Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Ch 7 Atomic Structure

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Rutherford Model

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3

4 Figure 7.2 Classification of Electromagnetic Radiation

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Electromagnetic Radiation Radiant energy that exhibits wavelength-like behavior and travels through space at the speed of light in a vacuum.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 Light is a wave Wavelength Distance between 2 similar points (meters) Frequency Number of waves in a second (frequency (s  1 ) or hertz (Hz)) Speed v (m/s) Light: energy that travels like a wave through space Wave properties:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Wavelength and Frequency

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Light is a wave All light travels the same speed: high, has short low, has long high = high energy

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 Light is a wave Created by movement of electric charge An electric field and magnetic field perpendicular to each Self-propagating

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Wavelength and frequency can be interconverted. = c/ (C =  = frequency (s  1 ) or hertz (Hz) = wavelength (m) c = speed of light (m s  1 ) ( x 10 8 m/s)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Figure 7.2 Classification of Electromagnetic Radiation

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Matter is not what it appears to be. Before 1900: Matter particle Light a wave Max Planck: not all energies were emitted from objects heated incandescence

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Planck’s Constant  E = change in energy, in J h = Planck’s constant,  10  34 J s = frequency, in s  1 = wavelength, in m Energy gained or lost only in whole number multiples. Transfer of energy is quantized: occur in discrete units, called quanta. Energy gained or lost only in whole number multiples. Transfer of energy is quantized: occur in discrete units, called quanta.   E = nh nhc  n = 1, 2,3,..

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 Light is a particle Einstein: theorizes that light made of photons. Gets Nobel Prize

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Photoelectric Effect

Copyright©2000 by Houghton Mifflin Company. All rights reserved High photon = high E photon - One photon hits one electron - If photon E not = to e - E nothing happens (even if bright) -Giant example Light made of photons

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 Photoelectric Effect

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 Light made of photons Einstein: electromagnetic radiation is quantized: E photon = h = hc

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 Energy and Mass Einstein’s special theory of relativity: (1905) Energy has mass: E = mc 2 Or m=E/c 2 E = energy m = mass c = speed of light

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 Energy and Mass (Hence the dual nature of light.) Does a photon have mass? for a photon with wavelength m = E = hc/  m  h c 2 c 2  c E photon = hc/ 

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 Figure 7.4 Dual Nature of light

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 If light can be a particle can a particle (e - ) be a wave? m = h /h / m  m = h /  v  h / m v = wavelength, in m h = Planck’s constant,  10  34 J s = kg m 2 s  1 m = mass, in kg v = velocity in m/s Louis de Broglie’s Equation 1920

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 Wave interference

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 25 Water wave interference

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Interference in water waves

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 Water interference

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 Interference patterns

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 Light interference

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 Interference Pattern

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 Electron interference Diffraction patterns caused by interference X-rays passing through NaCl crystal are diffracted. Electrons passing through NaCl crystal are diffracted x-rays

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 Diffraction using NaCl crystal

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 33 So, Debroglie was right: all matter show both wave like and particle like behavior

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 34 How does all this stuff relate to the e - and the atom? Bohr model: electrons are at set distances from the nucleus (energy levels)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 How do we know this?: we can use Einstein’s ideas to explain bright-line spectra Observe the light coming from the hydrogen emission tube. “Excited” atoms only emit certain frequencies (colors) of light Why not all frequencies of light? Look at hydrogen emission. Each line is one frequency of light.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 Figure 7.6 A Continuous Spectrum (a) and A Hydrogen Line Spectrum (b)

Absorption, emission, and energy Absorption photon Emission photon

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 Emission Spectrum The emitted photons are seen as light of specific frequency (i.e. colors ). What color emitted if electron could go anyway? A: the contiuous spectrum

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 What is the energy difference between levels? Must be equal to the energy of the photon emitted. Energy levels are quantized: ∆E = h  = h  c  

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40 Figure 7.7 A Change between Two Discrete Energy Levels

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 Figure 7.8 Electronic Transitions in the Bohr Model for the Hydrogen Atom

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 The Bohr Model E = energy of the levels in the H-atom z = nuclear charge (for H, z = 1) n = an integer The electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits. The energy of each level is given by this equation:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 43 The Bohr Model Ground State: The lowest possible energy state for an atom (n = 1).

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 Energy Changes in the Hydrogen Atom  E = E final state  E initial state

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 45 Standing waves in spring: - exist only at specific and (.5, 1, 1.5, 2.0, 2.5, ect.) - are quantized. Electron waves: - exist only at certain and (and energy) - only form certain distances from nucleus. - are quantized

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 46 Figure 7.10 The Hydrogen Electron Visualized as a Standing Wave Around the Nucleus

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 47 Quantum Mechanics Schrodinger’s equations: Based on the wave properties of the atom  = wave function = mathematical operator E = total energy of the atom A specific wave function is often called an orbital.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 48 What is the orbital for H when n=1? 1s orbital Orbitals are not the Bohr orbits. Where is the electron in the orbital?

