Atomic and Molecular Structure Michael Abosch Brian Pflaum 2nd Period

Unit Outline Atomic and Electronic Structure, and Quantum Mechanics Periodic Trends Molecular Structure Bonding Theory

The Wave Nature of Light Electromagnetic Radiation- all visible light, radio waves, infrared, X-rays etc. Electromagnetic Spectrum- shows radiation arranged in order of increasing wavelength Visible light is only a small portion of spectrum. http://steve.files.wordpress.com/2006/03/Elcetromagnetic%20spectrum.jpg

The Wave Nature of Light fλ= c (frequency)(wavelength)= Speed of light (2.9979x108 m/s) Frequency measured in s-1 (often Hz) Wavelength measured in meters (often nm,μm)

The Quantization of Energy Quantum=The smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation. Energy, E, of a single quantum equals a constant times the frequency of radiation. E=hf h=planck’s constant=6.626X10-34Joule-seconds.

Photoelectric Effect When photons of sufficiently high energy (greater than the individual metal’s threshold energy) strike a metal surface, electrons are emitted from the metals Energy of Photon, E=hf (planck’s constant)(frequency) Kinetic Energy of ejected electrons: KEe=Ephoton-Ethreshold of metal

Wave Behavior of Matter Dual nature of radiant energy: both particle and wave-like properties DeBroglie wavelength: wavelength=(Planck’s constant)/(momentum)=(h)/(mv) Mass in Kg, Velocity in m/s

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

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

The Azimuthal Quantum Number The azimuthal quantum number, l, can have integral values from 0 to n-1 for each value of n. This quantum number defines the shape of an orbital. The value of l is generally designated by the letters s,p,d, and f. Value of l 1 2 3 Letter Used s p d f http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9020&vid=9884

The Magnetic Quantum Number The magnetic quantum number, ml, can have integral values between -l and l. Describes orientation of orbital in space.

Relationship amongst Quantum Numbers http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9020&vid=9884

Spin Magnetic Quantum Number and Pauli Exclusion Principle The Spin Magnetic Quantum Number, ms, has two possible values: +1/2, -1/2. No two electrons in an atom can have the same set of four quantum numbers n, l, ml, and ms Thus, an orbital can hold a maximum of two electrons, and they must have opposite spins.

Electron Configurations Electron Configuration=A particular arrangement of electrons in the orbitals of an atom. The orbitals are filled in order of increasing energy, with no more than two electrons per orbital.

Orbital Diagrams Each orbital is denoted by a box, and each electron by a half arrow (which represents spin-up or spin-down) Electrons having opposite spins are said to be paired when they are in the same orbital An unpaired electron is one not accompanied by a partner of opposite spin. http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9020&vid=9887

Hund’s Rule Hund’s Rule=For orbitals of the same energy level, the lowest energy is attained when the number of electrons with the same spin is maximized. Note how in the diagram below, all three p orbitals are filled singularly before an electron is paired http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9020&vid=9887

The Periodic Table and Electron Filling Order http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9020&vid=9888

Condensed Electron Configurations http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9020&vid=9887 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 denoted As 4s1

Ions Start by writing the electron configuration for the normal element Then remove (or add) electrons as necessary, always taking (or adding) from the highest principle quantum number first (ignoring the filling order). Fe=[Ar]4s23d6 Fe(II)=[Ar]3d6

Anomalous Electron Configurations Electron configurations of certain elements appear to violate the “rules” Frequently occurs when there are enough electrons to lead to precisely half-filled sets of degenerate (same energy-level) orbitals, or to completely fill an orbital. This conserves Energy No universal pattern or predictability Ex: Chromium is [Ar]4s13d5 instead of [Ar]4s23d4

Practice What’s the electron configuration for Lead? Answer: [Xe]6s24f145d106p2 Assign Quantum numbers to it’s last filled electron. Answer: n=6, l=1, ml=0, ms=+1/2 http://www.elementsdatabase.com/

Periodic Trends Atomic Size Ionic Size Ionization Energies Electronegativity campus.ru.ac.za/full_images/ img05206111510.jpg

Atomic Size Within each group, size increases from top to bottom, results primarily from the increase in principle quantum number of electrons In each period, atomic radius tends to decrease from left to right. Increase in the effective nuclear charge as we move across a row steadily draws valence electrons closer to nucleus Exceptions: The addition of a paired electron produces increased repulsion that sometimes leads to an increase in size (Like from a p3 to a p4 element.)

