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Chapter 4

Chapter 4. Elements, Atoms, and Ions. Ch#4 Contents. 4.1 The Elements 4.2 Symbols for the Elements 4.3 Dalton’s Atomic Theory 4.4 Formulas of Compounds 4.5 The Structure of the Atom 4.6 Modern Concept of Atomic Structure 4.7 Isotopes 4.8 Introduction to the Periodic Table

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Chapter 4

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  1. Chapter 4 Elements, Atoms, and Ions

  2. Ch#4 Contents 4.1 The Elements 4.2 Symbols for the Elements 4.3 Dalton’s Atomic Theory 4.4 Formulas of Compounds 4.5 The Structure of the Atom 4.6 Modern Concept of Atomic Structure 4.7 Isotopes 4.8 Introduction to the Periodic Table 4.9 Natural States of the Elements 4.10 Ions 4.11 Compounds that Contain Ions

  3. Elements A substance that cannot be broken down by chemical means. Elements are defined by the number of protons they possess. • 115 known: 88 found in nature, others are man made. • Most of the know elements are found on the Periodic Chart handing on the wall. • One or two letters are used to represent an element and they are called the symbol • Some elements are solids (black), some are liquids (blue) and some are gases (red)

  4. Element Distribution in Earth’s Surface

  5. Symbols • Each element has a unique one- or two-letter symbol. • First letter is always capitalized and the second is not. • The symbol usually consists of the first one or two letters of the element’s name. • Examples: Oxygen O Krypton Kr • Sometimes the symbol is taken from the element’s original Latin, Greek, English, or German name. • Examples: Gold Au aurum Lead Pbplumbum Wolfram W tungsten

  6. Symbols Located on the front page of your textbook

  7. Dalton’s Atomic Theory In 1808 Dalton published A New System of Chemical Philosophy, where he presented his theory of atoms: • Each element is made up of tiny particles called atoms. • 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 of different elements combine with each other. A given compound always has the same relative numbers and types of atoms. • Chemical changes involve reorganization of the atoms, to different ratios.

  8. Dalton’s Atomic Theory Which of the following statements regarding Dalton’s atomic theory are still believed to be true? • Elements are made of tiny particles called atoms. • All atoms of a given element are identical. • A given compound always has the same relative numbers and types of atoms. • IV. Atoms are indestructible.

  9. Dalton’s Atomic Theory Which of the following statements regarding Dalton’s atomic theory are still believed to be true? • Elements are made of tiny particles called atoms. • All atoms of a given element are identical. • A given compound always has the same relative numbers and types of atoms. • IV. Atoms are indestructible.

  10. Dalton’s Atomic Theory Which of the following statements regarding Dalton’s atomic theory are still believed to be true? • Elements are made of tiny particles called atoms. • All atoms of a given element are identical. • A given compound always has the same relative numbers and types of atoms. • IV. Atoms are indestructible.

  11. Formulas • Compound – distinct substance that is composed of the atoms of two or more elements and always contains exactly the same whole number ratio of those elements. • Chemical Formulas – expresses the types of atoms and the number of each type in each formula unit, or molecule of a given compound. • Formula Examples • H2O (molecule) • C6H12O6 (molecule) • NaCl (Formula Unit)

  12. Metals

  13. The Periodic Table Metallic Properties • Shiny luster • Conduct, electricity and heat • Malleable • Ductile • Lose electrons easily

  14. The Periodic Table Nonmetallic Properties • Dull luster like charcoal • Do not conduct, electricity and heat • Brittle • Do not lose electrons easily

  15. Atomic Structure J. J. Thomson (1898—1903) • Postulated the existence of electrons using cathode-ray tubes. • The atom must also contain positive particles that balance exactly the negative charge carried by particles that we now call electrons.

  16. Cathode-ray Tube

  17. Atomic Structure William Thomson (Plum Pudding Model) Reasoned that the atom might be thought of as a uniform “pudding” of positive charge with enough negative electrons scattered within to counterbalance that positive charge.

  18. Rutherford's Gold Foil Experiment • Alpha particles strike gold foil • Most pass through • Some bounce back at sharp angles • Rutherford postulated that • Atoms are mostly empty space • Used our solar system as and example • His model is referred to as the planetary model

  19. Atomic Structure Ernest Rutherford (1911) • Explained the nuclear atom. • Atom has a dense center of positive charge called the nucleus. • Electrons travel around the nucleus at a relatively large distance. • A proton has the same magnitude of charge as the electron, but its charge is positive.

