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  1. Hg(NO3)2(aq) + 2KI(aq) HgI2(s) + 2KNO3(aq) Conservation of Mass and the Law of Definite Proportions Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions. 3.25 g + 3.32 g = 6.57 g 4.55 g + 2.02 g = 6.57 g

  2. Elements and Compounds • Elements can combine to form compounds • Compounds can be separated into elements • Red powder (HgO) is heated to give a colorless gas and a silver liquid

  3. Law of Conservation of mass Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions. Hg(NO3)2 + 2 KI HgI2 + 2 KNO3

  4. Conservation of Mass and the Law of Definite Proportions Law of Definite Proportions: Different samples of a pure chemical substance always contain the same proportion of elements by mass. By mass, water is: 88.8 % oxygen 11.2 % hydrogen

  5. Dalton’s Atomic Theory • Elements are made up of tiny particles called atoms. • Each element is characterized by the mass of its atoms. • Atoms of the same element have the same mass, but atoms of different elements have different masses.

  6. Dalton’s Atomic Theory • Atoms can join together to form compounds, combining in small whole-number ratios. • Chemical reactions only rearrange the way that atoms are combined; the atoms themselves don’t change.

  7. Dalton’s Law of Multiple Proportions Law of Multiple Proportions: Elements can combine in different ways to form different substances, whose mass ratios are small whole-number multiples of each other. nitric oxide: nitrous oxide: 8 grams oxygen per 7 grams nitrogen 16 grams oxygen per 7 grams nitrogen

  8. Dalton’s Atomic Theory and the Law of Multiple Proportions Law of Multiple Proportions: Elements can combine in different ways to form different substances, whose mass ratios are small whole-number multiples of each other.

  9. The Structure of Atoms: Electrons Cathode-Ray Tubes: J. J. Thomson (1856-1940) proposed that cathode rays must consist of tiny negatively charged particles. We now call them electrons.

  10. Rutherford Experiment: the Nuclear Atom Rutherford proposed that the atom must consist mainly of empty space with the mass concentrated in a tiny central core—the nucleus.

  11. The Structure of Atoms: Protons and Neutrons The charge of the proton is opposite in sign but equal to that of the electron The mass of the atom is primarily in the nucleus

  12. Atomic Number Atomic Number (Z): Number of protons in an atom’s nucleus. Equivalent to the number of electrons around the atom’s nucleus. Mass Number (A): The sum of the number of protons and the number of neutrons in an atom’s nucleus. Isotope: Atoms with identical atomic numbers but different mass numbers.

  13. Atomic Numbers and Isotopes Isotope: Atoms with identical atomic numbers but different mass numbers.

  14. 14 12 C C 6 6 Atomic Number and Mass Number carbon-12 mass number 6 protons 6 electrons 6 neutrons atomic number carbon-14 mass number 6 protons 6 electrons 8 neutrons atomic number

  15. Write isotopic symbol for an atom of nitrogen with 7 protons and 8 neutrons: Atomic Number = 7 Mass Number = 7 + 8 =15

  16. Calculate the number of protons, neutrons and electrons in an atom of 1223Mg Number of protons = atomic number = 12 Number of neutrons = mass # minus atomic # = 23-12=11 Uncharged atoms: # protons = # electrons So: 12 protons 11 neutrons 12 electrons

  17. How many protons, neutrons and electrons are there in an atom of 1122Na+1? Atomic # 11 = 11 protons Mass # - atomic # = 22-11 = 11 neutrons +1 charge implies one more proton than electron, so 10 electrons 11 protons 11 neutrons 10 electrons

  18. Most elements in nature are uniform mixtures of two or more kinds of atoms with slightly different masses Atoms of the same element with different masses are called isotopes For example: there are 3 isotopes of hydrogen and 4 isotopes of iron Chemically, isotopes have virtually identical properties Isotopes 20

  19. Atomic Mass and Isotopes Three forms of Carbon occur in nature: 12C, 13C and 14C Most carbon is 12C, but the others do exist Atomic Mass: The weighted average of the isotopic masses of the element’s naturally occurring isotopes.

  20. Atomic Mass Why is the atomic mass of the element carbon 12.01 amu? carbon-12: 98.89 % natural abundance 12 amu carbon-13: 1.11 % natural abundance 13.0034 amu mass of carbon = (12 amu)(0.9889) + (13.0034 amu)(0.0111) = 11.87 amu + 0.144 amu = 12.01 amu

  21. Atomic Masses and the Mole Avogadro’s Number (NA): One mole of any substance contains 6.022 x 1023 formula units. Molar Mass: The mass in grams of one mole of any element. It is numerically equivalent to its atomic mass.

