The Periodic Table

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The Periodic Table
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Concepts to Master

• Who was important in the development of the
  periodic table and why?
• What is the difference between a chemical and
  physical property?
• What are elements arranged on the periodic table
  horizontally?
• What are elements arranged on the periodic table
  vertically?
• How many electrons can each orbital contain?
• Why can the periodic table be subdivided into
  s,p,d, and f blocks?
• What are the other names for the s, p, d, and f
  blocks?
• Is atomic radius a chemical or physical property?
• Is electronegativity a chemical or physical
  property?
• Why are the noble gases inert?
• Why do the alkali metals react so readily with
  the halogens?
• Why do the transition metals have multiple
  oxidation numbers?
• What is the trend in metallic character as you go
  down the periodic table?
• What is the common theme as you go down the
  periodic table?
• What is the diagonal rule?

• Where are electrons located?
• What are the characteristics of metals?
• What are the characteristics of nonmetals?
• What are the characteristics of metalloids?
• Explain periodicity?
• How are sublevels and PELs related?
• How are orbitals and orbits related?
• Whats the formula for the maximum number of
  electrons allowed in a certain PEL?
• List the elements that exist as diatomics.
• What are the trends in atomic radius as you go
  down and across the periodic table? Why?
• What are the trends in electronegativity as you
  go down and across the periodic table? Why?
• What are the trends in ionization energy as you
  go down and across the periodic table? Why?
• What is the most reactive metal and the most
  reactive non-metal?
• What is the common theme as you go across the
  periodic table?

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Vocab

• Melting point
• Metalloids
• Metals
• non metals
• Modern Periodic Law
• Monatomic
• Orbitals
• Orbits
• Periods
• Periodic Law
• Polyatomic
• Principle energy level
• Principle quantum number
• Reactivity
• Reactivity
• Sublevels
• Transition Metals
• Valence electrons
• Valence shell

• Alkali Metals
• Alkaline Earth Metals
• Anion
• Atom
• Atomic Radius
• Boiling point
• Cation
• Density
• Diatomic
• Diatomic
• Ductile
• Electronegativity
• Groups
• Halogens
• Inert Gases
• Ionization Energy
• Kernel
• Lustrous
• Malleable

Labs Graphing Trends and Constructing a Table
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Cool Websites

• http//www.colorado.edu/physics/2000/applets/a2.ht
  ml
• http//www.youtube.com/watch?vDYW50F42ss8
• Element song

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Mendeleev facts

• Bunsen didnt invent burner (mid-1800s)
• just improved design and advertised well
• He did build the first spectroscope (instrument
  used to separate visible light film)
• Since unique to each element, some chemist
  thought it made sense to organize the periodic
  table by their spectra
• Six different people independently built their
  own PT
• Mendeleev was one of the 6.
• He became the PT hero due to his bio
• Born in Siberia
• 1 of 14 kids
• Dad died when he was 13
• Mom ran glass factory, but it burned down
• She had identified her son as brilliant so she
  rode with him on horseback 1200 miles to univ of
  moscow
• They rejected him because he wasn't from Moscow
• Rode another 400 miles to his father alma mater
  in St petersburg
• He graduated from their and then traveled to
  germany worked under Bunsen
• They didnt get along so he went back to univ of
  st petersburg to be prof
• More than the other 5 chemists, Mendeleev
  understood that certain traits among elements
  persists whereas others dont atomic weight
• Had deep knowledge of metals due to his extensive
  work in chem labs he knew them by smell,
  texture, taste, how they reacted so he could
  arrange them into columns and rows
• He was smart enough (and humble enough) to
  realize new elements would come to fruition.

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People and Periodic Table

• Dmitri Mendeleev
• Credited with organization of FIRST periodic
  table
• Mendeleevs greatest achievement was recognizing
  the fundamental rule that the chemical elements
  show an approximate repetition in their
  properties.
• Elements were arranged by increasing atomic mass.
• Elements were listed in columns so that those
  with similar properties were side by side.
• He predicted the existence and properties of new
  elements (blank spaces in the first periodic
  table).

