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Sec 3.1 Science History Democritus - ~ 400 BC Greek Philosopher o 1st particle theory of matter. He called nature’s basic particle an “atom”, which meant indivisible. Aristotle – “Einstein” of the Greeks o Did not believe in Democritus idea. o He believed that all matter was continuous. o Everyone believed him because of his prestige. Neither option was supported by experimental evidence, just speculation. For almost 2000 years, the idea of the “atom” was lost. Law of Conservation of Mass – we can not create or destroy mass during ordinary chemical reactions or by physical changes. Law of Definite Proportions – Compounds are always composed of a fixed proportion of elements. ex. NaCl always contains 1 Na and 1 Cl or 39.34% by mass Na 60.66% by mass Cl Law of Multiple Proportions – the same element may combine to form more than one compound. ex. CO (Carbon monoxide) and CO2 (Carbon Dioxide) O2 (Oxygen) and O3 (Ozone) Dalton’s Atomic Theory (In 1808, he explained the last 3 laws) 1. All matter is composed of extremely small particles 2. 3. 4. 5. called atoms. Atoms of a given element are identical in size, mass, and other properties. Atoms of different elements differ by size, mass, and other properties. Atoms can not be subdivided, created, or destroyed. Atoms of different elements combine in simple whole # ratios to form chemical compounds. In chemical reactions, atoms are combined, separated, or rearranged. ** By relating atoms to the measureable property of mass, Dalton turned Democritus’ idea into a scientific theory. Parts of Dalton’s Atomic Theory no longer hold true, but it has been modified to explain new observations - Modern Atomic Theory ex. Atoms are divisible into smaller particles (protons, neutrons, e-) but the conservation of mass still holds true. ex. Atoms of the same element can have different mass ( C12& C14 ) Sec 3-2 Atom – the smallest particle of an element that retains the chemical properties of that element. The proton, neutron, and electron are called subatomic particles. Joseph Thomson (1897) – discovered the e- when he passed electric current through a various gases in a cathode ray tube. Robert Millikan (1909) – showed that the mass of e- was 1/1837 the mass of the hydrogen atom or 9.109 x 10-31 kg and that e- carry a charge. Because of these discoveries, 2 inferences were made about the atomic structure. 1. Because atoms are electrically neutral, they must contain a (+) charge to balance the (-) charge of the e- . 2. Because e- have so much less mass than atoms, atoms must contain other particles that account for most of the mass. Ernest Rutherford (1911) – discovered the atomic nucleus. He bombarded a thin foil with alpha particles. ** If the atom were equal to the size of a football field, the nucleus would be the size of a marble. (Mostly empty space) All nuclei (except the simplest type of H) contains protons (p+) and neutrons (N). Mass of p+ = 1.673 x 10 -27 kg or 1836 times the mass of e-. Mass of n = 1.675 x 10 -27 kg All About Atoms - What are atoms? Sec 3-3 Atoms of different elements have different # of p+. Atomic Number = tells us the # of p+ in the nucleus of each atom of that elements. Also tells us the # of e-. Ex. How many protons does Ca have? How many electrons does Ca have? Ex. What is the element that has an atomic #7? Nitrogen The identity of an atom is determined by the # of p+, not eor N Isotopes – atoms of the same element that have different masses (because they differ by the number of neutrons). Isotopes of hydrogen are Protium, Deuterium, and Tritium. Most elements consists of mixtures of isotopes: Tin has 10 isotopes. Mass # = p+ + n We don’t include the e- because their mass is negligible. p+ n Mass # Protium 1 0 1 Deuterium 1 1 2 Tritium 1 2 3 Normally, we name the isotope by its mass #. ex. Hydrogen – 1 or H - 1 Uranium – 235 or U - 235 How many protons, neutrons, and electrons does Na – 23 have? 23 = p+ + n 23 = 11 + n n = 12 Another way to write Na – 23 is How many p+, n, and e- are in ? How many p+, n, and e- are in ? And what is the name of x? ***Do not confuse the mass # with the average atomic mass on the periodic table. The masses for atoms are extremely small. ex. O atom = 2.657 x 10-29 g Therefore we use the relative atomic mass, which is relative to the Carbon – 12 atom. C-12 has been assigned a mass of exactly 12 atomic mass units or amu’s. (not kg or grams) Using this standard, the p+ and n have ≈ masses = 1 amu. Mass # ≈ atomic mass in amu’s. Ave atomic mass – is a weighted average of all the isotopes that occur naturally. Ex. Your grades are weighted by: 20% Homework 40% Quizzes 40% Tests Ex. C-12 (98.9%) C-13 (1.1%) C-14 (very very small %) It takes an enormous amount of H2O molecules to get enough for us to see. Too many for us to count. However, in chemistry, we need to keep track of all of these molecules, so chemists relate the # of molecules to its mass. We call this a “mole”. 1 mole of any substance = 6.022 x 10 23 Avagadro’s Number Avagadro’s Number was found by experiment. Exactly 12 g of C-12 atoms contained 6.022 x 1023 C-12 atoms or we say 12g/mole. That is 602 sextillion – an extremely huge # 1 oxygen atom has a mass of 15.999 amu 1 mole of oxygen atoms has a mass of 15.999g or can be written as 15.999 g/mole. Gram to Mole Conversions . What is the mass in grams of 3.50 mol of Cu? 1 2. What is the mass of 2.25 mol of Fe? 3. A chemist produced 11.9 g of Al. How many moles of Al were produced? 4. How many moles of Cu are in 3.22 g Cu? Practice problems on pg 83: top of the page 2-4 bottom of the page 1 & 2 6. How many atoms of Ca are in 4 moles of Ca? 7. How many moles of K are in 4.52 x 1023 atoms of K? 8. How many grams of Ag are in 3.01 x 1023 atoms of Ag? 9. How many grams of C are in 2.25 x 1022 atoms of C? 10. How many atoms are in 2.5 g of Fe? 11. What mass of gold, contains the same # of atoms as 9.0 g of Aluminum?