Transcript Slide 1

PowerPoint to accompany
Chapter 1:
Part 2
Atomic theory
Introduction:
Matter,
Measurement and
Molecules
Atomic Theory
The theory that atoms are the fundamental
building blocks of matter came into being
during the period 1803 to 1807 in the work
of an English schoolteacher, John Dalton.
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Dalton’s Postulates

Each element is composed of extremely small particles called
atoms.

All atoms of a given element are identical to one another in mass
and other properties, but the atoms of a particular element are
different from the atoms of all other elements.

Atoms of an element are not changed into atoms of a different
element by chemical reactions; atoms are neither created nor
destroyed in chemical reactions. This is the basis of the law of
conservation of mass (or law of conservation of matter) which states
that the total mass of substances present at the end of a chemical
process is the same as the mass of substances present before the
process took place.

Compounds are formed when atoms of more than one element
combine; a given compound always has the same relative number
and kind of atoms. This is the basis of the law of constant
composition (or law of definite proportions) which states that the
relative numbers and kinds of atoms are constant, i.e. the elemental
composition of a pure substance never varies.
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The Law of Multiple
Proportions

Was deduced by Dalton from the preceding four
postulates and states that:

If two elements A and B combine to form more
than one compound, the masses of B that can
combine with a given mass of A are in the ratio
of small whole numbers.
Examples
 H2O consists of 2 hydrogens to 1 oxygen
 H2O2 consists of 1 hydrogen to 1 oxygen
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Cathode Rays and Electrons



Streams of negatively charged particles
were found to emanate from cathode
tubes.
J. J. Thompson is credited with its
discovery in 1897.
He determined the charge/mass ratio of the
electron to be 1.78 x 108 Cg-1.
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Cathode Ray Tubes (TV’s!)
Cathode-  Anode+
electons
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Millikan Oil Drop Experiment
Once the charge/mass ratio of the electron
was known, determination of either the
charge or the mass of an electron would
yield the other.
Figure 1.19
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Millikan Oil Drop Experiment
In 1909, Robert Millikan at the
University of Chicago determined the
charge on the electron.
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Radioactivity
The spontaneous emission of radiation by
an atom was first observed by Henri
Becquerel. It was also studied by Marie
and Pierre Curie.
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Radioactivity
Three types of radiation were discovered by
Ernest Rutherford



 particles
 particles
 rays
Figure 1.21
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Discovery of the Nucleus
Ernest Rutherford shot  particles at a thin
sheet of gold foil and observed the pattern
of scatter of the particles.
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The Nuclear Atom
Some particles were
deflected at large
angles. This led
Rutherford to postulate
that the atom had a
nucleus.
Figure 1.22
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The Nuclear Atom


Rutherford postulated a very small,
dense nucleus with the electrons around
the outside of the atom.
Most of the volume of the atom is empty
space.
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Subatomic Particles



Protons and electrons are the only particles that
have a charge.
Protons and neutrons have essentially the same
mass.
The mass of an electron is so small we ignore it.
Table 1.5
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Atomic Numbers, Mass
Numbers and Isotopes
Mass Number (A)
12
6
C
Symbol of element
Atomic Number (Z)
The number of protons in the nucleus of an atom of any
particular element is called that element’s atomic number (Z).
The mass number (A) of an atom is the total number of
protons plus neutrons in that atom.
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Isotopes
Atoms with identical atomic numbers (Z) but
different mass numbers (A), or atoms with
the same number of protons which differ only
in the number of neutrons are called
isotopes.
Examples:
11
6C
12
6C
carbon-12
isotope
13
6C
14
6C
carbon-14
isotope
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Which Isotopes make up carbon in
diamonds, graphite, petroleum and coal?
Where does 14C originate and why is
there so little in nature except for living
things?
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Atomic Mass
Atomic and molecular masses can be
measured with great accuracy with a mass
spectrometer.
Figure 1.23
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Average Atomic Mass
(commonly called Atomic
Mass)

We use average masses in calculations,
because we use large amounts of atoms and
molecules in the real world.

Average atomic mass is calculated from the
fractional abundance of each isotope and mass
of that isotope.
For example, the average atomic mass of C made up mostly of 12C (98.93%) and 13C
(1.07%) - is 12.01 u.
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Mass Spectrum of Chlorine Atoms
How many protons
does Chlorine have?
a) What makes the rest
of the masses?
b) Why are there
2 different isotopes?
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Mass Spectrum of Chlorine Atoms
What would the
mass spectrum
of Chlorine
Molecules look
like?
a) What are the masses?
b) Which is least common?
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Periodicity
Figure 1.25
When one looks at the chemical properties
of elements, one notices a repeating
pattern of reactivities.
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Periodic Table


A systematic catalogue of elements.
Elements are arranged in order of atomic
number.
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Aluminum wrench,
Copper pipe, Lead
shot, Gold nuggets
Bromine L+V vial,
Iodine crystals,
Carbon black +
Diamond
+ Graphite pencil
Sulfur
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Periodic Table

The rows are called periods.


Elements in each row are similar sizes.
The columns are called groups.


