History of the Atom

The idea of an Atom  Democritus and Leucippus Greek philosophers who thought maybe small particles like grains of sand could be made of something smaller…but they were over- ruled.  The Early Greeks thought of everything as being made up of four basic elements because that is what Aristotle told them Earth. Earth. Air. Air. Fire. Fire. Water. Water.

John Dalton  Dalton – thought an atom was a tiny indivisible sphere of matter.  Dalton’s Atomic theory…  All matter is made of tiny indivisible particles called atoms.  Atoms of the same element are identical, those of different atoms are different.  Atoms of different elements combine in whole number ratios to form compounds  Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.

The problems with Dalton’s theory  following scientists did find that the atom contained smaller pieces…  and we’ve learned that an atom can be created and destroyed

J.J. Thomson

Voltage source Vacuum tube Metal Disks + - Thomson’s cathode ray tube ....

Thomson’s Experiment Voltage source +-

Thomson’s Experiment Voltage source +-

Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

Thomson’s Experiment  By adding an electric field

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field he found that the moving pieces were negative + -

JJ Thompson  Thompson – concluded that the atom a divisible sphere that was positively charged and embedded with negative particles.  The problem: was not supported by Rutherford’s experiment

Thomson’s Model  Found the electron  Couldn’t find positive (for a while)  Said the atom was like plum pudding  A bunch of positive stuff, with the electrons able to be removed

 Other pieces  Proton was found to be a positively charged particle… 1840 times heavier than the electron  Neutron - no charge but the same mass as a proton.

Ernest Rutherford  Rutherford, a New Zealand scientist, had studied radiation and wanted to work with Thompson in investigating the make-up of the atomic structure.  Rutherford had discovered three types of radiation and named them alpha particles, beta particles and gamma particles.  Rutherford designed an experiment to probe the atom using radiation (alpha particles) as bullets.

Lead block Uranium Gold Foil Florescent Screen

He Expected  The alpha particles to pass through without changing direction very much  Because  The positive charges were spread out evenly. Alone they were not enough to stop the alpha particles

What he expected

Because, he thought the mass was evenly distributed in the atom

What he got

How he explained it +  Atom is mostly empty  Small dense, positive piece at center positive piece at center  Alpha particles are deflected byit if they deflected byit if they get close enough get close enough

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Rutherford’s gold foil experiment  Conclusion – every atom must have a nucleus containing a dense positive mass which has a lot of empty space around it in which tiny negative particles are found.

 As technology, chemistry and physics advanced the model of the atom was again questioned and altered by a Danish scientist named Neils Bohr.  The problem with the Rutherford model was … if the nucleus is positive and opposite charges attract, why didn’t the negative electrons just crash in to the nucleus?

Modern View  The atom is mostly empty space  Two regions  Nucleus- protons and neutrons  Electron cloud- region where you might find an electron

Density and the Atom  Since most of the particles went through, it was mostly empty.  Because the pieces turned so much, the positive pieces were heavy.  Small volume, big mass, big density  This small dense positive area is the nucleus

Neils Bohr  Bohr’s model was compared to our solar system. While the Sun attracts the planets with great gravitational force, the plants move in an orbit due to the speed of their motion.  The Bohr model describes the electrons in rapid motion with set amounts of energy around the nucleus. Where the electron is depends on how much energy it has. There are energy levels around the nucleus, like orbits, and due to their energy the electrons move around the nucleus in a set path.

 Bohr’s Theory Allowed Orbitals Allowed Orbitals An electron can only orbit around an atom in specific orbitsAn electron can only orbit around an atom in specific orbits Radiationless Orbits Radiationless Orbits An electron in an allowed orbit does not emit radiant energy as long as it remains in the orbit.An electron in an allowed orbit does not emit radiant energy as long as it remains in the orbit. Quantum Leaps Quantum Leaps An electron gains or loses energy only by moving from one allowed orbit to another.An electron gains or loses energy only by moving from one allowed orbit to another. The lowest energy state is known as the ground stateThe lowest energy state is known as the ground state Higher states are known as excited statesHigher states are known as excited states

 Each time an electron males a "quantum leap," moving from a higher energy orbit to a lower energy orbit, it emits a photon of a specific frequency and energy value.

Counting the pieces  The number of protons in an atom determines what atom it is…this is referred to as the atom’s atomic number.  For neutral atoms, the number of protons (positive charges) and the number of electrons (negative charges) are equal.  The mass number is the mass of the nucleus (the number of protons plus the number of neutrons).

Atomic Mass  The mass of an atoms comes from the particles in the nucleus.  Each proton and neutron has a mass of 1.66 x g. Comparably the mass of an electron is negligible.  The mass of an individual atom is too small to measure, therefore we base the mass on ratios – comparing elements to each other.

 Carbon-12 is used as the reference element for atomic mass.  Carbon-12 has 12 protons and 12 neutrons… therefore, defining one atomic mass unit as 1/12 of a carbon-12 atom, means that therefore, defining one atomic mass unit as 1/12 of a carbon-12 atom, means that 1 proton has a mass of 1 amu and 1 neutron has a mass of 1 amu.

Isotopes  We now know that a number of the postulates of Dalton’s atomic theory were wrong.  We know that particles smaller than the atom exist.  We also know that atoms of the same element can have different masses, because they can have different numbers of neutrons. These are called isotopes.

Isotope Symbols  To symbolize the composition of an isotope, the superscript to the left is used to represent the mass number, and the subscript to the left is used to represent the atomic number.  Some common isotopes can be found on page 78 in your textbook. 18O8

Average atomic mass  The average atomic mass of an element is based on the weighted average of its isotopes and their percent abundance

Quantum Mechanics  The Quantum Concept. In 1900 Max Plank introduced the idea that matter emits and absorbs energy in discrete units called quanta. In 1900 Max Plank introduced the idea that matter emits and absorbs energy in discrete units called quanta. In 1905 Albert Einstein extended the quantum concept to include light and that light consist of discrete units called photons. In 1905 Albert Einstein extended the quantum concept to include light and that light consist of discrete units called photons. The energy of a photon is directly proportional to the frequency of vibration. The energy of a photon is directly proportional to the frequency of vibration. E=hfE=hf where E = energy where E = energy h = Plank’s constant = 6.63 X J  s h = Plank’s constant = 6.63 X J  s f = frequency f = frequency

Quantum Mechanics  Quantum mechanics states that light and matter, including electrons, have a dual nature of both particles and waves.

 (A) Light from incandescent solids, liquids, or dense gases, produces a continuous spectrum as atoms interact to emit all frequencies of visible light (B) Light from an incandescent gas produces a line spectrum as atom emit certain frequencies that are characteristic of each element.

 Atomic hydrogen produces a series of characteristic line spectra in the ultraviolet, visible, and infrared parts of the total spectrum. The visible light spectra always consist of two violet lines, a blue-green line, and a bright red one.

 These fluorescent lights emit light as electrons of mercury atoms inside the tube gain energy from the electric current. As soon as they can, the electrons drop back to their lower-energy orbit, emitting photons with ultraviolet frequencies. Ultraviolet radiation strikes the fluorescent chemical coating inside the tube, stimulating the emission of visible light.