Thompson’s Experiment

Rutherford’s Experiment:

Explanation

Diffraction:- the change in direction of a wave as it passes the edge of an object

The Wave Nature of Light Light as electromagnetic waves: polarization, interference, diffraction, reflection, and refraction

Electromagnetic Spectrum

Atomic Line Spectra (Line Emission Spectra) Hydrogen has the simplest atomic emission spectrum ( 1880s).

The Spectrum of Atomic Hydrogen Compare the absorption spectrum to the emission spectrum of H If we pass light from a continuous source (eg from a hot object) through a substance, then the continuous spectrum has some of the wavelengths removed! -- absorption spectrum

Light as Particles

Bohr Atom

Hidrojen İçin Bohr Atom Modeli (1913) Elektronlar belirli yörüngelerde bulunabilirler

Bohr Model of Hydrogen Atoms Assumptions Quantized energy levels – total energy for each level is the sum of the kinetic energy of the electron plus its potential energy. Electrons do not radiate while in their orbits, but do when they move from one orbit to another.

Energy E = 0 eV Paschen Series (IR) n = 3 Balmer Series (visible) Lyman Series (ultraviolet) E1 = -13.6 eV n = 1 Lyman Balmer Paschen Example Data

Line Spectra The Lyman and Balmer series of lines in the hydrogen spectrum correspond to transitions that the electrons make between higher and lower energy levels. The Bohr model only has one quantum number, n, which represents the energy level. http://www.walter-fendt.de/ph11e/bohrh.htm

Electron in the Hydrogen Atom

Elektronlar gerçekten de çekirdek etrafında belirli yörüngelerde mi dönerler? Bunu tespit etmek mümkün mü?

Elektronun konumu ve momentumunun ölçülmesi Işının her çarpışında elektronun da konumu değişir. Bu sebeple ışıkla bir elektronunun konumu ve hızı hassas bir şekilde belirlenemez.

Heisenberg, Werner 1901–76, Alman Fizikçi 1932 Nobel Fizik Ödülü Eğer bir taneciğin nerede olduğunu kesin olarak biliyorsak, aynı anda taneciğin nereden geldiğini ve nereye gittiğini kesin şekilde bilemeyiz.

Peki elektronlar nerede ve nasıl hareket ediyorlar?

Erwin Schrödinger 1927 yılında Elektronlar, zamanlarının büyük bir çoğunluğunu orbital denen bölgelerde geçirirler. Değişik şekillerde orbitaller mevcuttur.

Atomik Orbitaller Şekli Orbital sayısı e sayısı s küresel 1 2 p halter 3 6 d karışık 5 10 f karışık 7 14 Herbir orbital 2 elektron içerir

Quantum Mechanics

The First Shell The innermost shell (1) only contains an s orbital.   1s

Atomic Orbitals, s-type

S orbitalleri

Atomic Structure

Quantum Mechanics

Electron Configuration in p Orbital

Atomic Orbitals, p-type

p Orbitals Rather than being a sphere, the "p" orbital has two lobes pointed in opposite direction away from the nucleus.   One p orbital points along each the x, y, and z axis. There are three p orbtitals in every shell except the first.

The Second Shell 2py 2px 2pz 2s The second shell contains an s type orbital as well as a new kind of orbital called a "p" orbital. 2py 2px 2pz 2s

Atomic Orbitals, d-type

d Orbitals There are 5 types of d orbitals. Four of the five have four lobes at 90o to one another. The fifth looks like a donut around a p-orbital Image from: HMChem

The Third Shell The 3rd shell has: 1 s orbital (3s) 3 p orbitals (3px, 3py, 3pz) 5 d orbitals (3dxy, 3dyz, 3dxz, 3dx2-y2, 3dz2)

f orbitals f-orbitals have 6 lobes and are very challenging to envision Image from: HMChem

The First Two Shells Picture from and more info on How Atoms Work

The Fouth Shell The 4th shell has: 1 s orbital (4s) 3 p orbitals (4px, 4py, 4pz) 5 d orbitals (4dxy, 4dyz, 4dxz, 4dx2-y2, 4dz2) 7 f orbitals What pattern do you see in the relationship between shells and orbitals? Electrons are rarely if ever observed in orbitals beyond f so, we most chemists are not too worried about them.

Principle quantum number n = 1, 2, 3,….. describes orbital size and energy Angular momentum quantum number l = 0 to n-1 describes orbital shape Magnetic quantum number ml = l, l-1…-l describes orientation in space of the orbital relative to the other orbitals in the atom Spin quantum number ms = +1/2 or -1/2 describes the direction of spin of the e- on its axis Pauli Exclusion Principle: "no two electrons in an atom can have the same set of quantum numbers", or, only two electrons (of opposite spin) per orbital.

