atomic and molecular structure n.
Skip this Video
Loading SlideShow in 5 Seconds..
Atomic and Molecular Structure PowerPoint Presentation
Download Presentation
Atomic and Molecular Structure

Atomic and Molecular Structure

270 Vues Download Presentation
Télécharger la présentation

Atomic and Molecular Structure

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

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

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

  3. 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.

  4. 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)

  5. 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.

  6. 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

  7. 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

  8. 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

  9. 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

  10. 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.

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

  12. Relationship amongst Quantum Numbers

  13. 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.

  14. 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.

  15. 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.

  16. 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

  17. The Periodic Table and Electron Filling Order

  18. Condensed Electron Configurations 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

  19. 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

  20. 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

  21. 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

  22. Periodic Trends • Atomic Size • Ionic Size • Ionization Energies • Electronegativity img05206111510.jpg

  23. 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.)

  24. Atomic Size

  25. 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.

  26. Ionic vs. Atomic Size

  27. 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+

  28. 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.

  29. 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.

  30. 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.

  31. Periodic Trends in First Ionization Energy

  32. 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.

  33. 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

  34. 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 • . . . .

  35. 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.

  36. 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

  37. 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

  38. Formal Charge Count all unshared electrons • Example- CN- 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 0

  39. 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.

  40. Linear Trigonal Bipyramidal Trigonal Planar Octahedral Tetrahedral • Shape based on number of electron domains in the molecule Molecular Shapes5 Basic Shapes All Pictures:

  41. Linear • One or Two electron Domains • 1 or 2 bonding domains • Bond angles = 180˚ • Example- CO2

  42. 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

  43. Tetrahedral 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

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

  45. 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

  46. Dodecahedral • Just Kidding

  47. 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

  48. 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

  49. No Yes Yes No Polar Bonds Present? Nonpolar Molecule Polar Bonds arranged symmetrically? Polar Molecule Molecular Polarity When is a molecule Polar?

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