The Periodic Table

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The Periodic Table

• Chapter 6
• Chemistry 112

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6.1 Development of the Modern Periodic Table

• The properties of the elements in the table
  repeat in a periodic way which is why the term
  periodic is used to describe the table.
• Example periods of the moon

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• Antoine Lavoisier was credited with
• Compiling a list of 23 elements including silver,
  gold, carbon and oxygen

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• In 1864, Newlands noticed if he arranged the
  elements by atomic mass, their properties
  repeated every eight element. He named this
  relationship the Law of Octaves, after the
  musical notes that repeat every eight tone. The
  acceptance of this law was hampered because the
  law did not work for all elements and octave
  was criticized because many thought the musical
  analogy was unscientific.

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• Lothar Meyer and Dmitri Mendeleev were the two
  scientists that demonstrated a relationship
  between atomic mass and elemental properties in
  1869. Mendeleev received the most credit
  because he published first.

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• Mendeleev is known as the father of the periodic
  table

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• Mendeleevs table was widely accepted because he
  was able to predict the existence and properties
  of undiscovered elements

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• Henry Mosley is credited with arranging the
  Periodic Table by atomic number, instead of by
  atomic mass.

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• Periodic Law is a periodic repetition of chemical
  and physical properties of the elements when they
  are arranged by increasing atomic number.

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• Group elements in the same column (also known
  as family)
• Periods elements in the same horizontal row

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• A Representative elements
• B transition elements

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• The three main classifications of elements are
  Metals, Non-metals, and metalloids.

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• Some common properties of metals, nonmetals, and
  metalloids are
• Metals hard, shiny, conduct electricity and
  heat, malleable, ductile.
• Nonmetals brittle, dull, poor conductors
• Metalloids properties of metals and nonmetals

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• Metals are to the left of the stairstep (not
  including the metalloids), except Hydrogen.
• Non-metals are to the right.
• All elements touching the line are metalloids
  EXCEPT for Aluminum

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• These definitions will be discussed more
  completely later in your notes
• Atomic radii - Atomic radius measures size of
  atoms and is defined by how closely an atom lies
  to a neighboring atom.
• Ionization energy - Energy required to remove an
  electron from an atom which is used to overcome
  the attraction between the positive nucleus and
  the negative electron.
• Electronegativity - The relative ability of an
  atom to attract electrons in a chemical bond.

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Advancement of Chemistry

• There are three factors that led to the
  advancement of chemistry in the 1800s
• The advent of electricity
• Development of a spectrometer
• The industrial revolution

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Oxidation Numbers

• The positive or negative charge of an ion
• Equal to the number of electrons transferred from
  an atom to form an ion
• Sodium loses one electron 1
• Chlorine gains one electron -1
• Why do they act like this????? Think about how
  many valence electrons all atoms want.

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Oxidation Numbers of Groups

• Alkali metals have 1 valence electron that they
  want to LOSE. They all form a 1 charge.
• Alkaline earth metals have 2 valence electrons
  that they want to LOSE. They all form a 2
  charge.
• Transition metals for ions in more than one way.
  We will skip the d-block for this chapter and
  revisit it next unit.

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• Boron Group has three valence electrons they want
  to LOSE. They all form a 3 charge.
• Carbon Group has four valence electrons. It is
  just as easy to lose four as it is to gain four
  electrons. Carbon group forms a 4 or -4 charge.
  (there are some exceptions that we will talk
  about next unit).

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• Nitrogen group has five valence electrons. They
  want to GAIN three more to make a total of eight.
  They all form a -3 charge.
• Oxygen group has six valence electrons and they
  want to GAIN two more. They all form a -2
  charge.
• Halogens all have seven valence electrons and
  want to GAIN one more. They all form a -1
  charge.
• Nobel gases have a full valence shell already and
  do not form ions.

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Practice Problems What are the oxidation
numbers of the following ions?

