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**Review Models of the Atom**

Dalton proposes the indivisible unit of an element is the atom. Thomson discovers electrons, believed to reside within a sphere of uniform positive charge (the “plum-pudding model). Review Models of the Atom Rutherford demonstrates the existence of a positively charged nucleus that contains nearly all the mass of an atom. Bohr proposes fixed circular orbits around the nucleus for electrons. In the current model of the atom, electrons occupy regions of space (orbitals) around the nucleus determined by their energies. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

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**Rutherford’s Gold Foil Expt.**

Observation Inference Atom is mostly empty space. Positively charged part of atom must be small. Core of atom is very small, dense and positively charged. Most α particles pass through undeflected. Some α particles are slightly deflected. Few particles are reflected back (i.e. do not pass through).

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**Bohr Model and Emission Line Spectra of Hydrogen**

Observation 4 distinct bands of colour Inference e- can only posses certain energies (quantized) Certain wavelengths of light are absorbed by atom as e- transitions from lower to higher energy level (orbit). As e- returns to lower energy level, photons of light are released. If energy released corresponds to a wavelength of visible light, a specific band of colour will be observed.

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**Quantum Mechanical Model**

Niels Bohr & Albert Einstein Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).

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Electrons as Waves Energy is quantized. photons –> particles of light (E=hν Planck) Matter has wavelike properties. (deBroglie) Louis de Broglie ~1924 Standing waves n = 4 Standing Wave 200 150 100 50 - 50 -100 -150 -200 200 150 100 50 - 50 -100 -150 -200 Forbidden n = 3.3 Adapted from work by Christy Johannesson

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**Schrodinger’s Wave Equation**

Used to determine the probability of finding the electron at any given distance from the nucleus Orbital “electron cloud” -region in space where the electron is likely to be found 90% probability of finding the electron Orbital

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**Heisenberg Uncertainty Principle**

p. 200 Crazy Enough? Heisenberg Uncertainty Principle Impossible to know both the velocity and position of an electron at the same time Werner Heisenberg ~1926 g Microscope In order to observe an electron, one would need to hit it with photons having a very short wavelength. Short wavelength photons would have a high frequency and a great deal of energy. If one were to hit an electron, it would cause the motion and the speed of the electron to change. Lower energy photons would have a smaller effect but would not give precise information. Electron

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**Modern View The atom is mostly empty space Two regions Nucleus**

protons and neutrons Electron cloud region where you might find an electron

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Models of the Atom e + + - Dalton’s model (1803) Greek model (400 B.C.) Thomson’s plum-pudding model (1897) Rutherford’s model (1909) Bohr’s model (1913) Charge-cloud model (present) 1803 John Dalton pictures atoms as tiny, indestructible particles, with no internal structure. 1897 J.J. Thomson, a British scientist, discovers the electron, leading to his "plum-pudding" model. He pictures electrons embedded in a sphere of positive electric charge. 1911 New Zealander Ernest Rutherford states that an atom has a dense, positively charged nucleus. Electrons move randomly in the space around the nucleus. 1926 Erwin Schrödinger develops mathematical equations to describe the motion of electrons in atoms. His work leads to the electron cloud model. 1913 In Niels Bohr's model, the electrons move in spherical orbits at fixed distances from the nucleus. 1924 Frenchman Louis de Broglie proposes that moving particles like electrons have some properties of waves. Within a few years evidence is collected to support his idea. 1932 James Chadwick, a British physicist, confirms the existence of neutrons, which have no charge. Atomic nuclei contain neutrons and positively charged protons. 1904 Hantaro Nagaoka, a Japanese physicist, suggests that an atom has a central nucleus. Electrons move in orbits like the rings around Saturn. Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125

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**Homework Describe the current model of the atom.**

Describe the contributions of deBroglie, Schrodinger and Heisenberg to the development of the modern view of the atom.

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**ΔE is released as a photon of light.**

Explaining emission spectrum: ΔE is released as a photon of light.

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Sommerfeld (1915) -main lines of H spectrum are actually composed of more lines -introduced the idea of sublevels which is linked to the orbital shape

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Zeeman effect - Single lines split into multiple lines in the presence of a magnetic field. Orbits can exist at various angles and their energies are different when the atom is near a strong magnet.

