Anatomy of an atom In the 1910s Ernest Rutherford and Niels Bohr devised a model for the atom. Electrons – electrically negative – surround a positive.

Anatomy of an atom

• In the 1910s Ernest Rutherford and Niels Bohr devised a model for the atom.

• Electrons – electrically negative – surround a positive nucleus.

•The nucleus contains protons (+) and (neutral) neutrons (James Chadwick 1932).

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Problems with the “Astronomical”model of the atom.

• Classical theory requires that a circulating charged particle emits radiation.

• So the circulating electron would lose energy continuously.

• And quickly spiral down into the nucleus.• Solution to the problem demands that

energy should not be emitted continuously.• Another opportunity for Quantum science!• If one young man could revolutionise

science, why not another?• Enter Niels Bohr.

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In the meantime....Bohr-Rutherford atom.• Niels Bohr (again) made the quantum hypothesis the

centre of his new theory (1913) of the structure of the atom.

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Bohr bolted together Newtonian dynamics and Quantum Theory.

Proposed that there are stable orbits – which do not spiral into the nucleus as classical physics demands.

Transitions between orbitals obey Planck’s law. Results agree with spectrum for Hydrogen.

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Incoming radiation causes ejection of an electron to a higher quantum orbit.

Electron falling to a lower quantum orbit releases radiation

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Beyond the Bohr atom

• Bohr’s hypothesis agreed with the spectrum of Hydrogen.

• H is the lightest atom - 1 proton in the nucleus and 1 orbiting electron,

• Bohr’s method failed for larger elements containing many electrons.

• But the key ingredient - “Quantisation of stable electron states and

• changes in energy was established.

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Louis de BroglieAfter detailed reflection on Einstein’s work,

“I suddenly had the idea, during the year 1923, that the discovery made by Einstein in 1905 should be generalised by extending it to all material particles and notably to electrons.”

E = mc2 (Einstein)

E = hν (Planck);

So simply coupling the two

equations we get:

E = hν = mc2

So electron “mass” has an

associated wave motion.

Particles become “Wavicles”22

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DeBroglie’s Harmonics

De Broglie had in mind wave patterns for a stretched string.• Rather than electrons occupying “orbits”- as Bohr had suggested, he thought that the lowest level electron state would be the fundamental mode of “vibration”.

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• Higher orbitals could be represented by harmonics.

De Broglie’s thesis supervisor and examiners had doubts about the validity of the work and consulted Einstein who was supportive: “I believe it is a first feeble ray of light on this worst of our physics enigmas.”

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Some reactions to De Broglie’s proposal: hν =E = mc2

• Roger Penrose in “The Emperor’s New Mind”:• “Thus, according to De Broglie’s proposal, the

dichotomy between particles and fields that had been a feature of classical theory is not respected by nature.

• Somehow, Nature contrives to build a consistent world in which particles and field-oscillations are the same thing.”

• Or, rather, her world consists of some more subtle ingredient, the words ‘particle’ and ‘wave’ conveying but partially appropriate pictures”.

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(Section of the) Periodic Table

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Alkalimetals

Inert gases

Halogens

• He, Ne , Ar, Kr, Xe are called “inert gases”.

• Why?? Because: they don’t react!

• Because: Helium has 2 electrons in a filled 1s Shell.

• Neon has 10 electrons in filled shells with principal quantum number 1,2 &3

• Filled shells are the most stable arrangement of electrons.

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Atoms - Where are the electrons?

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Electrons make elements

For Principal Q.N. N=2: there are up to 2 electrons in 2s,

and 2 electrons in each of three polar (2p) orbitals,

pointing in 3 directions at right angles. Total = 8

Shells fill up in an orderly manner: 1s before 2s before 2p, before 3s etc.

Atomic orbitals• Electrons surround nucleus in “orbitals” with well-defined “shapes”.

• Orbital shape and size is specified by Quantum numbers.

• Size: specified by the “Principal Quantum number”, N.

• A second Q.N. specifies the shape and symmetry:

• Shape can be labelled: s(spherical); & for N greater than 1: p (polar).

• For N=1, there are 2 elements:H (1s1), He (1s2)

• Each type of orbital can contain 2 electrons – and no more – as it was later found that electron can “spin” in 2 ways.

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Erwin Schrödinger (1887-1962)

Schrödinger set to work and over a Christmas holiday in the Alps, devised “Schrödinger’s Wave Equation”.This reproduced energies of atomic orbitals without Bohr’s arbitrary assumptions. The numbers fell out naturally as the ‘standing wave’ patterns.

Also, the shapes of the electron probability clouds closely resemble those presented in earlier work following Bohr’s model.

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Shapes of the first 5 atomic “orbitals”: 1s, 2s, 2px, 2py, 2pz

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• Cross-section views of electron probability clouds

Hydrogen lowest level H next higher 2s stateCarbon, 1s(orange), 2s (red) and one 2p (green)

2p orbitals

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“Ionic” Bonding of elements

• Ne, filled 1s and filled 2s & 2p shells (10 in all).

• H, Li & Na have 1 electron more than the filled shell configuration, which is easily removed so it is very reactive.

• F, Cl have one electron less than the filled shell and could gain one. So they too are very reactive.

• If, for example, Na donates its extra electron to Cl , then both achieve the stable filled shell “rare gas” configuration.

• The result is a positive Na+ ion and a negative Cl - ion

• So the compound Na+ Cl - is formed with an electrostatic “ionic” bond between the two ions.

Bonding by electron transfer - “IONIC” bonding

Chlorine (a few more inner electrons) is chemically similar to Fluorine.

Bonding by electron sharing

Also called “covalent” bonding

Electron sharing: “Covalent” bonds

Electron “shape” now wraps around all 3 atoms in water

Carbon bonds•Carbon has a filled 1s shell, 2 electrons in the 2s shell and 2 electrons in the 2p shell.

This suggests that carbon should form 2 bonds using 2p electrons.

•In fact, 2s and 2p electrons become more stable if they “hybridise” to form four sp3 bonds, directed towards the corners of a tetrahedron.

This now becomes the basis of the structure of diamond and much of organic chemistry.

“Hydrogen bonding” between water molecules

Hydrogen bonding in ice crystals

Pauli Exclusion principle (or why we don’t fall through the floor)

• Wolfgang Pauli: Austrian, worked with Bohr (1922-23);

• Famous for caustic criticism.• 1925 Proposed Exclusion Principle:

“Never two (or more) electronswith identical quantum numbers.”

• Two electrons can occupy the same atomic orbital (because electrons have 2 “spin” directions and therefore Q.N.s.)

• Vastly important - otherwise any number of electrons could crowd into any of the orbitals.

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Pauli’s Exclusion Principle• Key point is that ‘identical’ electrons cannot exist

together - in the same quantum orbits.

• By identical: having the same ‘Quantum numbers’.

• Quantum numbers specify: energy of the orbit, ‘shape’ (& ‘spin quantum number’ of the electron).

• Electrons enter the available orbits starting with the lowest energy level.

• Normally, only the outer orbits are involved in chemistry.

• If the Exclusion Principle did not apply, all the electrons would enter the lowest level so all atoms would be the same.

• Without the Pauli principle, “matter would collapse in on itself….” Roger Penrose.

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Pauli Continued• Without the Pauli Exclusion principle Chemistry -

formation of molecules, etc. - would not exist.

• An atom could pass through other atoms with full shells.

• We would fall through the floor!

Atoms in your foot

Atoms in the floor