Ch. 1: Atoms
Dr. Namphol Sinkaset
Chem 200: General Chemistry I
I. Chapter Outline
I.
II.
III.
IV.
V.
VI.
VII.
Introduction
Particulate View of the World
The Scientific Approach
Measurement in Science
History of the Atom
Subatomic Particles
Atomic Mass
I. Real-Life Legos®
• Everything is comprised of small parts
connected into a complex whole.
• The structure of the whole determines
its properties.
II. Particles
•
We will approach chemistry with two
key principles in mind:
1. Matter is particulate.
2. Structure of particles determines
properties of matter.
• Chemistry seeks to understand
properties of matter by studying the
structure of particles that compose it.
II. Matter
• Matter is anything that occupies space
and has mass.
• Everything around you is composed of
matter – desk, book, air.
• Remember: matter is particulate.
II. Atoms and Molecules
• Atoms are the basic particles that
compose ordinary matter.
• Atoms can bind to one another in
specific arrangements to yield
molecules.
• For example, a water molecule is
comprised of 1 oxygen atom and 2
hydrogen atoms.
II. Structure and Properties
• Boils at 30 °C
• Feels like gasoline
• Doesn’t dissolve salt
• Boils at 100 °C
• Feels like water
• Dissolves salt
II. Classifying Matter
• Any sample of matter is called a
substance.
• Matter can be classified by state or by
composition.
• State determined by relative positions
and interactions of particles.
• Composition determined by types of
particles.
II. States of Matter
II. States of Matter
• solid: strong particle attractions, pack in fixed
locations, only vibrate in place, not
compressible
• liquid: slightly weaker particle attractions,
pack in non-fixed locations, fixed volume,
assume shape of container
• gas: weak particle attractions, free to move,
large distances between particles,
compressible
II. Composition of Matter
• Can also classify matter by the kinds of
particles out of which it is comprised.
• If there is only one type of particle, then
it is a pure substance.
• If there is more than one type of
particle, then it is a mixture.
II. Types of Matter
III. The Scientific Approach
• a.k.a. scientific method, is a flexible process
of creative thinking and testing aimed at an
objective
III. Differences Between
Hypothesis and Theory
•
•
•
•
Hypothesis not thoroughly tested
Theory more “developed”
Hypothesis does not predict
Experiments on hypothesis test
hypothesis itself
• Experiments on theory test predictions
of theory
• Theory can be expanded to many
related situations
III. Differences Between
Hypothesis and Theory
• Compare the two statements below.
• “Methane reacts w/ oxygen to form
carbon dioxide and water.”
• “Hydrocarbons undergo a combustion
reaction w/ oxygen to form carbon
dioxide and water.”
IV. Measurement in Science
• Observations in the lab can be
qualitative or quantitative.
• Science is powerful because many
observations can be assigned an
accurate number.
• e.g. Hot/cold vs. 262 °C/12 °C
IV. Units
• All measured quantities have a number
and a unit!!!!
• Without a unit, a number has no
meaning in science.
• e.g. a person’s height is 6 feet, 4
inches, not 6-4.
• ANY ANSWER GIVEN W/OUT A UNIT
WILL BE GRADED HARSHLY.
IV. SI Units
• Scientists have agreed to use the Système
International d’Unités, a.k.a. SI units
IV. Derived SI Units
• Combinations of fundamental SI units
are used to describe other quantities.
• e.g. speed is distance per time, so its SI
unit is m/s
V. History of the Atom
• The Greeks were the first to wonder
about matter.
• Greek philosophers around 430 B.C.E.
debated what made up the world
around them.
• Leucippus and Democritus vs. Plato
and Aristotle
V. Atomos vs. Fire, Air, Earth,
Water
V. Revival of the Atom
• The idea of the atom was discarded and
forgotten about for almost 2000 years.
• In the late 18th and early 19th centuries,
three natural laws baffled everyone.
• John Dalton resurrected the idea of the
atom to explain what was observed.
V. Law of Mass Conservation
• In a reaction, matter
is neither created
nor destroyed.
• Credit Antoine
Lavoisier (1789).
V. Law of Mass Conservation
V. Law of Definite Proportions
• All samples of a given compound have
the same proportions of constituent
elements.
 Credit Joseph Proust (1797)
• e.g. Ammonia has 14.0 g N for every
3.0 g of H:
14.0 𝑔 𝑁
𝑀𝑎𝑠𝑠 𝑟𝑎𝑡𝑖𝑜 =
= 4.7 𝑜𝑟 4.7: 1
3.0 𝑔 𝐻
V. Law of Multiple Proportions
• In 1804, John Dalton found that when two
elements (A and B) form two different
compounds, the masses of element B that
combine with 1 g of element A can be
expressed as a ratio of small whole numbers.
