The Periodic Table

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The Periodic Table

• chapter 6

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Howd They Come Up With That?

• Our current society takes for granted all of the
  hard work, research, chance,
• and luck that has gone into creating
• and discovering the materials that are used in
  the products we utilize every day.
• For example, who was the first person to set or
  find a random black rock (coal) on fire and
  discover that it provided a good, constant source
  of heat?
• Who was the first person to discover that a
  substance found in some rocks was capable of
  being the ultimate explosive (uranium)?

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• In nature and in the lab we have discovered over
  100 different elements.
• Weve organized the elements into a table based
  on their PHYSICAL and CHEMICAL PROPERTIES
• It took us almost 2000 years to figure out the
  properties of the elements currently in the
  Periodic Table of Elements and arrange them.

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Developing the Periodic Table

• By the early 1800s, enough information was known
  about the elements that scientists wanted an easy
  way to categorize the Earths ingredients.
• Many methods of organization were tried before
  scientists found the most effective way of
  grouping the elements

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Mayan Periodic Table, named for its similarity
to the Mayan calendar.
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Johann Dobereiner 1780 - 1849

• In 1829, he classified some elements into groups
  of three, which he called triads.The elements in
  a triad had similar chemical properties and
  orderly physical properties.
• (ex. Cl, Br, I and Ca, Sr, Ba)
• Model of triads

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John Newlands 1838 - 1898

• In 1863, he suggested that elements be arranged
  in octaves because he noticed (after arranging
  the elements in order of increasing atomic mass)
  that certain properties repeated every 8th
  element.
• Law of Octaves

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Dmitri Mendeleev(1834 1907)

• Russian chemist, Dmitri Mendeleev organized
  elements into a table based on atomic mass and
  similar properties.
• Mendeleev stated that the properties of elements
  are a periodic function of their atomic masses.

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Mendeleevs Periodic Table
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Mendeleevs Prediction

• Mendeleevs table had several missing elements.
  When these elements were discovered, they were
  almost exactly as Mendeleev predicted.
• The following is an example of the element we
  know as Germanium.

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Germanium is located below silicon. Mendeleev
predicted its properties based on this location
in his table.
Ekasilicon (Es) Germanium (Ge)
1. Atomic mass 72 1. Atomic mass 72.61
2. High melting pt. 2. Melting pt 945 C
3. Density 5.5g/cm3 3. Density 5.323g/cm3
4. Dark gray metal 4. Gray metal
5. Will obtain from K2EsF6 5. Obtain from K2GeF6
6. Will form EsO2 6. Forms oxide (GeO2)
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• However, in spite of Mendeleevs great
  achievement, problems arose when new elements
  were discovered and more accurate atomic weights
  determined. By looking at our modern periodic
  table, can you identify what problems might have
  caused chemists a headache?
• 18Ar, 39.95 amu and 19K, 39.10 amu
• 27Co, 58.93 amu and 28Ni, 58.69 amu

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• Modern Periodic Law

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Henry Moseley 1887 - 1915

• In 1913, through his work with X-rays, he
  determined the actual nuclear charge (atomic
  number) of the elements. He rearranged the
  elements in order of increasing atomic number.
• There is in the atom a fundamental quantity
  which increases by regular steps as we pass from
  each element to the next. This quantity can only
  be the charge on the central positive nucleus.

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• Increasing atomic number is the basis for our
  current periodic law.

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His research was halted when the British
government sent him to serve as a foot soldier in
WWI. He was killed in the fighting in Gallipoli
by a snipers bullet, at the age of 28. Because
of this loss, the British government later
restricted its scientists to noncombatant duties
during WWII.
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Glenn T. Seaborg 1912 - 1999

• After co-discovering 10 new elements, in 1944 he
  moved 14 elements out of the main body of the
  periodic table to their current location below
  the Lanthanide series. These became known as the
  Actinide series.

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• He is the only person to have an element named
  after him while still alive.
• 106Sg- Seaborgium
• "This is the greatest honor ever bestowed upon me
  - even better, I think, thanwinning the Nobel
  Prize."

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Periodic Table
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Periodic Table Review

• Rows on the periodic table are called PERIODS
• Columns on the periodic table are called GROUPS
  or FAMILIES

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Periodic Table Review

• There are 7 periods and 18 groups.
• Electron arrangements are repeated in periods.
• Elements with similar e- configurations are
  placed in the same group.
• Elements in groups are also listed in order of
  their increasing principal quantum numbers.

