1 / 15

CHAPTER 7 AND PART 16

CHAPTER 7 AND PART 16. AP CHEMISTRY. LIGHT. Wavelength = λ nm/ wave, m/wave, or nm or m λ f= c; c = 2.998 x 10 8 m/s Frequency = f Waves/s or 1/s = c/ λ High frequency = short wavelength Low frequency = long wavelength. E = hf Planck’s constant h = 6.626 x 10 -34 J . s

astro
Télécharger la présentation

CHAPTER 7 AND PART 16

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. CHAPTER 7 AND PART 16 AP CHEMISTRY

  2. 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

  3. 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 PHOTON ENERGY

  4. 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

  5. QUANTUM NUMBERS • Most probable place an electron can be found • Ψ amplitude (height) of the electron wave • Ψ2 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.

  6. 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

  7. 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

  8. 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

  9. 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

  10. 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

  11. IONIC RADIUS • Ionic radius • Ionic radius increases as you move down the group • Cations decrease left to right • Anions decrease left to right

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

  13. 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

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

  15. 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

More Related