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People and Periodic Table

• Henry Moseley
• With the discovery of isotopes of the elements,
  it became apparent that atomic mass was not the
  significant player in the periodic law as
  Mendeleev, had proposed.
• Moseley used X-rays to determine the atomic
  number of the known elements and then arranged
  them according to increasing atomic number.
• Because of Moseley's work, the modern periodic
  table is based on the atomic numbers of the
  elements.

pg186
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Mosely

• Student in Rutherford s lab at the univ of
  Manchester
• Remember Rutherford was playing with radioactive
  sources to do gold-foil exp so he had access to
  xrays
• Found a mathematical relationship between the amt
  of energy in the xray beam (wavelength) and the
  of protons in nucleus.
• This was huge because at the time Rutherfords
  idea of a nucleus wasnt proven (not enough
  scientists had repeated it) so it was too
  tentative for other scientists to accept.
• Moselys work confirmed Rutherfords nucleus
  conclusion .
• His work could be repeated by anyone
• Linked the order of the elements with a physical
  characteristic based on the atoms structure
• Now scientists knew what to look for when
  searching for new elements
• He died in battle during WWI he was 27

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Periodic Law

• Mendeleev - The properties of the elements are a
  periodic function of their atomic masses.
• Moseley - The properties of the elements are a
  periodic function of their atomic numbers.
• Modern Periodic Law states that many of the
  physical and chemical properties of the elements
  tend to recur in a systematic manner with
  increasing atomic number.
• Periods are the horizontal rows in the table.
• Progressing from the lightest to the heaviest
  atoms, certain properties of the elements
  approximate those of precursors at regular
  intervals of 2, 8, 18, and 32 (periodicity).
• Examples
• The 2d element (helium) is similar in its
  chemical behavior to the 10th (neon), as well as
  to the 18th (argon), the 36th (krypton), the 54th
  (xenon), and the 86th (radon).
• The chemical family called the halogens, composed
  of elements 9 (fluorine), 17 (chlorine), 35
  (bromine), 53 (iodine), and 85 (astatine), is an
  extremely reactive family.

Pg 187
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Electron Location

• Kernel of an atom is the nucleus and all the
  electrons but the valence electrons.
• The elements are arranged vertically in columns
  of the periodic table called GROUPS or FAMILIES.
• Group indicates the number of valence
  electrons.
• Because of Periodicity, the elements with the
  same of valence electrons are in the same
  group.
• These electrons influence the chemical and
  physical properties of elements the most.
• Electron Configuration shows the location of all
  the electrons for the atom.

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Physical and Chemical properties

• Because of Periodicity, the elements with the
  same of valence electrons are in the same group
  so they share similar chemical and physical
  properties.
• Chemical properties of matter describe its
  "potential" to undergo some chemical change or
  reaction by virtue of its composition. What
  elements, electrons, and bonding are present to
  give the potential for chemical change. The
  result of the change is the formation of a new
  substance.
• Toxicity
• Flammability
• Reactivity
• Electronegativity
• Ionization Energy
• Physical properties can be observed or measured
  without changing the composition of matter. They
  are used to observe and describe matter.
• Atomic radius
• Density
• Melting Point
• Boiling Point
• Color
• Solubility
• Odor
• Conductivity

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Because of periodicity.
Noble Gases or Inert Gases
Halogens
Alkali Metals
Alkaline Earth Metals
Transition Metals
Lanthanide Series
Actinide Series
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Alkali Metals
1A
NOT

• The alkali metals are silver-colored
• Soft solids (Fr and Cs are liquids)
• The first three are biologically important
• low-density metals
• react readily with halogens
• react readily with water
• one valence electron
• so they want to lose an electron and achieve a
  noble configuration (which is?)
• form 1 cations

Alkali Metals
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Alkaline Earth Metals
1A

• silvery colored
• Soft solids
• Ca and Mg have biological functions
• low-density metals
• react readily with halogens
• react readily with water - though not as rapidly
  as the alkali metals
• Beryllium is an exception It does not react with
  water
• two valence electron
• so they want to lose 2 electrons and achieve a
  noble configuration (an octet)
• form 2 cations

2A
Alkali Metals
Alkaline Earth Metals
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Transition Metals

• They often form colored compounds.
• They are often good catalysts
• lowers activation energy so rxns are faster
• not used up in the rxn
• enzymes
• They are silvery-blue at room temperature (except
  copper and gold) - lustrous.
• Malleable
• They are solids at room temperature (except Hg)
• Partly filled d sublevel
• They can have a variety of different charged
  cations
• 4s fills before 3d (clouds are more apparent and
  overlapping occurs)
• chromium
• Iron
• Vanadium
• Silver (5s fills before 4f)
• Good conductors of electricity (Why?)