  22. Avogadro’s Number and the Mole Avogadro’s Number (NA): One mole of any substance contains 6.022 x 1023 formula units. One mole of any substance is equivalent to its molecular or formula mass. HCl: 1 mole = 36.5 g C2H4: 1 mole = 28.0 g

  23. How many moles of carbon are there in36.0 grams of carbon? 1 mole carbon = 12.0 grams carbon

  24. How many moles of sodium are there in 56.25 grams? 1 mole sodium = 23.0 grams sodium

  25. How much will 3.25 moles of silver weigh?

  26. Mass in grams Number of atoms Number of moles Avogadro’s Number 6.02 X 1023 Atomic mass in grams Conversion between grams, and moles, and atoms

  27. Mass # of atoms # of moles 6.02 X 1023 Atomic mass How many nitrogen atoms are there in 5.25 grams of nitrogen atoms?

  28. Matter Can Be Classified By Its Properties: • Matter is either a pure substance or a mixture • Mixtures may be separated using physical methods such as chromatography, filtration, sieving

  29. What Is An Element? • Elements - substances that cannot be decomposed into simpler substances • shown on the periodic table as symbols: “K” for potassium and “Na” for sodium • made of identical atoms, either singly or in groups

  30. What Is A Compound? • Compounds - formed from two or more atoms of different elements combined in a fixed proportion • Have different characteristics than the elements that compose them • Can be broken down into elements by some chemical changes

  31. Mixtures • mixtures consist of varying amounts of two or more elements or compounds • Homogeneous mixtures or “solutions”- have the same properties throughout the sample • Brass, tap water • Heterogeneous mixtures-consist of two or more phases • Salad dressing, pepperoni pizza

  32. Iron, Sulfur and Iron Sulfide

  33. Compounds and Mixtures Variable composition Same composition Variable properties Similar properties Copyright © 2008 Pearson Prentice Hall, Inc.

  34. Compounds and Mixtures The same number of protons Cannot be separated physically Copyright © 2008 Pearson Prentice Hall, Inc.

  35. Molecules, Ions, and Chemical Bonds Covalent Bond: Results when two atoms share several (usually two) electrons. Typically a nonmetal bonded to a nonmetal.

  36. When elements for compounds, the properties can change drastically

  37. Molecules, Ions, and Chemical Bonds Ionic Bond: A transfer of one or more electrons from one atom to another. An electrostatic attraction between charged particles.Typically a metal bonded to a nonmetal. Ion: A charged particle. Cation: A positively charged particle. Metals tend to form cations. Anion: A negatively charged particle. Nonmetals tend to form anions.

  38. Na + Cl Na1+ + Cl1- Molecules, Ions, and Chemical Bonds In the formation of sodium chloride, one electron is transferred from the sodium atom to the chlorine atom. 11 protons 11 electrons 17 protons 17 electrons 11 protons 10 electrons 17 protons 18 electrons

  39. Molecules, Ions, and Chemical Bonds

  40. Which of the following compounds would be ionic and which would be molecular (covalent)? • BaF2 ionic • SF4 covalent • NO2 covalent • MgO ionic

  41. Which of the following drawings most likely represents a molecular (covalent) compound?

  42. Write a molecular formula for each of the following compounds: White = H Red = O Black = C Blue = N C3H7NO2 C2H6O2 C3H4O2

  43. Naming Chemical Compounds Charges for Typical Main-Group Ions 1+ 2+ 3+

  44. Naming Chemical Compounds Ionic Compound: A neutral compound in which the total number of positive charges must equal the total number of negative charges. Binary Ionic Compounds sodium chloride: Na1+ Cl1- NaCl magnesium oxide: Mg2+ O2- MgO aluminum sulfide: Al3+ S2- Al2S3 Copyright © 2008 Pearson Prentice Hall, Inc.

  45. Calculating Charge of Transition Elements FeCl3 CoO Work backwards from nonmetal Metal charge must equal total negative Iron (III) chloride cobalt (II) Oxide

  46. Calculate the charge on the following transition metals

  47. Naming Chemical Compounds Stock System: Use Roman numerals in parentheses to indicate the charge on metals that form more than one kind of cation. Binary Ionic Compounds iron(III) oxide: Fe3+ O2- Fe2O3 tin(II) chloride: Sn2+ Cl1- SnCl2 lead(II) fluoride: Pb2+ F1- PbF2

  48. Write formulas for the following: Magnesium oxide MgO Cobalt (III) sulfide Co2S3 Aluminum chloride AlCl3 Titanium (IV) oxide TiO2

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