond order increases, greater

repulsive forces exist between repulsive forces exist between adjacent electron domains, creating adjacent electron domains, creating bigger anglebigger angle

As Bond order increases, more energy As Bond order increases, more energy is needed to break the bondis needed to break the bond

Bond PolarityBond Polarity Happens when electrons are shared unevenly Happens when electrons are shared unevenly

between atomsbetween atoms Therefore does not happen between like atoms Therefore does not happen between like atoms

(i.e. H-H)(i.e. H-H) Generally, electronegativity differences of .4 or Generally, electronegativity differences of .4 or

higher are considered polarhigher are considered polar When electronegativity difference is great enough, When electronegativity difference is great enough,

the bond is considered ionic, not polar covalentthe bond is considered ionic, not polar covalent Ex. H-H CEx. H-H C≡N Na-Cl≡N Na-Cl 2.1-2.1-0 2.5-3.0=.5 .9-3.0= 2.12.1-2.1-0 2.5-3.0=.5 .9-3.0= 2.1 0<.4 .5>.4 2.1>>.40<.4 .5>.4 2.1>>.4 Nonpolar Polar IonicNonpolar Polar Ionic

Molecular PolarityMolecular Polarity

No

Yes Yes

No

Polar Bonds Present?

Polar Bonds arranged symmetrically?

Nonpolar Molecule

Polar Molecule

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When is a molecule Polar?

Molecular PolarityMolecular Polarity

Symmetrical Molecules Asymmetrical MoleculesSymmetrical Molecules Asymmetrical Molecules

Linear Trig. Planar

Tetrahedral Trig. Bipyramidal

Octahedral Square Planar

SeesawPyramidal T-Shaped

Square Pyramidal

Bent

Valence Bond Theory Valence Bond Theory (hybrid orbitals)(hybrid orbitals)

Bonds occur when electron shells overlapBonds occur when electron shells overlap Since electrons are simultaneously Since electrons are simultaneously

attracted to both nuclei, bonds occurattracted to both nuclei, bonds occur Valence bond theory alone does not explain Valence bond theory alone does not explain

polyatomic molecules. For this, Hybrid polyatomic molecules. For this, Hybrid orbitals are neededorbitals are needed

sp orbitalssp orbitalsConsider the linear non-polar BeF2 molecule

1s 2s 2pAs it is, the molecule would not bond, since it has a full 2s shell

After promotion…

1s 2s 2p

The molecule can now bond, however it would make non identical polar bonds

By Hybridizing the 2s and 2p shells…

1s sp 2p

Now the Be molecule can make 2 identical bonds

The remaining 2p orbitals end up unhybridized

Additionally, the bigger lobes produced by the sp orbital allow for more overlap, which means stronger bonds

All 2p orbitals can be hybridized. When this occurs, the same amount of orbitals must be created. Ex.

1s 2s 2p promote 1s 2s 2p hybridize 1s sp2 2p

More Hybrid OrbitalsMore Hybrid Orbitals

sp makes 2 180sp makes 2 180˚ orbitals˚ orbitals spsp2 2 makes 3 120makes 3 120˚orbitals˚orbitals spsp33 makes 4 109.5 tetrahedron- makes 4 109.5 tetrahedron-

arranged orbitalsarranged orbitals

http://upload.wikimedia.org/wikipedia/commons/thumb/9/9f/Sp3-Orbital.svg/290px-Sp3-Orbital.svg.png

Sigma(Sigma(σσ)) and Pi( and Pi(ππ)) BondsBonds

σσ bondsbonds occurring on the internuclear bondsbonds occurring on the internuclear axisaxis

ππ bondsbonds occurring between two p bondsbonds occurring between two p orbitals oriented perpendicularly to the orbitals oriented perpendicularly to the internuclear axis. internuclear axis.

ππ bonds produce a sideways overlap, which is bonds produce a sideways overlap, which is not as substantial, and therefore, not as strong, not as substantial, and therefore, not as strong, as a as a σσ bond bond

Single bonds are Single bonds are σσ bonds, double bonds consist bonds, double bonds consist of one of one σσ bond and one bond and one ππ bond, triple bonds bond, triple bonds have one have one σσ bond and 2 bond and 2 ππ bonds bonds


Molecular Orbital TheoryMolecular Orbital Theory

Better explains excited states of Better explains excited states of moleculesmolecules

Each Molecular Orbital (MO) holds Each Molecular Orbital (MO) holds up to two electrons, of opposing up to two electrons, of opposing spinsspins

Associated with the entire molecule, Associated with the entire molecule, not just a single atomnot just a single atom

Molecular Orbital TheoryMolecular Orbital Theory•Easiest way to analyze is through an energy level diagram

•Bottom half of each shell is the bonding molecular orbital, and is lowest energy.

•Top half is the antibonding molecular orbital, and is higher in energy•As the name suggests, antibonding orbitals cancel out the bonding orbitals

•Because energy increases as the position on the chart increases, slots are filled in from the bottom up http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?

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Molecular Orbital TheoryMolecular Orbital Theory

Additionally, based on the positioning in Additionally, based on the positioning in the diagram, the bonds can be analyzed the diagram, the bonds can be analyzed as either as either σσ or or ππ bonds bonds

The diagrams can be used to determine The diagrams can be used to determine whether or not an atom would form whether or not an atom would form naturallynaturally


Bond OrderBond Order Bond order = ½(# of bonding electrons- # Bond order = ½(# of bonding electrons- #

of antibonding electrons)of antibonding electrons) Bond order of HBond order of H22 = ½(2-0) = 1, Therefore = ½(2-0) = 1, Therefore

HH22 exists exists Bond order of HeBond order of He22 = ½(2-2) = 0. Therefore, = ½(2-2) = 0. Therefore,

Helium is not diatomic in natureHelium is not diatomic in nature

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Atomic and Molecular Structure Michael Abosch Brian Pflaum 2 nd Period.

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Transcript of Atomic and Molecular Structure Michael Abosch Brian Pflaum 2 nd Period.