The Periodic Table

1
The Periodic Table
2
The Modern Periodic Table

• The modern periodic table is based on the atomic
  numbers of the elements.

3
The Modern Periodic Table

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

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The Periodic Table

• Elements in vertical columns showed similar
  properties.

For example, alkaline earth metals have high
melting points and low density and are silver in
color, ductile, and malleable.
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Periodicity

• This repeated pattern is an example of
  periodicity in the properties of elements.
• Periodicity is the tendency to recur at regular
  intervals.

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The Periodic Table

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

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The Periodic Table

• There are 7 periods in the table.

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The Periodic Table

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

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Metals, Nonmetals, and Metalloids

• There are 3 main regions of the table metals,
  nonmetals and metalloids.

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Metals

• The metals are in blue.

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Metals

• Metals are elements that have luster (are shiny),
  conduct heat and electricity, and usually bend
  without breaking (malleable).
• Metals are also ductile (can be drawn out into a
  wire).

12
Metals

• Most metals have one, two, or three valence
  electrons.
• Metals tend to lose electrons in order to achieve
  the stability of a filled octet.

13
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.
• Reactivity for metals increases as you go down a
  
  group and left
  across a period.

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Metal Reactivity
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Metal Reactivity

• 1. 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

• The nonmetals are in yellow.

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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.

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Nonmetals

• Many are gases at room temperature those that
  are solids lack the luster of metals.
• Their melting points tend to be lower than those
  of metals.

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Nonmetals

• With the exception of carbon, nonmetals have
  five, six, seven, or eight valence electrons.
• Nonmetals tend to gain electrons in order to
  achieve the stability of a filled octet.

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Nonmetal Reactivity

• A nonmetals reactivity is its ability to react
  with another substance.
• Reactivity for nonmetals increases as you go left
  to right and up the periodic table.

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Nonmetal Reactivity

• 2. Consult the Activity Series of Halogens 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

• The metalloids are in pink.

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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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Group Names

• Groups 1, 2, and 13 - 18 (Group A elements) are
  called representative (main group) elements.

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Group Names

• Groups 3 - 12 (Group B elements) are called
  transition elements.

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Group 1 The Alkali Metals
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Group 1 The Alkali Metals

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

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Group 2 The Alkaline Earth Metals
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Group 2 The Alkaline Earth Metals

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

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Group 17 The Halogens
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Group 17 The Halogens

• Group 17 elements have seven valence electrons.
• They form 1- ions after gaining one more electron.

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Group 18 The Noble Gases
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Group 18 The Noble Gases

• Group 18 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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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)
(6)
c) Tin (Sn)
(4)
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Question

• 2. 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
2. chromium
d
3. iodine
a, c
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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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Periodic Trends
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Atomic Size (Atomic Radius)

• The atomic radius of a chemical element is a
  measure of the size of its atoms, usually the
  mean or typical distance from the nucleus to the
  boundary of the surrounding cloud of electrons.

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Trends in Atomic Size (Radii)

• Atomic size is influenced by two factors.
• Energy Level A higher energy level is farther
  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
3. (a) State why atoms get bigger as you go down
a group on the periodic table. (b) State why
the radius decreases across a period.
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Question
4. 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
5. 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)

• Ionic radius is the radius of an atom's ion.
• When an atom gains or loses one or more
  electrons, it becomes an ion.

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Ionic Size (Ionic Radius)

• Recall that metals tend to lose electrons in
  order to achieve the stability of a filled octet.
• As a result, metals tend to form cations which
  are positive ions.

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Ionic Size (Ionic Radius)

• A cation has a smaller radius than its neutral
  atom.

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Ionic Size (Ionic Radius)

• Nonmetals tend to gain electrons in order to
  achieve the stability of a filled octet.
• As a result, nonmetals tend to form anions which
  are negative ions.

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Ionic Size (Ionic Radius)

• An anion has a larger radius than its neutral
  atom.

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Question
6. 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
7. For each of the following pairs, predict which
atom is larger.
a) Mg, Sr
(Sr)
(Ge)
d) Ge, Br
(Sr)
b) Sr, Sn
(W)
e) Cr, W
(Sn)
c) Ge, Sn
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Question
8. 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 (IE) 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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Ionization Energy
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What Determines Ionization Energy (IE)

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

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Ionization Energy

• As you go down a group, first IE decreases.
• This is because the electron is farther 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 because of the increasing number of
  protons.
• Therefore, IE generally increases from left to
  right.

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Question

• 9. (a) State why ionization energy decreases as
  you go down a group.
• (b) State why ionization energy increases
  across a period.

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Question
10. 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
11. 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 a pair of electrons to itself when it is
  chemically combined with another element.
• Large electronegativity means the atom pulls the
  electron toward it.

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Electronegativity
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Electronegativity

• Electronegativity decreases down a group.
• The farther you go down a group, the farther the
  electron is away from the nucleus and the more
  electrons an atom has.

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Electronegativity

• It is harder to attract extra electrons if the
  available energy level is far from the nucleus,
  so electronegativity decreases.

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Electronegativity

• As you go across a row, electronegativity
  increases.
• Remember the radius of the atoms decreases across
  the periodic table.
• With the smaller size, there is a greater
  attraction for the nucleus to electrons.

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Question

• 12. (a) State why electronegativity decreases as
  you go down a group.
• (b) State why electronegativity increases
  across a period.

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Question
13. 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
14. Choose the element from the pair with the
smaller electronegativity.
a) magnesium (Mg) or calcium (Ca)
(calcium)
b) nitrogen (N) or oxygen (O)
(nitrogen)
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Summary of the Periodic Trends

• Moving Left ? Right
• Atomic Radius Decreases
• Ionization Energy Increases
• Electronegativity Increases
• Moving Top ? Bottom
• Atomic Radius Increases
• Ionization Energy Decreases
• Electronegativity Decreases

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Electron Affinity

• Electron affinity is the energy change that
  accompanies a gaseous atom when it gains an
  electron to form a gaseous ion.

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Metallic Character andPeriodic Trends

• Recall again that metals tend to lose electrons
  in order to achieve the stability of a filled
  octet.
• Therefore as metallic character increases toward
  francium (Fr) ionization energy,
  electronegativity, and electron affinity
  decrease.