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Stern-Gerlach experiment – a beam of neutral atoms is separated into two groups by passing them through a magnetic field

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**Orbitals s px pz py carbon y z x 2s 2p (x, y, z)**

Using Computational Chemistry to Explore Concepts in General Chemistry Mark Wirtz, Edward Ehrat, David L. Cedeno* Department of Chemistry, Illinois State University, Box 4160, Normal, IL Mark Wirtz, Edward Ehrat, David L. Cedeno*

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Y21s Y22s Y23s r r r r r r Distance from nucleus (a) 1s (b) 2s (c) 3s

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p-Orbitals px pz py Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 335

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**Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.**

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**Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.**

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d Orbitals Courtesy Christy Johannesson

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**Shapes of s, p, and d-Orbitals**

s orbital p orbitals d orbitals

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Atomic Orbitals

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**Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.**

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**Principal Energy Levels 1 and 2**

Note that p-orbital(s) have more energy than an s-orbital.

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**Homework Describe the shape of s and p orbitals. Define orbital.**

How does the size of the orbital change as n increases?

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**Quantum Numbers (used to represent different energy states)**

Pauli Exclusion Principle No two electrons in an atom can have the same 4 quantum numbers. Wolfgang Pauli 1. Principal # (n) 2. Secondary # (l) 3. Magnetic # (ml) 4. Spin # (ms) energy level sublevel (s,p,d,f), shape tilt or orientation of orbital electron spin Key expt work l hydrogen spectral lines made up of closely spaced lines ml splitting of spectral lines caused by a magnetic field ms magnetism Courtesy Christy Johannesson

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**Quantum Numbers 2s 2px 2py 2pz l = 0 (s) l = 1 (p) l = 2 (d) l = 3 (f)**

Courtesy Christy Johannesson

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**Quantum Numbers n (principal quantum #) n = 1,2,3…**

l (secondary quantum #) l = 0,1…n-1 ml (magnetic quantum #) ml = -l to +l ms (spin quantum #) ms = ½, -½

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**Quantum Numbers n = # of sublevels per level**

Principal level n = 1 n = 2 n = 3 Sublevel s s p s p d Orbital px py pz px py pz dxy dxz dyz dz2 dx2- y2 An abbreviated system with lowercase letters is used to denote the value of l for a particular subshell or orbital: l = Designation s p d f • The principal quantum number is named first, followed by the letter s, p, d, or f. • A 1s orbital has n = 1 and l = 0; a 2p subshell has n = 2 and l = 1(and contains three 2p orbitals, corresponding to ml = –1, 0, and +1); a 3d subshell has n = 3 and l = 2 (and contains five 3d orbitals, corresponding to ml = –2, –1, 0, –1, and +2). n = # of sublevels per level n2 = # of orbitals per level Sublevel sets: 1 s, 3 p, 5 d, 7 f Courtesy Christy Johannesson

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**Maximum Capacities of Subshells and Principal Shells**

n n l Subshell designation s s p s p d s p d f Orbitals in subshell An abbreviated system with lowercase letters is used to denote the value of l for a particular subshell or orbital: l = Designation s p d f • The principal quantum number is named first, followed by the letter s, p, d, or f. • A 1s orbital has n = 1 and l = 0; a 2p subshell has n = 2 and l = 1(and contains three 2p orbitals, corresponding to ml = –1, 0, and +1); a 3d subshell has n = 3 and l = 2 (and contains five 3d orbitals, corresponding to ml = –2, –1, 0, –1, and +2). Relationships between the quantum numbers and the number of subshells and orbitals are 1. each principal shell contains n subshells; – for n = 1, only a single subshell is possible (1s); for n = 2, there are two subshells (2s and 2p); for n = 3, there are three subshells (3s, 3p, and 3d); 2. each subshell contains 2l + 1 orbitals; – this means that all ns subshells contain a single s orbital, all np subshells contain three p orbitals, all nd subshells contain five d orbitals, and all nf subshells contain seven f orbitals. Subshell capacity Principal shell capacity n2 Hill, Petrucci, General Chemistry An Integrated Approach 1999, page 320

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**Practice Decide whether each set of quantum #s is allowed.**

n = 1 l = 0 ml = 0 ms = ½ n = 2 l = 1 ml = -1 ms =- ½ n = 3 l = 3 ml = 2 ms = ½ n = 0 l = 0 ml = 0 ms = ½ 2. Give the complete set of quantum #s for a lone e- in the third orbital of a 2p sublevel.

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Homework p. 184 #1-8

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The Quantum Mechanical Model of the Atom. Niels Bohr In 1913 Bohr used what had recently been discovered about energy to propose his planetary model of.

The Quantum Mechanical Model of the Atom. Niels Bohr In 1913 Bohr used what had recently been discovered about energy to propose his planetary model of.

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