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  • 1. The Periodic Table and Periodic Law Chemistry Chapter 6
  • 2. Main Ideas
    • Periodic trends in the properties of atoms allow us to predict physical and chemical properties.
    • The periodic table evolved over time as scientists discovered more useful ways to compare and organize elements.
    • Elements are organized into different blocks in the periodic table according to their electron configurations.
    • Trends among elements in the periodic table include their size and their ability to lose or attract electrons
  • 3. Development of the Modern Periodic Table
    • Objectives:
    • Trace the development of the periodic table
    • Identify key features of the periodic table
  • 4. Development of the Periodic Table
    • In the 1700’s, Lavoisier compiled a list of all the known elements of the time.
      • 33 elements
  • 5. Development of the Periodic Table
    • The 1800s brought large amounts of information and scientists needed a way to organize knowledge about elements.
    • Advent of electricity – break down compounds
    • Development of the spectrometer – identify newly isolated elements
  • 6. Development of the Periodic Table
    • The 1800s brought large amounts of information and scientists needed a way to organize knowledge about elements.
    • Industrial revolution – new chemistry based ingredients and compounds.
    • 70 known elements by 1870
  • 7. Development of the Periodic Table
    • The 1800s brought large amounts of information and scientists needed a way to organize knowledge about elements.
    • John Newlands proposed an arrangement where elements were ordered by increasing atomic mass.
  • 8. Law of Octaves
    • Newlands (1864) noticed when the elements were arranged by increasing atomic mass, their properties repeated every eighth element.
  • 9. Law of Octaves
    • Octaves was used due to the musical analogy, but was widely dismissed.
    • Some elements didn’t follow the pattern
  • 10. The Periodic Table
    • Meyer and Mendeleev both demonstrated a connection between atomic mass and elemental properties.
  • 11. The Periodic Table
    • Mendeleev’s Table – A Russian scientist – gets the most credit because he published first.
      • Arranged elements by increasing mass and columns with similar properties.
      • Predicted the existence and properties of undiscovered elements.
      • Still some inconsistencies.
  • 12. The Periodic Table
    • Moseley discovered that each element had a distinct number of protons.
      • Once rearranged by increasing atomic number, the table resulted in a clear periodic pattern.
  • 13. The Periodic Table
    • Periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number is called periodic law.
  • 14. Development of the Periodic Table
  • 15. The Modern Periodic Table
    • The modern periodic table contains boxes which contain the element's name, symbol, atomic number, and atomic mass.
  • 16. The Modern Periodic Table
    • Rows of elements are called periods. (total of 7)
    • Columns of elements are called groups. (total of 18)
    • Elements in groups 1,2, and 13-18 possess a wide variety of chemical and physical properties and are called the representative elements.
    • Elements in groups 3-12 are known as the transition elements .
  • 17. Types of Elements
    • Elements are classified as metals, non-metals, and metalloids.
    • Metals are made up of most of the representative elements and all of the transition elements.
      • They are generally shiny when smooth and clean, solid at room temperature, and good conductors of heat and electricity.
      • Most are Ductile and Malleable –
        • Ductile – the ability to be drawn into wire.
        • Malleable – the ability to be pounded into sheets
  • 18. Types of Elements
    • Elements are classified as metals, non-metals, and metalloids.
    • Alkali metals are all the elements in group 1 except hydrogen, and are very reactive.
    • Alkaline earth metals are in group 2, and are also highly reactive.
  • 19. Types of Elements
    • The transition elements (groups 3 - 12) are divided into transition metals and inner transition metals.
      • The two sets of inner transition metals are called the lanthanide series and actinide series and are located at the bottom of the periodic table.
      • Lanthanides are phosphors – elements that emit light when struck by electrons.
  • 20. The Modern Periodic Table
    • Non-metals are elements that are generally gases or brittle, dull-looking solids, and poor conductors of heat and electricity.
    • Group 17 is composed of highly reactive elements called halogens.
    • Group 18 gases are extremely unreactive and commonly called noble gases.