Atomic Size http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9021&vid=9896#oa217306

Ionic vs. Atomic Size Cations: Compared to its neutral atom, cations are smaller because electrons have vacated the biggest orbital Anions: Compared to its neutral atom, anions are larger because adding electrons increases repulsions, which leads to more space.

Ionic vs. Atomic Size http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9021&vid=9896#oa217306

Isoelectronic Series Isoelectronic Series=A group all containing the same number of electrons. As the atomic number increases, the radius decreases. Ex: Cl-, Ar, K+ Size: Cl->Ar>K+

Ionization Energy Ionization Energy=The minum energy required to remove an electron from the ground state of the isolated gaseuous atom or ion The Greater the ionization energy, the more difficult it is to remove an electron.

Variations in Successive Ionization Energies I1>I2>I3 etc… It’s more difficult to pull away an electron from an increasingly more-positive ion There is a sharp increase in ionization energy to remove a core electron, as they are closer to the nucleus. http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9021&vid=9897

Periodic Trends in First Ionization Energy Within each period, ionization energy generally increases with increasing atomic number.(Smaller atomic radius) Within each group, Ionization generally decreases from top to bottom (Larger atomic radius). Irregularities: Added “p” orbital sometimes leads to decrease in ionization energy because the “p” orbitals have more space than the “s” orbitals. Adding a paired electron can also lead to a decrease in ionization energy, as there is increased electron-electron repulsion.

Periodic Trends in First Ionization Energy http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9021&vid=9897

Electronegativity Electronegativity=An order of an atom’s overall ability to attract electrons. It combines atomic size and ionization energy into a single summary number. http://www.green-planet-solar-energy.com/images/PT-small-electroneg.gif

Covalent Bonding Created when two atoms share electrons Strive to fulfill the Octet rule- “atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons” Many covalent bonds are exceptions to the octet rule www.ider.herts.ac.uk/.../covalent_bonding.gif http://academic.brooklyn.cuny.edu/biology/bio4fv/page/covalent-hydrogen.jpeg

Lewis Symbols Consists of the Atom’s chemical symbol, plus one dot for every valence electron it has Anions have extra dots, cations fewer dots Examples: . . . . H• : Ar : :F: • C •

Drawing Lewis Structures Write the Chemical symbols for every atom in the molecule The atom that makes the most bonds is generally the central atom Determine the total amount of Valence Electrons in the molecule Place single bonds between all atoms in the molecule that bond With remaining electrons, fill up octets on all the atoms If extra electrons exist, place them on the central atom If too few electrons exist, create double, or triple bonds, keeping the octet rule in mind.

Drawing Lewis Structures Write all Chemical symbols Example- CO2 Carbon makes more bonds (4) than oxygen (2) O+O+C = 6+6+4= 16 Place single bonds O C O - - = = Fill all Octets Not Enough! Must make double bonds This Creates 16 electrons, while satisfying the octet rule

Formal Charge Formal charge= the charge the atom would have if each bonding electron pair were shared evenly between its two atoms To determine formal charge draw Lewis structure, and Count all unshared electrons per atom Add half of the single, double, or triple bonds electrons to the total (either 1,2, or 3 electrons) Subtract this number from that atom’s usual amount of valence electrons

[:C≡N:]- Formal Charge Example- CN- -1 Count all unshared electrons Add half of bond total [:C≡N:]- Subtract from Atom’s usual amount of valence electrons 2 2 +(6/2) +(6/2) 4- = 5 5- = 5 -1

Electron Domains Any Bond (only single bonds) plus electron pairs (or last unpaired electron) counts as an electron domain. Electron Domains are important in understanding molecular shape Shapes are categorized by the amount of total electron domains, then described further by the amount of bonding domains If an atom has 5 electron domains, but only 3 are bonding domains, the other 2 are considered non bonding domains, and are lone pairs.

