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					CHAPTER 6 THE PERIODIC TABLE

6.1 Organizing the Elements
6.2 Classifying the Elements
6.3 Periodic Trends
6.1 ORGANIZING THE ELEMENTS

Key Concepts

1.   How did chemists begin to organize the known elements?
2.   How did Mendeleev organize his periodic table?
3.   How is the modern periodic table organized?
4.   What are three broad classes of elements?
Searching for an Organizing Principle:

1. Only about 13 elements had been identified by 1700, these
   included copper, silver, and gold followed by hydrogen, nitrogen,
   and oxygen.
2. The problem was not only how to isolate the elements but how to
   organize them to establish some type of pattern.
3. In 1829, J.W. Dobereiner published a classification system in
   which elements were grouped in triads.
4. A triad is any of several sets of three chemically similar
   elements, the atomic weight of one of which is approximately
   equal to the mean of the atomic weights of the other two.
5. Not all the elements fit into the triad system.
6. Such triads—including chlorine-bromine-iodine, calcium-
   strontium-barium, and sulfur-selenium-tellurium
Mendeleev’s Periodic Table

1. In 1869 Mendeleev and then Meyer published nearly identical charts
of elements.
2. Mendeleev used cards to represent elements and began moving the
cards so that the elements were grouped according to their increasing
atomic mass.
3. Mendeleev noticed periodic trends in the elements present and he
predicted properties of missing elements.
4. As predicted elements were isolated it convinced scientists that
Mendeleev’s periodic table was a powerful tool.
The Periodic Law
1. Mendeleev experienced some problems with the organization od
    his chart based on his methodology – arrangement by mass.
2. In 1913 Henry Moseley determined the atomic number of each
    element based on the protons in the atom.
3. The modern periodic table is arranged in order of increasing
    atomic number.
4. There are seven rows or periods on the periodic table.
5. Each period corresponds to a principle energy level.
6. There are more elements in higher number periods because there
    are more orbitals in higher energy levels.
7. Elements are also arranged in groups or columns of elements with
    similar properties.
8. The properties of elements within a period change as you move
    across the periodic table.
9. The pattern of properties that develops in the elements as they
    cross the chart repeats itself for each energy level.
10.This periodic repetition of properties is know as periodic law.
11. Periodic law states; “When elements are arranged by increasing
    atomic number, there is a periodic repetition of their physical and
    chemical properties.
Metals, Nonmetals, and Metalloids
1. Each group in the periodic table has three labels
   a. In the U.S. Numbers and letters signify the column – 1A or 2B
   b. In Europe, Roman numerals and letters are used – VB or IIA
    c. The International Union of Pure and Applied Chemistry number
   the groups 1 to 18 from left to right as you cross the chart.
2. Elements are can be grouped into three broad classes based on their
   general properties.
   a. Metals
    b. Nonmetals
   c. Metalloids
3. Metals
   a. 80 % of the elements found mostly on the left half of the chart
   b. Good conductors of heat and electricity
   c. Have a shiny luster or sheen
   d. Good ability to reflect light
   e. Solid at room temperature except Hg
   f. Ductile – can be drawn into wire
   g. Malleable – can be hammered into sheets
4. Nonmetals
    a. Found in the upper right portion of the periodic table.
    b. A greater variation in physical properties
    c. Most are gasses are room temperature.
    d. Some are solid like Sulfur and Phosphorus
    e. Bromine is a dark red liquid
    f. Poor conductors except for Carbon
    g. Solids tend to be brittle
    h. Vary in color
     i. Not malleable or ductile
5. Metalloids
   a. Form a stairway from Boron to Astatine
   b. Properties reflect those of metals and nonmetals
   c. The properties of metalloids can be controlled giving them
       many technological applications

Homework – Questions p 160, Handout 6.1, and Journal
Questions
6.2 Classifying the Elements

Key Concepts

1. What type of information can be displayed in a
   periodic table?
2. How can elements be classified based on their electron
   configurations?
Squares in the Periodic Table

1. The periodic table displays the symbols and names of the
   elements as well as the atomic number, atomic mass, electron
   configuration, crystalline structure and group based on its
   background color.
2. Groups that are generally color coded include;
   a. 1A Alkali Metals
   b. 2A Alkali Earth Metals
   c. 7A Halogens
   d. 1B to 8B the Transition Metals
   e. 8A Noble Gases
Electron Configurations in Groups

1. Elements can be sorted into noble gases, representative
   elements, transition elements, or inner transition metals based
   on their electron configuration
2. Noble Gases
   a. Are Group A8 elements
   b. Nonmetals
   c. Inert
   d. The s and p sublevels are completely filled except He which is
   full with an s sublevel.
3. The representative Elements
   a. Is the portion of the table from 1A to 7A.
   b. They are referred to representative elements because they
   display a wide range of physical and chemical properties.
   c. They include metals, nonmetals, and metalloids
   d. Most are solids, some gases, and a liquid
   e. The s and p sublevels of the highest occupied energy levels are
   not filled
   f. For any representative element, its group number equals the
   number of electrons in the highest occupied level.
Transition Elements

