Textbook Notes (369,072)
Canada (162,367)
Chemistry (43)
CHEM 112 (33)
Chapter Final

Chapter 8, 9, 11, 12, 26 and 27 December Exam Review

8 Pages
145 Views

Department
Chemistry
Course Code
CHEM 112
Professor
John Carran

This preview shows pages 1,2 and half of page 3. Sign up to view the full 8 pages of the document.
Description
Chapter 8 – Quantum mechanics Electromagnetic radiation – form of energy transmission in which electric and magnetic fields as are propagated as waves through empty space, or through a medium Wave – disturbance that transmits energy through space or a material medium • Energy is discontinuous (energy changes by a jump, or quantum) • Small particles of matter may at times display wave-like properties Photoelectric effect – phenomenon explaining that when light strikes the surface of metals, electrons are ejected Heisenberg uncertainty principle – cannot measure position and momentum accurately , simultaneously Classical mechanics • Objects = particles • Know exact positions and velocities • Energy is continuous Quantum mechanics – accounts for experiments related to atoms, electrons and light that cannot be explained by classical mechanics • Objects = particles or waves • Cannot know exact positions and velocities • Energy is transferred in discrete units  Lyman series, n=1  Balmer series, n=2  Pachen series, n=3 Standing wave – wave that remains in a constant position Assigning Quantum Numbers 1. Principle quantum number, n. Only positive, non-zero number (n=1, 2, 3) - ENERGY LEVEL/MOST PROBABLE DISTANCE OF FINDING e FROM NUCLEUS 2. Orbital angular momentum quantum number, l. Not larger than n-1 (l=0; s, 1; p, 2;d, 3; f) SHAPE OF ORBITAL 3. Magnetic quantum number, m. Negatlve or positive integer, incl 0 and ranging from –l to l (where l is the orbital angular momentum quantum number (m = -l, (-l l+1)….-1, 0, 1, (l-1), +l)) For example, if l = 2, then m can be +/- : -2, -1, 1, 2 and 0. ORIENTATION • ***NOT involved in orbital designation*** • S  spherical, pdumb-bell, d  clover leaf 4. Electron spin, M (+s- 1/2) • All orbitals with same n and l value are in the same sub-shell • Cannot have two electrons in an atom with all 4 quantum numbers the same (Pauli Excl. Princ.) • Only two electrons may occupy the same orbital, and these must have opposing spins • For degenerate orbitals, electrons initially occupy them singly (atom tends to have as many unpaired electrons as possible) Hund’s rule • count the superscripts of the electron configuration of a species to find the atomic number Degenerate – orbitals at the same energy level • For electron configuration, only Cr and Cu are exceptions to the usual filling rule Reference http://answers.yahoo.com/question/index?qid=20090804201613AAGReAq Chapter 9 – Periodic table properties Effective nuclear charge – true nuclear charge minus charge screened out by electrons: Z = Zeff S • Increases across the row, and is constant in each family Diamagnetic atom/ion – all electrons are paired and the individual magnetic effects cancel out Paramagnetic atom/ion – unpaired electrons present and individual magnetic effects don’t cancel out Isoelectric – when atoms or ions have the same number of electrons (same position on periodic table) Metals - more associated to small atomic radius, high ionization energy and negative electron affinities Nonmetals -more associated to greater atomic radius and lower ionization energy • If two ions have the same electronic configuration, in general, the ion with the higher nuclear charge is smaller • Elements within the same row are more related than those in the same family • Atomic radii: o Covalent radius o Metallic radius o Ionic radius • Ionization energy- energy change for the removal of an electron o inverse of radii • Electron affinity- energy change for the addition of an electron, or enthalpy change as atom in gas phase gains an electron. Chapter 12 – Chapter 10 - Chemical Bonding I • Double/triple bonds are more likely to happen with S, N, C, O, H • Formal charge = valence-lone-(1/2)bonded • Expanded valence – nonmetals in period 3 or later (central atom expand to 10 or 12 electrons) • Coordinate covalent bond – when both electrons in a covalent bond are provided by one of the atoms Hybridization (count regions of high electron density; could be lone pair, double or triple bond) • 4 regions = sp  s + p + p + p • 3 regions = sp  s + p + p •
More Less
Unlock Document

Only pages 1,2 and half of page 3 are available for preview. Some parts have been intentionally blurred.

Unlock Document
You're Reading a Preview

Unlock to view full version

Unlock Document

Log In


OR

Join OneClass

Access over 10 million pages of study
documents for 1.3 million courses.

Sign up

Join to view


OR

By registering, I agree to the Terms and Privacy Policies
Already have an account?
Just a few more details

So we can recommend you notes for your school.

Reset Password

Please enter below the email address you registered with and we will send you a link to reset your password.

Add your courses

Get notes from the top students in your class.


Submit