Download Unit 2 PPT

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The Periodic Table
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Dmitri Mendeleev (1869, Russian)
 Organized elements by increasing atomic mass.
 Elements with similar properties were put in the
same row and/or column.
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Mendeleev
 In mass order, some properties didn’t match up.
 But he predicted properties of undiscovered elements.
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Moseley
 Henry Mosely (1913, British)
 Organized elements by increasing atomic number.
 Properties vary in regular ways up and down and across
the periodic table.
 Resolved discrepancies in Mendeleev’s arrangement.
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Tellurium & Iodine
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Periodic Law
 When the elements are arranged in order of increasing
atomic number, there is a periodic repetition of their
physical and chemical properties.
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Modern Periodic Table
 Family/Group
 Numbered 1-18 from left to right.
 Elements in each column have similar chemical and
physical properties.
 Period
 Numbered 1-7 from top to bottom.
 Closer elements have more similar properties.
 Chemical and physical properties change somewhat
regularly across a row.
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The Modern Periodic Table
 Ninety naturally occurring elements plus an additional
twenty-eight synthetic elements for a total of one
hundred & eighteen elements.
 Most elements are solids at room temperature.
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The Modern Periodic Table
 Lanthanide Series & Actinide Series
 Metallic elements that fit into the table just after
elements 56 & 88
 Placed at the bottom to keep the table from being too
wide.
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Nonmetals
Metalloids
Metals
?
?
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METALS
1. Good conductor of heat and electricity.
2. Ductile (drawn into wire)
3. High tensile strength
(ability to resist breaking when pulled)
4. Malleable (hammered into a thin sheet)
5. Most have a silver or grayish-white luster.
6. Solid at room temperature, except Hg
NONMETALS
 Poor conductor of heat and electricity.
 Often gases, but not all of them.
 Solids tend to be brittle
METALLOIDS
 Have some characteristics of both metals and
nonmetals.
 All are solids at room temperature
 Tend to be semiconductors of electricity
 Do not unconditionally conduct electricity like
metals but do not impede electric flow like nonmetals
Periodic Groups
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Groups of the Periodic Table
Alkali Metals – Group 1
 Lithium, Sodium, Potassium, Rubidium, Cesium, and
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Francium
As pure elements: Soft, silver, metals. So soft that the
metal can easily be cut with a dull knife.
But… they are SO reactive that they are never found in
nature as pure elements.
React vigorously with most non-metals
React strongly with water (even as vapor) to produce
hydrogen
Usually stored in kerosene to keep them from reacting
Melt at low temperatures.
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Alkaline Earth Metals – Group 2
 Beryllium, Magnesium, Calcium, Strontium, Barium,
and Radium
 Harder, denser, stronger than the Alkali Metals and
have higher melting points
 Less reactive than the Alkali Metals – not found in
nature as free elements but when purified in the lab
they can be kept pure in a jar.
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Transition Metals – Group 3-12
 Mercury, Gold, Silver, Copper, Zinc, Nickel, etc.
 1 or 2 valance electrons
 Less reactive then Alkali or Alkaline Earth metals
 Good conductors of electricity
 High luster
 Posses metallic properties
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Halogens – Group 17
 Fluorine, Chlorine, Bromine, Iodine, and Astatine
 Are the most reactive non-metals. React vigorously
with metals.
 Do not exist in nature as single atoms of the pure
element. F2, Cl2, Br2, and I2 exist because they’ve
formed molecules.
 F2 and Cl2 are gases at room temperature, Br2 is a
liquid, and I2 is a dark purple solid. Astatine is a
synthetic element.
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Noble Gases – Group 18
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Helium, Neon, Argon, Krypton, Xenon, and Radon
All are gases at room temperature
He is less dense than air (It “floats” in air)
Ne, Ar, Kr, and Xe are used in lighting
Rn is radioactive
All have 8 electrons in the valence shell without
forming compounds, so …
 Considered to be un-reactive (or “inert”) so natural
compounds of these elements do not exist
 In 1962 Xenon tetrafluoride was produced at extremely
high temperatures and pressures.
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Periodic Trends
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Valence Electrons
 The number of valence electrons in an atom is equal to
the number of electrons in the highest numbered
energy level.
 In any period, the number increases moving from
group 1 to 2, and then from group 13 to 18.
 In any group, all atoms have the same number of
valence electrons.
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Valence Electrons
1ve-
8ve2ve-
3ve- 4ve- 5ve- 6ve- 7ve-
1ve- or 2ve-
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Trends in Atomic Radius
 The atomic radius is ½ the distance between the
centers of the nuclei of two identical atoms bonded
together.
 SIZE OF THE ATOM
 Done this way because the “edge” of an atom is
somewhat “fuzzy”
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Trends in Atomic Radius
 PERIOD TREND
 Decreases going across left to right
 Why? More positively charged nucleus (more protons)
pulls electrons in tighter.
 GROUP TREND
 Increases going down
 Why? Increasing number of energy levels (more
electrons)
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Fig. 14.10, p401
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Ionic Radii
 Positive ions (cations) are SMALLER than the
corresponding neutral atom
 The removal of the outer layer of electrons makes it
smaller
 There are the same number of protons, but fewer
electrons. Each electron is pulled closer by the
unbalanced charge
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Ionic Radii
 Negative ions (anions) are LARGER than the
corresponding neutral atom
 There are more electrons, but no more protons. Each
electron is not drawn to the nucleus as strongly,
electrons in the cloud repel one another
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Trends in Ionization Energy
 The amount of energy needed to remove an electron from a
neutral atom is the ionization energy.
 It may be easy to remove the first electron, but it takes more
energy to remove each one after that.
 Ionization energy generally decreases as you move down a
group because the electrons that you are removing are
further away from the positive nucleus – the “pull” isn’t as
great.
 Generally, Ionization energy increases as you move across a
period because the valence level is getting more full.
Increasingly, it would rather gain electrons.
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Trends in Electron Affinity
(NOT in your book – but it is on the PA Keystone Standards)
 Electrons can also be gained. The energy change that occurs
when a neutral atom gains an electron is the electron affinity.
 Trends are opposite the trends of Ionization Energy.
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Decrease across a period / Increase down a group
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Trends in Electronegativity
 Valence electrons hold the atoms together in a compound,
but when this happens it causes the electrons to gather closer
to one atom than to the other.
 The electronegativity is the ability of an atom to attract the
electrons.
 The most electronegative elements are in the upper right if
you ignore the noble gases
 The least electronegative elements are in the lower left, the
values only decrease slightly going down a group.
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Trends in Chemical Reactivity
(Reactivity = ability to form compounds via reactions)
 Noble Gases do NOT naturally react with anything
 Elements in Groups 1 and 17 (Alkali metals and
Halogens) are HIGHLY reactive. They do not exist in
nature as single atoms.
 Elements in Groups 2 and 16 are reactive, but not as
reactive as groups 1 + 17. They can exist as pure
elements but only after we artificially purify them.
 Transition metals can exist as pure elements (but often
exist as compounds
 Metals are more reactive at the bottom of a group
 Non-metals are more reactive at the top of the group
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