Write a valid set of quantum numbers for each of the following sub-shells: (a) 2 s n = 2, l = 0, ml = 0, ms = - 1/2 n = 2, l = 0, ml = 0, ms = ± 1/2 2 combinations

Write a valid set of quantum numbers for each of the following sub-shells: (a) 2 s n = 2, l = 0, ml = 0, ms = - 1/2 n = 2, l = 0, ml = 0, ms = ± 1/2 2 combinations (b) 2 p n = 2, l = 1, ml = -1, ms = - 1/2 n = 2, l = 1, ml = -1, 0 or 1, ms = ± 1/2 6 combinations

Write a valid set of quantum numbers for each of the following sub-shells: (a) 2 s n = 2, l = 0, ml = 0, ms = - 1/2 n = 2, l = 0, ml = 0, ms = ± 1/2 2 combinations (b) 2 p n = 2, l = 1, ml = -1, ms = - 1/2 n = 2, l = 1, ml = -1, 0 or 1, ms = ± 1/2 6 combinations (c) 3 d n = 3, l = 2, ml = -2, ms = - 1/2 n = 3, l = 2, ml = -2, -1, 0, 1, or 2, ms = ± 1/2 10 combinations

How many orbitals in a subshell? l = 1, px, py, pz 3 l = 2, dxy,, dxz,, dyz ,, dx2-y2, dz2 5

How many orbitals in a subshell? l = 1, px, py, pz 3 l = 2, dxy,, dxz,, dyz ,, dx2-y2, dz2 5 2 l + 1 orbitals per subshell

How many orbitals in a subshell? l = 1, px, py, pz 3 l = 2, dxy,, dxz,, dyz ,, dx2-y2, dz2 5 2 l + 1 orbitals per subshell How many orbitals in a shell? n = 1, 1s 1 n = 2, 2s, 2px, 2py, 2pz 4 n = 3, 3s, 3px, 3py, 3pz, 3dxy,, 3dxz,, 3dyz ,, 3dx2-y2, 3dz2 9

How many orbitals in a subshell? l = 1, px, py, pz 3 l = 2, dxy,, dxz,, dyz ,, dx2-y2, dz2 5 2 l + 1 orbitals per subshell How many orbitals in a shell? n = 1, 1s 1 n = 2, 2s, 2px, 2py, 2pz 4 n = 3, 3s, 3px, 3py, 3pz, 3dxy,, 3dxz,, 3dyz ,, 3dx2-y2, 3dz2 9 n2 orbitals per principal quantum level

Hydrogen atom- all orbitals within a shell have the same energy electrostatic interaction between e- and proton

Hydrogen atom- all orbitals within a shell have the same energy electrostatic interaction between e- and proton Multi-electron atoms- the energy level of an orbital depends not only on the shell but also on the subshell electrostatic interactions between e- and proton and other e-

Orbital Energies 3dxy 3dxz 3dyz 3dx2-y2 3dz2 3px 3py 3pz 3s Energy 2px 2py 2pz 2s 1s

Electronic Configuration: Filling-in of Atomic Orbitals Rules: 1. Pauli Principle

Electronic Configuration: Filling-in of Atomic Orbitals Rules: 1. Pauli Principle 2. Fill in e-'s from lowest energy orbital upwards (Aufbau Principle)

Electronic Configuration: Filling-in of Atomic Orbitals Rules: 1. Pauli Principle 2. Fill in e-'s from lowest energy orbital upwards (Aufbau Principle) 3. Try to attain maximum number of unpaired e- spins in a given sub-shell (Hund's Rule)

Electronic Configuration: Filling-in of Atomic Orbitals Rules: 1. Pauli Principle 2. Fill in e-'s from lowest energy orbital upwards (Aufbau Principle) 3. Try to attain maximum number of unpaired e- spins in a given sub-shell (Hund's Rule) H (Z = 1) 1s1 2s 2p Energy 1s

Electronic Configuration: Filling-in of Atomic Orbitals Rules: 1. Pauli Principle 2. Fill in e-'s from lowest energy orbital upwards (Aufbau Principle) 3. Try to attain maximum number of unpaired e- spins in a given sub-shell (Hund's Rule) N (Z = 7) 1s2, 2s2, 2p3, 2p 2s Energy 1s

Electronic Configuration: Filling-in of Atomic Orbitals Rules: 1. Pauli Principle 2. Fill in e-'s from lowest energy orbital upwards (Aufbau Principle) 3. Try to attain maximum number of unpaired e- spins in a given sub-shell (Hund's Rule) B (Z = 5) 1s2, 2s2, 2p1 2p 2s Energy 1s

