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The Periodic Properties of the Elements. By Lauren Querido, Chris Via, Maggie Dang, Jae Lee. The Founders of the Periodic Table. Luthar Meyer. Dmitri Mendeleev.

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the periodic properties of the elements

The Periodic Properties of the Elements

By Lauren Querido, Chris Via, Maggie Dang, Jae Lee

the founders of the periodic table
The Founders of the Periodic Table

Luthar Meyer

Dmitri Mendeleev

7 1 developing the periodic table
7.1 Developing the Periodic Table
  • Dmitri Mendeleev (1869)- and Luthar Meyer Published very similar documents to classify the elements. And were the first to make the modern periodic table
    • Used chemical and physical properties to classify
  • Henry Moseley (1887-1915)- Developed concept of atomic numbers
    • Found that frequency increases as the atomic mass increases
7 2 electron shells and size of atoms
7.2 Electron Shells and Size of Atoms
  • Electron Shells in Atoms
    • Gilbert N. Lewis – electrons are arranged in shells surrounding the nucleus.
  • Atomic sizes-
Bonding Atomic Radius- the distance between the center of two bonding atoms

practice problem 1
Practice Problem #1
  • Predict the lengths of C-S, C-H, and S-H bonds in this molecule
  • Radius of C = 0.77 Å
  • Radius of S = 1.02 Å
  • Radius of H = 0.37 Å
  • When determining the bonding radius, you add the radius of the bonding atoms together
answer to practice problem 1
Answer to Practice Problem #1
  • C-S bond length = radius of C + radius of

= 0.77 Å + 1.02 Å = 1.79 Å

  • C-H bond length = 0.77 Å + 0.37 Å = 1.14 Å
  • S-H bond length = 1.02 Å +0.37 Å = 1.39 Å
7 2 continued
7.2 continued
  • When moving across a row, the number of core electrons stay the same but the nuclear charge increases
  • The effective nuclear charge increases even though the quantum number remains the same
  • Shielding is the process of blocking the protons effective charge on the outermost electrons
7 3 ionization energy
7.3 Ionization Energy
  • Ionization Energy – to remove an electron from the ground state
  • Second Ionization – removing the 2nd electron from the ground state
  • I1<I2<I3 and so forth; It increases in magnitude
  • The greater effective nuclear charge, the greater the ionization energy
7 3 cont
7.3 cont..
  • There is a sharp increase in ionization energy when an inner shell electron is removed
  • Periodic Trends
  • Within each row, the ionization energy increases with atomic number
  • Within a group, the ionization energy generally decreases with increasing atomic number
7 3 cont1
7.3 cont..

3. The ionization energy of transition elements & f-block metals increase slowly as you read from left to right.

  • The transition in ionization energy are affected by how strong an electron is attracted to an atom
  • It is affected by the effective nuclear charge and the average distance from the nucleus.
  • The irregularities are explained through the periodic table
  • Electrons in the s orbital are more effective at shielding than in the p orbital
7 4 electron affinities
7.4 Electron Affinities
  • Positive ionization energy = energy put into atom in order to remove electrons
  • Electron affinity = attraction of change in energy when the electron is added
  • Most atoms = energy is released when electron is added
  • A positive electron affinity, an ion will not form
7 4 cont
7.4 cont..
  • On the periodic table, electron affinity becomes negative towards halogen (closest to being stable)
  • The electron affinity does not change when they move down a group (noble gases)
element classification http www elementsdatabase com images periodic table1 gif
Element Classification
7 5 metals nonmetals and metalloids
7.5 Metals, Nonmetals, and Metalloids
  • Metals
    • Tend to have low ionization energies and lose electrons when they undergo chemical reaction
    • Most metal oxides are basic oxides that dissolve in water react to for metal hydroxides

Metal Oxide + Water  Metal Hydroxide

    • Metal oxides show their basicity by reacting with acids to form water and salts

Metal Oxide + Acid  Salt + Water

  • Characteristics of Metals
    • Have a shiny luster
    • Various colors
    • Solids are malleable and ductile
    • Good conductors of heat and electricity
    • Most metal oxides are ionic solids that are basic
  • Nonmetals
    • Tend to gain electrons and become anions

Metal + Nonmetal  Salt

    • Most nonmetal oxides are acidic oxides that dissolve in water react to form acids

Nonmetal Oxide + Water  Acid

    • The acidity of nonmetal oxides is shown by the fact they dissolve in basic solutions to form salts

Nonmetal Oxide+ Base  Salt + Water

  • Characteristics of Nonmetals
    • Do not have a luster
    • Various colors
    • Solids are usually brittle; some are hard, and some are soft
    • Poor conductors of heat and electricity
    • Most nonmetallic oxides are molecular substances that form acidic solutions
  • Metalloids
    • Have properties intermediate between nonmetals and metals

7 6 group trends for the active metals
7.6 Group Trends for the Active Metals
  • Group 1A: The Alkali Metals (most active)
    • Metallic Characteristics
      • Silvery
      • Metallic luster & high thermal
      • Electrical conductivities
      • Have low densities & melting points
      • Very reactive b/c they want to lose 1 electron to form ions with a 1+ charge so it becomes more stable
7 6 cont
7.6 cont..
  • As you move down a group
    • Atomic radius increases
    • 1st ionization energy decreases
7 6 cont1
7.6 cont..
  • Group 2A: The Alkaline Earth Metals
    • Properties of Alkaline Earth Metals
      • Harder
      • More Dense
      • Melt at higher temps
      • Highly Reactive
    • Compared to alkali Metals, Alkaline Earth metals..
      • Have lower 1st ionization energies
      • Are less reactive
  • Group 6A
  • Oxygen is a colorless gas at room temperature while all the other elements in this group are solid.
  • Oxygen has two main forms: 02=“oxygen” and 03=“ozone”.
  • This is an example of an allotrope, it has different forms of the same element.
  • The most stable form of sulfur is S8, It is a yellow solid.
  • All of the elements in this group have the tendency to gain electrons form other elements.
  • Group 7A: Halogens
  • Halogens is named Greek words, “halos” and “gennao” meaning salt formers.
  • Fluorine and Chlorine are gases, bromine is a liquid, and iodine is a solid at room temperature.
  • These elements melting and boiling points increase with atomic number.
  • These elements have highly negative electron affinities because they have the need to gain electrons from other elements.
the fluorine atom is very reactive
The Fluorine atom is very reactive!

  • Group 8A: Noble Gases
  • All of the elements are nonmetals at room temperature and they are monatomic
  • They are very unreactive because they have completely filled s and p orbitals.
  • They also have very large 1st ionization energies.