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Law of Conservation of Mass:

Law of Conservation of Mass:. Mass is neither created nor destroyed in a chemical reaction. Therefore in a chemical reaction, Mass of Reactants must equal Mass of Products. Example:. Mercury (II) Sulfide Makes Mercury + Sulfur.

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Law of Conservation of Mass:

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  1. Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction. Therefore in a chemical reaction, Mass of Reactants must equal Mass of Products Example: Mercury (II) Sulfide Makes Mercury + Sulfur Mercury (II) Sulfide → Mercury + Sulfur (reactant) (products) If 216.0 g of Mercury (II) sulfide are used, and 200.0 g of Mercury are made, how many grams of Sulfur must have been made? 216.0 g = 200.0 g + x g x g = 216 g – 200.0 g = 16.0 g of S made Make special note of all problems! Expect to have to do this on a quiz or test!

  2. Law of Definite Proportions: A pure substance always contains the same proportions by mass of each atom in the substance regardless of where the sample of the substance was found or how much of the substance was analyzed. Percent by Mass (or Mass Percent) Mass of Element x 100 Mass % = Mass of Compound Example: sodium chloride (table salt) contains 39.34 % sodium ion and 60.66 % chloride ion by mass no matter where the pure sample is found and no matter how much sodium chloride is used during the analysis. A 100.00 g sample of sodium chloride would contain how many grams of sodium ion and how many grams of chloride ion? A 28.43 g sample of sodium chloride would contain how many grams of sodium ion and how many grams of chloride ion? Make special note of all problems! Expect to have to do this on a quiz or test!

  3. Law of Multiple Proportions: If two or more compounds are composed of the same two elements, then the mass ratio(s) of one of the elements will always be a ratio of small whole numbers. Example: at least three different compounds containing just chromium and chlorine are known. Data for these three compounds is given in the table below. Using the data in the table, the law of multiple proportions says that the ratios of the masses of the chlorine will be small whole numbers. Ratio of chlorine in B to A is 1.4998 or 1.500 which is 3/2. Ratio of chlorine in C to B is 2.000 which is 2/1. Ratio of chlorine in C to A is 3.000 which is 3/1. What does this do for us in terms of our understanding of chemistry?

  4. Law of Conservation of Mass Practice N2 (g) + 3 H2 (g) → 2 NH3 (g) 1) If 28.0 g of N2 react exactly with 6.0 g of H2, how many grams of NH3 are produced? N2 (g) + 3 H2 (g) → 2 NH3 (g) 2) If 53.5 g of NH3 are produced by the reaction of 42.42 g of N2 with excess H2, how many grams of are H2 consumed? 2 NH3 (g) → N2 (g) + 3 H2 (g) 3) If 128.4 g of NH3 are reacted to produce 98.4 g of N2, how many grams of are H2 produced?

  5. Mass % Practice 1) If 28.0 g of a sample of compound are found to contain 8.00 g of calcium and 20.00 g of iodine, what is the mass % of calcium in the sample? What is the mass % of iodine in the sample? 2) If a compound is known to be 22.00% carbon and 78.00% chlorine by mass, how many grams of carbon would there be in 3.721 g of the compound? How many grams of chlorine?

  6. Atomic Theories (theories of what atoms are like) John Dalton proposed a new atomic theory based on the three Laws we have been discussing and the ideas of Democritus. See page 64 for a discussion of Daltons ideas. Dalton’s Atomic Theory (1808 A.D.)-very similar to ideas of Democritus (~400 B.C.) • Each element is made up of tiny particles called atoms that can not be broken down into smaller particles. • The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. • Chemical compounds are formed when atoms combine with each other in simple whole number ratios. A given compound always has the same relative numbers and types of atoms. • Chemical reactions involve reorganization of the atoms - changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction. Was Dalton 100% correct?

  7. Cathode Ray: stream of charged particles produced by an electric field. The green “glow” in the picture. The particles move from the cathode (negatively charged plate) to the anode (positively charged plate).

  8. Cathode “rays” are influenced by both magnetic and electric fields (electric field shown in the diagram). Opposite charges are known to attract each other. Since the cathode “ray” was deflected towards the positively charged plate, what charge do the particles in the cathode “ray” have?

  9. J. J. Thompson and Millikan demonstrated that the mass of the particles in a Cathode Ray are much less than the mass of the smallest atom. How did this change our understanding of Dalton’s theory? J. J. Thompson proposed the Plum Pudding Model Plum Pudding Model: Electrons are evenly distributed throughout a uniform positive charge equal to the negative charge of all of the electrons. The part with the positive charge is where most of the mass of an atom is located.

  10. Ernest Rutherford (1911) “gold foil” experiment Large, high energy particles (alpha particles, +2 charge) do not always pass straight through a thin sheet of matter. The alpha particle being deflected or reflected is like having a 30-06 bullet bounce off of a sheet of tissue paper.