  20. Atomic Structure Rutherford and Chadwick (1932) • Most nuclei also contain a neutral particle called the neutron. • A neutron is slightly more massive than a proton but has no charge.

  21. Atomic Structure • Electrons – found outside the nucleus; negatively charged • Protons– found in the nucleus; positive charge equal in magnitude to the electron’s negative charge • Neutrons– found in the nucleus; no charge; virtually same mass as a proton

  22. Atomic Structure • The nucleus is: • Small compared with the overall size of the atom. • Extremely dense; accounts for almost all of the atom’s mass.

  23. Atomic Structure

  24. Atomic Structure Why do different atoms have different chemical properties? • The chemistry of an atom arises from its electrons. • Electrons are the parts of atoms that “intermingle” when atoms combine to form molecules. • It is the number of electrons that really determines chemical behavior.

  25. Isotopes • Atoms with the same number of protons but different numbers of neutrons. • Show almost identical chemical properties; chemistry of atom is due to its electrons. • In nature most elements contain mixtures of isotopes.

  26. Two Isotopes of Sodium

  27. Isotopes • X = the symbol of the element • Z = the atomic number (# of protons) • A = the mass number (# of protons and neutrons)

  28. Isotopes • C = the symbol for carbon • 6 = the atomic number (6 protons) • 14 = the mass number (6 protons and 8 neutrons) • C = the symbol for carbon • 6 = the atomic number (6 protons) • 12 = the mass number (6 protons and 6 neutrons)

  29. Isotopes

  30. Isotopes

  31. Isotopes

  32. Isotopes

  33. Isotopes

  34. Isotopes

  35. Isotopes

  36. Natural State of Elements • Most elements are very reactive. • Elements are not generally found in uncombined form. • Exceptions are: • Noble metals – gold, platinum and silver • Noble gases – Group 8

  37. Diatomic Molecules

  38. Allotropes • Different forms of a given element. • Example: • Solid carbon occurs in three forms. • Diamond • Graphite • Buckminsterfullerene

  39. Carbon Allotropes

  40. Ions • Atoms can form ions by gaining or losing electrons. • Metals tend to lose one or more electrons to form positive ions called cations and are named by using the name of the parent atom. • Nonmetals tend to gain electrons to form negative ions called anions and are named by using the root of the atom name followed by the suffix –ide.

  41. Common Oxidation States

  42. Common Oxidation States 1+ 2+ 1-

  43. Common Oxidation States By Group Number

  44. Sample Problem An ion with a 3+ charge contains 23 electrons. Which ion is it?

  45. Sample Problem An ion with a 3+ charge contains 23 electrons. Which ion is it? Fe3+ A certain ion X+ contains 54 electrons and 78 neutrons. What is the mass number of this ion?

  46. Sample Problem An ion with a 3+ charge contains 23 electrons. Which ion is it? Fe3+ A certain ion X+ contains 54 electrons and 78 neutrons. What is the mass number of this ion?

  47. Sample Problem An ion with a 3+ charge contains 23 electrons. Which ion is it? Fe3+ A certain ion X+ contains 54 electrons and 78 neutrons. What is the mass number of this ion? 133

  48. Ionic Compounds • Ions combine to form ionic compounds. • The simplest ratio between the ions is called the empirical formula or a formula unit. • Properties of ionic compounds • High melting points • Conduct electricity • If melted • If dissolved in water • Ionic compounds are electrically neutral. • The charges on the anions and cations in the compound must sum to zero.

  49. Ionic Compounds Ionic compounds do notexist as discrete molecules. Instead they exist as crystals where ions of opposite charges occupy positions known as lattice sites. Ions combine in the ratio that results in zero charge to form ionic compounds. Which ions are the smaller ones? Crystal Lattice of NaCl

  50. Formula Writing Rules Step 1 Write the symbols of the elements. Step 2 Assign oxidation numbers (O.N. = P – e) Step 3 Slide with Clyde! (number only) Step 4 Reduce if the compound is ionic Examples Write formulas for the following two elements Sodium and oxygen Sodium and nitrogen Calcium and oxygen Calcium and phosphorus

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