Transcript CHAPTER 7 AND PART 16

CHAPTER 7 AND PART 16
AP CHEMISTRY
LIGHT
Wavelength = λ
– nm/ wave, m/wave, or nm or m
– λf= c; c = 2.998 x 108 m/s
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 = 6.626 x 10-34 J.s
Continuous spectrum
– All wavelengths
Line spectrum
– 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
– where a photon and electron are in
completely separate energy states within an
atom
Ground state
– lowest amount of energy an electron has
Excited state
– 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 = 1, 2, 3, ..... etc.
Sublevels
– s, p, d, and f
– l = 0, 1, 2, 3, ...... etc.
CONTINUE
Orbitals or orientations
– m = l to -l
– l = 0 then m = 0
– l = 1 then m = 1, 0, -1
– l = 2 then m = 2, 1, 0,
-1, -2
Spin
– s = +1/2 , -1/2
Pauli Exclusion Principle
– 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
– 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 = 10-10 m
QUANTUM NUMBERS
Polyelectronic atoms
– Treat electrons as if they have nuclear attraction and
average repulsion from other electrons
– Effective nuclear charge
– Zeff = Zactual - (electron repulsion) Z = atomic
number
Periodic table can predict the filling of the sublevels
– Elements in groups
– Elements in groups
– Transition metals in groups
– The two sets of 14 elements at the bottom of the
table fill
ORBITAL
DIAGRAM
Hund’s rule
– Within a given sublevel, the order of filling is such
that there is the maximum number of half-filled
orbitals
Monoatomic ion electron configuration
– 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
–Decreases as you move left to right
–Increases as you move down
Effective nuclear charge
–Zeffective = Z - S
–Z = number or protons
–S = number of core electrons that are
shielding the outer electrons from the
nucleus
IONIC RADIUS
Ionic radius
– Ionic radius increases as you move down the
group
– Cations decrease left to right
– 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
– No luster, poor conductors
– 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
– Some properties of both, brittle, semi-conductor
Trends in metallic characteristics
– The more an element shows physical and chemical properties
characteristic for metals, increase right to left and top to
bottom