B group
Transition Metals
Thus they are transitioning between the filling
of their outermost orbitals.
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Transition Metal Colors
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Lanthanide

• Rare Earth Metals
• Silvery-white metals that tarnish when exposed to
  air
• Relatively soft metals
• Very reactive
• Many rare earth compounds fluoresce strongly
  under ultraviolet light

Lanthanide Series
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Actinide

• All are radioactive
• The metals tarnish readily in air
• Actinides are very dense metals
• Actinides combine directly with most nonmetals

Actinide Series
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Metallic Characteristics

• Conduct electricity and heat
• Dense
• Malleable (bendable to form shapes jewelry)
• Ductile (able to be drawn out into wires)
• Lustrous (shiny)
• Reactive
• High Melting point
• Solid at RT

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Metalloids

• Share properties with both metals and nonmetals
• Solids
• Semi-conductors (between a conductor and an
  insulator)
• Form cations or anions. of valence electrons
  varies.
• 4 sit on steps and 2 are beneath.

8A
Less metallic metalloid
7A
5A
3A
6A
4A
Poor Metals
More metallic metalloid
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NonMetals

• poor conductors of heat and electricity
• in solid form, they are dull and brittle
• usually have lower densities than metals
• most of the crust, atmosphere and oceans are made
  up of nonmetals.
• Bulk tissues of living organisms are composed
  almost entirely of nonmetals.
• Many nonmetals (hydrogen, nitrogen, oxygen,
  fluorine, chlorine, bromine, and iodine) are
  diatomic, and most of the rest are polyatomic
• Prefer to form anions - gaining electrons to
  achieve an octet.

8A
4A
6A
7A
5A
Halogens
Noble gases
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Halogens

• This high reactivity is due to their atoms being
  one electron short of a full outer shell.
• Halogens are highly reactive
• harmful or lethal to biological organisms in
  sufficient quantities.
• Fluorine is the most reactive element in
  existence
• This high reactivity is due to their atoms being
  one electron short of a full outer shell of
  electrons. They form -1 anions.
• Both chlorine and bromine are used as
  disinfectants
• They form diatomic molecules (F2, Cl2, Br2, I2)
• All three states of matter are represented
• fluorine and chlorine are gases
• bromine is a liquid
• iodine and astatine are solids

8A
7A
Noble Gases
Halogens
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Diatomics

• hydrogen, nitrogen, oxygen, fluorine, chlorine,
  bromine, and iodine
• Diatomic elements are nonmetal elements that form
  a covalent bond between two atoms.
• As elements they always travel in pairs of atoms
  and therefore you must write then as
• H2, N2, O2, F2, Cl2, Br2, I2

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Noble Gases or Inert Gases

• odorless, colorless, monatomic gases.
• Lighting (Ne), welding and space technology
  (processes performed under Ar so that no unwanted
  chem rxns occur) .
• Stable or unreactive
• They have the maximum number of valence electrons
  their outer shell can hold.

8A
Noble Inert Gases
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PEL / PQN

• Probable Location according to the wave
  mechanical model
• Principle Energy Levels (PEL) or Principle
  Quantum Number (PQN) n
• the total number of orbits around the nucleus
• Period n
• Max number of e in that PEL, PQN 2n2
• When n 1, max of e 2
• When n 2, max of e 8
• When n 3, max of e 18
• When n 4, max of e
• When n 5, max of e
• When n 6, max of e
• When n 7, max of e

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Sublevels

• Sublevels exist in each orbit (PEL)
• Shape of the electron cloud that is created by
  fast moving electrons.
• s p d f g
• The number of sublevels present in each PEL also
  n, so PEL 5 contains 5 sublevels. They are 5s,
  5p, 5d, 5f, and 5g

Element (Neutral) Galium Galium Galium Galium Galium Galium Galium Galium Galium Galium
E shown on the periodic table 2 8 8 18 18 18 3 3 3 3
PQN (4th period) 1 2 2 3 3 3 4 4 4 4
Sublevels 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f
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Orbitals

• Each sublevel has orbitals
• The possible orientations of the shapes around
  the x,y, and z axis
• s sublevel has 1 orbital
• p sublevel has 3 orbitals
• d sublevel has 5 orbitals
• f sublevel has 7 orbitals

Element (Neutral) Galium Galium Galium Galium Galium Galium Galium Galium Galium Galium
E shown on the periodic table 2 8 8 18 18 18 3 3 3 3
PQN 1 2 2 3 3 3 4 4 4 4
Sublevels 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f
Orbitals 1 1 3 1 3 5 1 3 5 7
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s sublevel orbitals
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p sublevel orbitals
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d sublevel orbitals
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Electrons in Orbitals

• Each orbital can contain 2 electrons maximum
• 2n2

Element (Neutral) Galium Galium Galium Galium Galium Galium Galium Galium Galium Galium
E shown on the periodic table 2 8 8 18 18 18 3 3 3 3
PQN 1 2 2 3 3 3 4 4 4 4
Sublevels 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f
Orbitals 1 1 3 1 3 5 1 3 5 7
of E possible 2 2 6 2 6 10 2 6 10 14
Actual of E 2 2 6 2 6 10 2 1    
E configuration 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p1    
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Determining spdf configuration using THE Table
The number of columns present in the block
equals the number of possible electrons for
that sublevel.
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f Block location if inserted
35
Really know the location of your valence electrons
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Using PT to determine detailed electron
configuration
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A short cut?