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Electron Configuration

• Sublevel / e- capacity
• s 2
• p 6
• d 10
• f 14

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S block (sublevel)

• Contains elements in Group 1, Group 2, and He
  from Group 18.
• Electrons are added to the s orbitals.
• EX H 1s1
• He 1s2
• Li 1s22s1
• Be 1s22s2

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P block (sublevel)

• Contains elements in Group 13, Group 14, Group
  15, Group 16, Group 17, and the remaining
  elements from Group 18 (except He)
• Electrons are added to the p orbitals.
• Ex B 1s22s22p1
• C 1s22s22p2
• N 1s22s22p3

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D block (sublevel)

• Contains elements from the center of the periodic
  table.
• These elements are called transition metals.
• Electrons are added to the d orbitals of the
  transitions metals as well as La and Ac of the
  inner transition elements (rare earth).

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F block (sublevel)

• Contains elements from the inner transition
  metals (rare earth elements)
• Electrons are added to the f orbitals.
• Ex Ce ? Lu
• Th ? Lr

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Octet Rule

• Atoms with full outer levels are stable (less
  reactive)
• For elements (except He) this stable
  configuration would have eight e-.
• (two in the outer s sublevels and six in the
  outer p sublevels)
• These outer eight e- (valence electrons) are
  called an octet.

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Octet Rule

• Eight electrons in an outer level render an atom
  unreactive.
• This is referred to as the Octet Rule.
• When atoms react with one another, they do so to
  obtain a stable config.
• Some atoms gain or lose e- (ions) and some share
  e- (molecules).

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Organizing Information on the Periodic Table

• Use a pen to label the following
• Group 1 Alkali metals
• Group 2 Alkaline earth metals
• Group 16 Chalcogens
• Group 17 Halogens
• Group 18 Noble gases
• Sc Uub Transition metals
• La Lu Lanthanoids
• Ac Lr Actinoids

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Organizing Information on the Periodic Table

• Draw a stair step dark line starting between B
  and Al.
• Label the right side metals
• Label the left side nonmetals
• Write METALLOID along stair step line.
• Label the valence e- (outer electrons).
• Use colored pencils to shade each group or
  category a different color.

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Basic Properties of Metals, Nonmetals, and
Metalloids

• Metals
• 1. Dense and shiny (luster).
• 2. Conduct heat and electricity well.
• 3. Have high melting/boiling points (high
  densities).
• 4.Malleable and ductile.

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• Nonmetals
• 1. Generally gases or brittle solids.
• 2. If solid, dull surface.
• 3. Good insulators.
• 4. Have low melting/boiling points (low
  densities)

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• Metalloids
• 1. Properties of both metals and nonmetals.
• 2. Some semiconductors.
• EX Silicon, for example, possesses a metallic
  luster, yet it is an inefficient conductor
  (semiconductor) and is brittle.

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Properties of Alkali Metals

• Group 1 metals
• Soft silver metals.
• Less dense than other metals and lower melting
  points.
• Very reactive due to large size and one loosely
  held valence electron.
• Too reactive to be found free in nature.

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Properties of Alkaline Earth Metals

• Group 2 Metals
• Shiny silvery-white metals
• Have 2 valence electrons
• Not as reactive as alkali metals but very
  reactive
• All found in the Earths crust in mineral form
• Too reactive to be found in free element form

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Properties of Halogens

• Group 17 nonmetals
• All diatomic gases at room temperature EX F2,
  (Br2 -liquid at room temp)
• Too reactive to be found as free elements in
  nature
• Most important group to be used in industry

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Properties of Chalcogens

• Group 16 nonmetals
• Diverse group that includes nonmetals,
  metalloids, and metals

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Properties of Noble Gases

• Group 18 nonmetals
• Complete octet of valence electrons
• s2p6
• Largely unreactive
• Monotomic gases

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Periodic Trends
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Using the Periodic Table to Predict Properties of
Elements

• The basis of the periodic table is the atomic
  structures of the elements.
• Position on the table and properties of these
  elements arise from the e- configurations of the
  atoms.
• Properties such as density, atomic radius,
  oxidation numbers, ionization energy, and
  electronegativity can be predicted.

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Trends in Oxidation Numbers

• Our knowledge of e- configurations and the
  stability of noble gases allows us to predict
  oxidation numbers for elements.
• Oxidation numbers represent the charge an ion
  obtains after losing or gaining valence electrons.

