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```									Atoms, Molecules, and Ions

De La Salle College Oaklands

Chapter 2

Chapter 2          1
The Atomic Theory of Matter
• John Dalton:
–   Each element is composed of atoms
–   All atoms of an element are identical.
–   In chemical reactions, the atoms are not changed.
–   Compounds are formed when atoms of more than one
element combine.
• Dalton’s law of multiple proportions: When two
elements form different compounds, the mass ratio of
the elements in one compound is related to the mass
ratio in the other by a small whole number.

Chapter 2                     2
The Discovery of Atomic Structure
The ancient Greeks were the first to postulate that
matter consists of indivisible constituents.
Later scientists realized that the atom consisted of
charged entities.

Cathode Rays and Electrons
A cathode ray tube (CRT) is a hollow vessel with an
electrode at either end.
A high voltage is applied across the electrodes.

Chapter 2                    3
The Discovery of Atomic Structure
Cathode Rays and Electrons
The voltage causes negative particles to move from the
negative electrode to the positive electrode.
The path of the electrons can be altered by the presence
of a magnetic field.

Chapter 2                  4
The Discovery of Atomic Structure
Cathode Rays and Electrons
Consider cathode rays leaving the positive electrode
through a small hole.
•If they interact with a magnetic field perpendicular to
an applied electric field, the cathode rays can be
deflected by different amounts.
•The amount of deflection of the cathode rays depends
on the applied magnetic and electric fields.
•In turn, the amount of deflection also depends on the
charge to mass ratio of the electron.

Chapter 2                  5
The Discovery of Atomic Structure
Cathode Rays and Electrons

Chapter 2       6
The Discovery of Atomic Structure
Cathode Rays and Electrons
In 1897, Thomson determined the charge to mass ratio
of an electron to be 1.76  108 C/g.
Goal: find the charge on the electron to determine its
mass.

Chapter 2                  7
The Discovery of Atomic Structure
Cathode Rays and Electrons
Consider the following experiment:
•Oil drops are sprayed above a positively charged plate
containing a small hole.
•As the oil drops fall through the hole, they are given a
negative charge.
•Gravity forces the drops downward. The applied
electric field forces the drops upward.
•When a drop is perfectly balanced, the weight of the
drop is equal to the electrostatic force of attraction
between the drop and the positive plate.

Chapter 2                   8
The Discovery of Atomic Structure
Cathode Rays and Electrons

Chapter 2       9
The Discovery of Atomic Structure
Cathode Rays and Electrons
Using this experiment, Millikan determined the charge
on the electron to be 1.60  10-19 C.

Knowing the charge to mass ratio, 1.76  108 C/g,
Millikan calculated the mass of the electron:
9.10  10-28 g.

With more accurate numbers, we get the mass of the
electron to be 9.10939  10-28 g.

Chapter 2                  10
The Discovery of Atomic Structure
Consider the following experiment:
•A radioactive substance is placed in a shield containing
a small hole so that a beam of radiation is emitted from
the hole.
•The radiation is passed between two electrically
charged plates and detected.
•Three spots are noted on the detector:
•a spot in the direction of the positive plate,
•a spot which is not affected by the electric field,
•a spot in the direction of the negative plate.

Chapter 2                  11
The Discovery of Atomic Structure

Chapter 2         12
The Discovery of Atomic Structure
A high deflection towards the positive plate corresponds
to radiation which is negatively charged and of low
mass. This is called b-radiation (consists of electrons).

No deflection corresponds to neutral radiation. This is

Small deflection towards the negatively charged plate
corresponds to high mass, positively charged radiation.

Chapter 2                  13
The Discovery of Atomic Structure

The Nuclear Atom
we conclude that the atom
consists of neutral, positively, and
negatively charged entities.
Thomson assumed all these
charged species were found in a
sphere.

Chapter 2   14
The Discovery of Atomic Structure
The Nuclear Atom
Rutherford carried out the following experiment:
A source of a-particles was placed at the mouth of a
circular detector.
The a -particles were shot through a piece of gold foil.
Most of the a-particles went straight through the foil
without deflection.
Some a-particles were deflected at high angles.
If the Thomson model of the atom was correct, then
Rutherford’s result was impossible.

Chapter 2                   15
The Discovery of Atomic Structure
The Nuclear Atom
Rutherford’s a-particle experiment:

Chapter 2     16
The Discovery of Atomic Structure
The Nuclear Atom
In order to get the majority of a-particles through a
piece of foil to be undeflected, the majority of the atom
must consist of a low mass, diffuse negative charge -
the electron.