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 49 Heisenberg There is a limit to how precisely we can both the position and momentum of a an electron at a given time

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 50 Heisenberg Uncertainty Principle x = position mv = momentum h = Planck’s constant The more accurately we know a particle’s position, the less accurately we can know its momentum.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 51 What is an orbital? If we don’t know the motion of an electron, what is an orbital? 4 square of the wave function gives the probability of finding an electron at a given position. --> (a) probability distrib. for H 1s orbital

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 52 Figure 7.12 Radial Probability Distribution (the probability distribution in each spherical shell.)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 53 Bohr radius Turns out that for H 1s orbital, the max radial probability is 5.29x10 -2 nm, = Bohr’s innermost “orbit”.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 54 How big is the 1s? Probab. Decreases with radius, but never goes to zero. Size definition: Size of the orbital is the radius of the sphere that encloses 90% of the electron’s probability.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 55 Quantum Numbers (QN) Schrodinger equation has many wave function solutions. Each are described by quantum numbers 1.Principal QN (n = 1, 2, 3,...) - related to size and energy of the orbital. (this gives the”rings” or “shells) 2. Angular Momentum QN (l = 0 to n  1) - relates to the shape of the orbital. (ex. s p d f also called subshells) 3.Magnetic QN (m l = l to  l ) - relates to orientation of the orbital in space relative to other orbitals. Gives you the number of each type of orbital. (ex.: p x p y p z ) 4. Electron Spin QN (m s = + 1 / 2,  1 / 2 ) - relates to the spin states of the electrons. (see pg 310, table 7.2)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 56 Orbitals when n=1 (principal quantum #): ( l = 0 to n-1) (m l = l to - l ) : l =0 m l =0 1s orbital only. Only 1 or 2 electrons are described by an orbital. Total electrons at n=1? :.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 57 When n= 2, ( l = 0 to n-1) m l = l to - l ): l = 0 and l = 1 When l = 0, m l = 0 2s orbital When l = 1 m l = -1, 0, 1 giving three 2p orbitals: 2p x 2p y 2p z Total orbitals: Total electrons:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 58 The P orbitals (energy level 2 and up)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 59 When n=3 ( l = 0 to n-1) m l = l to - l ): l = 0, 1, 2 When l = 0, m l = 0 giving one 3s orbital When l = 1 m l = -1, 0, 1 giving three 3p orbitals: 3p x 3p y 3p z When l = 2 m l = -2,-1, 0, 1,2 giving 5 d orbitals Total orbitals level 3: Total electrons level 3:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 60 When n=4 l = 0, 1, 2, 3 when l = 0, 1, 2 : 4s 4p x 4p y 4p z 4 d’s (5 of them) when l =3 4f orbitals (7 of them) Total orbitals: Total electrons:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 61 electron spin quantum number: m s = +1/2 or - 1/2

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 62 Pauli Exclusion Principle In a given atom, no two electrons can have the same set of four quantum numbers (n, l, m l, m s ). Therefore, an orbital can hold only two electrons, and they must have opposite spins.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 63 Quantum Model static.howstuffworks.com/ gif/atom-quantum.jpg

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 64 Hydrogen orbitals are degenerate All H orbitals with the same n have the same energy.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 65 Polyelectronic Atoms How does it work after Hydrogen? Shielding (e - repel, feel less attraction to +) Hydrogen orbitals: degenerate Hydrogenlike orbitals: NOT degenerate! E ns < E np < E nd < E nf …etc. Why?

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 66 Figure 7.20 A Comparison of the Radial Probability Distributions of the 2s and 2p Orbitals 2p appears closer to nucleus? Less energy? No, look at small 2s hump. “2s penetrates to the nucleus” penetration effect

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 67 Figure 7.21 The Radial Probability Distribution for the 3s, 3p, and 3d Orbitals so, E 3s <E 3p <E 3d

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 68 History of the Periodic Table 7.10 Dobereiner: triads Newlands: octaves Meyer / Mendeleev: arrangements by atomic masses. Theory → prediction

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 69 Figure 7.23 Mendeleev’s Early Periodic Table, Published in 1872 Prediction: Ga, Ge (see table 7.3, pg318 to see how cool Mendeleev was)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 70 Aufbau Principle As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals. “ an electron occupies the lowest energy orbital that can receive it”

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 71 Hund’s Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals. “in the p, d, f, orbitals, spread out before you pair up”

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 72 The Electron Configurations in the Type of Orbital Occupied Last for the First 18 Elements Note: elements in same group have number valence electrons The electrons in the outermost principle quantum level of (valence: The electrons in the outermost principle quantum level of an atom. an atom. ( (Core electron: other than valence) Ex:Cl 1s 2 2s 2 2p 6 3s 2 3p 5 or [Ne] 3s 2 3p 5 # valence = ?