Molecular Shapes 5 Basic Shapes Shape based on number of electron domains in the molecule Linear Trigonal Bipyramidal Trigonal Planar Octahedral Tetrahedral All Pictures: chemlab.truman.edu/.../MM1Files/Linear3.gif

Linear One or Two electron Domains 1 or 2 bonding domains Bond angles = 180˚ Example- CO2 www.renewacycle.com/2007_02_01_archive.html

Trigonal Planar Three Electron Domains Bond angle = 120˚ 3 bonding domains- trigonal planar Ex. BF3 2 bonding domains- bent molecule Ex. bent- NO2 Trigonal Planar Bent

Tetrahedral Four Electron Domains Bond Angle109.5˚ 4 bonding domains- Tetrahedral ex. CH4 3 bonding domains- trigonal pyramidal ex. NH3 2 bonding domains- bent Ex. H2O Trigonal pyramidal Bent

Trigonal Bipyramidal Five Electron Domains Bond Angles- Equatorial 120˚ Polar 180˚ 5 bonding domains- trigonal bipyramidal- ex. PCl5 4 bonding domains- Seesaw-ex. SF4 3 bonding domains- T-shaped- ex. ClF3 2 bonding domains- Linear- ex. XeF2 Linear See-Saw Trigonal Bipyramidal T-Shaped

Octahedral 6 Electron Domains Bond Angles- Equatorial- 90˚, Polar 180˚ 6 bonding domains- Octahedral Ex. SF6 5 bonding domains- Square Pyramidal Ex. BrF5 4 bonding domains- Square Planar Ex. XeF4 Octahedral Square Pyramidal Square Planar

Dodecahedral Just Kidding

Bond Order & Length Double bond= bond order of 2 Triple bond = bond order of 3 As Bond order increases, bond length decreases As Bond order increases, greater repulsive forces exist between adjacent electron domains, creating bigger angle As Bond order increases, more energy is needed to break the bond

Bond Polarity Happens when electrons are shared unevenly between atoms Therefore does not happen between like atoms (i.e. H-H) Generally, electronegativity differences of .4 or higher are considered polar When electronegativity difference is great enough, the bond is considered ionic, not polar covalent Ex. H-H C≡N Na-Cl 2.1-2.1-0 2.5-3.0=.5 .9-3.0= 2.1 0<.4 .5>.4 2.1>>.4 Nonpolar Polar Ionic

Molecular Polarity When is a molecule Polar? No Yes Yes Polar Bonds Present? Polar Bonds arranged symmetrically? Nonpolar Molecule Polar Molecule http://bitesizebio.com/wp-content/uploads/2007/12/picture-1.png

Molecular Polarity Symmetrical Molecules Asymmetrical Molecules Linear Trig. Planar Bent Square Pyramidal Tetrahedral Trig. Bipyramidal Pyramidal Seesaw T-Shaped Square Planar Octahedral

Valence Bond Theory (hybrid orbitals) Bonds occur when electron shells overlap Since electrons are simultaneously attracted to both nuclei, bonds occur Valence bond theory alone does not explain polyatomic molecules. For this, Hybrid orbitals are needed

sp orbitals Consider the linear non-polar BeF2 molecule 1s 2s 2p As it is, the molecule would not bond, since it has a full 2s shell After promotion… The molecule can now bond, however it would make non identical polar bonds 1s 2s 2p By Hybridizing the 2s and 2p shells… Now the Be molecule can make 2 identical bonds 1s sp 2p 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 Orbitals sp makes 2 180˚ orbitals sp2 makes 3 120˚orbitals sp3 makes 4 109.5 tetrahedron-arranged orbitals http://upload.wikimedia.org/wikipedia/commons/thumb/9/9f/Sp3-Orbital.svg/290px-Sp3-Orbital.svg.png

Sigma(σ) and Pi(π) Bonds σ bonds- bonds occurring on the internuclear axis π bonds- bonds occurring between two p orbitals oriented perpendicularly to the internuclear axis. π bonds produce a sideways overlap, which is not as substantial, and therefore, not as strong, as a σ bond Single bonds are σ bonds, double bonds consist of one σ bond and one π bond, triple bonds have one σ bond and 2 π bonds http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9030&vid=9925#oa220020

Molecular Orbital Theory Better explains excited states of molecules Each Molecular Orbital (MO) holds up to two electrons, of opposing spins Associated with the entire molecule, not just a single atom

Molecular 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?vbcid=9030&vid=9927

Molecular Orbital Theory Additionally, based on the positioning in the diagram, the bonds can be analyzed as either σ or π bonds The diagrams can be used to determine whether or not an atom would form naturally http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9030&vid=9927

Bond Order Bond order = ½(# of bonding electrons- # of antibonding electrons) Bond order of H2 = ½(2-0) = 1, Therefore H2 exists Bond order of He2 = ½(2-2) = 0. Therefore, Helium is not diatomic in nature http://myphlip.pearsoncmg.com/phproducts/student/ab2page.cfm?vbcid=9030&vid=9926