Development of the Atomic Model Essay

Custom Student Mr. Teacher ENG 1001-04 14 January 2017

Development of the Atomic Model

460 – 370? BC – Democritus – first theory of atom

– All matter is composed of particles called atoms which can’t be subdivided – different materials had different properties because their atoms were different – atoms have different sizes, regular shape, are in constant motion, and have empty space

450 BC – Empedocles – matter is composed of four elements – earth, air, fire, water

384 – 322 Aristotle – no voids! Opposed Democritus’ theory – 4 elements earth, fire, air water with dry, hot, moist and cold

500 – 1600 A.D. – age of alchemy

Late 1700’s – law of conservation of mass – mass doesn’t change during a chemical reaction

1799 – Proust – law of constant composition – compounds always have same proportion by mass of their elements

1766-1844 John Dalton (English) postulates atoms as a billiard ball model
– all matter is made of particles called atoms
– all atoms of an element are identical
– atoms of different elements have different properties
– atoms combine to form compounds
– atoms are neither created nor destroyed during a chemical reaction

Late 1800’s – Sir William Crookes and others
– used sealed glass tubes to generate a glow
– Cathode rays were attracted to positive plates
– therefore negatively charged
– Rays could be blocked – therefore a particle

– Negatively charged particles were called electrons

1897 JJ Thomson – used cathode ray tube and developed raisin bun model – Electrons randomly distributed through positive mass – told not to touch – broke everything but could see what was wrong with equipment

1904 Hantaro Nagaoka – developed Saturn model

1911 Earnest Rutherford – Thomson’s research assistant – testing Thomson’s theory – gold foil experiment – surprised – like shooting a cannon ball at a piece of tissue paper and having the cannon ball bounce back at you! – Most of atom is empty space, positively charged nucleus – Electrons in a cloud around the nucleus

– had hands of gold and knew how to use them to get answers
– didn’t mention electrons because he didn’t know what they did
– he did know they weren’t in orbits because the energy degenerates and in the atom, it doesn’t

1886 – Goldstein – discovery of the proton (shown to be a fundamental particle 20 years later)
– 1837 times heavier than an electron

1932 James Chadwick – discovered neutrons by bombarding Be with alpha particles
– Gave off rays which weren’t deflected by outside fields
– Neutron had mass approximately equal to a proton

1900 Max Planck – energy is absorbed and released in chunks called quantum (compare playing a piano vs a violin)

Einstein – radiant energy – energy packets called photons ; described photoelectric effect from observing that radiant energy on metal releases electrons

1913 Niels Bohr (worked first with JJ Thomson then with Rutherford) developed model for hydrogen where the electron orbits the nucleus. – He explained
the H emission spectra and the explanation was the foundation for n, the principle quantum number – the concept of energy levels – Mathematical predictions of lines only worked for hydrogen – won a Nobel prize for looking at the solar system and comparing it to the atom

1924 Louis de Broglie showed that if radiant energy could act like a stream of particles, then matter could act like a wave – the wave property of electrons

1927 Werner Heisenberg – developed uncertainty principle – impossible to know both exact momentum and location of an electron due to dual nature of matter

1926 Erwin Schodinger – Schodinger’s wave equation
– quantum mechanics (advanced calculus needed) takes into account the wave and particle nature of electrons. – equation (2 gives info on location of electron in terms of probability density – wave functions are called orbitals – [pic], where E is energy, e2 is electric potential, r is orbital radius and h is Planck’s constant

1925 Wolfgang Pauli – each orbital has only 2 electrons is now explained due to direction of spin of electrons. Spinning electrons create magnetic field. Only 2 electrons of opposite spin in an orbital referred to as Pauli exclusion principle

Hund’s rule – half fill each orbital before adding second electron

Aufbau principle – energy sublevel must be filled before moving onto next higher sublevel

Principle Quantum Number, n
– integer that Bohr used to label the orbits and energy levels – a main shell of electrons
– seen in low resolution spectra
– still used today although we now use orbitals instead of orbits

Secondary Quantum Number, l
– Arnold Sommerfeld (1915) extended Bohr’s theory.
– H has 3 elliptical orbitals for n = 2
– Explained the observed line splitting seen for H in high resolution line spectra – Introduced l to describe sublevels
– l has values 0 to n-1
– relates energy levels to shape of electron orbital and explains regions of the periodic table – l=0, s orbital – sharp
– l=1, p orbital – principle
– l=2, d orbital – diffuse
– l=3, f orbital – fundamental

Magnetic Quantum Number, ml
– from experimentation with emission line spectra
– place a gas discharge tube near a strong external magnet, and some single lines split into new lines not initially seen – done by Pieter Zeeman in 1897 – called normal Zeeman Effect – Zeeman Effect explained by Sommerfeld and Peter Debye (1916) – They proposed that the orbits could exist at various angles – If orbits in space are in different planes, the energies of the orbits are different when the atom is near a strong magnet – For each value of l, ml can vary from –l to +l

– If l = 1, ml can be -1, 0, 1 suggesting 3 orbits with the same energy and shape but with a different orientation in space (degenerate orbitals)

Spin Quantum Number, ms
– to explain more and new evidence, ie the additional line splitting seen in a magnetic field – student of Bohr and Sommerfeld – Pauli – suggested each electron spins on its axis and is like a tiny magnet. – Could only have one of two spins equal in magnitude, opposite in direction (vector) – Values + ½ or – ½

– Opposite pair is a stable arrangement like bar magnets stored in pairs arranged opposite to each other (produce no magnetism) – If single unpaired electrons present, magnetism is present and atom is affected by magnetic fields

Overall – each electron in an atom is described by a set of 4 quantum numbers – fits perfectly arrangement of electrons and the structure of the periodic table

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