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The Periodic Table

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Title: 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.

4
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.
5
Periodicity
  • This repeated pattern is an example of
    periodicity in the properties of elements.
  • Periodicity is the tendency to recur at regular
    intervals.

6
The Periodic Table
  • On the periodic table a period, sometimes also
    called a series, consists of the elements in a
    horizontal row.

7
The Periodic Table
  • There are 7 periods in the table.

8
The Periodic Table
  • A group, sometimes also called a family, consists
    of the elements in a vertical column.

9
Metals, Nonmetals, and Metalloids
  • There are 3 main regions of the table metals,
    nonmetals and metalloids.

10
Metals
  • The metals are in blue.

11
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.

14
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.

15
Metal Reactivity
16
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)
17
Nonmetals
  • The nonmetals are in yellow.

18
Nonmetals
  • Although the majority of the elements in the
    periodic table are metals, many nonmetals are
    abundant in nature.

19
Nonmetals
  • Most nonmetals dont conduct electricity, are
    much poorer conductors of heat than metals, and
    are brittle when solid.

20
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.

21
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.

22
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.

23
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)
24
Metalloids
  • The metalloids are in pink.

25
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.

26
Group Names
  • Groups 1, 2, and 13 - 18 (Group A elements) are
    called representative (main group) elements.

27
Group Names
  • Groups 3 - 12 (Group B elements) are called
    transition elements.

28
(No Transcript)
29
Group 1 The Alkali Metals
30
Group 1 The Alkali Metals
  • Group 1 elements have one valence electron.
  • They form 1 ions after losing the one valence
    electron.

31
Group 2 The Alkaline Earth Metals
32
Group 2 The Alkaline Earth Metals
  • Group 2 elements have two valence electrons.
  • They form 2 ions after losing the two valence
    electrons.

33
Group 17 The Halogens
34
Group 17 The Halogens
  • Group 17 elements have seven valence electrons.
  • They form 1- ions after gaining one more electron.

35
Group 18 The Noble Gases
36
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.

37
Valence Electrons
38
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)
39
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.

40
Question
Column A
1. strontium
2. chromium
3. iodine
41
Answers
1. strontium
b, c
2. chromium
d
3. iodine
a, c
42
Periodic Trends
43
  • Because the periodic table relates group and
    period numbers to valence electrons, its useful
    in predicting atomic structure and, therefore,
    chemical properties.

44
Periodic Trends
45
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.

46
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.

47
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
48
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.

49
Period Trend for Atomic Radii
Na
Mg
Al
Si
P
S
Cl
Ar
50
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.
51
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)
52
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)
53
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.

54
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.

55
Ionic Size (Ionic Radius)
  • A cation has a smaller radius than its neutral
    atom.

56
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.

57
Ionic Size (Ionic Radius)
  • An anion has a larger radius than its neutral
    atom.

58
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)
59
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
60
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)
61
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.

62
Ionization Energy
63
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.

64
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.

65
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.

66
Question
  • 9. (a) State why ionization energy decreases as
    you go down a group.
  • (b) State why ionization energy increases
    across a period.

67
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)
68
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)
69
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.

70
Electronegativity
71
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.

72
Electronegativity
  • It is harder to attract extra electrons if the
    available energy level is far from the nucleus,
    so electronegativity decreases.

73
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.

74
Question
  • 12. (a) State why electronegativity decreases as
    you go down a group.
  • (b) State why electronegativity increases
    across a period.

75
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)
76
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)
77
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

78
Electron Affinity
  • Electron affinity is the energy change that
    accompanies a gaseous atom when it gains an
    electron to form a gaseous ion.

79
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.
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