Chemistry 1A - Foothill College
Chapter 2: Atoms, Molecules and Ions
The topics in this chapter should be review from a previous
course. It is expected that you are able to review and master this
material quickly and somewhat independently.
From this Chapter you should:
• Understand the atomic theory of matter
• Understand and be able to discuss/describe the history of the
discovery of subatomic particles
• Understand the structure of the atom, atomic number and mass
number
• Understand atomic mass, isotopes and be able to calculate average
atomic mass.
• Become more familiar with the Periodic table: Families (columns)
• Understand compound formulas: molecular, empirical, structural
• Understand the difference between molecular and ionic compounds
• Become proficient in inorganic nomenclature: names and formulas of
compounds (LAB)
• Be able to name and write formulas for a few types of simple organic
compounds.
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Dalton’s Atomic Theory (1803-1806) Postulates
1. Each element is composed of extremely small, indivisible particles
called atoms.
2. All atoms of a given element are identical to one another in mass and
other properties, but the atoms of one element are different from the
atoms of all other elements.
3. 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.
4. Compounds are formed when atoms of more than one element
combine; a given compound always has the same relative number
and kind of atoms.
•
Which of these postulated is consistent with the Law of Conservation
of Mass?
•
Which of these postulated is consistent with the Law of Definite
Proportions?
•
Which of these postulates have since been modified?
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History of Modern Atomic Theory
In the late 19th and early 20th century, a series of experimentation
allowed scientists to establish a model of the atom that remains
today the foundation of modern atomic theory.
The Discovery of the Electron-J.J. Thomson
Cathode ray tubes (CRT) - mid 1800’s - electron beam
Thompson concluded that cathode rays are streams of negatively
charged particles based upon their behavior. He also measured
the charge/mass ratio of the electron to be 1.76 × 108 C/g in
1897. The exact mass of the electron was still unknown. (C is the
symbol for coulomb, the SI unit for electric charge.)
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History of Modern Atomic Theory
Millikan Oil Drop Experiment (1909)
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.
Robert Millikan (University of
Chicago) determined the
charge on the electron in
1909.
Electron Charge: 1.602x10-19 C
Electron Mass = ? (calculate it)
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History of Modern Atomic Theory
Radioactivity-evidence that the atom is divisible!
•
•
•
•
The spontaneous emission of radiation by an atom.
First observed by Henri Becquerel (1896).
Also studied by Marie and Pierre Curie.
Three types of radiation were identified by Ernest Rutherford:
– α particles (He nucleus, + charge and heavy)
– β particles (electron, - charge and light)
– γ rays (high energy electromagnetic radiation: no charge, no mass)
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Discovery of the Nucleus
In 1910, Ernest Rutherford
shot beam of α particles at a
thin sheet of gold foil and
observed the pattern of scatter
of the particles.
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Rutherford is credited with the “discovery” of the
nuclear atom.
• Based on the alpha particle scatter pattern, Rutherford
postulated in 1911 that the atom contains a very small,
dense nucleus with the electrons surrounding the
nucleus.
• Most of the volume of the atom is empty space.
• Protons were later “discovered” by Rutherford in 1919.
(They were first detected by Eugen Goldstein in 1886, emitted
in the opposite direction compared to electrons from a CRT.)
• Neutrons were finally discovered by James Chadwick in
1932.
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Model of the Atom
•
•
•
•
Electrons (-), e–
Protons (+), p+
Neutrons (0), n0
Neutral atoms: number
of protons=number of
electrons
Summary of Subatomic Particles
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Mass Spectrometry: The Discovery of Isotopes!
• A mass spectrometer measures the relative mass and
abundance of an element’s isotopes.
Mass Spectrum of Cl
The two isotopes of Cl are clearly defined.
Determine their approximate percent abundance.
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Some Examples of Isotopic Abundances and Masses
Determined Using Mass Spectrometry
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Atomic Number, Mass Number and Isotopic
Notation
• Atomic Number, Z: The number of protons in the nucleus of
an element. Each specific element has a unique atomic
number.
• Mass number, A: The number of protons + neutrons in the
nucleus of an element. Mass number is NOT unique.
Isotopes
• Atoms of the same element with different masses.
• Isotopes have different numbers of neutrons.
