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Modern Atomic Theory Notes

Modern Atomic Theory Notes. Electromagnetic radiation – energy that travels through space as waves. Waves have three primary characteristics: Wavelength (  - lambda) – distance between two consecutive peaks or troughs in a wave. Unit = meter

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Modern Atomic Theory Notes

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  1. Modern Atomic Theory Notes

  2. Electromagnetic radiation – energy that travels through space as waves. Waves have three primary characteristics: • Wavelength (- lambda) – distance between two consecutive peaks or troughs in a wave. Unit = meter • Frequency (  = nu) – indicates how many waves pass a given point per second. Unit = Hertz (Hz) • Speed – velocity (c = speed of light = 3 x 108 m/sec) - indicates how fast a given peak moves in a unit of time c = 

  3. Electromagnetic radiation (light) is divided into various classes according to wavelength.

  4. Wave- Particle Theory – Light as waves – Light as photons (de Broglie) Photon/quantum – packet of energy – a “particle” of electromagnetic radiation

  5. Energy - (E – change in energy) – Unit Joules (J) Planck’s Constant – (h = 6.626 x 10-34 J * s) Ephoton = h Change in Energy of a photon = (Planck’s Constant) x (frequency) c =  + Ephoton = h = Ephoton = hc  Ex: What is the wavelength of light with a frequency of 6.5 x 1014 Hz? What is the change in Energy of the photon? E = hc  ΔE = 4.3 x 10-19 J • = c  • Given • = 6.5 x 1014 Hz  = ? ΔE = ? = 3 x 108 m/sec 6.5 x 1014 Hz • = (6.626 x 10-34 J x s)(3 x 108 m/s) • 4.6 x 10-7 m λ= 4.6 x 10-7 m

  6. Wrap – up So with light waves, you can convert between wavelength, frequency, and energy with two equations: l = c E = h And two constants: c = 3 * 108 m/s h = 6.626 * 10-34 J s In the visible part of the spectrum, different colors correspond to different frequencies, wavelengths and energies. Blue light has a short wavelength, high frequency and high energy. Red light has a long wavelength, low frequency, and low energy.

  7. Learning Check

  8. Excited State – atom with excess energy Ground State – lowest possible energy state Wavelengths of light carry different amounts of energy per photon Only certain types of photons are produced (see only certain colors) Quantized – only certain energy levels (and therefore colors) are allowed

  9. Emission and Absorption Spectra Intensity Color Emission Spectrum – bright lines on a dark background. Produced as excited electrons return to a ground state – as in flame tests. Absorption Spectrum – dark lines in a continuous spectrum. Produced as electrons absorb energy to move into an excited state, only certain allowable transitions can be made. Energy absorbed corresponds to the increase in potential energy needed to move the electron into allowed higher energy levels. The frequencies absorbed by each substance are unique.

  10. An Element’s Fingerprint • When excited by heat or electricity, gases glow with characteristic colors. • A prism can be used to spread out the light from these hot gases. • This reveals a series of discrete lines, the element’s fingerprint. • Chemists use these fingerprints (called spectral lines) to identify elements both in the lab and in space.

  11. Here are some spectral lines

  12. Learning Check Now, try matching each of the spectra from column A with its corresponding line plot from column B. A B

  13. Bohr Model – suggested that electrons move around the nucleus in circular orbits Only Correct for Hydrogen Wave Mechanical Model – Described by orbitals gives no information about when the electron occupies a certain point in space or how it moves *aka – Heisenberg's Uncertainty Principle

  14. 1 2 3 4 5 6 7 s p d f Parts of the Wave Mechanical Model 1. Principle Energy Level (n) – energy level designated by numbers 1-7. -called principle quantum numbers 2. Sublevel – exist within each principle energy level -the energy within an energy level is slightly different -each electron in a given sublevel has the same energy -lowest sublevel = s, then p, then d, then f

  15. Parts of the Wave Mechanical Model cont. 3. Orbital – region within a sublevel or energy level where electrons can be found s sublevel – 1 orbital p sublevel – 3 orbitals d sublevel – 5 orbitals f sublevel – 7 orbitals - ** No more than two electrons can occupy an orbital** -an orbital can be empty, half-filled, filled

