Unit 3 History of the Atom

Democritus about 460 bc stated that the “ultimate particle” is atmos He believed that atoms were indivisible and indestructible was not based on the scientific method – but just philosophy

John Dalton - 1803 English school teacher the atom was a small, indivisible particle like a “small marble”

Dalton’s Atomic Theory (experiment based!) All matter is made of tiny indivisible particles called atoms Atoms of the same element are chemically the same. Atoms of different elements are chemically different. John Dalton (1766 – 1844) Individual atoms of an element have slightly different masses. We use average mass. Different elements have different ave masses Atoms can’t be divided in normal chemical reactions.

Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron

Modern Cathode Ray Tubes Television Computer Monitor Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

Thomson’s Atomic Model J. J. Thomson Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

Ernest Rutherford’s Gold Foil Experiment - 1911 Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil Particles that hit on the detecting screen (film) are recorded

Rutherford’s Findings Most of the particles passed right through A few particles were deflected VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions: The nucleus is small The nucleus is dense The nucleus is positively charged

The Rutherford Atomic Model Based on his experimental evidence: atom - mostly empty space all positive charge, and most mass is in the center called a “nucleus” nucleus = protons and neutrons (neutrons confirmed by Chadwick in 1932) electrons distributed around the nucleus, and occupy most of the volume His model was called a “nuclear model”

Niels Bohr - 1913 Niels Bohr wondered why horseshoes heated in a forge changed color, but did not react Nucleus? Electrons? http://www.askaboutireland.ie/_internal/gxml!0/m6s6aicjogoampxnh1a9dzzq7k4g1nn$o0srg946q1ltpizssbw14g5qaft4vmx

Neils Bohr All atoms contain energy – the energy has to do with the electrons 1913 Bohr said that electrons are in definite areas with definite amounts of energy

Planetary Model Electrons are in the “ground state” – normal lowest energy state *excited state electrons gain energy in fixed amounts (photons) and go to higher levels

So, Energy = Planck’s constant x frequency E = hf E = energy h = Planck’s constant f = frequency So, Energy = Planck’s constant x frequency

Electrons don’t stay in the excited state, but fall back to the ground state and give off energy in the form of light called a spectrum (which is unique for each element)

http://www. mhhe. com/physsci/astronomy/applets/Bohr/applet_files/Bohr http://www.mhhe.com/physsci/astronomy/applets/Bohr/applet_files/Bohr.html http://www.colorado.edu/physics/2000/quantumzone/bohr.html

Speed of light (constant) = 2.998 x 108 m/sec c = f λ wavelength, lambda speed of light frequency

c = f λ frequency=speed/wavelength 2.998 x 108 m/s divided by 7.6 x 10-7 = 3.9 x 1014 1/s = 3.9 x 1014 Hz Energy of an electron = 2.179 x 10-18 J/n2

Schrodinger - 1925 Wave Mechanical Model – electrons exist in areas of probability called “space orbitals” (Rutherford and Bohr support this)

Quantum Mechanics – each electron is identified by 4 quantum numbers: 1. Principle quantum # - level 2. Orbital quantum # - shape 3. Magnetic quantum # - orientation 4. Spin quantum # - clockwise or counter http://www.colorado.edu/physics/2000/quantumzone/schroedinger.html

Werner Karl Heisenberg -1927 Uncertainty Principle – you can’t locate the exact position of an electron at any given time (too small, too fast)