11
6
C
12
6
C
13
6
C
14
6
C
Isotopic notation : AZ X
carbon -12 : 126 C
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Atomic Mass
• Mass values on the periodic table are relative to carbon-12.
The mass values are given as amu (atomic mass units).
• 1 amu = 1/12 the mass of an atom of carbon-12 (defined).
• We convert to absolute units (g, kg, lbs. etc.) by using the
conversions:
1 amu = 1.661x10-24 g
1 g = 6.02214x1023 amu (look familiar?)
• The atomic masses listed in the periodic table are average
masses. These are weighted averages based upon the
naturally occurring isotopic abundances.
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Calculating Atomic Masses
• Textbook Problem 2.35: Naturally occurring magnesium has
the following isotopic abundances.
(a)What is the average atomic mass of magnesium?
(b)Sketch the mass spectrum of magnesium.
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The Periodic Table
• Mendeleev is given
credit for first proposal of
the Periodic Table,
published in 1896.
• Groups or families;
these are the columns of
the periodic table and
contain elements with
similar properties.
• Rows; these are called
periods.
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Navigating the Periodic Table
• Make sure you can locate and identify the various classifications
of the elements.
•
•
•
•
•
Alkali metals: Group 1A
Alkaline earth metals: Group 2A
Halogens: Group 7A
Nobel gases : Group 8A
The 7 diatomic elements MUST
be memorized!
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Representing Molecular Compounds
•
Molecule: smallest identifiable unit of a pure
COVALENT (not ionic) compound.
Molecules are primarily formed from the nonmetal
elements.
Molecular compounds can be gases, liquids or solids
at room temperature and pressure.
•
There are various ways to represent molecules
structurally:
Perspective drawings show the three-dimensional
arrangement of atoms.
Structural formulas show the order in which atoms
are bonded.
Example:
Condensed formulas (condensed structural) group
atoms that are bonded together.
Example:
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Types of Chemical Formulas
•
Molecular formulas give the actual number of atoms of each
element in one molecule of molecular a compound. (Used only for
covalent compounds.)
Example:
Different compounds can have the same molecular formula with a
different arrangement of the atoms.
•
Empirical formulas give the lowest whole-number ratio of atoms of
each element in a compound.
Example:
Compounds with different molecular formulas can have the same
empirical formula.
Empirical formulas are always used for ionic compounds. Do you
know why?
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Molecular Compounds-More Examples
• Let’s write the structural, condensed, molecular and empirical
formulas for acetic acid.
• Let’s compare ethanol (b.p. = 78.5°C) with dimethyl ether
(b.p. = –23.6°C).
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Molecular Models
• Colors are often used to represent specific atoms.
• What is the molecular formula of cysteine?
Cysteine
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Ionic Compounds: Cations + Anions
CATIONS are formed by the loss of ELECTRONS
• Cations are formed when an atom LOSES electrons (oxidation).
lost e– (to anion)
• Cations are (+) charged ions.
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Common monatomic CATIONS formed from the
group 1A, 2A and 3A METALS
•
•
The monatomic cations
are all metals.
They lose valence shell
electrons = column #.
– Group 1A :
lose ? e–
– Group 2A:
lose ? e–
– Metals of group 3A:
lose ? e–
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Ionic Compounds: Cations + Anions
ANIONS are formed by the GAIN of ELECTRONS
• Anions are formed when an atom GAINS electrons (reduction).
gained e–
(from cation)
• Anions are (–) charged ions.
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Common monatomic ANIONS are formed from
the NONMETAL elements
•
•
The monatomic anions are all
nonmetals.
They gain electrons to fill their
outer valence shell (octet rule).
– Group 4A nonmetal:
gain ? e –
– Group 5A nonmetals:
gain ? e –
– Group 6A nonmetals:
gain ? e –
– Group 7A:
gain ? e –
– H when bonded to a metal:
gains ? e–
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Common ions YOU SHOULD MEMORIZE!
Formulas of ionic compounds are empirical formulas, they reflect
the smallest whole number ratio of cations to anions that results
in a net charge of zero.
Examples:
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Coulomb’s Law for Charged Particles: IONIC
Attractions!
• The force of attraction between two charged particles is given
by Coulomb’s Law: (Electrostatic force)
Force = k
•
•
•
•
(n+e)(n!e)
d2
k is a constant
n+ and n- are the magnitude of the (+) and (–) charges
e is the charge of an electron
d is the distance between the atoms
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Coulomb’s Law for Charged Particles: IONIC
Attractions!