1. Elements of the B Group separate the A groups
2. Elements in the B groups provide a connection between the two
   sets of representative elements.
3. There are two type of transition elements based on their
   electron configuration.
   a. Transition Metals
   b. Inner Transition Metals
4. Transition Metals
   a. Are group B elements displayed in the main body of the
   periodic table
   b. Examples – Cu, Au, Ag, Fe
   c. The highest occupied s sublevel and a nearby d sublevel contain
   electrons.
   d. There are some exceptions in writing electron configurations
   found in 6B and 1B
5. Inner Transition Metals
   a. Appear below the main body of the periodic table
   b. the highest occupied s sublevel and the nearby f sublevel
   generally contain electrons.
   c. Are characterized by f orbitlas
Blocks of Elements

1. If electron configuration and the position of the element are
   considered simultaneously, another pattern emerges.
2. The periodic table can be divided into sections or blocks that
   correspond to the highest occupied sublevel.
3. S block – 1A and 2A
4. P block – 3A, 4A, 5A, 6A, 7A, 8A
5. D block – Transition elements
6. F block – Inner transition elements
7. D sublevels have a principal energy level that is one less than the
   period number.
8. F sublevels have a principal energy level that is two less than the
   period number.

Homework – Question #10 to 15 p. 167, Review sheet 6.2, and
  Journal questions
6.3 PERIODIC TRENDS

Key Concepts
1. What are the trends among the elements for atomic
   size?
2. How do ions form?
3. What are the trends among the elements for the
   first ionization energy, ionic size, and
   electronegativity?
4. What is the underlying cause of periodic trends?
Trends in Atomic Size

1. Molecules –a neutral group of atoms joined together by covalent
   bonds
   a. examples – H2 , O2, N2, F2, Cl2, Br2, I2
2. Atoms in each molecule are identical
3. Distances between the nuclei of molecules can be used to estimate
   the size of the atoms
4. Atomic Radius – Is one half the distance between the nuclei of two
   atoms of the same element when the atoms are joined.
5. Atomic radius is measured in picometers
6. In general, atomic size increases from top to bottom within a
   group and decreases from left to right across a period.
Group trends in atomic size:

1. The atomic radius within a group increases as the atomic number
   increases.
2. As the atomic number increases in a group so does the charge.
3. As the atomic number increases in a group so does the number
   of occupied energy levels
4. How atomic size can be affected;
   a. An increase in positive charge draws electrons closer to the
   nucleus.
   b. An increase in occupied orbitals shields the outer electrons
   from being drawn to the nucleus and creates a larger atom
Periodic trend in atomic size:

1. Atomic size decreases from left to right
   a. Each element has one more proton and one more electron
   b. Electrons are added to the same principal energy level
   c. The shielding effect is constant for all the elements in the period
   d. The increasing nuclear charge pulls the electrons in the highest
   energy level closer to the nucleus and decreases atomic size.
Ions:

1. Some compounds are composed of particles called ions.
2. An ion is an atom or group of atoms that has a positive or
   negative charge.
3. A neutral atom has equal numbers of electrons and protons
4. A cation is a positively charged atom with more protons then
   electrons
5. Cations are atoms that tend to lose electrons
6. An anion is a negatively charged atom with more electrons then
   protons
7. Anions are atoms that tend to gain electrons
Trends in Ionization Energy:

1. Ionization energy – is the energy needed to remove an electron
   from an atom.
2. The energy is measured when an element is in its gaseous state.
3. First ionization energy – the energy required to remove the first
   electron from the atom.
4. The cation produced by removing one electron is +1
5. First Ionization energy tends to devrease from top to bottom
   within a group and increases from left to right across a period.
6. Ionization energy can help predict what ions elements will form.
7. Ionization energy increases as more electrons are removed to
   produce cations with +2 or +3 charge.
Group Trends in Ionization energy:

1. In general first ionization energy decreases from top to
   bottom within a group
2. As the size of an atom increases, the nucleus has a smaller
   effect on the higher level electrons, therefore it requires
   less energy to remove the electrons.

Periodic trends in Ionization energy:

1. Ionization energy tends to increase from left to right.
2. Because the nuclear charge increases while the shielding
   effect remains constant, there is an increase in the
   attraction of the nucleus for an electron. Therefore, there
   is an increase in the energy to remove the electrons from
   the atom.
Trends in Ionic Size:

1. In ionic bonds between metals and nonmetals, metals
   tend to lose electrons while nonmetals gain electrons.
2. Cations are always smaller than the atoms from which
   they form.
3. Anions are always larger than the atoms from which they
   form.
4. Periodic trend reveals a decrease in size of cations
   followed by a decrease in size of anions.
5. Group trends reveal that there is a increase in ionic size
   down the group.
Trends in electronegativity:

1. Electronegativity – is the ability of an atom of an element to
   attract electrons when the atom is in a compound.
2. Ionization energy is used to calculate electronegativity.
3. Electronegativity can be used to determine the type of bond that
   will form during a reaction.
4. Electronegativity values can be represented in a table, however the
   noble gases do not appear because they are not very reactive.
5. The values are expressed in units called Paulings.
6. Electronegativity values decrease from top to bottom within a
   group.
7. For representative elements, the values tend to increase from left
   to right across a period.
8. Electronegativity values among the transition elements are not
   regular.
9. The least electronegative element is Cesium at 0.7
10.The most electronegative element is Fluorine at 4.0

Homework – Questions 16 to 22 p. 178, Handout 6.3, Journal
http://www.meta-synthesis.com/webbook/24_complexity/complexity2.html

				
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