  • 21. The Modern Periodic Table
    • Metalloids have physical and chemical properties of both metals and non-metals, such as silicon and germanium. They are found along the stair step of the table starting with Boron
  • 22.  
  • 23. Questions What is a row of elements on the periodic table called? A. octave B. period C. group D. transition
  • 24. Questions What is silicon an example of? A. metal B. non-metal C. inner transition metal D. metalloid
  • 25. Practice Problems
    • Page 181 #1-7
  • 26. Classification of the Elements
    • Objectives:
    • Explain why elements in the same group have similar properties.
    • Identify the four blocks of the periodic table on their electron configuration.
  • 27. Organizing the Elements by Electron Configuration
    • Electron configuration determines the chemical properties of an element.
    • Recall electrons in the highest principal energy level are called valence electrons.
  • 28. Organizing the Elements by Electron Configuration
    • All group 1 elements have one valence electron.
    • All group 2 elements have two valence electrons.
  • 29. Organizing the Elements by Electron Configuration
  • 30. Organizing the Elements by Electron Configuration
    • The energy level of an element’s valence electrons indicates the period on the periodic table in which it is found.
    • The number of valence electrons for elements in groups 13-18 is ten less than their group number.
    • After the s-orbital is filled, valence electrons occupy the p-orbital.
    • Groups 13-18 contain elements with completely or partially filled p orbitals.
  • 31. Organizing the Elements by Electron Configuration
  • 32. Organizing the Elements by Electron Configuration
  • 33. Organizing the Elements by Electron Configuration
    • The d-block contains the transition metals and is the largest block.
      • There are exceptions, but d-block elements usually have filled outermost s orbital, and filled or partially filled d orbital.
      • The five d orbitals can hold 10 electrons, so the d-block spans ten groups on the periodic table.
  • 34. Organizing the Elements by Electron Configuration
    • The f-block contains the inner transition metals.
      • f-block elements have filled or partially filled outermost s orbitals and filled or partially filled 4f and 5f orbitals.
      • The 7 f orbitals hold 14 electrons, and the inner transition metals span 14 groups.
  • 35. Practice Problems
    • Page 186 #8-15
  • 36. Periodic Trends
    • Objectives:
    • Compare period and group trends of several properties.
    • Relate period and group trends in atomic radii to electron configuration
  • 37. Atomic Radius
    • Atomic radius – is determined by the amount of positive charge in the nucleus and the number of valence electrons of an atom. It is usually measured in picometers (10 -12 ).
    • For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element.
    • For nonmetals, the atomic radius is the distance between nuclei of identical atoms.
  • 38. Atomic Radius
  • 39. Atomic Radius
    • The periodic trend: decreases from left to right (periods) and increases top to bottom (groups) due to the increasing positive charge in the nucleus.
  • 40. Atomic Radius
  • 41. Atomic Radius
    • Atomic radius generally increases as you move down a group.
    • The outermost orbital size increases down a group, making the atom larger.
    • Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons.
  • 42. Ionic Radius
    • Ions – atom(s) that gain or lose one or more electrons to form a net charge.
    • Ionic radius is the radius of a charged atom.
    • When atoms lose electrons and form positively charged ions, they always become smaller.
      • Lost electrons are usually valence electrons and could leave the outer orbital empty and therefore smaller.
      • Electrostatic repulsion between remaining electrons decreases and pulls closer to nucleus.
  • 43. Ionic Radius
    • When atoms gain electrons and forms a negatively charged ion, they become larger.
      • Increased electrostatic repulsion increases distance of outer electrons.
  • 44. Ionic Radius
    • Periodic Trend: radius of an ion decreases from left to right (periods) until charge changes and then the radii increases dramatically. After the change, the radius continues to decrease. Ionic radii increases top to bottom (groups) until change in charge.
  • 45. Ionic Radius
  • 46. Ionization Energy
    • Ionization energy is the energy needed to remove an electron from the positive charge of the nucleus of a gaseous atom. (how strongly a nucleus holds on to an electron.)
    • First ionization energy is the energy required to remove the first electron.