  16. Electron Configuration – arrangement of the electrons among the various orbitals of the atom Ex: 1s22s22p6 = Neon Sulfur = 1s2 2s2 2p6 3s2 3p4 1s2 2p6 Cd = 2s2 4p6 5s2 3s2 3p6 4s2 3d10 4d10 2s2 1s2 2p6 3s1 Na = Ne Na

  17. Learning Check Write electron configurations: Oxygen Chromium

  18. Shapes of orbitals All s orbitals are spherical as the principle energy level increases the diameter increases. All p orbitals are dumbbell or figure-8 shaped – all have the same size and shape within an energy level

  19. 4 of the d orbitals are 4-leaf clover shaped and the last is a figure-8 with a donut – all have the same size and shape within an energy level

  20. f orbitals are complicated!!!!!

  21. Electron Spin Spin – motion that resembles earth rotating on its axis– clockwise or counterclockwise Pauli Exclusion Principle – two electrons in the same orbital must have opposite spins Hund’s Rule – All orbitals within a sublevel must contain at least one electron before any orbital can have two Orbital Diagram – describes the placement of electrons in orbitals • use arrows to represent electrons with spin • line represents orbital (s=1, p=3, d=5, f=7) ____ full ____ half-full ____ empty

  22. Orbital Diagrams Neon = 1s__ 2s__ 2p__ __ __ Ex: Carbon = 1s__ 2s__ 2p__ __ __ 2p__ __ __ 3s__ 3p__ __ __ Zinc = 1s__ 2s__ 4s__ 3d__ __ __ __ __ 2p__ __ __ 3s__ Gallium = 1s__ 2s__ 3p__ __ __ 4s__ 3d__ __ __ __ __ 4p__ __ __

  23. Learning Check Draw orbital diagrams: Oxygen Chromium

  24. Noble Gas Configuration – Shorthand configuration that substitutes a noble gas for electrons Ex: • Valence Electrons – Electrons in the outermost (highest) principle energy level in an atom • Core Electrons – innermost electrons – not involved in bonding • Valence Configuration – shows just the valence electrons Ex: Na = 1s22s22p63s1 or [Ne]3s1 Sn = 1s22s22p63s23p64s23d104p65s24d105p2 or [Kr]5s24d105p2 Na = 1s22s22p63s1 1 Valence Sn = 1s22s22p63s23p64s23d104p65s24d105p2 4 Valence Na = 3s1 3rd Shell/1valence electron Sn = 5s25p2 5th Shell/4 valence electrons

  25. Learning Check Write the noble gas configuration, valence configuration, and number of valence electrons: Oxygen Chromium

  26. Periodic Table Dimitri Mendeleev-1869- developed the first version of the periodic table. He expressed the regularities as a periodic function of the atomic mass. Henry Moseley- revised Mendeleev periodic table by describing regularities in physical and chemical properties as periodic functions of the atomic number

  27. Periods – horizontal rows • Period number corresponds to the principal quantum number of valence electrons Groups (family) – vertical column Elements with similar valence electrons configurations Group 1 – alkali metals – reactive Group 2 - alkaline earth metals – reactive Group 3-12 – transition metals Group 15 – nitrogen family Group 16 – oxygen family – reactive Group 17 – halogens – very reactive Group 18 – noble gases

  28. Periodic Trends 1. Atomic Radius/Size – size of an atom Increases – down a group Decreases – across a period Size of ions Cation Ca+2/Ca Ca larger because Ca+2 lost 2 electrons Anion S-2/S S-2 larger because S-2 gained 2 electrons

  29. 2. Ionization Energy – energy required to remove an electron from an individual atom in a gas phase M(g) M+(g) + e- (energy to make a positive ion) • Metals lose electrons to non-metals so relatively low energy is needed • High ionization energy means an electron is hard to remove Decreases – down a group Increases - across a period

  30. 3. Electron Affinity – Electron affinity is the energy involved when an electron is added to a gaseous atom. • Negative values of energy mean that energy was released during the process. Atoms with negative values of electron affinity have a very strong attraction for electrons. • Positive values of electron affinity have very little attraction for electrons. (energy involved in negative ions) Decreases – down a group Increases - across a period

