Periodic Properties of the Elements - PowerPoint PPT Presentation

Loading...

PPT – Periodic Properties of the Elements PowerPoint presentation | free to download - id: 76ffd9-NTllM



Loading


The Adobe Flash plugin is needed to view this content

Get the plugin now

View by Category
About This Presentation
Title:

Periodic Properties of the Elements

Description:

Chapter 6 Periodic Properties of the Elements Mn Oxides Hg and Br Mn Oxides Hg and Br PowerPoint to accompany * Power Point Slides edited and annotated by Tark ... – PowerPoint PPT presentation

Number of Views:54
Avg rating:3.0/5.0
Slides: 33
Provided by: Pears178
Learn more at: http://faculty.camosun.ca
Category:

less

Write a Comment
User Comments (0)
Transcript and Presenter's Notes

Title: Periodic Properties of the Elements


1
Chapter 6
  • Periodic Properties of the Elements

2
Development of the Periodic Table
  • Elements in the same group generally have similar
    chemical properties.
  • Properties are not identical, however.

3
O2 gas In a gas burette
S8 rings stack to make Orthorhombic Crystals
4
Development of the Periodic Table
  • Dmitri Mendeleev and Lothar Meyer independently
    came to the same conclusion about how elements
    should be grouped.

Figure 6.2
5
Development of the Periodic Table
Mendeleev, for instance, predicted the discovery
of germanium,
Table 6.1
  • (which he called eka-silicon) as an element with
    an atomic weight between that of zinc and
    arsenic, but with chemical properties similar to
    those of silicon.

6
Effective Nuclear Charge
  • In a many-electron atom, electrons are both
    attracted to the nucleus and repelled by other
    electrons.
  • The nuclear charge that an electron experiences
    depends on both factors.

7
Effective Nuclear Charge
  • The effective nuclear charge, Zeff, is found this
    way
  • Zeff Z - S
  • where Z is the atomic number and S is a
    screening constant, usually close to the number
    of inner electrons.

Figure 6.3
8
Sizes of Atoms and Ions
  • The bonding atomic radius is defined as one-half
    of the distance between covalently bonded nuclei.

Figure 6.4
9
Sizes of Atoms
  • Bonding atomic radius
  • tends to
  • decrease from left to right across a row
  • due to increasing Zeff
  • increase from the top to the bottom of a column
    due to increasing value of n

Figure 6.5
10
Sizes of Ions
  • Ionic size depends upon
  • Nuclear charge
  • Number of electrons
  • Orbitals in which electrons reside

Figure 6.6
11
Sizes of Ions
  • Cations are smaller than their parent atoms
  • The outermost electron is removed and repulsions
    are reduced.

Figure 6.6
12
Sizes of Ions
  • Anions are larger than their parent atoms
  • Electrons are added and repulsions are increased.

Figure 6.6
13
Sizes of Ions
  • Ions increase in size as you go down a column
  • Due to increasing value of n.
  • In an isoelectronic series, ions have the same
    number of electrons.
  • Ionic size decreases with an increasing nuclear
    charge.

14
Ionisation Energy
  • Amount of energy required to remove an electron
    from the ground state of a gaseous atom or ion
  • First ionisation energy is that energy required
    to remove the first electron.
  • Second ionisation energy is that energy required
    to remove the second electron, etc.

15
Ionisation Energy
  • It requires more energy to remove each successive
    electron.
  • When all valence electrons have been removed, the
    ionisation energy takes a quantum leap.

Table 6.2
16
(No Transcript)
17
Trends in First Ionisation Energies
  • As one goes down a column, less energy is
    required to remove the first electron.
  • For atoms in the same group, Zeff is essentially
    the same, but the valence electrons are farther
    from the nucleus.

Figure 6.8
18
Trends in First Ionisation Energies
  • Generally, as one goes across a row, it gets
    harder to remove an electron
  • As you go from left to right, Zeff increases.

Figure 6.7
19
Electron Affinities
  • Energy change accompanying the addition of an
    electron to a gaseous atom
  • Cl e- ??? Cl-
  • ?E -349 kJ/mol

20
Trends in Electron Affinity
  • In general, electron affinity becomes more
    exothermic as you go from left to right across a
    row.

Figure 6.9
21
Metal, Nonmetals and Metalloids
Figure 6.10
22
Metals versus Nonmetals
Table 6.3
  • Differences between metals and nonmetals tend to
    revolve around these properties.

23
Metals versus Nonmetals
  • Metals tend to form cations.
  • Nonmetals tend to form anions.

Figure 6.12
24
Metals
  • Tend to be lustrous, malleable, ductile, and good
    conductors of heat and electricity.
  • Figure 6.11

25
Metals
  • Compounds formed between metals and nonmetals
    tend to be ionic.
  • Metal oxides tend to be basic.

Figure 6.15
26
Reactivity of Metals
  • Metal reactivity varies with the type of metal.
  • Alkali metals are very reactive because it is
    very easy to remove the 1s1 electron. All alkali
    metals react vigorously with water to produce
    metal hydroxide and hydrogen gas
  • 2M(s) 2H2O(l) -gt 2MOH(aq) H2(g)

Figure 6.13
27
Reactivity of Metals
  • Alkaline earth metals are less reactive than
    alkali metals.
  • Mg reacts only with steam, Be does not react with
    water, but others react readily with water.
  • Reactivity tends to increase as you move down
    groups.

Figure 6.14
28
Nonmetals
  • Dull, brittle substances that are poor conductors
    of heat and electricity.
  • Tend to gain electrons in reactions with metals
    to acquire noble gas configuration.
  • Substances containing only nonmetals are
    molecular compounds.
  • Most nonmetal oxides are acidic.

Figure 6.16
29
Metalloids
  • Have some characteristics of metals and some of
    nonmetals.
  • For instance, silicon looks shiny, but is brittle
    and a fairly poor conductor.

30
Mn Oxides
31
Hg and Br
32
(No Transcript)
About PowerShow.com