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1
2 or 4
0
2
Tend to have more than one oxidation number
3
3-
2-
1-
3
3 or 4
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• Two hydrogen atoms are walking down the road. One
  said, I think I lost an electron!.
• Really, the other replied, Are you sure?.
• Yes, Im positive.

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Atomic Radius

• Simply put, this is a measurement of the size of
  an atom
• (its determined by finding ½ the bond distance
  between two atoms of the same element).

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• 1. Group trends
• As we increase the atomic number (or go down a
  group). .
• each atom has another energy level,
• so the atoms get bigger

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• 2 - Period Trends
• Going from left to right across a period, the
  size gets smaller.
• Electrons are in the same energy level.
• But, there is more nuclear charge.
• Outermost electrons are pulled closer which
  reduces the volume of the electron cloud.

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Predicting Atomic Radius

• General rule atomic size increases as you move
  diagonally from top right corner to bottom left
  corner.

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When graphed, atomic radii demonstrates a
periodic trend
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Radii of ions Ions are atoms that have gained or
lost e- from the outer orbitals.

• Cations ()
• Become smaller
• 1. Positive charged nucleus attracting fewer e-
  so pulls electron cloud in tighter.
• 2. Reduced the number of energy levels.

• Sodium atom is much larger than the positive
  sodium ion.

Na1 11p 10e-
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• The pull on each electron is reduced expanding
  the electron cloud.

• Anions (-)
• Become larger
• 1. Positive charged nucleus attracting more e-
  expands electron cloud.
• 2. Add more energy levels.

S-2 16p 18e-
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Ionization Energy

• The energy required to remove an e- from an atom.
• The larger the atom, the less energy is required
  because the e- are farther from the positive
  center.
• As atoms get larger ionization energy decreases
  because of the shielding effect (which says that
  the farther an electron is from the nucleus, the
  less tightly the positive nucleus grabs it).

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• Remove the most loosely held e- is first
  ionization energy.
• Measured in kilojoules per mole
• kJ/mol

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Ionization energy increases diagonally from
bottom left corner to top right corner.
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Classification based on First Ionization Energy

• METAL
• 1. Low 1st ionization energy.
• 2. Located on left side of Periodic Table.
• 3. Form positive ions.

• NONMETAL
• 1. High 1st ionization energy.
• 2. Located on the right side of Periodic Table.
• 3. Form negative ions.

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Multiple Ionization Energies

• Additional e- can be lost from an atom and the
  ionization energies can be measured.

IONIZATION ENERGIES (kilojoules per
mole) Element 1st 2nd 3rd
4th 5th H 1312.0 He 2372.3
5220 Li 520.2 7300 11750 Be 899.5
1760 14850 20900 B 800.6 2420 3660
25020 32660
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Electronegativity

• Electronegativity is the ability of an atom to
  capture an electron.
• The smaller the atom the stronger its ability to
  take electrons from other atoms.
• Electronegativity is a unitless value.
• Fluorine is highest at 3.98
• Francium is the lowest at 0.7

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• It increases from bottom left to top right
  corners.

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Electron Affinity

• e- affinity is a measure of an atoms attraction
  for an e-.
• Metals have low e- affinities.
• Nonmetals have high e- affinities.
• Chemical reactions occur between atoms with high
  e- affinity and those with low e- affinity.
• EX Al Br ? Al2Br3
• (low) (high) (more stable)

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Review
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Review

• Based on our trends
• The most reactive metal element would be
• Francium
• The most reactive nonmetal element would be
• Fluorine

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In Summary

• Periodic table is a chart of elements in which
  the elements are arranged based on their e-
  configurations which dictates their properties.
• Moving down a group in the periodic table, atomic
  radii becomes larger because more energy levels
  are needed for more e-.

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In Summary

• As the size becomes larger, the e- are located
  farther away from the positive center.
• This decreases the affinity of that atom to hold
  on to these outer e-, thus decreasing e-
  affinity.
• Ionization energy is low because it is easy for
  the atom to lose these outer e-.

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In Summary

• Moving across a period in the periodic table,
  atomic radii becomes smaller because the energy
  levels of periods are the same but the positive
  centers of atoms increase. This pulls the e-
  cloud closer to the nucleus, making the atom
  smaller.
• Ionization energy and e- affinity increases for
  these smaller atoms.

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THE END