The Periodic Table

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The Periodic Table
2
History

• In 1829, the German chemist J.W. Döbereiner
  classified some elements into groups of three,
  which he called triads.
• The elements in a triad had similar chemical
  properties, and their physical properties varied
  in an orderly way according to their atomic
  masses.

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History

• The Russian chemist, Dmitri Mendeleev was
  studying the properties of the elements and
  realized that the chemical and physical
  properties of the elements repeated in an orderly
  way when he organized the elements according to
  increasing atomic mass.

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History

• Mendeleev later developed an improved version of
  his table with the elements arranged in
  horizontal rows.
• This arrangement was the forerunner of todays
  periodic table.

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History

• Patterns of changing properties repeated for the
  elements across the horizontal rows.
• Elements in vertical columns showed similar
  properties.

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History

• Mendeleev grouped elements in columns by similar
  properties in order of increasing atomic mass.
• He found some inconsistencies and felt that the
  properties were more important than the mass, so
  he switched order.

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History

• Mendeleev left some gaps in his periodic table,
  deciding there must be undiscovered elements.
• He predicted their properties before they were
  found.

• Mendeleev is considered to be the Father of the
  Periodic table.

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Periodicity

• This repeated pattern (when Mendeleev grouped
  elements in columns by similar properties) is an
  example of periodicity in the properties of
  elements.
• Periodicity is the tendency to recur at regular
  intervals.

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History

• By 1860, scientists had already discovered 60
  elements and determined their atomic masses.

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The Modern Table

• Fifty years after Mendeleev, the British
  scientist Henry Moseley discovered that the
  number of protons in the nucleus of a particular
  type of atom was always the same.

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The Modern Table

• When atoms were arranged according to increasing
  atomic number, the few problems with Mendeleev's
  periodic table disappeared.
• Because of Moseley's work, the modern periodic
  table is based on the atomic numbers of the
  elements.

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The Modern Table

• The statement that the physical and chemical
  properties of the elements repeat in a regular
  pattern when they are arranged in order of
  increasing atomic number is known as the periodic
  law.

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The Modern Table

• On the periodic table a period, sometimes also
  called a series, consists of the elements in a
  horizontal row.

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• There are 7 periods.

1
2
3
4
5
6
7
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The Modern Table

• A group, sometimes also called a family, consists
  of the elements in a vertical column.

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• Elements are placed in columns by similar
  properties.

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• The elements in the A groups are called the
  representative elements.

8A0
1A
2A
3A
4A
5A
6A
7A
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• The B groups are called the transition elements.

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• Group 1A elements are the alkali metals.

1A
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Group 1A The Alkali Metals

• Group 1A elements have one valence electron.
• They form 1 ions after losing the one valence
  electron.

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• Group 2A elements are the alkaline earth metals.

2A
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Group 2A The Alkaline Earth Metals

• Group 2A elements have two valence electrons.
• They form 2 ions after losing the two valence
  electrons.

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• Group 3A is called the boron group.

3A
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Group 3A The Boron Group

• Group 3A elements have three valence electrons.
• They form 3 ions after losing the three valence
  electrons.

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• Group 4A is called the carbon group.

4A
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Group 4A The Carbon Group

• Group 4A elements have four valence electrons.
• They form 4 ions after losing the four valence
  electrons or 4- ions after gaining four
  additional electrons.

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• Group 5A is called the nitrogen group.

5A
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Group 5A The Nitrogen Group

• Group 5A elements have five valence electrons.
• They form 3- ions after gaining three more
  electrons.

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• Group 6A is called the oxygen group.

6A
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Group 6A The Oxygen Group

• Group 6A elements have six valence electrons.
• They form 2- ions after gaining two more
  electrons.

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• Group 7A is called the halogens.

7A
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Group 7A The Halogens

• Group 7A elements have seven valence electrons.
• They form 1- ions after gaining one more electron.

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• The word halogen is from the Greek words for
  salt former so named because the compounds that
  halogens form with metals are saltlike.

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• Group 8A elements are the noble gases.

8A
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Group 8A The Noble Gases

• Group 8A elements have eight valence electrons
  except for helium which only has two.
• The noble gases, with a full complement of
  valence electrons, are generally unreactive.