    • Removing the second electron requires more energy, and is called the second ionization energy.
  • 47. Ionization Energy
    • Atoms with large ionization energies have a strong hold of its electrons and are less likely to form positive ions.
    • Atoms with low ionization energies lose their outer electrons easily and readily form positive ions.
    • The ionization at which the large increase in energy occurs is related to the number of valence electrons.
  • 48. Ionization Energy
    • Periodic Trend : First ionization energy increases from left to right across a period. First ionization energy decreases down a group because atomic size increases and less energy is required to remove an electron farther from the nucleus.
  • 49. Ionization Energy
    • The octet rule states that atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons. The octet rule is useful for predicting what types of ions an element is likely to form.
  • 50. Ionization Energy
  • 51. Electronegativity
    • Electronegativity of an element indicates its relative ability to attract electrons in a chemical bond. Measured in Paulings: numbers 4 and less.
  • 52. Electronegativity
    • Periodic Trend: electronegativity decreases down a group and increases left to right across a period.
  • 53. Questions The lowest ionization energy is the ____. A. first B. second C. third D. fourth
  • 54. Questions The ionic radius of a negative ion becomes larger when: A. moving up a group B. moving right to left across period C. moving down a group D. the ion loses electrons
  • 55. Practice Problems
    • Page 189 #16-19; Page 194 #20-23
  • 56. Key Concepts
    • The elements were first organized by increasing atomic mass, which led to inconsistencies. Later, they were organized by increasing atomic number.
    • The periodic law states that when the elements are arranged by increasing atomic number, there is a periodic repetition of their chemical and physical properties.
    • The periodic table organizes the elements into periods (rows) and groups (columns); elements with similar properties are in the same group.
  • 57. Key Concepts
    • Elements are classified as either metals, nonmetals, or metalloids.
    • The periodic table has four blocks (s, p, d, f).
    • Elements within a group have similar chemical properties.
    • The group number for elements in groups 1 and 2 equals the element’s number of valence electrons.
    • The energy level of an atom’s valence electrons equals its period number.
  • 58. Key Concepts
    • Atomic and ionic radii decrease from left to right across a period, and increase as you move down a group.
    • Ionization energies generally increase from left to right across a period, and decrease as you move down a group.
    • The octet rule states that atoms gain, lose, or share electrons to acquire a full set of eight valence electrons.
    • Electronegativity generally increases from left to right across a period, and decreases as you move down a group.
  • 59. Chapter Questions
    • The actinide series is part of the
    • A. s-block elements.
    • B. inner transition metals.
    • C. non-metals.
    • D. alkali metals.
  • 60. Chapter Questions
    • In their elemental state, which group has a complete octet of valence electrons?
    • A. alkali metals
    • B. alkaline earth metals
    • C. halogens
    • D. noble gases
  • 61. Chapter Questions
    • Which block contains the transition metals?
    • A. s-block
    • B. p-block
    • C. d-block
    • D. f-block
  • 62. Chapter Questions
    • An element with a full octet has how many valence electrons?
    • A. two
    • B. six
    • C. eight
    • D. ten
  • 63. Chapter Questions
    • How many groups of elements are there?
    • A. 8
    • B. 16
    • C. 18
    • D. 4
  • 64. Chapter Questions
    • Which group of elements are the least reactive?
    • A. alkali metals
    • B. inner transition metals
    • C. halogens
    • D. noble gases
  • 65. Chapter Questions
    • On the modern periodic table, alkaline earth metals are found only in ____.
    • A. group 1
    • B. s-block
    • C. p-block
    • D. groups 13–18
  • 66. Chapter Questions Bromine is a member of the A. noble gases. B. inner transition metals. C. earth metals. D. halogens.
  • 67. Chapter Questions
    • How many groups does the d-block span?
    • A. two
    • B. six
    • C. ten
    • D. fourteen
  • 68. THE END
  • 69. Chapter Questions
  • 70. Chapter Questions
  • 71. Chapter Questions
  • 72. Chapter Questions
  • 73. Chapter Questions