Review Unit 1 (Chp 6,7): Atoms, Electrons, & Periodicity Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Review Unit 1 (Chp 6,7): Atoms, Electrons, & Periodicity John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

Development of Atomic Models 1803 Dalton Atomic Theory 1904 Thomson Plum Pudding + 1911 Rutherford Nuclear Model These illustrations show how the atomic model has changed as scientists learned more about the atom’s structure.

same # of protons (& electrons), but different # of neutrons Isotopes element: same or different mass: same or different why? same # of protons (& electrons), but different # of neutrons 1 H 2 1 H 3 1 H protium deuterium tritium

Average Atomic Mass Avg. Mass = (Mass1)(%) + (Mass2)(%) … average atomic mass: calculated as a weighted average of isotopes by their relative abundances. lithium-6 (6.015 amu), which has a relative abundance of 7.50%, and lithium-7 (7.016 amu), which has a relative abundance of 92.5%. amu = 1/12 mass of C-12 1amu = 1.66053X10^-24 grams (6.015)(0.0750) + (7.016)(0.925) = 6.94 amu Avg. Mass = (Mass1)(%) + (Mass2)(%) …

isotopes separated by difference in mass Mass Spectrometry isotopes separated by difference in mass

Development of Atomic Models 1803 Dalton Atomic Theory 1904 Thomson Plum Pudding + What evidence? 1911 Rutherford Nuclear Model These illustrations show how the atomic model has changed as scientists learned more about the atom’s structure. 1913 Bohr Shell Model 1926 Quantum Mechanical Model

Atomic Emission Spectra 5.3 Atomic Emission Spectra elements give discrete lines of E & f. (only specific colors of energy & frequency)

Bohr’s Shell Model (1913–Niels Bohr) electrons occupy only specific levels (or shells) of “quantized” energy (& wavelength & frequency) Electrons as Waves quantized into specific multiples of wavelengths, but none in between.

Bohr’s Shell Model ∆E EXCITED state e–’s emit (–) energy, move back to inner levels (n=5 to n=2) e–’s absorb (+) energy, move to outer levels (n=2 to n=5) GROUNDstate 5 4 3 2 ∆E 2 2 Which transition shows a light wave of the greatest energy? n=5 to n=2

Electromagnetic Spectrum Y G B I V Electromagnetic Spectrum Lowest Energy Highest Energy (higher ) (shorter ) All EM radiation travels at the speed of light (c), 2.998  108 m/s. c =  E = h (of 1 photon)

+ 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 nucleus Aufbau: Fill lowest energy orbitals first. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 Hund: 1 e– in equal orbitals before pairing () (3d fills after 4s) ? Pauli Exclusion: -Where should we start placing electrons first? -Opposite spin alleviates repulsion. -All single before double. no e–’s same props (opp. spin) (↑↓) nucleus +

species are attracted by a magnet (caused by unpaired electrons). Paramagnetic: species are attracted by a magnet (caused by unpaired electrons). Fe: [Ar] ↑↓ ↑↓ ↑ ↑ ↑ ↑ 4s 3d Diamagnetic: species are slightly repelled by magnets (caused by all paired electrons) Zn: [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

d block metals lose their outer s electrons before any core d electrons to form ions. Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Fe2+ 1s2 2s2 2p6 3s2 3p6 3d6 Fe3+ 1s2 2s2 2p6 3s2 3p6 3d5 d block (trans. metals) have colored ions due to light excited e– movement in d orbitals

Other Aspects Ar, K+, Ca2+, S2–, Cl– List 3 species isoelectronic with Ca2+ & S2–. P3– , Cl– , Ar, K+ , Sc3+ , Ti4+, V5+, Cr6+, Mn7+ Arrange the following species by increasing size: Ar, K+, Ca2+, S2–, Cl– Ca2+ < K+ < Ar < Cl– < S2–

SPECTROSCOPIC TECHNIQUE Spectroscopy SPECTROSCOPIC TECHNIQUE EM REGION APPLICATION Molecular Structure by molecular Rotation Microwave Microwave Types of bonds by bond Vibration IR Infrared Vis/UV Atomic Emission Spectra (lines of frequencies/colors) Visible & Ultraviolet Transition of e–’s between energy levels Ionization of e–’s shows e– configuration PES (Photoelectron Spectroscopy) X-ray WATCH this 6 min Video Explanation of PES at HOME.

Photoelectron Spectroscopy (PES) Which peak is H and which is He? higher peak = more e–’s 1s2 Relative # of e–’s He 1s1 H 6 5 4 3 2 1 0 Binding Energy ...or Ionization Energy (required to remove e–’s) (MJ/mol)  further left = more energy required (stronger attraction due to more protons)

Photoelectron Spectroscopy (PES) Which peak is H and which is He? 2p6 higher peak = more e–’s ? Ne 1s2 Identify the element & e-config Relative # of e–’s He 1s1 1s2 2s2 H 6 5 4 3 2 1 0 Binding Energy ...or Ionization Energy (required to remove e–’s) (MJ/mol)  further left = more energy required (stronger attraction due to more protons)

PES (A) PES (B) 3d10 2p6 3p6 1s2 2s2 3s2 4s2 4p2 Ge n = 1 n = 2 n = 3 Identify element (A) 3d10 2p6 3p6 1s2 2s2 3s2 4s2 4p2 Ge n = 1 n = 2 n = 3 n = 4 WS #1,5 PES (B) Identify element (B) 4s1 ? K

Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Review Unit 1 (Chp 7): Periodicity …or… Periodic Trends in Atomic Properties John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

(explains all periodic trends and properties) We will explain observed trends in Atomic (and Ionic) Radius Ionization energy Electronegativity size lose e– attract e– Zeff & shielding (explains all periodic trends and properties)

Zeff & Shielding attraction shielding Zeff Na atom effective nuclear charge, (Zeff): Zeff = Z − S Z = nuclear charge (+proton’s) S = shielding (core e–’s) attraction shielding Zeff shielding, (S): inner core e–’s shield valence e–’s from nuclear attraction. Z = +11 +11 Zeff = +1 Na atom

decreases across a period Atomic Radius att. =shield Zeff decreases across a period -due to increasing Zeff (more protons) -due to increasing shielding (more energy levels) increases down a group att. shield =Zeff

Ionic Radius Why? Cations are smaller than atoms. loses a shell core shell closer to nucleus new valence e– ‘s less shielded (greater Zeff) Anions are larger than atoms. electrons are added and repulsions are increased (same Zeff & same shielding)

Ionization Energy (IE) energy required to remove an electron more energy to remove each electron IE1 < IE2 < IE3, … look for a huge jump in IE once all valence e–’s are removed, the next e– is on an inner level with attraction (shielding & Zeff). huge jump in IE4 b/c 4th e– on inner level (must have 3 valence e–’s)

increases across a period Trends in First IE IE tends to… -due to increasing Zeff (more protons) increases across a period att. =shield Zeff -due to increasing shielding (more energy levels) decreases down a group att. shield =Zeff

Trends in Electronegativity (EN) -ability of an atom to attract electrons when bonded with another atom. increases across a period -greater Zeff -more shielding (more energy levels) decreases down a group

Periodic Trends (Summary) Electronegativity Can you explain all of this in terms of p’s and e’s? Zeff & shielding Electronegativity WS #2,3 Atomic radius