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Chapter 3 – Atoms: The Building Blocks of Matter I. Laws A. 1790s = Study of matter changed by analysis of chemical reactions B. Atomic Theory led to discovery of several basic laws (also in book) 1. Law of Conservation of Mass = States that mass is neither created nor destroyed during ordinary chemical reactions or physical changes 2. 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 of source of the compound. 3. 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 combine with a certain mass of the first element is ALWAYS a ratio of small WHOLE numbers. II. Development of the Atomic Theory A. John Dalton (1808) 1. English Schoolteacher who proposed explanation for the three laws 2. Dalton’s Atomic Theory (also in book) a. All matter is composed of extremely small particles called atoms. b. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. c. Atoms cannot be subdivided, created, or destroyed. d. Atoms of different elements combine in simple whole number ratios to form compounds. e. In chemical reactions, atoms are combined, separated, or rearranged. 3. Problems with Dalton’s Model = atoms do have smaller parts (subatomic particles/quarks), atoms of the same element can have different mass, atoms of different elements can have the same mass (isotopes) 4. Other parts of his model are correct III. Discovery of subatomic particles A. J.J. Thomson concluded that the atom had negatively charged particles using cathode-­‐ray tubes 2. He concluded that subatomic particles can be found in atoms. 3. He is credited with discovering the electron 4. His model looked like “plum pudding” with a large positive pudding and small, evenly distributed negative charges (plums) B. Ernest Rutherford tested Thomson’s Model in 1909 1. Used alpha particles to bombard thin sheets of gold/metals 2. Determined that the atom contained a small, dense center of positive charge = the nucleus 3. Very small volume compared to size of the atom 4. Contained protons – positively charged particles 5. 99.9% of mass and all positive charge in nucleus 6. Suggested that electrons moved around like bees around a hive and thought that neutral particles existed C. Irene and Frederic Joilot-­‐Curie and James Chadwick credited with discovering neutrons. 1. Joilot-­‐Curies discovered the neutron, but did not know what it was 2. James Chadwick determined it was a new particle called a neutron 3. Same mass as a proton, but no charge, found in the nucleus D. Millikan discovered electron mass was very small compared to proton IV. The Atom A. The smallest particle of an element that retains chemical properties of element. B. 2 regions: 1. Nucleus = small region located at the center of an atom made up of two kinds of particles (protons and neutrons) a. Proton i.
Magnitude of charge = negative charge of an electron
ii.
Mass = 1.673 X 10-­‐27 kg b. Neutrons i.
Mass = 1.675 X 10-­‐27 kg ii.
No charge
2. Electron Cloud = region surrounding the nucleus a. Very large volume compared to nucleus, but no/little mass b. Contains electrons which have a negative charge c. Neutral atoms contain equal numbers of protons and electrons 3. Forces in the Nucleus = Same electric charge, particles should repel one another, instead there is a strong attraction between them V. Atomic Number, Mass Number, Ions and Isotopes A. Atomic number = the number of protons in the nucleus of one atom of the element 1. Also equals the number of electrons in neutral atoms, but not in ions 2. No 2 elements have the same atomic number and thus atomic number defines the element! B. Mass Number 1. These numbers increase by amounts greater than 1 2. The Mass number is equal to the protons + neutrons. 3. Written after the element name. (Carbon-­‐12) 4. Mass number – atomic number = neutrons C. Ions = charged particles/the charge is determined by the electrons, protons do not change D. Isotopes = Atoms of the same element that have different masses
1. Atoms of the same element that has the same number of protons and electrons but differ in their # of neutrons. 2. Carbon has different isotopes: carbon-­‐12, carbon-­‐13,carbon-­‐14 E. Nuclear Symbol – Mass Number on top and atomic number on the bottom. F. You can calculate all three atomic particles with this number. G. Element Symbol on the periodic table has the atomic number (generally on top of symbol), but does NOT have the mass number 1. Average atomic mass is given instead as mass numbers may differ 2. Weighted average of the atomic masses of the naturally occurring isotopes of an element 3. When using the average atomic mass in calculations, we will round this number to TWO decimal places. 4. Example: Carbon’s average atomic mass is 12.0107 = 12.01