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Atoms : The Building Blocks of Matter Chapter 3 The Atom: From philosophical idea to scientific theory 500 B. C. Aristotle: Matter was continuous and there were no basic particles. 400 B. C. Democritus: Matter was composed of particles, or atoms(Greek for indivisible) 1700’s: Elements could not be broken down by ordinary chemical means. Elements combined to form compounds that had properties different from the original elements. Such a transformation was called a chemical reaction. Foundations of Atomic Theory 1790’s Quantitative analysis came about and several basic laws were discovered. Law of conservation of mass: Mass is neither created nor destroyed in an ordinary chemical reaction or physical change. (Lavoisier) Law of definite proportions: A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. (Proust) Law of multiple proportions: If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a small whole number ratio. (Dalton) Law of conservation of mass Law of multiple proportions Dalton’s Atomic Theory 1808 John Dalton proposed an explanation for the 3 laws. He stated the first atomic theory based on the teachings of Democritus. All matter is composed of small particles called atoms. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. Atoms cannot be subdivided, created, or destroyed. Atoms of different elements combine in simple whole number ratios to form chemical compounds. In chemical reactions, atoms are combined, separated, or rearranged. Dalton’s Theory Explains the Laws The law of conservation of mass is explained by the fact that chemical reactions involve combination, separation, or rearrangement of atoms and that during the process, no atoms are lost, created, or subdivided. The law of definite proportions results from the fact that a given chemical compound is always composed of the same combination of atoms. As for the law of multiple proportions, compounds of the same elements have the same amount of one element and a multiple of the second element. Modern Atomic Theory Not all aspects of Dalton’s atomic theory have proved correct but Dalton is still known as the father of atomic theory for his beginning thoughts. We now know that atoms are composed of protons, neutrons, and electrons. They do not have the same properties as the atom however. We also know that atoms of the same element exist that differ in mass. Such atoms are called isotopes. Atoms of different elements can have the same mass, too. The structure of the atom An atom is the smallest particle of an element that retains the chemical properties of the element. All atoms consist of two regions---the tiny, dense nucleus and the large electron cloud. The nucleus contains protons(p+) and neutrons(n0) while the electron cloud contains electrons(e- ). These are subatomic particles. The structure of the atom http://ippex.pppl.gov/interactive/matter/ato ms.html http://www.pbs.org/wgbh/aso/tryit/atom/ Discovery of the electron Late 1800’s: The first subatomic particle, the electron, was discovered in an evacuated cathode ray tube by J. J. Thomson. When an electric discharge was passed through various gases at very low pressures in cathode ray tubes, the surface of the tube opposite the negative electrode(cathode) glowed. Scientists hypothesized that the glow was due to cathode rays which traveled from the cathode to the positive electrode(anode). Cathode Ray Tube Cathode ray tube experiments An object placed between the cathode and the opposite end of the tube cast a shadow on the glass. A paddle wheel placed on rails between the electrodes rolled along the rails from the cathode toward the anode. Cathode rays were deflected by a magnetic field. Cathode rays were deflected away from a negatively charged object. Results of Cathode Ray Experiments All particles of the cathode ray are negative. Thomson measured the charge to mass ratio. He found that the ratio was always the same, no matter what the cathode metal was or what the gas was. Thus Thomson concluded that all cathode rays are composed of identical negatively charged particles. Millikan’s Oil Drop Experiment Millikan(1909) showed that the mass of the electron is about 1/2000th the mass of a hydrogen atom or 9.11 x 10-31 kg. He also found that the electron carries a charge of 1.6 x 10-19C. Because all kinds of elements produced the same cathode rays, he concluded that electrons were fundamental to all atoms. Because atoms are neutral, if there is a negative charge, there must also be a positive one. There must be other particles that help account for mass of the atom. Rutherford’s Gold Foil Experiment Rutherford(1911) bombarded thin gold foil with alpha particles, which are like helium nuclei. Most the alpha particles passed through the foil; others were deflected at angles. Still others bounced back from the foil. This experiment did more to characterize the atom than any other. Rutherford’s Experiment The atom was mostly space since most particles passed through the foil. The center of the atom was a small, very dense core (nucleus) because there were so few particles deflected. The core was positively charged because it was able to deflect the positive alpha particles. He assumed that the electrons were outside the nucleus orbiting the nucleus like planets around the sun. Rutherford’s Gold Foil Experiment Discovery of the proton(1919) Protons were also discovered in a cathode ray tube but the cathode in it was perforated. Rays were found to come from the back of the cathode. They were called canal rays. They were found to be opposite in charge from the cathode rays but equal in magnitude. Rutherford and Goldstein concluded that a positive ion may be formed from any neutral atom so each atom must have positive particles. Discovery of Neutrons There was extra mass in the atom that could not be explained by protons and electrons. In 1932, Chadwick found that there were neutral particles in the nucleus that were given off as a result of radioactive decay when Be atoms were bombarded with alpha particles. The mass of a neutron is approximately 1 g and it may decompose into a proton and an electron. The electron may be emitted as β rays while the proton is reabsorbed by the nucleus. Composition of the Nucleus All atomic nuclei are made up of protons and neutrons except an isotope of hydrogen called protium. It has no neutrons. Protons are positive and electrons are negative and equal numbers of them cause an object to be neutral. The number of protons determines an atom’s identity. Charges of the same type generally repel one another but in the nucleus, protons are attracted to one another and are held together with gluons. These are short range nuclear forces. Atomic Number The proton number is the atomic number, Z. Elements on the periodic table are ordered by atomic number. The atomic number is an identifying characteristic. Moseley formulated the atomic number concept. Isotopes Isotopes are atoms of the same element that have different numbers of neutrons. The three isotopes of hydrogen are protium, deuterium, and tritium. Since there are different numbers of neutrons, the masses of the isotopes will be different. Isotopes do not differ significantly in their chemical behavior. Mass Number Identifying an isotope requires knowing both the name or atomic number and the mass of the isotope. The mass number, A, of the isotope is the sum of the protons and the neutrons. Specifying isotopes may be done in two ways. P+ + N0 = A Write the name of the element, then a hyphen and then the mass number. Carbon-13 Write the nuclear symbol. 235U is uranium-235. Nuclide is the general term for any isotope of any element. Relative Atomic Masses Masses of atoms in grams are extremely small so one atom has been arbitrarily chosen as the standard and assigned a relative mass value. All other masses are expressed in relation to this standard. The standard is the carbon-12 nuclide. It has been assigned a mass of 12. Thus a hydrogen-1 atom will have a mass of approximately 1/12th of carbon12 atom One amu is exactly 1/12th the mass of a carbon-12 atom. Average Atomic Masses of Elements Most elements occur naturally as mixtures of isotopes. The natural abundance of an isotope is taken into consideration when determining the average atomic mass of an element. Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. Weighted Average Do these contribute equally to the sample mass? Relating Mass to Numbers of Atoms The mole is the SI unit for amount of substance. A mole is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12. It is a collective term like dozen is. The number of particles in a mole is 6.022 x 1023, which is Avogadro’s number. Another definition of the mole is the amount of substance in grams that contains Avogadro’s number of particles. The mass of one mole is the molar mass or gram formula mass(gfm).