Electronic Configuration: Filling-in of Atomic Orbitals Rules: 1. Pauli Principle 2. Fill in e-'s from lowest energy orbital upwards (Aufbau Principle) 3. Try to attain maximum number of unpaired e- spins in a given sub-shell (Hund's Rule) F (Z = 9) 1s2, 2s2, 2p5 2p 2s Energy 1s

Hydrogen 2s 3s 4s 1s 2p 3p 4p 3d 4d 4f Multi-electron atoms 1s 2s 3s 4s 5 s 2p 3p 4p

1s 2s 2px 2py 2pz H 1s1 He 1s2 Li 1s2, 2s1 Be 1s2, 2s2 B 1s2, 2s2, 2px1 C 1s2, 2s2, 2px1, 2py1 N 1s2, 2s2, 2px1, 2py1, 2pz1 O 1s2, 2s2, 2px2, 2py1, 2pz1 F 1s2, 2s2, 2px2, 2py2, 2pz1 Ne 1s2, 2s2, 2px2, 2py2, 2pz2

H 1s1 He 1s2 Li [He], 2s1 Be [He], 2s2

H 1s1 He 1s2 Li [He], 2s1 Be [He], 2s2 B [He], 2s2, 2p1 Ne [He], 2s2, 2p6 Na [He], 2s2, 2p6, 3s1  [Ne], 3s1

H 1s1 He 1s2 Li [He], 2s1 Be [He], 2s2 B [He], 2s2, 2p1 Ne [He], 2s2, 2p6 Na [He], 2s2, 2p6, 3s1  [Ne], 3s1 Mg [He], 2s2, 2p6, 3s2  [Ne], 3s2 Al [Ne], 3s2, 3p1 Si [Ne], 3s2, 3p2

H 1s1 He 1s2 Li [He], 2s1 Be [He], 2s2 B [He], 2s2, 2p1 Ne [He], 2s2, 2p6 Na [He], 2s2, 2p6, 3s1  [Ne], 3s1 Mg [He], 2s2, 2p6, 3s2  [Ne], 3s2 Al [Ne], 3s2, 3p1 Si [Ne], 3s2, 3p2 P [Ne], 3s2, 3p3 S [Ne], 3s2, 3p4 Cl [Ne], 3s2, 3p5 Ar [Ne], 3s2, 3p6

H 1s1 He 1s2 Li [He], 2s1 Be [He], 2s2 B [He], 2s2, 2p1 Ne [He], 2s2, 2p6 Na [He], 2s2, 2p6, 3s1  [Ne], 3s1 Mg [He], 2s2, 2p6, 3s2  [Ne], 3s2 Al [Ne], 3s2, 3p1 Si [Ne], 3s2, 3p2 P [Ne], 3s2, 3p3 S [Ne], 3s2, 3p4 Cl [Ne], 3s2, 3p5 Ar [Ne], 3s2, 3p6 outermost shell - valence shell most loosely held electron and are the most important in determining an element’s properties

K [Ar], 4s1 Ca [Ar], 4s2 Sc [Ar], 4s2, 3d1 Ti [Ar], 4s2, 3d2

K [Ar], 4s1 Ca [Ar], 4s2 Sc [Ar], 4s2, 3d1 Ca [Ar], 4s2, 3d2 Zn [Ar], 4s2, 3d10 Ga [Ar], 4s2, 3d10, 3p1 Kr [Ar], 4s2, 3d10, 3p6

K [Ar], 4s1 Ca [Ar], 4s2 Sc [Ar], 4s2, 3d1 Ca [Ar], 4s2, 3d2 Zn [Ar], 4s2, 3d10 Ga [Ar], 4s2, 3d10, 3p1 Kr [Ar], 4s2, 3d10, 3p6 Anomalous electron configurations d5 and d10 are lower in energy than expected Cr [Ar], 4s1, 3d5 not [Ar], 4s2, 3d4 Cu [Ar], 4s1, 3d10 not [Ar], 4s2, 3d9

Electron Configuration of Ions Electrons lost from the highest energy occupied orbital of the donor and placed into the lowest unoccupied orbital of the acceptor (placed according to the Aufbau principle)

Electron Configuration of Ions Electrons lost from the highest energy occupied orbital of the donor and placed into the lowest unoccupied orbital of the acceptor (placed according to the Aufbau principle) Examples: Na [Ne], 3s1 Na+ [Ne] + e- Cl [Ne], 3s2, 3p5 + e- Cl- [Ne], 3s2, 3p6 Mg [Ne], 3s2 Mg2+ [Ne] O [He], 2s2, 2p4 O2- [He], 2s2, 2p6

Modern Theories of the Atom - Summary Wave-particle duality of light and matter Bohr theory Quantum (wave) mechanical model Orbital shapes and energies Quantum numbers Electronic configuration in atoms