  31. 4. Electronegativity is the tendency of an atom to draw electrons to itself when in a covalent bond. Consequently, the trends are the same as for electron affinity. The atoms with the highest electronegativity are fluorine, then oxygen, then nitrogen. It is also important to know that the electronegativity of hydrogen is slightly less than that of carbon. Decreases – down a group Increases - across a period

  32. 5. Metallic Character Increases – down a group Decreases – across a period Electronegativity Electronegativity

  33. Learning Check Put the following elements in order of increasing atomic radius: a. Ge, Se, Fe, Ca b. C, Pb, Sn, Si, Ge Put the following elements in order of increasing electronegativity: a. Ge, Se, Fe, Ca b. C, Pb, Sn, Si, Ge

  34. Notes- Chemical Bonding

  35. Bond- force that holds groups of two or more atoms together and makes them function as a unit bond energy- energy required to break the bond (tells the bond strength) Ionic bonding- between ionic compounds which contain a metal and a nonmetal • Atoms that lose electrons relatively easily react with an atom that has a high affinity for electrons • Transfer of electrons Covalent bonding- between two nonmetals • Electrons are shared by nuclei Polar Covalent bonding- unequal sharing of electrons • positive end attracted to the negative end • (delta) indicates partial charge

  36. electronegativity-(p. 362) relative ability of an atom in a molecule to attract shared electrons to itself • The higher the atom’s electronegativity value, the closer the shared electrons tend to be to that atom when it forms a bonds Increases – across a period Decreases- down a group H-H = 2.1 - 2.1 = 0 Ex. List the following in order of increasing polarity. H-H, O-H, Cl-H, S-H, F-H O-H = 3.5 - 2.1 = 1.4 Cl-H = 3.0 - 2.1 = .9 H-H, S-H, Cl-H, O-H, F-H S-H = 2.5 - 2.1 = .4 F-H = 4.0 - 2.1 = 1.9

  37. Dipole moment- has a center of positive charge and a center of negative charge • Represented by an arrow • Arrow points toward the negative charge

  38. Chemical Formula – type of notation made with numbers and chemical symbols • indicates the composition of a compound • indicates the number of atoms in one molecule Molecule - Bonded collection of two or more atoms of the same element or different elements - monatomic molecule – one atom molecules - diatomic molecule – two atom molecules (seven) MEMORIZE Br, I, N, Cl, H, O, F

  39. Semi-metals Nonmetals METALS Metals Location: Left side of Periodic Table Properties: Ductile – drawn into wires Malleable – hammered into sheets Metallic Luster – shine Good Conductors of Heat and Electricity Nonmetals Location: Right side of Periodic Table Properties: Brittle Lack Luster – not shiny Poor Conductors of Heat and Electricity Semi-metals Location: Along Stair-step Properties: Have properties of metals and nonmetals also called METALLOIDS Si, Ge, As, Sb, Te, Po, At

  40. Molecular Nomenclature Molecular Compounds (molecules) – compounds made from two nonmetals - electrons are shared by two atoms Naming Molecular Prefixes: (MEMORIZE) Mono-1 tetra-4 hepta-7 deca-10 di-2 penta-5 octa-8 tri-3 hexa-6 non-9 prefixes are used with both the first named and second named element. Exception: mono- is not used on the first word second word ends in –ide If a two syllable prefix ends in a vowel, the vowel is dropped before the prefix is attached to a word beginning with a vowel monooxide Writing molecular formulas Translate prefixes Examples: N2O dihydrogen monoxide Si8O5 tetrasulfur hexachloride NH3 carbon monoxide P3I10 carbon dioxide = Dinitrogen monoxide = H2O = Octasilicon pentoxide = S4Cl6 = Nitrogen trihydride = CO = Triphosphorus deciodide = CO2

  41. Learning Check Write the name: • C2O4 • P2O5 Write the formula: a. Dihydrogen monoxide b. Phosphorus trihydride

  42. Valence electrons are used in bonding. 8 1 6 2 4 7 3 5 • Stable elements want to achieve 8 electrons similar to the noble gases • If it’s a metal it wants to achieve the configuration for the noble gas before. • If it’s a nonmetal it wants to achieve the configuration for the noble gas after. 2

