CHAPTER 7 AND PART 16 AP CHEMISTRY

LIGHT Wavelength = λ –nm/ wave, m/wave, or nm or m –λf= c; c = x 10 8 m/s Frequency = f Frequency = f –Waves/s or 1/s = c/ λ High frequency = short wavelength Low frequency = long wavelength

PHOTON ENERGY E = hf Planck’s constant h = x J.s Continuous spectrum –A–A–A–All wavelengths Line spectrum –D–D–D–Discrete energies light going through a prism ROY G BIV

HYDROGEN ATOM Bohr model used hydrogen atom Electron orbits around nucleus Energy levels: Zero energy –w–w–w–where a photon and electron are in completely separate energy states within an atom Ground state –l–l–l–lowest amount of energy an electron has Excited state –e–e–e–electron getting extra energy E = hf = Ehi - Elo Electrons can only give off certain energies

QUANTUM NUMBERS Most probable place an electron can be found Ψ amplitude (height) of the electron wave Ψ 2 directly proportional to the probability of finding the electron Quantum numbers Principal energy level –n–n–n–n = 1, 2, 3,..... etc. Sublevels –s–s–s–s, p, d, and f –l–l–l–l = 0, 1, 2, 3, etc.

CONTINUE Orbitals or orientations ––m––m = l to -l ––l––l = 0 then m = 0 ––l––l = 1 then m = 1, 0, -1 ––l––l = 2 then m = 2, 1, 0, -1, -2 Spin ––s––s = +1/2, -1/2 Pauli Exclusion Principle ––N––No two electrons can have the same four quantum numbers Capacities Each energy level (n) has n sublevels Each sublevel (l) has 2l + 1 orbitals Each orbital (m) has 2 electrons

CONTINUE Heisenberg principle –s–s–s–state that there are limitations in knowing what the position and momentum are at any given time Probability distribution intensity of color -electron density map Most probable place to find a hydrogen electron is.529Ǻ from nucleus 1Ǻ = m

Polyelectronic atoms –T–T–T–Treat electrons as if they have nuclear attraction and average repulsion from other electrons –E–E–E–Effective nuclear charge –Z–Z–Z–Zeff = Zactual - (electron repulsion) Z = atomic number Periodic table can predict the filling of the sublevels –E–E–E–Elements in groups –E–E–E–Elements in groups –T–T–T–Transition metals in groups –T–T–T–The two sets of 14 elements at the bottom of the table fill

ORBITAL DIAGRAM Hund’s rule –W–W–W–Within a given sublevel, the order of filling is such that there is the maximum number of half-filled orbitals Monoatomic ion electron configuration –I–I–I–Ions with noble gas configuration Anions are formed from atoms with Cations are formed from atoms with

PERIODIC TRENDS Atomic radius In general atomic radii –D–D–D–Decreases as you move left to right –I–I–I–Increases as you move down Effective nuclear charge –Z–Z–Z–Zeffective = Z - S –Z–Z–Z–Z = number or protons –S–S–S–S = number of core electrons that are shielding the outer electrons from the nucleus

IONIC RADIUS Ionic radius –I–I–I–Ionic radius increases as you move down the group –C–C–C–Cations decrease left to right –A–A–A–Anions decrease left to right

IONIZATION ENERGY Energy required to remove an electron Increase as you move across left to right Decrease as you move down

ELECTRONEGATIVITY Electron Affinity Energy released from an atom as it acquires another electron Atomic radius Decreases going across because the valence electrons are being pulled in, due to an increase of proton attraction. Shielding remains constant across the period Electronegativity How much an atom wants an electron

METALS Metals Metallic luster, ductile, malleable, good conductor of heat and electricity Compounds formed from metals and nonmetals tend to be ionic

METALS AND NONMETALS Nonmetals –N–N–N–No luster, poor conductors –N–N–N–Nonmetals tend to gain electrons to become anions When compounds are made up chiefly of nonmetals they are molecular compounds Most nonmetal oxides are Semimetals –S–S–S–Some properties of both, brittle, semi-conductor Trends in metallic characteristics –T–T–T–The more an element shows physical and chemical properties characteristic for metals, increase right to left and top to bottom