• K
• P
• I
• Ar
• Al
• S

• 1
• 3-
• 1-
• 0
• 3
• 2-

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Write the ion with the proper oxidation number or
charge for each of the following

• Mg
• Cs
• Al
• N
• O
• Br
• Kr

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Trends on the Periodic Table

• You will be responsible for five trends.
• 1. Atomic radius
• 2. Metal reactivity
• 3. Non-metal reactivity
• 4. Ionization Energy
• 5. Electronegativity

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Atomic Radius

• All trends on the periodic table relate to atomic
  radius! They all have to do with how easy it is
  to lose or gain electrons.
• Atomic radius measures size of atoms
• Defined by how closely an atom lies to a
  neighboring atom.
• Trend for within a period and within a
  group.graph to find out!!

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Trends for Atomic Radius

• Atomic Radius INCREASES as you go down the P.T.
  because there are increasing number of energy
  levels (think of it as layers of fatgetting
  bigger).
• Atomic Radius DECREASES across the P.T. because
  there is more of a positive pull on the
  electrons. More of an attraction pulls the
  electrons a bit closer. (think about having four
  people pull on a rope versus five. Five people
  are able to pull harder and will therefore bring
  the rope closer)

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Hands on a clock

• A way to remember the trends is to relate them to
  hands on a clock.
• You point an arrow in the direction where the
  trend INCREASES.
• For atomic Radius, you would have a hand pointing
  down and a hand pointing left.
• This looks like 930 on a clock.

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Practice Problems Which atomic radius is
bigger? (Hint, write he symbols how they appear
on the periodic tablewhich ever one the arrow
points to is bigger)

• Na or S
• Ba or Ca
• K or Ra
• Cl or Au
• As or P
• Fr or Cs

• Na
• Ba
• Ra (period matters more than group)
• Au
• As
• Fr

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Ion Radius

• Atoms that lose electrons to become cations,
  always become smaller. (when you lose weight, you
  become smaller!) Since there are less electrons
  that are repelling each other, they can settle in
  a little closer to each other.
• Atoms that gain electrons to become anions,
  always become larger. (when you gain weight, you
  become bigger). More electrons mean more
  repulsion. They have to spread out even further!

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Practice Problems Which is SMALLER! (pay
attention to the questions)

• Ca atom or Ca ion (hintwhat oxidation number
  does Ca form?)
• F atom or F ion
• Fe2 or Fe3

• Ca2 ion is smaller because it loses 2 e-
• F atom is smaller because the ion gains 1e-
• Fe3 is smaller because it loses 3 e- compared to
  Fe2 that loses 2 e-

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Metal Reactivity

• When Metals react, they lose electrons due to the
  fact they have a low amount of valence electrons.
  
• Would it be easier for an electron to be lost if
  it was close to the nucleus or far away?
• Electrons are lost easier when they are farther
  away from the nucleus since there is less of a
  positive pull on them.
• Therefore, the greater the Atomic Radius, the
  more reactive the metal is. (easier to lose
  more reactive)

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Metal Reactivity Trendhttp//video.google.com/vid
eoplay?docid-2134266654801392897qbrainiachlen
Blocked by Cobb County but you can watch this at
home!

• INCREASES DOWN a Group due to the fact that the
  electrons are further from the positive pull of
  the nucleus and they are easier to remove.
• DECREASES ACROSS a period due to the fact that
  there is more of a positive pull holding onto the
  electrons and they are harder to remove.

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Which metal is more reactive?

• Na or Cs
• Ni or Fe
• Au or Cu
• Ag or K

• Cs
• Fe
• Au
• K (group always matters morealkali metals are
  more reactive than transition metals! Good thing
  or we wouldnt be able to wear jewerly!!)

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Non Metal Reactivity

• When Non-metals react, they gain electrons due to
  the fact that they have a high number of valence
  electrons.
• Would it be easier to gain electrons if they were
  closer to the nucleus or farther away?
• Electrons are gained easier if they are closer to
  the nucleus due to the positive attraction.
• The trend is exactly opposite of Metal Reactivity
  and Atomic Radius.
• Noble Gases are excludedthey hardly ever react!

