Atomic Structure and Periodicity Chapter 2 Atomic Structure and Periodicity

2.1 Electromagnetic Radiation 2.2 The Nature of Matter 2.3 The Atomic Spectrum of Hydrogen 2.4 The Bohr Model 2.5 The Quantum Mechanical Model of the Atom 2.6 Quantum Numbers 2.7 Orbital Shapes and Energies 2.8 Electron Spin and the Pauli Principle 2.9 Polyelectronic Atoms 2.10 The History of the Periodic Table 2.11 The Aufbau Principle and the Periodic Table 2.12 Periodic Trends in Atomic Properties 2.13 The Properties of a Group: The Alkali Metals Copyright © Cengage Learning. All rights reserved

Different Colored Fireworks Copyright © Cengage Learning. All rights reserved

Eletromagnetic radiation (The way that energy travels through space) Properties of Light Eletromagnetic radiation (The way that energy travels through space) (a) ex: Sun light, microwave. X-ray, radiant heat (b) Wavelike behavior: l n = c wavelength ( l ) : m frequency (v): s-1 (= hertz, Hz) velocity ( c ) : m/s

Classification of Electromagnetic Radiation Copyright © Cengage Learning. All rights reserved

Pickle Light Copyright © Cengage Learning. All rights reserved

Classic theory Quantum theory matter : particle energy is same mass as matter ! energy: continuous , wavelike

(1) Blackbody radiation radiation depend on Temp 3 paradoxes : (1) Blackbody radiation radiation depend on Temp Plank : energy is quantized (quanta)  only certain values allowed (2) Photoelectric effect Einstein : light has particulate behavior  photon (3) Atomic line spectra Bohr : energy of atoms is quantized photon emitted when electron change orbit

Observation : solid body (metal) is heat T : 750℃ T > 1200℃ Planck’s eqn (1900): Observation : solid body (metal) is heat T : 750℃ T > 1200℃ metal → dull red → brighter → brilliant white light Classical physics : atoms & molecules could emit or absorb any arbitrary amount of E  continue Planck proposal : energy , like matter , is discontinuous  quantum of energy & the energy DE=nhn n : positive integer h : 6.6210-34 JS  An atom can emit only certain amounts of energy E = hn , 2hn , 3hn ,  is not continuous but quantized

Einstein & the photoelectric effect (1905 (1) Photons : particles of light n > n0  photon current n < n0  no e－ ejected classical : energy associate intensity weak blue light & intense red light

b) light : dual nature Einstein & the photoelectric effect (1905) (2) Ephoton = hn = hc/l E = mc2  a) energy is quantized light wave photon (particle) mass Speed of light energy b) light : dual nature

Bohr’s postulation (for hydrogen atom) a) e moving around the nucleus in a circular orbit  Planetary model b) only a limited number of orbits with certain E are allow  orbits are quantized c) E of electrons in orbit  its distance from nucleus E = -2.178  10-18 (Z2 / n2) d) Electrons can pass from one allowed orbit to another. Fig. 2.9

lower E → higher E ni < nf DE > 0 absorption spectrum higher E → lower E ni > nf DE < 0 emission spectrum (fire works) Niel Bohr had tied the unseen (interior of the atom) to the seen (spectrum) But the model is only good for one e atom: H , He+ , Li2+

p. 70, Fig. 2-9

wave mechanics or quantum mechanics (A) Louis de Broglie (1982 – 1987) light wave photon How about matter ? matters have both wave & particle behavior 2pr = nl mnr = n(h/2p) h  l = ── mn wave properties particle properties

(B) Schrödinger’s model of H atom & wave function () = f (x, y, z) an electron in an atom could be described by equation for wave motion  wave function ()  characterize the e as a matter wave.

 E is quantized : only certain  are allowed & each  with allowed E. (2) Schrödinger’s theory choose to define the E of e precisely, i.e. can only describe the probability of electron.  E is quantized : only certain  are allowed & each  with allowed E.  electron density = probability of finding the e = 2  orbitals : specific wave functions for a given e. () The matter waves for the allow E states. orbits : Bohr’s model , was a path supposedly followed by the electron. Copyright © Cengage Learning. All rights reserved 18 18

h/2 ( h= h/2) (C) The uncertainty principle (1927) Heisenberg: It’s impossible to know simultaneously both the momentum & the position of a particle at a given time with certainty. only probability of finding an e with a given energy a given space.  (DX)(DP)  (DX)(Dmν)  h/4p h/2 ( h= h/2)

Probability Distribution for the 1s Wave Function Copyright © Cengage Learning. All rights reserved

Radial Probability Distribution Copyright © Cengage Learning. All rights reserved

(1)The principle quantum number (n) Quantum Numbers 3 quantum numbers are required to describe the distribution of e in atoms (1)The principle quantum number (n) (a) n = 1, 2, 3,……..∞ (shell) (b) related to the size & energy of the orbital. (c) the bigger the n, the larger the orbital, the less stable the orbital.

(2) The angular momentum quantum number ( l ) (a) l = 0,1, 2, 3,……., n－1 (subshells) (b) tell the orbital shapes or types. (c) (3) The magnetic quantum number （ml） a) ml = ﹣l , －l +1 , …, 0 , … , ( l-1 ) , l  ( 2l + 1 ) integral values b) relates to the orientation of the orbital in space l 1 2 3 4 name of orbital s p d f g

Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom Copyright © Cengage Learning. All rights reserved

Two Representations of the Hydrogen 1s, 2s, and 3s Orbitals Copyright © Cengage Learning. All rights reserved

The Boundary Surface Representations of All Three 2p Orbitals Copyright © Cengage Learning. All rights reserved

The Boundary Surfaces of All of the 3d Orbitals Copyright © Cengage Learning. All rights reserved

Representation of the 4f Orbitals in Terms of Their Boundary Surfaces Copyright © Cengage Learning. All rights reserved

(1) n, l, ml : define the orbital for an electron (2) for muti-electron atom : we need one more quantum number：electron spin ( ms ) ms = + ½ , －½ (3) the Pauli exclusion principle no two electrons in an atom can have the same set of four quantum numbers (n, l, ml, ms)  no atomic orbital can contain more than two electrons (opposite spins) He : 1s2 , (n, l, ml, ms) = (1, 0, 0, ½) (1, 0, 0, －½) 

(4) Paramagnetic : can be attracted by a magnetic fied  atoms contain upaired e. Diamagnetic : e spin are paired with partners  magnetic effects cancel out. odd / even e ？

(1) For polyelectronic atoms, we need electron configuration to understand electrons behavior. (2) electron configuration : how the electrons are distributes among the various atomic orbitals. (3) order of subshell E － depend on n & l a) E ↑ with “ n＋l ” value ↑ b) if same value of （n＋l）, then lower n lower E

(4) Effective nuclear charge (Zeff) a) polyelectronic atoms with two type of interactions nucleus － electron attraction, Zeff ↑ electron － electron repulsions, Z eff ↓ b) ∵ atomic E has a value, stronger attractions, lower E but repulsions, higher E -

(1) ground state electron configuration (2) Hund’s rule : the most stable arrangement of e in subshell (p, d, f) is that with the maximum number of unpaired e. H 1S1 1S  n l # of e in the orbital (subshell)

The Orbitals Being Filled for Elements in Various Parts of the Periodic Table Copyright © Cengage Learning. All rights reserved

The Values of First Ionization Energy for the Elements in the First Six Periods Copyright © Cengage Learning. All rights reserved

Atomic Radii for Selected Atoms Copyright © Cengage Learning. All rights reserved

Special Names for Groups in the Periodic Table Copyright © Cengage Learning. All rights reserved