Theories of Matter Composition
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- Ferdinand Pierce
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1 Chapter 2
2 Theories of Matter Composition Democritus (5 th 4 th century BC) ATOMISM Aristotle (4 th 5 th century BC) CONTINUOUS MATTER FOUR ELEMENTS Earth, Air, Fire, Water Boyle (17 th century) Reintroduced ATOMISM in modern times. Dalton (19 th century) Atomic Theory to explain results of EXPERIMENTS. F-2!
3 Development of the Modern Atomic Theory In 1803, John Dalton proposed an atomic theory that is still the basis for many of our theories about the atom. 1. All matter is composed of atoms, which are tiny, indivisible particles. 2. A chemical reaction is a rearrangement of atoms to form different compounds. Atoms are neither created nor destroyed in a chemical reaction (the law of conservation of mass). 3. Atoms of one element cannot be converted into another element. Atoms of an element are identical in mass and other properties, and are different from every other element. 4. A compound is a combination of atoms of two or more elements in specific ratios (the law of definite composition). F-3!
4 Cathode Ray Tube: Electron: Very small, subatomic, negatively charged particle. From his experiments, Thompson calculated the ratio of the electron s mass, m e, to its electric charge, e. F-4!
5 F-5!
6 The work of Thomson and Millikan combined provided us with the electron s mass of x kg, which is more than 1800 times smaller than the mass of the lightest atom (hydrogen). These experiments also showed that the electron is a subatomic particle. F-6!
7 JJ Thomson favored the plum pudding model. He believed that the atom is a large mass of positive charge, with tiny electrons moving around in circles on the surface. Ernest Rutherford, a student of Thomson, was working with alpha particles, large positively charged particles. He and his students (Geiger & Morrison) found that the behavior of alpha particles was not consistent with Thomson s model. F-7!
8 Rutherford used alpha particles to further study the atomic model. He directed alpha particles from a source at thin pieces of metal foil that were a relatively small number of atoms thick. F-8!
9 F-9!
10 Rutherford s Experimental Design
11 Rutherford s Experimental Design
12 Rutherford s Experimental Design
13 Rutherford s Experimental Design
14 Rutherford s Experimental Design
15 An atom is made up of a small, central, positively charged nucleus, surrounded by even smaller, negatively charged electrons that are moving around the atom. NUCLEUS: 1) Subatomic particle: Core of the atom. 2) Positively charged. 3) Takes up very little space in the atom. 4) Contains most of the mass of the atom. F-15!
16 Atoms are made up of three subatomic particles: Proton (p + ): Positively charged subatomic particle found in the nucleus of the atom. Neutron (n 0 ): Neutral subatomic particle found in the nucleus of the atom. Electron (e ): Negatively charged subatomic particle found outside the nucleus. F-16!
17 The identity of an element is determined by the number of protons in the nucleus. For example, any atom with 6 protons in the nucleus is a Carbon atom. Number of protons in the nucleus = atomic number (Z) In a neutral atom, # electrons = # protons. The symbol for an element is simply its 1, 2, or 3 letter abbreviation from the periodic table. Elements are arranged in the periodic table by their atomic number (Z) F-17!
18 F-18!
19 Isotopes are atoms of the same element with different numbers of neutrons. Most elements have more than one isotope. Isotopes of an element are have nearly identical properties. The number of protons and electrons, which are the same in all isotopes of an element, has much more to do with the chemical and physical properties of an element. F-19!
20 F-20!
21 F-21!
22 Can be used to separate isotopes of an element. F-22!
23 F-23!
24 F-24!
25 Consider the isotope Phosphorus-31 symbol: 31 P 15 A = 31 (mass number = the sum of protons and neutrons) Z = 15 (atomic number = protons ) In order to determine how many neutrons are in the nucleus of an atom simply subtract: # neutrons = mass number atomic number = = 16 neutrons F-25!
26 Isotope Name Isotope Symbol Atomic Number (Z) Mass Number (A) protons electrons neutrons Rhenium Np
27 Atoms are very small. For example, 18 ml (~ 4 teaspoons) of water contains ~ 602,200,000,000,000,000,000,000 molecules of H 2 O. A single water molecule has a mass of 2.99 x grams. Working in grams to describe the mass of single atoms or molecules is not convenient. Instead, we typically express the mass of atom in terms of atomic mass units, amu, or simply u. F-27!
28 The atomic mass unit is defined as 1 / 12 of the mass of an atom of the carbon-12 isotope. 1 amu = 1/12 mass of C-12 Both the neutron and the proton have a mass of approximately 1 amu: 1 amu 1 proton 1 neutron Electrons are very small: 1 electron amu Because the mass number is the sum of protons and neutrons, the mass number is a whole number approximation of the atomic mass of an isotope. F-28!
29 Most elements are a mixture of two or more isotopes. The percentage of an isotope in a naturally occurring sample of an element is called the isotopic abundance (or percentage abundance) of that that isotope. The isotopic mass is the mass of a single atom of an isotope. The average atomic mass is the weighted average of the masses of isotopes of an element. F-29!
30 (Fractional Abundance Isotope Mass) Consider the element Magnesium, composed of three isotopes of the following percentage abundances: Isotope % Abundance Isotope Mass (amu) Mg % Mg % Mg % F-30!
31 F-31!
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