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Non-Metal Reactivity Trend

• DECREASES DOWN a group due to the fact that the
  atomic radius is getting larger and it is harder
  for the positive nucleus to attract electrons.
• INCREASES ACROSS a period due to the fact that
  there is more of a positive pull and is easier to
  attract the electrons.
• What time would this be on a clock?

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Which non-metals is more reactive?

• F or Br
• S or Cl
• N or He

• F
• Cl
• N (remember Nobel gases dont react!)

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Ionization Energy

• Energy required to remove an electron from an
  atom.
• Energy is needed to overcome the attraction
  between the positive charge in the nucleus and
  the negative charge of the electron.
• What type of atoms want to lose electrons?
• Metalstherefore, they have lower I.E.

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• 1st I.E. Energy to remove the 1st electron
• 2nd I.E. Energy to remove the 2nd electron
• 3rd I.E. etc.

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Trends for I.E.

• I.E. DECREASES down a group. It becomes easier
  to remove an electron because there are more
  energy levels blocking the positive pull.
• I.E. INCREASES across a period. There is more of
  a positive pull holding on to the electrons,
  therefore, it is harder to remove them. You are
  also approaching the Non-metals, which gain
  electrons!

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When will you see a drastic increase in I.E.???

• After the element has lost its valence electrons.
• Example Potassiums 1st I.E. is low. Its 2nd
  is very high.
• Calciums 1st and 2nd I.E. is low. Its 3rd is
  high.
• See page 168 Table 6-2

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Practice problems Which atom has a lower 1st
I.E.?

• Ca or Br
• Cl or I
• Au or Cu
• Ag or Rb
• Mg or I

• Ca
• I
• Au
• Rb
• Mg (metals always have a lower I.E. than
  nonmetals.metals want to lose, nonmetals want to
  gain)

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Electronegativity

• The relative ability of its atoms to attract
  electrons in a chemical bond.
• What type of elements want to ATTRACT electrons??
• Nonmetals! Therefore, nonmetals have a high E.N.
  value.
• Numerical value from 0.0 (least) to 4.0 (most)
• Unit is a Pauling
• Noble Gases are excluded!!

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Trends for E.N.

• Down a group, E.N. DECREASES. More energy levels
  makes it harder to attract electrons
• Across a period, E.N. INCREASES. More and more
  valence electrons, and less for the atoms to have
  a full octet. More positive pull also. Also,
  pointing towards the non-metals.

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Which element is more electronegative?

• F or Be
• S or Po
• Sr or Br
• Kr or As

• F (highest E.N. value on P.T.)
• S
• Br (Nonmetals always higher than metals)
• As (Nobel gases are are excluded)

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Using E.N. values to calculate bond types

• Ionic Bond electrostatic force holds oppositely
  charged particles together in an ionic compound.
  Formed by losing or gaining electrons.
• Polar Covalent Bond formed when electrons are
  not shared equally.
• Non-polar Covalent bond formed from the equal
  sharing of valence electrons.

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Like Dissolves Like

• Water is a polar covalent solvent (meaning one
  end of the molecule has a slight charge).
• Only Ionic compounds and Polar compounds will
  dissolve in water because they also have a
  charge.
• Non-polar covalent compounds will not dissolve in
  water because it does not have a charge.

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E.N. differences

• Subtract the lowest E.N. element in a compound
  from the highest E.N. element (the subscripts do
  not matter)
• Difference greater than 1.70 IONIC
• Difference 0.5 to 1.70 Polar
• Difference less than .5 Non-polar

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Use your book pg. 403 to find the E.N. values of
each element. What kind of bond will the
compounds likely have?

• SO2
• BaCl2
• H2S
• Ca2N3
• As2O3

• .94 polar covalent
• 2.27 ionic
• 0.3 non-polar
• 2.04 ionic
• 1.26 polar covalent