Elements in columns bond alike and
Have similar chemical properties.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Groups
Table 1.7
The above five groups are
known by their names.
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Periodic Table
Nonmetals are on the right side of the
periodic table (with the exception of H).
Metalloids border the stair-step line (with the
exception of Al and Po).
Metals are on the left side of the chart.
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Molecules and Chemical
Formulae
The subscript to the right of
the symbol of an element tells
the number of atoms of that
element in one molecule of
the compound.
Notice how the composition of
each compound is given by
its chemical formula.
Figure 1.29
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Diatomic Molecules
Figure 1.28
These seven elements occur naturally as
molecules containing two atoms. Which are
gases?
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Molecular Compounds
Molecular compounds are composed of
molecules and almost always contain
only nonmetals.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Types of Formulae
Empirical formulae give the lowest whole-number
ratio of atoms of each element in a compound, e.g.
HO.
Molecular formulae give the exact number of atoms
of each element in a compound, e.g. H2O2.
Structural formulae show which atoms are attached
to which within the molecule, e.g. H-O-O-H.
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Picturing Molecules
Different
representations of
the methane (CH4)
molecule.
Figure 1.30
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Common Ionic Charges
Figure 1.31

When atoms lose or gain electrons, they become
ions.

Cations are positive and are formed by
elements on the left side of the periodic chart.

Anions are negative and are formed by
elements on the right side of the periodic chart.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Ionic Compounds
Ionic compounds (such as NaCl) are
generally formed between metals and
nonmetals.
Figure 1.32
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Writing Ionic Formulae

Because compounds are electrically neutral, one
can determine the formula of a compound by:
 writing the value of the charge on the cation
as the subscript on the anion.
 writing the value of the charge on the anion
as the subscript on the cation.
Note: if the subscripts are not in the lowest
whole number ratio, simplify it, e.g. Ca2O2
would become CaO.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical Nomenclature
Positive Ions (Cations)
a) Cations formed from metal atoms have the same
name as the metal, e.g. Na+ is the sodium ion.
b) If a metal can form different cations, the positive
charge is indicated by a Roman numeral in
parentheses following the name of the metal,
e.g. Au+ is the gold(I) ion and Au3+ is the gold(III)
ion.
c) Cations formed from nonmetal atoms have
names that end in -ium, e.g. NH4+ is the
ammonium ion.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical Nomenclature
Common Cations
Table 1.8
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Chemical Nomenclature
Negative Ions (Anions)
a) The names of the monatomic anions are formed by
replacing the ending of the name of the element with ide, e.g. O2- is the oxide ion, OH- is the hydroxide ion,
N3- is the nitride ion.
b) Polyatomic anions containing oxygen (called
oxyanions) have names ending in -ate for the most
oxidized form e.g. SO42- is the sulfate ion or –ite for the
more reduced form, e.g. SO32- is the sulfite ion.
c) Anions derived by adding H+ to an oxyanion are named
by adding the prefix hydrogen e.g. HCO3- is the
hydrogen carbonate ion (bicarbonate ion) or
dihydrogen as in dihydrogen phosphate H2 PO4-2 .
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical Nomenclature
Common Anions
Table 1.9
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Common Oxy-anion
Names, Formulae & Charges
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K-Iron-CN
salts:
Left:
KFe2+(CN)3
(Potassium
ferrocyanide)
Right:
KFe3+(CN)4
(Potassium ferricyanide)
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Chemical Nomenclature Figure 1.34
More on naming oxyanions
Examples:
ClO4 perchlorate ion (one more O atom than chlorate)
ClO3- chlorate (most common oxidized form)
ClO2- chlorite ion (one less O atom than chlorate)
ClO- hypochlorite ion (one O atom less than chlorite)
Which is the most oxidized ion? Which is most reduced?
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Chemical Nomenclature Figure 1.36
Name and Formulae of Acids
1.
Acids containing anions
whose names end in -ide
are named by changing the
-ide ending to -ic, adding
the prefix hydro- to this
anion name, and then
following with the word
acid.
2.
Acids containing anions
whose names end in -ate
or -ite are named by
changing the -ate ending to
-ic and the -ite ending to
-ous and then adding the
word acid.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Chemical Nomenclature
Ionic Compounds
Names of ionic compounds consist of the
cation followed by the anion name, e.g.
Cu(ClO4)2 is copper(II) perchlorate, and
CaCO3 is calcium carbonate.
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Table 1.10
Chemical Nomenclature
Binary Molecular Compounds
1. The name of the (more reduced) element farther to
the left in the periodic table is written first (usually a
metal), eg. MnO2 , NaCl, FeFe2 S4 , K(MnO4) .
It’s all Greek prefixes to me!
2. If both elements are in the same group in the periodic
table, the one having the higher atomic number (more
reduced) is written first eg. SO2 , SiC .
3. The name of the second element is given an -ide ( or
–ite or –ate) ending.
4. Greek prefixes are used to indicate the number of
atoms of each element.
Examples
N2O4 is dinitrogen tetroxide
P4S10 is tetraphosphorus decasulfide.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia