History of the Atom

Lecture #9 History of the Atom

John Dalton's Atomic Theory You all know Dalton’s Atomic Theory…that all matter is made of atoms; atoms cannot be broken down; and atoms of one element are different from atoms of another element, but the same as atoms of the same element. However, John Dalton also came up with an atomic model. From his experiments and observations, he suggested that atoms were like tiny, hard balls. Each chemical element An element is a substance made from only one type of atom. An element cannot be broken down into any simpler substances. element had its own atoms that differed from others in mass. Dalton believed that atoms were the fundamental building blocks of nature and could not be split. In chemical reactions, the atoms would rearrange themselves and combine with other atoms in new ways. Dalton did not recognize that there were parts of an atom. He assumed that the atom was the smallest part. He just thought atoms differed by their size and mass.

Joseph John Thompson's Atomic Theory Joseph John Thompson believed that electrons were imbedded in a positively charged sphere. This was essentially the first model of an atom containing positive and negative particles. Thompson recognized the (-) electron, but he didn’t know that the nucleus has 2 parts.

J.J. Thomson and others were performing experiments with cathode ray tubes. A cathode ray tube is an evacuated tube that contains a small amount of gas between two metallic plates. When a potential is placed between the cathode (the negatively charged plate) and the anode (the positively charged plate) a "ray" of electric current passes from one plate to the other. Thomson discovered that this ray was actually composed of particles. When a second set of plates is placed around the tube, the ray is bent toward the positive plate indicating that the ray is composed of negatively charged particles. By varying the potential on the plates, Thomson was able to determine the mass to charge ratio of these particles.

In further experiments he varied what metal was used to make the electrodes and what gas was used to filled the tube. In each case, the properties of the ray particles were exactly the same. He concluded that the negatively charged particles were subatomic particles that were part of every atom. He further surmised that, since atoms were electrically neutral, the atom must also contain some positive charge. Based on these conclusions Thomson proposed that an atom was composed of a spherical ball of positive charge with "corpuscles" of negative charge imbedded in it. The corpuscles would later become known as electrons.

Ernest Rutherford's Atomic Theory Ernest Rutherford believed that an atom had a positive nucleus with electrons surrounding it. This model was fairly close to correct, but Rutherford believed that all the electrons were imbedded in a sphere around the nucleus.

Thomson had identified that the atom was composed of positive and negative charges and had proposed that the atom was a solid mass of these particles. If this model were true, any particles shot at the atom should be deflected by it. If the negative and positive charges were in some arrangement that left empty space in the atom, particles shot at the atom might be able to pass through them. In 1909, Rutherford set a fellow scientist, Hans Geiger, and a student, Ernest Marsden, to work on this problem. They devised a system that allowed alpha particles (the nuclei of helium atoms) to be shot at a very thin piece of gold foil and the trajectory of the particles monitored. They observed that while most of the particles passed through the foil with little or no deflection, some were deflected to a great degree.

Since the gold film was so thin, Rutherford proposed that all of the deflections observed were from single encounters of alpha particles with the atom. In order to deflect the relatively large and swiftly moving alpha particles to such a large extent, a large force was required. This force, he contended, could only be caused by a large concentration of positive charge within the atom. This large concentration of charge was located at the center of the atom and became known as the nucleus. The negative charge, in the form of electrons, was then distributed throughout the rest of the space occupied by the atom.

In order to account for the fact that many of the alpha particles passed through the gold film, Rutherford discounted Thompson's solid ball model of the atom, and believed that the central positive charge of the atom represented only a small fraction of the atom's size, and that the remainder was primarily empty space. He calculated that, while an individual atom was about 1x10-10 meters in diameter, the nuclear diameter was only about 1x10-14 meters. While solving the problem of the observed alpha particle deflection, Rutherford's model created another. If the positive charge was located at the center of the atom, why were the negatively charged electrons not immediately drawn into it (opposite charges attract). Rutherford was not unaware of this problem but his model so adequately (and mathematically) explained the scattering results that it became widely accepted.

Rutherford's Gold Foil Experiment • Gold is a metal that can be pounded into a foil only a few atoms thick. • Rutherford’s lead-shielded box contained radioactive polonium. As polonium decays, it gives off a helium nucleus (2 protons & 2 neutrons). The nucleus is positively charged because it has no electrons. • When the beam hits the gold, most of the particles passed through the gold. • But some of the particles bounced back, kind of like a baseball hitting tissue paper and bouncing back. Rutherford had discovered that the nucleus has 2 parts!

