chapter 6 notes - Chemistry-WEB VERSION

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					Chemistry – Chapter 6 – The Periodic Law
Section 6-1: History of the Periodic Table
Objectives:
     Explain the roles of Mendeleev and Moseley in the development of the periodic table.
     Describe the modern periodic table.
     Explain how the periodic law can be used to predict the physical and chemical properties of elements.
     Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic
        number.



Russian chemist Dmitri Mendeleev – 1869 – created a periodic table of the elements and grouped the elements in
order of increasing atomic mass.


English scientist Henry Moseley – 1911 (while working with Ernest Rutherford-not pictured) – discovered that
elements in the periodic table fit into patterns better when they were arranged in increasing atomic number.


1) Periodic law – the physical and chemical properties of the elements are periodic functions of their atomic
   numbers.
2) Periodic table – an arrangement of the elements in order of their atomic numbers so that elements with similar
   properties fall in the same column, or group.
3) Noble gases – unreactive gasses –group 18 elements.
4) Lanthanides – the 14 elements with atomic numbers from 58-71.

Common Properties of the Lanthanides

Lanthanides share the following common properties:

     Silvery-white metals that tarnish when exposed to air, forming their oxides.
     Relatively soft metals. Hardness increases somewhat with higher atomic number.
     High melting points and boiling points.
     Very reactive.

5) Actinides – the 14 elements with atomic numbers from 90-103

Common Properties of the Actinides

Actinides share the following common properties:

     All are radioactive.
     The metals tarnish readily in air.
     Actinides are very dense metals with distinctive structures.
     They react with boiling water or dilute acid to release hydrogen gas.

6) Periodicity – repeating patterns on the periodic table
Section 6-2: Electron Configuration and the Periodic Table
Objectives:
     Describe the relationship between electrons in sublevels and the length of each period of the periodic table.
     Locate and name the four blocks of the periodic table. Explain the reasons for these names.
     Discuss the relationship between group configurations and group numbers.


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     Describe the locations in the periodic table and the general properties of the alkali metals, the alkaline-earth
      metals, the halogens, and the noble gases.
1) Periods – horizontal rows on the periodic table.
2) Groups – vertical columns on the periodic table.




s-block elements:
1) They are chemically reactive metals (groups 1 and 2)
2) Alkali metals – the elements of group 1 of the periodic table (lithium, sodium, potassium, rubidium, cesium, and
    francium)
    a) They have a silvery appearance
    b) They are soft enough to be cut with a knife
    c) Because they are so reactive, they are not found in nature as free elements
    d) They combine vigorously with most nonmetals
    e) They react strongly with water to produce hydrogen gas and aqueous solutions of substances known as alkalis
    f) Since they are so reactive with air or moisture, they are usually stored in kerosene.
3) Alkaline-earth metals – elements of group 2 of the periodic table (beryllium, magnesium, calcium, strontium,
    barium, and radium)
    a) They contain 2 electrons in their outermost s sublevel
    b) They are harder, denser, and stronger than the alkali metals and also have higher melting points
    c) They are less reactive than the alkali metals, but they are also too reactive to be found in nature as free
        elements.
4) d-Block Elements:
    a) groups 3-12
    b) transition metals – d-block elements that are metals with typical metallic properties.
        i) Good conductors of heat and electricity
        ii) They have a high luster
        iii) They are typically less reactive than the alkali metals and the alkaline-earth metals
        iv) Palladium, platinum, and gold are among the least reactive of all the elements
5) p-Block Elements:
        i) groups 13-18
    b) main-group (or representative) elements – the p-block elements together with the s-block elements
        i) generally harder and denser than the s-block metals, but softer and less dense than the d-block metals
        ii) found in nature only in the form of compounds (exception = bismuth)


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   c) metalloids (semiconductors) – boron, silicon, germanium, arsenic, antimony, and tellurium (6 elements) –
       sometimes polonium and astatine are also considered to be metalloids.
       i) Brittle solids with some properties of metals and some of nonmetals
       ii) They have electrical conductivity between that of metals and nonmetals
   d) Halogens – the elements of group 17 – (fluorine, chlorine, bromine, iodine, and astatine)
       i) Halogens are the most reactive nonmetals and react vigorously with most metals to form salts.
       ii) The have seven electrons in their outer energy levels (they need one more electron to be stable).
6) f-Block Elements: Lanthanides and Actinides
   a) between groups 3 and 4 (the two rows at the bottom of the periodic table)
   b) shiny metals similar in reactivity to the group 2 alkaline-earth metals
   c) The actinides are all radioactive and thorium, protactinium, uranium and neptunium have been found
       naturally on Earth. The other actinides are only known as lab-made elements.

Section 6-3: Electron Configuration and Periodic Properties
Objectives:
     Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity.
     Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for
        these variations.
     Define valence electrons, and state how many are present in atoms of each main-group element.
     Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of
        the main-group elements.
                                                                1) Atomic radii – one-half the distance between the
                                                                    nuclei of identical atoms that are bonded together.
                                                                    Atomic radii decrease from left to right across a
                                                                    period and increase down a group.
                                                                2) Periodic trends – the trend to smaller atoms
                                                                    across a period is caused by the increasing positive
                                                                    charge of the nucleus.
                                                                3) Group trends – in general, the atomic radii of the
                                                                    main-group elements increase down a group
                                                                4) Ionization energy (IE) – the energy required to
                                                                    remove one electron from a neutral atom of an
                                                                    element. In general, first ionization energies
                                                                    increase across a period and decrease down a
                                                                    group.
                                                                    a) Second and third ionization energies – the
                                                                        energies for removal of additional electrons
                                                                        form an atom.
    b) Each successive electron removed from an ion feels an increasingly stronger effective nuclear charge (the
        nuclear charge minus the electron shielding).
5) Ion – an atom or group of bonded atoms that has a positive (cation), or negative (anion) charge.
6) Ionization – any process that results in the formation of an ion.
7) Electron affinity – the energy change that occurs when an electron is acquired by a neutral atom.
8) Ionic radii – cations are smaller and anions are larger than the atoms from which they are formed. There is a
    gradual increase of ionic radii down a group.
9) Valence electrons – the electrons available to be lost, gained, or shared in the formation of chemical compounds.
10) Electronegativity – a measure of the ability of an atom in a chemical compound to attract electrons. Fluorine is
    the most electronegative element. Electronegativities tend to increase across each period, although there are
    exceptions, and they tend to decrease down a group or remain about the same.




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