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 73 Figure 7.25 Electron Configurations for Potassium Through Krypton Why doe the 4s fill before the 3d? Penetration effect Notice Cr, Cu columns.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 74 Figure 7.26 The Orbitals Being Filled for Elements in Various Parts of the Periodic Table After lathanum [Xe] 6s 2 5d 1, go to lathanide series, fill the 4fs (fig 7.27: note anomalies) After Actinium [Rn]7s 2 6d 1, fill actinide series with 5fs.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 75 Figure 7.27 The Periodic Table With Atomic Symbols, Atomic Numbers, and Partial Electron Configurations

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 76 Figure 7.36 Special Names for Groups in the Periodic Table Main-group elements or representative elements: 1A 2A 3A 4A 5A 6A 7A 8A or 1,2, M-g e: each group has same #valence

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 77 Broad Periodic Table Classifications Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O) Transition Elements: filling d orbitals (Fe, Co, Ni) Lanthanide and Actinide Series (inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 78 Figure 7.30 The Positions of the Elements Considered in Sample Exercise 7.7

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 79 Ionization Energy The quantity of energy required to remove one electron from the gaseous atom or ion. X (g) → X + (g) + e -

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 80 Periodic Trends First ionization energy: increases from left to right across a period (increase +, no shielding) decreases going down a group. ( increase n, more shielding)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 81 Figure 7.31 The Values of First Ionization Energy for the Elements in the First Six Periods

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 82 Trends in Ionization Energies for the Representative Elements Who has the highest IE? Who has the lowest? Do metals or nonmetals have higher IE? What does IE tell us about metal reactivity

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 83 Periodic Trends: IE Al(g) --> Al + + e - I 1 = 580 kJ/mol Al + (g) --> Al 2+ + e - I 2 = 1815 kJ/mol Al 2+ (g) --> Al 3+ + e - I 3 = 2740 kJ/mol Al 3+ (g) --> Al 4+ + e - I 4 = 11,600 kJ/mol Why the differences? Indicates something about the electron structure.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 84 Practice: 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 Which atom has the largest first I.E.? Which one has the smallest second I.E.? Explain.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 85 More Practice IE increases across the period. Check IE of P and S on pg 329. Explain the anomaly

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 86 Electron Affinity The energy change associated with the addition of an electron to a gaseous atom or ion. X(g) + e   X  (g) If change is exothermic, then E.A. is negative.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 87 Figure 7.33 The Electronic Affinity Values for Atoms Among the First 20 Elements that Form Stable, Isolated X - Ions

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 88 Electron Affinity Left to right mostly more neg kJ, more energy released. (note the missing elements: why C, but not N? write o.d. for both and t.t.y.n) Down a group, usually less neg kJ, less energy released. (see table 7.7. Notice anomaly. T.t.y.n.)

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 89 E.A. Who has more neg EA, metals or nonmetals? What does this say about the reactivity of nonmetals?

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 90 Radius trend Data usually from distance between nuclei in a compound.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 91 Radius trend Data usually from distance between nuclei in a compound.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 92 Figure 7.35 Atomic Radii for Selected Atoms Period trend? Group trend? Why?

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 93 Periodic Trends Atomic Radii: decrease going from left to right across a period: increase + draws in valence. increase going down a group: increase in orbital sizes with n.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 94 Alkali Metals Trends look at pg 335. Note trends down the group in: 1. IE, radius 2. Density. Why? 3. mp/bp

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 95 Alkali metals Reaction with water. Write the balanced equation for Na(s) with water. List alkali from most to least reactive (think I.E.) Cs > Rb > K > Na > Li But in water, Li > K > Na, even though K loses electrons the easiest. why? Hydration energy (see table 7.9) Li is small, higher charge density, better at attracting water. But there is more:

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 96 Alkali metals What we observe when they react with water: K > Na > Li Look at mp. K and Na melt, increasing reaction rate.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 97 Waves Waves have 3 primary characteristics: 1.Wavelength: distance between two peaks in a wave. 2.Frequency: number of waves per second that pass a given point in space. 3.Speed: speed of light is 3.00  10 8 m/s.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 98 Figure 7.1 The Nature of Waves