• The electrostatic forces that hold cations and anions together
in an extended array (a lattice) are STRONG!
• As a result of these strong attractive forces, ALL ionic
compounds are solids at room temperature. (Ionic
Compounds have high melting points.)
• Solid ionic compounds DO NOT conduct electricity.
• In the molten state ionic compounds DO conduct electricity.
• Water solutions containing dissolved ionic compounds DO
conduct electricity.
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Polyatomic Ions
• Polyatomic ions consist of more than
one atom.
• Polyatomic ions are usually groups of
NONMETAL elements covalently
bonded together. This group as a
whole has a net charge (+ or –).
The naming of ionic and molecular
compounds will be covered in lab.
However, you are expected to know
the names for the lecture exams!
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Chemical Nomenclature: Names and Formulas
•
•
Inorganic Compounds (Done in LAB, should be review for you!)
Some simple Organic (carbon based) Compounds:
Alkanes: contain only C and H, all bonds are single
Alcohols: contain the –OH functional* group, covalently bonded to
a carbon atom
Carboxylic acids: contain the –COOH functional group
*Functional Group: An atom or group of atoms that imparts characteristic
chemical properties to an organic compound.
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Questions and Problems
Complete the following problems from the textbook:
2.1
A charged particle is caused to move between two electrically
charged plates, as shown below.
a)
Why does the path of the charged particle bend?
b)
c)
What is the sign of the electrical charge on the particle?
As the charge on the plates is increased, would you expect the bending to
increase, decrease, or stay the same?
As the mass of the particle is increased while the speed of the particles
remains the same, would you expect the bending to increase, decrease, or
stay the same?
d)
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Questions and Problems
2.4
Does the following drawing represent a neutral atom or an ion?
Write its complete chemical symbol including mass number, atomic
number, and net charge (if any).
2.5
Which of the following diagrams is most likely to represent an ionic
compound, and which a molecular one? Explain your choice.
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Questions and Problems
2.6
Write the chemical formula for the following compound. Is the compound
ionic or molecular? Name the compound.
2.22
a) Which two of the following are isotopes of the same element:
b) What is the identity of the element whose isotopes you have selected?
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Questions and Problems
2.49
Fill in the gaps in the following table:
2.53
Using the periodic table to guide you, predict the chemical formula and
name of the compound formed by the following elements:
(a) Ga and F, (b) Li and H, (c) Al and I, (d) K and S.
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Questions and Problems
2.75
a) What is a hydrocarbon?
b) Butane is the alkane with a chain of four carbon atoms. Write a
structural formula for this compound, and determine its molecular and
empirical formulas.
2.77
a) What is a functional group?
b) What functional group characterizes an alcohol?
c) With reference to Exercise 2.75, write a structural formula for 1-butanol,
the alcohol derived from butane, by making a substitution on one of the
end carbon atoms.
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Questions and Problems
2.84 A cube of gold that is 1.00 cm on a side has a mass of 19.3 g. A
single gold atom has a mass of 197.0 amu.
(a) How many gold atoms are in the cube?
(b) From the information given, estimate the diameter in Å of a single
gold atom.
(c) What assumptions did you make in arriving at your answer for
part (b)?
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Questions and Problems
2.86 (a) Assuming the dimensions of the nucleus and atom shown in
Figure 2.12 (Shown on slide 7 of these notes), what fraction of the
volume of the atom is taken up by the nucleus?
(b) Using the mass of the proton of 1.67 x 10–24 g and assuming its
diameter is 1.0 x 10-15 m, calculate the density of a proton in g/cm3.
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Questions and Problems
2.93 There are two different isotopes of bromine atoms. Under normal
conditions, elemental bromine consists of two atoms (a diatomic
molecule) and the mass of a Br2 molecule is the sum of the masses
of the two atoms in the molecule. The mass spectrum of Br2 consists
of three peaks:
Mass (amu)
Relative Peak Size
157.836
0.2569
159.834
0.4999
161.832
0.2431
(a) What is the origin of each peak (of what isotopes does each
consist)?
(b) What is the mass of each isotope?
(c) Determine the average molecular mass of a Br2 molecule.
(d) Determine the average atomic mass of a bromine atom.
(e) Calculate the abundances of the two isotopes.
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