Unit 5: Periodicity

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Unit 5 Periodicity
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History of the Periodic Table

• Dimitri Mendeleev (1836-1907)
• Russian chemistry professor
• Noticed that as he was arranging elements
  according to their properties that they generally
  were organized in order of increasing atomic
  mass.

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Genius of Mendeleevs Work

• Left spaces for elements not yet discovered.
• He predicted that some still-unknown elements
  must exist to fit in the holes.

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Mendeleevs Periodic Table

• Contained 1 inconsistency.
• He placed the elements in order of atomic mass
• Forced to break pattern a couple times to
  preserve the patterns he had discovered.

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Henry Moseley

• Once the proton was discovered by Rutherford,
  Moseley worked with X-rays to determine how many
  protons elements had.
• He arranged the elements in order of increasing
  atomic number.
• The eliminated the inconsistencies seen in
  Mendeleevs table.
• He was killed at the age of 28 in WWI.

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Periodic Law

• Periodic Law The properties of the elements are
  periodic functions of their atomic numbers.
• What this means is that if we arrange the
  elements in order of increasing atomic number, we
  will periodically encounter elements that have
  similar chemical and physical properties.
• These elements appear in the same vertical column
  (group).

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Elements

• Science has come along way since Aristotles
  theory of Air, Water, Fire, and Earth.
• Scientists have identified 92 naturally occurring
  elements, and created about 28 others.

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Periodic Table

• We can learn a lot about an element based on
  where it is located.
• You can predict the physical and chemical
  properties and how an element will react.

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LABEL THESE ON YOUR PERIODIC TABLE
Nonmetals
Metalloids
Metals
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Properties of Metals

• Good conductors of heat and electricity.
• Shiny
• Ductile (can be stretched into thin wires).
• Malleable (can be pounded into thin sheets).
• Reacts with water which results in corrosion.

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Properties of Non-Metals
Sulfur

• poor conductors of heat and electricity.
• not ductile or malleable.
• Solid non-metals are brittle and break easily.
• dull
• Many non-metals are gases.

Bromine
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Properties of Metalloids

• Metalloids (metal-like) have properties of both
  metals and non-metals.
• They are solids that can be shiny or dull.
• They conduct heat and electricity better than
  non-metals but not as well as metals.
• They are ductile and malleable.

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• Columns of elements are called groups or
  families.
• Elements in each family have similar but not
  identical properties.
• For example, lithium (Li), sodium (Na), potassium
  (K), and other members of family 1 are all soft,
  white, shiny metals.
• All elements in a family have the same number of
  valence electrons.

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• Each horizontal row of elements is called a
  period.
• The elements in a period are not alike in
  properties.
• In fact, the properties change greatly across
  even given row.
• The first element in a period is always an
  extremely active solid. The last element in a
  period, is always an inactive gas.

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Hydrogen

• The hydrogen square sits atop Family 1, but it is
  not a member of that family. Hydrogen is in a
  class of its own.
• Its a gas at room temperature.
• It has one proton and one electron in its one and
  only energy level.
• Hydrogen only needs 2 electrons to fill up its
  valence shell.

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Alkali Metals

• first column of the periodic table
• 1 valence electron
• They are shiny, have the consistency of clay, and
  are easily cut with a knife

That is Sodium!
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Alkali Metals

• most reactive metals
• react violently with water
• never found as free elements in nature. They are
  always bonded with another element.

That is sodium on top of the water.
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Alkaline Earth Metals

• never found alone in nature, always bonded with
  something else.
• two valence electrons

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Transition Metals

• elements in the d-block (Groups 3-12)
• These are the metals you are probably most
  familiar

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Transition Elements

• Transition elements have properties similar to
  one another and to other metals, but their
  properties do not fit in with those of any other
  family.
• Many transition metals combine chemically with
  oxygen to form compounds called oxides.

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Boron Family

• The Boron Family is named after the first element
  in the family.
• 3 valence electrons
• This family includes a metalloid (boron), and the
  rest are metals.
• This family includes the most abundant metal in
  the earths crust (aluminum).

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Carbon Family

• 4 valence electrons
• This family includes a non-metal (carbon),
  metalloids, and metals.
• The element carbon is called the basis of life.
  There is an entire branch of chemistry devoted to
  carbon compounds called organic chemistry.

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Nitrogen Family

• The nitrogen family is named after the element
  that makes up 78 of our atmosphere.
• This family includes non-metals, metalloids, and
  metals.
• 5 valence electrons. They tend to share electrons
  when they bond
• Other elements in this family are phosphorus,
  arsenic, antimony, and bismuth.

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Oxygen Family

• 6 valence electrons
• Most elements in this family share electrons when
  forming compounds.
• Oxygen is the most abundant element in the
  earths crust. It is extremely active and
  combines with almost all elements.

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Halogen Family

• 7 valence electrons, which explains why they are
  the most active non-metals. They are never found
  free in nature.
• Halogen atoms only need to gain 1 electron to
  fill their outermost energy level.
• They react with alkali metals to form salts.

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Noble Gases

• Noble Gases are colorless gases that are
  extremely un-reactive.
• One important property of the noble gases is
  their inactivity. They are inactive because their
  outermost energy level is full.
• Because they do not readily combine with other
  elements to form compounds, the noble gases are
  called inert.
• All the noble gases are found in small amounts in
  the earth's atmosphere.

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Rare Earth Elements

• The thirty rare earth elements are composed of
  the lanthanide and actinide series.
• One element of the lanthanide series and most of
  the elements in the actinide series are called
  trans-uranium, which means synthetic or man-made.

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Octet Rule

• Octet means to have 8 valence electrons.
• Is associated with the stability of the noble
  gases.
• Helium (He) is stable with 2 valence electrons
• Valence Electrons
• He 1s2 2
• Ne 1s2 2s2 2p6 8
• Ar 1s2 2s2 2p6 3s2 3p6 8
• Kr 1s2 2s2 2p63s2 3p6 4s2 3d10 4p6 8

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Electron Shielding
Shielding electrons are those electrons in the
energy levels between the nucleus and the valence
electrons. They are called "shielding" electrons
because they "shield" the valence electrons from
being pulled closer to the nucleus. Remember,
the nucleus has a positive charge!
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Periodic Trends

• Trends in properties of the elements that follow
  a pattern down a group and across a period in the
  periodic table.

down a group
across a period
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Trends in Atomic Radius (size)

• atomic radius distance from center of nucleus
  to edge of electron cloud

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Atomic Radius

• group trend increases going down a group.
• As you move down a group, you add energy levels,
  thus increasing the size of the electron cloud,
  so the atoms get larger.

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Atomic Radius

• periodic trend decreases going left to right
  across a period.
• Why? You are not adding energy levels, you are
  adding protons. As you do this, the nucleus gets
  a greater positive charge which pulls the
  electrons closer to it, decreasing the radius.

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Atomic Radius

• Atomic Radius

• Increases to the LEFT and DOWN

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Ionization Energy

• ionization energy the energy required to remove
  a valence electron.

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Trends in Ionization Energy
group trend decreases going down a group This is
due to the shielding effect - an electron in the
outer energy level of a large atom is easier to
remove because it is well-shielded from the pull
of the nucleus by the inner electrons. periodic
trend increases going across a period This is
due to nuclear charge - across a period, nuclear
charge increases, so they hold onto their
electrons tighter. metals have a much greater
tendency to lose electrons than nonmetals do.
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Ionization Energy

• Ionization Energy

• Increases UP and to the RIGHT

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Trends in Electron Affinity

• electron affinity (electron-liking) the energy
  change that accompanies the addition of an
  electron to an atom.
• group trend EA decreases going down a group.
• Why? It is harder to add an electron when it is
  so far away from the nucleus. The nucleus cannot
  grab onto it.
• periodic trend EA increases going across a
  period.
• Why? As the nuclear charge becomes more positive,
  it is able to hold onto electrons better.
• Note that this periodic trend supports the idea
  that nonmetals have a much greater tendency to
  gain electrons than metals do.

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Electron Affinity

• Electron Affinity

• Increases UP and to the RIGHT

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Trends in Electronegativity

• electronegativity the tendency of an atom to
  attract electrons to itself when it is chemically
  bonded with another element.
• Diagram of water molecule
• In H2O, oxygen is more electronegative than
  hydrogen, so it pulls the electrons closer, and
  thus obtains a partially negative charge.

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Trends in Electronegativity
group trend decreases down a group. Larger atoms
have more energy levels, so it is harder for them
to attract electrons to the nucleus (shielding
effect). periodic trend increases across a
period. The nuclear charge is greater with more
protons and can hold onto electrons closer to the
nucleus.
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Electronegativity

• Electronegativity

• Increases UP and to the RIGHT

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Ionic Radius

• Ion charged atom
• When atoms lose electrons, they become positive
  ions
• More p than e-
• When atoms gain electrons, they become negative
  ions.
• -more e- than p

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Positive Ions

• Example
• Sodium has 11 p and 11 e-. To become the sodium
  ion it loses an e-. Now it has 11 p and 10 e-
• Sodium atom Na Sodium ion Na1

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Negative Ions

• Example
• The oxygen atom has 8 p and 8 e-. The oxygen
  ion gains 2 e- to have a total of 8 p and 10e-.

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Ionic Radius

• Positive ions are smaller than the neutral atom.
• This is because you remove valence electrons
  and they no longer have the outer shell.

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Ionic Radius

• Negative ions are larger than the neutral atom.
• They get bigger because adding electrons to the
  outer shell causes them to want to be further
  apart (like charges repel).

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Trends in Ionic Radius

• Ionic radius increases going down and to the left.

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Summary of Periodic Trends

• Ionization Energy, Electron Affinity, and
  Electronegativity

• Increases UP and to the RIGHT

• Atomic Radius

• Increases to the LEFT and DOWN

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