• Diagonal rule - A guideline explaining the order
  in which electrons fill the orbital levels.
• Pros
• Easy to use if given to you
• Cons
• Memorization required
• There are exceptions to this rule when filling
  the orbitals of heavier elements.

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Using Diagonal Rule compared to using PT for
Determining detailed electron configuration
Use PT for determination of Pd
Use PT for determination of Cr
Use diagonal rule for determination of Pd
Use diagonal rule for determination of Cr
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Summary

• Elements are arranged by increasing atomic
  across the periodic table.
• Periods are the horizontal rows. The period
  PEL.
• Elements are grouped vertically by similar
  chemical and physical properties.
• Groups (or Families) are the vertical columns.
  The group of valence electrons.

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Textbook pg 186-191

• Practice problems 5.1-5.3

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Atomic Radius

• Half the distance between the nuclei of two like
  atoms in a diatomic molecule.
• Atom size vs ion size

Combine to form O2
Pg 187-188
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As you go down a group, Atomic Radius increases
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As you go across a period, Atomic Radius
decreases
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Why?

• As you go down a group the number of PELs
  increases, more electrons are present to fill
  these energy levels, so atomic radius increases.
• As you go across a period, atomic increases
  which means that the of protons in the nucleus
  increases, so nuclear charge is increasing and
  attracting electrons with a greater force.
  Opposite attract.

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Ionization Energy

• The amount of energy required to remove an outer
  electron
• The more difficult it is to remove an electron,
  the greater the ionization energy
• smaller atoms have greater ionization energy
  since the valence electrons are closer to the
  nucleus and more strongly attracted and,
  therefore, more difficult to remove
• X energy ? X e-
• First Ionization Energy
• X energy ? X e-
• Second Ionization Energy (greater than 1st)
• X energy ? X2 e-
• Third Ionization Energy (greater than 2nd)
• X2 energy ? X3 e-

Pg 189
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As you go down a group, Ionization Energy
decreases
If the ionization energy is high, that means it
takes a lot of energy to remove the outermost
electron. If the ionization energy is low, that
means it takes only a small amount of energy to
remove the outermost electron
Top trend is group 18 Bottom trend is group 2
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As you go across a period, Ionization energy
increases
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Why?

• As you go down a group the number of PELs
  increases. Attraction is less for the electrons
  furthest from the nucleus so it takes less energy
  for electrons to be pulled away.
• As you go across a period, nuclear charge is
  increasing and attracting electrons with a
  greater force. Since that force is increasing, it
  takes more energy for the electrons to be pulled
  away. Going towards inert gases which have a full
  valence shell and are extremely resistant to give
  up any electrons.

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Electronegativity

• An atoms affinity for electrons
• Arbitrary scale from 0-4
• 0 is least electronegative
• 4 is most electronegative

Fluorine is the most electronegative 4
Francium is the least electronegative 0.7
Neon and the other noble gases have
an Electronegativity of 0
Pg 189-191
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As you go down a group, Electronegativity
decreases
Electronegativity is a measure of the tendency of
an atom to attract electrons. The arbitrary scale
of 0-4 is the most commonly used. Fluorine (the
most electronegative element) is assigned a value
of 4.0.
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As you go across a period, Electronegativity
increases
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Why?

• As you go down a group the number of PELs
  increases. Electron attraction to the nucleus is
  less when they are farther from the nucleus.
• As you go across a period, nuclear charge is
  increasing and thus attracting electrons to a
  greater extent.
• Why do the inert gases have an electronegativity
  of 0?

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Chemical Reactivity

• Francium is the most reactive metal.
• Fluorine is the most reactive non-metal.

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Elements on opposite Sides of the Periodic Table
are attracted to each other.

• Sodium likes to combine with Chlorine Why?
• Atoms become cations due to less ionization
  energy.
• METALS
• Fr
• Atoms become anions due to high
  electronegativity.
• NONMETALS
• F