  43. • • • • F F • • • • • • • • Li Cl [Li]+1 + [ Cl ]-1 Lewis Structure- representation of a molecule Shows how the valence electrons are arranged among the atoms in the molecule. s pzXpx py • • • For an element: Oxygen 1s22s22p4 O • • • For a compound: +  For a molecule:

  44. H & He Duet rule- only two electrons in the full shell Octet rule- surrounded by eight electrons Bonding pair- electrons shared with other atom Lone pair or unshared pair- not involved in bonding Happy Eight!!!!! Line (-) = 2 electrons dots (••) = 2 electrons/each dot is one electron

  45. 5 Steps for Covalently Bonded Lewis Structures • Find the total number of valence electrons. • Calculate the number of “needed” electrons to give each atom 8 electrons, except for H which wants 2. • Subtract valence electrons from the “needed” electrons. This is the number of bonding electrons. • Divide the bonding electrons by 2, to find the number of bonds. • Subtract the bonding electrons from the valence electrons to find the non-bonding electrons or lone pairs. • Choose a central atom and assemble the pieces to make all atoms involved stable. • • • • Br • Ex. GeBr4 • • • • • • • Br Br Ge Valence = 1(4) + 4(7) = 32 • • • • • • Needed = 1(8) + 4(8) = 40 Bonding = 40 – 32 = 8 • • Br • • Bonds = 8/2 = 4 lines • • Lone e- = 32 – 8 = 24 dots Central atom = Ge

  46. Single bond- involves two atoms sharing one pair • Double bond- involves two atoms sharing two pairs • Triple bond- involves two atoms sharing three pairs Ex. CH4 C2H4 C2H2 1. 1(4) + 4(1) = 8 1. 2(4) + 4(1) = 12 1. 2(4) + 2(1) = 10 2. 1(8) + 4(2) = 16 2. 2(8) + 4(2) = 24 2. 2(8) + 2(2) = 20 3. 16 - 8 = 8 3. 24 - 12 = 12 3. 20 - 10 = 10 4. 8/2 = 4 lines 4. 12/2 = 6 lines 4. 10/2 = 5 lines 5. 8 – 8 = 0 dots 5. 12 – 12 = 0 dots 5. 10 – 10 = 0 dots Central atom = C Central atom = C Central atom = C H H H C C H C C H H H C H H H

  47. Resonance- more than one Lewis structure can be drawn for the molecule Ex. CO2 • • • • O C O • • 1. 1(4) + 2(6) = 16 • • 2. 1(8) + 2(8) = 24 3. 24 - 16 = 8 • • • • 4. 8/2 = 4 lines O C O • • 5. 16 – 8 = 8 dots • • Central atom = C • • • • O C O • • • •

  48. Exceptions to the Octet Rule 1. boron and beryllium- tend to be electron deficient • boron can hold 6 total electrons • beryllium can hold 4 total electrons ex. BF3 BeH2 • • • • F • • 1. 1(3) + 3(7) = 24 1. 1(2) + 2(1) = 4 2. 1(6) + 3(8) = 30 2. 1(4) + 2(2) = 8 B 3. 30 - 24 = 6 3. 8 - 4 = 4 • • • 4. 6/2 = 3 lines 4. 4/2 = 2 lines • F F • • 5. 24 – 6 = 18 dots 5. 4 – 4 = 0 dots • • • • • • Central atom = B Central atom = Be H Be H

  49. 2. Electrons are small spinning electric charges that create magnetic fields • Diamagnetic- substances which have paired electrons that cancel out the magnetic field • Paramagnetic- substances the have one or more unpaired electrons that show great attraction to the magnetic field Ex. O2 PH3 • • • • O O • • • • • • H P H H

  50. • • • • • • • • • • • • • • • • • • • • • • • • • • • • F • • • • • • • • F F S F F F 3. Odd number of electrons • You cannot write electron dot structures that fulfill the octet rule, when the total number of valence electrons is odd Ex. NO 4. Expanded Octet- expand the valence shell to include more than 8 electrons • Phosphorus and sulfur can expand to include 10 or 12 electrons • You will know you have an expanded octet when you don’t have enough bonds for the atoms present Ex. SF6 1. 1(5) + 1(6) = 11 No Drawing

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