To account for the small number of high deflections of
the a-particles, the center or nucleus of the atom must
consist of a dense positive charge.

Chapter 2                  17
The Discovery of Atomic Structure

The Nuclear Atom
Rutherford modified Thomson’s
model as follows:
assume the atom is spherical but
the positive charge must be
located at the center, with a
diffuse negative charge
surrounding it.

Chapter 2                 18
The Modern View of Atomic Structure
The atom consists of positive, negative, and neutral
entities (protons, electrons, and neutrons).

Protons and neutrons are located in the nucleus of the
atom, which is small. Most of the mass of the atom is
due to the nucleus.
There can be a variable number of neutrons for the same
number of protons. Isotopes have the same number of
protons but different numbers of neutrons.

Electrons are located outside of the nucleus. Most of
the volume of the atom is due to electrons.

Chapter 2                        19
The Modern View of Atomic Structure

Chapter 2        20
The Modern View of Atomic Structure

Chapter 2        21
The Modern View of Atomic Structure
Isotopes, Atomic Numbers, and Mass Numbers
Atomic number (Z) = number of protons in the nucleus.
Mass number (A) = total number of nucleons in the
nucleus (i.e., protons and neutrons).   A
By convention, for element X, we write Z X
Isotopes have the same Z but different A.

Chapter 2                22
The Modern View of Atomic Structure
Isotopes, Atomic Numbers, and Mass Numbers

Chapter 2           23
TO DO

• Answer Exercises 2.17-2.35 (red only).

Chapter 2              24
The Periodic Table
The Periodic Table is used to organize the 118 elements
in a meaningful way.
As a consequence of this organization, there are
periodic properties associated with the periodic table.

Chapter 2                  25
The Periodic Table
Columns in the periodic table are called groups
(numbered from 1A to 8A or 1 to 18).
Rows in the periodic table are called periods.
Metals are located on the left hand side of the periodic
table (most of the elements are metals).
Non-metals are located in the top right hand side of the
periodic table.
Elements with properties similar to both metals and
non-metals are called metalloids and are located at the
interface between the metals and non-metals.

Chapter 2                 26
The Periodic Table

Chapter 2   27
The Periodic Table
Some of the groups in the periodic table are given
special names.
These names indicate the similarities between group
members:
Group 1A: Alkali metals.
Group 2A: Alkaline earth metals.
Group 6A: Chalcogens.
Group 7A: Halogens.
Group 8A: Noble gases.

Chapter 2                  28
Molecules and Molecular Compounds
Molecules and Chemical Formulas
Molecules are assemblies of two or more atoms bonded
together.
Each molecule has a chemical formula.
The chemical formula indicates
which atoms are found in the molecule, and
in what proportion they are found.
Compounds formed from molecules are molecular
compounds.

Chapter 2             29
Molecules and Molecular Compounds
Molecular and Empirical Formulas
Molecular formulas
give the actual numbers and types of atoms in a molecule.
Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4.
Empirical formulas
give the relative numbers and types of atoms in a molecule.
That is, they give the lowest whole number ratio of atoms in a
molecule.
Examples: H2O, CO2, CO, CH4, HO, CH2.

Chapter 2                     30
Molecules and Molecular Compounds
Picturing Molecules
Molecules occupy three dimensional space.
However, we often represent them in two dimensions.

The structural formula gives the connectivity between
individual atoms in the molecule.
The structural formula may or may not be used to show
the three dimensional shape of the molecule.

Chapter 2                31
Molecules and Molecular Compounds
Picturing Molecules

Chapter 2   32
Molecules and Molecular Compounds

Picturing Molecules
If the structural formula does show the
shape of the molecule, then either a
perspective drawing, ball-and-stick
model, or space-filling model is used.

Chapter 2                 33
Ions and Ionic Compounds
When an atom or molecule loses electrons, it becomes
positively charged.
For example, when Na loses an electron it becomes Na+.
Positively charged ions are called cations.

Chapter 2                      34
Ions and Ionic Compounds
When an atom or molecule gains electrons, it becomes
negatively charged.
For example when Cl gains an electron it becomes Cl-.
Negatively charged ions are called anions.
An atom or molecule can lose more than one electron.