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• All transition elements have 2 valence electrons.

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Question
How many valence electrons are in an atom of each
of the following elements?
a) Magnesium (Mg)
(2)
b) Selenium (Se)
(2)
c) Tin (Sn)
(2)
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Metals

• Metals are elements that have luster, conduct
  heat and electricity, and usually bend without
  breaking.
• Most metals have one, two, or three valence
  electrons.

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Metals

• All metals except mercury are solids at room
  temperature in fact, most have extremely high
  melting points.

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Metal Reactivity

• A metals reactivity is its ability to react with
  another substance.

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Metal Reactivity

• Consult the Activity Series of Metals in the
  Chemistry Reference Tables to determine the more
  active metal.

a) cobalt (Co) or manganese (Mn)
(manganese)
b) barium (Ba) or sodium (Na)
(barium)
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Nonmetals

• Although the majority of the elements in the
  periodic table are metals, many nonmetals are
  abundant in nature.

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Nonmetals

• Most nonmetals dont conduct electricity, are
  much poorer conductors of heat than metals, and
  are brittle when solid.
• Many are gases at room temperature those that
  are solids lack the luster of metals.

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Nonmetals

• Their melting points tend to be lower than those
  of metals.
• With the exception of carbon, nonmetals have
  five, six, seven, or eight valence electrons.

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Nonmetal Reactivity

• A nonmetals reactivity is its ability to react
  with another substance.

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Nonmetal Reactivity

• Consult the Activity Series of Metals in the
  Chemistry Reference Tables to determine the less
  active nonmetal.

a) fluorine (F2) or chlorine (Cl2)
(chlorine)
b) chlorine (Cl2) or iodine (I2)
(iodine)
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Metalloids

• Metalloids have some chemical and physical
  properties of metals and other properties of
  nonmetals.
• In the periodic table, the metalloids lie along
  the border between metals and nonmetals.

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Metalloids
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Metalloids

• Some metalloids such as silicon, germanium (Ge),
  and arsenic (As) are semiconductors.
• A semiconductor is an element that does not
  conduct electricity as well as a metal, but does
  conduct slightly better than a nonmetal.

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• The metals are in pink.
• The nonmetals are in lime green.

• The metalloids are in white.

B
Si
As
Ge
Sb
Te
Po
At
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Question

• Match each element in Column A with the best
  matching description in Column B. Each Column A
  element may match more than one description from
  Column B.

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Question
Column A
1. strontium
2. chromium
3. iodine
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Answers
1. strontium
b, c
d
2. chromium
a, c
3. iodine
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Periodic Trends
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• Because the periodic table relates group and
  period numbers to valence electrons, its useful
  in predicting atomic structure and, therefore,
  chemical properties.

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Atomic Size (Atomic radius)

Radius

• Atomic Radius half the distance between two
  nuclei of a diatomic molecule.

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Trends in Atomic Size (Radii)

• Atomic size is Influenced by two factors.
• Energy Level A higher energy level is further
  away.
• Charge on nucleus - More charge (protons) pulls
  electrons in closer.

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Group Trend for Atomic Radii
H
Li
Na

• As you go down a group, each atom has another
  energy level so the atoms get bigger.

K
Rb
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Period Trend for Atomic Radii

• As you go across a period, the radius gets
  smaller.
• Atoms are in the same energy level, but as you
  move across the chart, atoms have a greater
  nuclear charge (more protons).
• Therefore, the outermost electrons are closer.

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Period Trend for Atomic Radii
Na
Mg
Al
Si
P
S
Cl
Ar
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Question

• Choose the element from the pair with the larger
  atomic radius.

a) lithium (Li) or beryllium (Be)
(lithium)
b) silicon (Si) or tin (Sn)
(tin)
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Question

• Choose the element from the pair with the smaller
  atomic radius.

a) silver (Ag) or gold (Au)
(silver)
b) cesium (Cs) or barium (Ba)
(barium)
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Ionic Size (Ionic radius)

• When an atom gains or loses one or more
  electrons, it becomes an ion.
• Because an electron has a negative charge,
  gaining electrons produces a negatively charged
  ion, an anion, whereas losing electrons produces
  a positively charged ion, a cation.