  11. a) Results of golf foil experiment expected if Plum Pudding Model were true. b) Results of golf foil experiment explained by new model. Rutherford proposed a nuclear model of the atom. The small, dense nucleus contains virtually all the mass of the atom and all of the positive charge while the negatively charged electrons exist apart from the nucleus. Rutherford did not know where the electrons were-they were just outside of the nucleus. How does this model fit Rutherford’s results?

  12. Nuclear Model of an Atom Small, very dense nucleus containing massive protons and neutrons, surrounded by small rapidly moving electrons If all atoms contain the same types of particles, what makes one atom hydrogen and another carbon? The number of protons in the nucleus determine what element an atom is. Remember that in normal atoms, the number of protons is equal to the number of electrons (they are neutral-a total charge of zero). How many times bigger is the atom compared to the nucleus? If the nucleus is very small compared to the overall size of the atom, then the atom consists of mostly empty space!

  13. Rutherford called the positive particle in the nucleus a proton. In 1932, James Chadwick discovered a neutral particle in the nucleus and called it a neutron. Summary of Subatomic Particles (mid 1900’s) Particle Symbols Relative Charge Mass # Relative Mass Actual Mass Electron -1 0 1/1837 9.11X10-31kg Proton +1 1 1 1.67X10-27 kg Neutron 0 1 1 1.68X10-27 kg Beta -1 0 1/1837 9.11X10-31kg Alpha a +2 4 4 6.64X10-27 kg p+ e- no 4 1 0 0 0 1 He e b e H n Other Related Particles (mid 1900’s) 2 -1 -1 1 -1 0 e-

  14. Atomic Number: Atomic number is the number of protons in the nucleus of an element. Since in neutral atoms the number of protons is equal to the number of electrons, atomic number also indicates the number of electrons in a single atom of an element. The Periodic Table is arranged according to atomic number. For example, hydrogen is element 1 and has 1 proton in its nucleus. Similarly, helium is element 2 and has 2 protons in its nucleus. How many protons does element number 25 contain? What is the name of this element? How many electrons does element 25 contain? What is the charge of one atom of element 25?

  15. Isotopes are atoms with the same number of protons (same element) but with a different number of neutrons. Na-24 (11 + 13 = 24) Na-23 (11 + 12 = 23)

  16. Identifying Specific Isotopes (Nuclides) Mass Number (total number of protons and neutrons in a nucleus) Element Symbol Atomic Number (total number of protons in a nucleus) Mass # = Atomic # + # Neutrons # Neutrons = Mass #-Atomic # or # Neutrons = A-Z

  17. Use a periodic table to help you fill in the missing information in the following tables.

  18. Unified AMU: unified atomic mass unit is a unit of mass based on 12C. The mass of one atom of 12C = 12 u so 1 u = 1/12 of the mass of one atom of 12C On this scale, 1 proton = 1.007276 u 1 neutron = 1.008665 u 1 electron = 0.0005486 u The Mass of “C” on the periodic table is 12.011 u because pure carbon is made up of both 12C (12 u exactly) and 13C (close to 13 u). The mass of “C” is a weighted average of all isotopes of C!

  19. Data for the known isotopes of Magnesium Isotope Abundance Mass 24Mg 78.99% 23.985 u 25Mg 10.00% 24.986 u 26Mg 11.01% 25.982 u Average Mass of Mg: (78.99%)*(23.985 u) + (10.00%)*(24.986 u) + (11.01%)*(25.982 u) 24.3049697 u Average Atomic Mass: the weighted average of the atomic masses of all naturally occurring isotopes of an element Average Mass = [(% abundance of isotope one)*(mass of isotope one) + (% abundance of isotope two)*(mass of isotope two)] abundance*atomic mass 24.30 u

  20. One Mole is the number of atoms in exactly 12 g of carbon-12. Avogadro’s Number is the number of particles in exactly one mole and has a value of 6.022X1023 The definition of “mole” and Avogadro’s number allows us to “count” atoms by weighing them when we know what the mass of one mole of a substance is. Molar Mass is the mass in grams of one mole of a pure substance. When the atomic masses on the periodic table are expressed in grams (instead of “u”), the atomic mass represents the mass of one mole of that element. Important Equivalence Statements: 1 mol = 6.022X1023 particles (atoms, molecules, ions, etc…) 1 mol = atomic mass in grams for each element on the periodic table for example: 24.3050 g Mg = 1 mol Mg and 55.845 g Fe = 1 mol Fe Remember that Equivalence Statements can be used to convert from one set of units to another!

  21. Example (like Sample Problem B on page80) What would be the mass of 0.3742 mol of aluminum?

  22. Example (like Sample Problem C on page 81) How many moles of copper are present in a sample of copper that has a mass of 127.32g?

  23. Example (like Sample Problem D on page 82) A chemist has a sample of gold that contains 4.37X1022 atoms of gold, how many moles of gold does she have in the sample?

  24. Example (like Sample Problem E on page 82) A chemist has a sample of iron that contains 5.27X1024 atoms of iron, what would the mass of the sample be?

  25. Example (extension-like conversion factor problem on page 84 and Sample Problem C on page 81) A chemist has a sample of silicon that has a mass of 84.39 mg. How many moles of silicon are in the sample?

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