Chapter 2                     35
Ions and Ionic Compounds
Predicting Ionic Charge
The number of electrons an atom loses is related to its
position on the periodic table.
Metals tend to form cations whereas non-metals tend to
form anions.

Chapter 2                 36
Ions and Ionic Compounds
Ionic Compounds
The majority of chemistry involves the transfer of
electrons between species.
Example:
To form NaCl, the neutral sodium atom, Na, must lose an
electron to become a cation: Na+.
The electron cannot be lost entirely, so it is transferred to a
chlorine atom, Cl, which then becomes an anion: Cl-.
The Na+ and Cl- ions are attracted to form an ionic NaCl
lattice which crystallizes.

Chapter 2                        37
Ions and Ionic Compounds
Important: note that there are no easily identified NaCl
molecules in the ionic lattice. Therefore, we cannot use
molecular formulas to describe ionic substances.

Chapter 2                 38
Ions and Ionic Compounds
Ionic Compounds
Consider the formation of Mg3N2:
Mg loses two electrons to become Mg2+
Nitrogen gains three electrons to become N3-.
For a neutral species, the number of electrons lost and
gained must be equal.
However, Mg can only lose electrons in twos and N can
only accept electrons in threes.
Therefore, Mg needs to lose 6 electrons (2  3) and N
gain those 6 electrons (3  2).

Chapter 2                 39
Ions and Ionic Compounds
Ionic Compounds
I.e., 3Mg atoms need to form 3Mg2+ ions (total 3x2+
charges) and 2 N atoms need to form 2N3- ions (total
2x3- charges).
Therefore, the formula is Mg3N2.

Chapter 2                 40
TO DO

• Answer: Visualizing Concepts 2.3, 2.4
• Answer Exercises 2.35 – 2.55 (red only)

Chapter 2              41
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Naming of compounds, nomenclature, is divided into
organic compounds (those containing C) and inorganic
compounds (the rest of the periodic table).
Cations formed from a metal have the same name as
the metal.
Example: Na+ = sodium ion.
If the metal can form more than one cation, then the
charge is indicated in parentheses in the name.
Examples: Cu+ = copper(I); Cu2+ = copper(II).
Cations formed from non-metals end in -ium.
Example: NH4+ ammonium ion.

Chapter 2              42
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds

Chapter 2            43
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Monatomic anions (with only one atom) are called
-ide.
Example: Cl- is chloride ion.
Exceptions: hydroxide (OH-), cyanide (CN-), peroxide (O22-).
Polyatomic anions (with many atoms) containing
oxygen end in -ate or -ite. (The one with more oxygen is
called -ate.)
Examples: NO3- is nitrate, NO2- is nitrite.

Chapter 2                     44
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Polyatomic anions containing oxygen with more than
two members in the series are named as follows (in
order of decreasing oxygen):

per-….-ate
-ate
-ite
hypo-….-ite

Chapter 2                45
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds

Chapter 2            46
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Polyatomic anions containing oxygen with additional
hydrogens are named by adding hydrogen or bi- (one
H), dihydrogen (two H), etc., to the name as follows:
CO32- is the carbonate anion
HCO3- is the hydrogen carbonate (or bicarbonate) anion.
H2PO4- is the dihydrogen phosphate anion.

Chapter 2                        47
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Name the anion then cation for the ionic compound.
Example: BaBr2 = barium bromide.

Chapter 2                   48
Naming Inorganic Compounds
Names and Formulas of Acids
The names of acids are related to the names of anions:
-ide becomes hydro-….-ic acid;
-ate becomes -ic acid;
-ite becomes -ous acid.

Chapter 2              49
Naming Inorganic Compounds
Names and Formulas of Acids

Chapter 2   50
Naming Inorganic Compounds
Names and Formulas of Binary Molecular
Compounds
Binary molecular compounds have two elements.
The most metallic element is usually written first (i.e.,
the one to the farthest left on the periodic table).
Exception: NH3.
If both elements are in the same group, the lower one is
written first.
Greek prefixes are used to indicate the number of
atoms.

Chapter 2                  51
Naming Inorganic Compounds
Names and Formulas of Binary Molecular
Compounds

Chapter 2            52
To Do

• Read Section 2.8 – 2.9 in textbook
• Answer Visualizing concepts 2.5, 2.6
• Answer Exercises 2.57-2.67 (red only); 2.67 –
2.72 (all)

Chapter 2                   53
Atoms, Molecules, and Ions

End of Chapter 2

Chapter 2                54

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