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Ionic Size (Ionic radius)

• As you might expect, the loss of electrons
  produces a positive ion with a radius that is
  smaller than that of the parent atom.
• Conversely, when an atom gains electrons, the
  resulting negative ion is larger than the parent
  atom.

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Ionic Size (Ionic radius)

• Practically all of the elements to the left of
  group 4A of the periodic table commonly form
  positive ions.
• As with neutral atoms, positive ions become
  smaller moving across a period and become larger
  moving down through a group.

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Ionic Radius Group Trend
Li1

• As you go down a group, you are adding an energy
  level.
• Ions get bigger as you go down.

Na1
K1
Rb1
Cs1
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Ionic Size (Ionic radius)

• Most elements to the right of group 4A (with the
  exception of the noble gases in group 8A) form
  negative ions.
• These ions, although considerably larger than the
  positive ions to the left, also decrease in size
  moving across a period.

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Ionic Size (Ionic radius)

• Like the positive ions, the negative ions
  increase in size moving down through a group.

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Ionic Radius Period Trend

• Across the period, nuclear charge increases so
  they get smaller.
• Energy level changes between anions and cations.

N-3
O-2
F-1
B3
Li1
C4
Be2
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Question

• Choose the element from the pair with the smaller
  radius.

a) silver (Ag) or the silver ion (Ag1)
(silver ion)
b) oxygen (O) or the oxygen ion (O2-)
(oxygen)
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Question
For each of the following pairs, predict which
atom is larger.
a) Mg, Sr
(Sr)
d) Ge, Br
(Ge)
(Sr)
b) Sr, Sn
e) Cr, W
(W)
(Sn)
c) Ge, Sn
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Question
For each of the following pairs, predict which
atom or ion is larger.
a) Mg, Mg2
(Mg)
(I-)
d) Cl, I
(S2-)
b) S, S2
(Na)
e) Na, Al3
c) Ca2, Ba2
(Ba2)
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Ionization Energy

• Ionization energy is the amount of energy
  required to completely remove an electron from a
  gaseous atom.
• Removing one electron makes a 1 ion. The energy
  required to do this is called the first
  ionization energy.

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What Determines Ionization Energy (IE)

• The greater the nuclear charge ( of protons),
  the greater IE.
• The distance from the nucleus increases IE.

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Ionization Energy

• As you go down a group, first IE decreases
  because the electron is further away, thus there
  is more shielding by the core electrons from the
  pull of the positive nucleus.

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Ionization Energy

• All the atoms in the same period have the same
  energy level.
• They have the same shielding, but as you move
  across the chart there is an increasing nuclear
  charge.
• Therefore, IE generally increases from left to
  right.

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Question

• Choose the element from the pair with the greater
  ionization energy.

a) silver (Ag) or iodine (I)
(iodine)
b) oxygen (O) or selenium (Se)
(oxygen)
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Question

• Choose the element from the pair with the smaller
  ionization energy.

a) chromium (Cr) or tungsten (W)
(tungsten)
b) sodium (Na) or magnesium (Mg)
(sodium)
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Electronegativity

• Electronegativity is the tendency for an atom to
  attract electrons to itself when it is chemically
  combined with another element.
• Large electronegativity means it pulls the
  electron toward it.

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Electronegativity

• The further you go down a group, the farther the
  electron is away from the nucleus and the more
  electrons an atom has.
• It is harder to attract extra electrons if the
  available energy level is far from the nucleus,
  so the electronegativity decreases.

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Electronegativity

• As you go across a row, electronegativity
  increases as the metallic character of the
  elements decreases.

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Question

• Choose the element from the pair with the greater
  electronegativity.

a) sodium (Na) or rubidium (Rb)
(sodium)
b) selenium (Se) or bromine (Br)
(bromine)
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Question

• Choose the element from the pair with the smaller
  electronegativity.

a) magnesium (Mg) or calcium (Ca)
(calcium)
b) nitrogen (N) or oxygen (O)
(nitrogen)