The idea that an elementary unit of charge should exist seems to have originated with Benjamin Franklin around 1850. It was not until the work of Robert Millikin that the number value of this charge could be determined. It was known that X-rays could be used to impart a negative charge to an oil droplet in a chamber that contains it. Like all objects of mass, the influence of gravity causes the oil droplet to fall. Millikin placed charged plates at the top and bottom of his chamber. By varying the potential between the plates, he discovered that he was able to suspend the droplet in mid-air. The droplet remained suspended when the downward force of gravity was exactly balanced by the upward electrical force caused by the charged plates. Since both gravitational and electrical equations were known to determine these forces, Millikin was able to calculate the charge on each of the droplets he tested (Note 1). The calculated charges on the droplets all turned out to multiples of a single number. Millikin therefore reasoned the elementary charge, or the smallest of charge, must be equal to this value. By combining his new information with the mass to charge ratio for the electron determined by Thomson, the mass of an electron was calculated for the first time. Robert Milikin

In 1886, Eugen Goldstein pierced several holes in the cathode of a cathode ray tube and noticed a stream of particles in the space behind the cathode. He called these streams canal rays. Using powerful magnetic fields in 1898, Wilhelm Wien verified that these particles were positively charged and had masses comparable to a hydrogen atom. Additionally, he concluded that, as the particles were not all influenced by the magnetic field to the same extent, they must have different masses. These two results were in stark contrast to the normal cathode ray that was composed solely of electrons (small and uniform in mass).

Niels Bohr's Atomic Theory Niels Bohr believed that each electron circled the nucleus in a fixed orbit. He believed that only one electron occupied each orbit and could not move from orbit to orbit.

Niels Bohr (1885-1962) refined Rutherford's model by introducing different orbits in which electrons spin around the nucleus. This model is still used in chemistry. Elements are distinguished by their "atomic number", which specifies the number of protons in the nucleus of the atom. Electrons are held in their orbits through the electrical attraction between the positive nucleus and the negative electron. Bohr argued that each electron has a certain fixed amount of energy, which corresponds to its fixed orbit. Therefore, when an electron absorbs energy, it jumps to the next higher orbit rather than moving continuously between orbits. The characteristic of electrons having fixed energy quantities (quanta) is also known as the quantum theory of the atom.

The Electron Cloud Model Today, it is believed that electrons move in a cloud around the nucleus. There are different levels to the cloud and electrons can move between the levels. Here is a picture of what the different electron orbital clouds actually look like, take with an electron microscope. Note, we still draw atoms using the Bohr model for ease.

Review of the Atom

The Atomic Model Remember that the inside ring (orbital) can hold only up to 2 electrons. The next ring can hold up to 8 electrons.

11 22.99 80 Atomic Number Element Symbol Element Name Atomic Mass Hg Na Mercury Sodium 200.59 Remember, each periodic table may be a little different. The upper left is from the poster, the upper right is from your agenda, the lower left is from the handout. Boron Element Name Atomic Number Element Symbol Atomic Mass 5 B 10.811

The atomic number equals the number of protons in an atom. atomic # = # protons Atomic # = 11 # Protons = 11 The atomic mass, minus the atomic number, equals the number of neutrons in an atom. atomic mass - atomic # = # neutrons Atomic mass = 22.99 = 23 23 - 11 = 12 = # neutrons In a neutral atom, the number of electrons is equal to the number of protons. # protons = # electrons 11 protons = 11 electrons

The atomic number equals the number of protons in an atom. atomic # = # protons The atomic mass, minus the atomic number, equals the number of neutrons in an atom. atomic mass - atomic # = # neutrons In a neutral atom, the number of electrons is equal to the number of protons. X # protons = # electrons Try: Ne He #p = #p = #n = #n = #e- = #e- =

The atomic number equals the number of protons in an atom. atomic # = # protons The atomic mass, minus the atomic number, equals the number of neutrons in an atom. atomic mass - atomic # = # neutrons In a neutral atom, the number of electrons is equal to the number of protons. X # protons = # electrons Try: Ne He #p = #p = #n = #n = #e- = #e- = 2 2 2 10 10 10

How do you figure out the number of electrons in an ion? If the atom is positively charged, the ion is a cation. cations are + If the atom is negatively charged, the ion is an anion. anions are - If the atom is an anion (-), the atomic number plus the charge equals the number of electrons. atomic # + X2- = # electrons If the atom is a cation (+), the atomic number minus the charge equals the number of electrons. atomic # - X2+ = # electrons

Try: Rb+ N3- Al3+ #p = #p = #p = #n = #n = #n = #e- = #e- = #e- = Cl- Mg2+ Sb4- #p = #p = #p = #n = #n = #n = #e- = #e- = #e- =

Try: Rb+ N3- Al3+ #p = #p = #p = #n = #n = #n = #e- = #e- = #e- = Cl- Mg2+ Sb4- #p = #p = #p = #n = #n = #n = #e- = #e- = #e- = 7 14-7=7 7+3=10 13 27-13=14 13-3=10 37 85-37=48 37-1=36 51 122-51=71 51+4=55 17 35-17=18 17+1=18 12 24-12=12 12-2=10

Valence Electrons Valence electrons are found on the outside ring of the atom. The easiest way to find out how many valence electrons there are on an atom is to look at the periodic table, find the element, and find the number at the top of the column. That number is the number of valence electrons in each element in that column. Ve-

Lewis Dot Structures Lewis Dot Structures are the visual representation of the number of valence electrons found on an atom. Each dot on a Lewis Dot Structure represents one valence electron. Look at the periodic table to determine how many valence electrons the atom has, and then start drawing dots around the element symbol. Remember, when you draw dots, start with one on each side, if there is more than 4, start doubling them up.

An examples of Lewis Dot Structures: Na Al Cl Try: Mg S Xe P Cs Te

History of the Atom