V. Dalton’s Atomic Theory
• John Dalton revived the idea of the
atom to explain the natural laws that
had everyone perplexed.
• His atomic theory (1808) worked so well
that it was quickly accepted.
V. Postulates of Dalton’s Theory
1. Each element is composed of tiny,
indestructible particles called atoms.
2. All atoms of a given element have the same
mass and other properties that distinguish
them from the atoms of other elements.
3. Atoms combine in simple, whole-number
ratios to form compounds.
4. Atoms of one element cannot change into
atoms of another element. In a chemical
reaction, atoms only change the way that
they are bound together with other atoms.
V. The Nuclear Atom
• Dalton’s theory treated atoms as
permanent, indestructible building
blocks that composed everything.
• A series of experiments were conducted
that led to a new view of the atom.
V. Cathode Rays
• What conclusions about cathode rays can be
made from these experiments?
V. Cathode Rays
• Using EM fields (late 1800s), J.J.
Thomson measured the cathode ray
particle’s mass to charge ratio.
• He estimated that cathode ray particles
were about 2000 times lighter than a
hydrogen atom.
• Result implies that atoms can be
divided into smaller particles.
V. Cathode Rays
• Using his famous oil drop experiment,
Robert Millikan (1909) calculated the
charge of a cathode ray particle.
• His value is w/in 1% of today’s accepted
value: -1.602 x 10-19 C.
• Mass was determined to be
9.109 x 10-28 g.
• Of course, cathode ray particles are
now known as electrons.
V. Plum Pudding
• If electrons are in all
matter, there must
be positivelycharged species as
well.
• J.J. Thomson
proposed the plum
pudding model of
the atom.
 Electron “raisins”
 “Pudding” of positive
charge
V. The Role of Radioactivity
• Henri Becquerel and Marie Curie
discovered radioactivity by accident.
• Ernest Rutherford used radium, an
alpha (a) particle emitter.
• These a-particles are dense and have a
positive charge.
V. Rutherford’s a-Particle
Experiment (1909)
V. Conclusions from
Rutherford’s Experiment
• Most of an atom’s mass and all of its
positive charge exists in a nucleus.
• Most of an atom is empty space,
throughout which electrons are dispersed.
• By having equal numbers of protons and
electrons, an atom remains electrically
neutral.
• Note: neutrons discovered 20 years later.
V. Rise of the Nuclear Atom
VI. Subatomic Particles
• Therefore, all atoms are made up of
protons, neutrons, and electrons.
VI. Atomic Number
• The atomic number (Z) of an element
equals the # of protons in the nucleus
 All atoms of an element have same, unique
atomic number!!
• Protons are responsible for an atom’s
identity.
• e.g. All carbon atoms have 6 protons
and all uranium atoms have 92 protons.
VI. Chemical Symbols
• Each element has a unique symbol.
• The symbol is either a 1 or 2
abbreviation of its name.
• e.g. carbon  C; nitrogen  N; chlorine
 Cl; sodium  Na; gold  Au
VI. Mass Number
• The mass number (A) is the total
number of protons and neutrons in the
nucleus.
• e.g. A carbon atom with 6 neutrons has
a mass number of 12.
VI. Isotopes
• The # of protons determines the identity
of the atom, but the # of neutrons has
no effect.
• Thus, atoms of the same element can
have different mass numbers.
• Since chemical properties are mainly
due to e-, isotopes are almost identical
chemically.
• Different isotopes of an element exist in
certain percentages – natural
abundances.
VI. Depicting an Isotope
VI. Sample Problem
1. What are the atomic number, mass
number, and symbol for the carbon
isotope with 7 neutrons?
2. How many protons and neutrons are
present in an atom of potassium-39?
VI. Ions
• Atoms can lose or gain electrons and
become ions.
• Ions are charged particles.
• Positively-charged particles = cations.
 e.g. Li  Li+ + 1e-
• Negatively-charged particles = anions.
 e.g. F + 1e-  F-
VII. Atomic Mass
• Postulate #2 of Dalton’s atomic theory
stated that all atoms have the same
mass.
• With the existence of isotopes, this can’t
be true, but we can calculate an
average mass.
• The average mass of an element is
called the atomic mass.
VII. Atomic Mass
• Atomic masses are calculated using a
weighted average of all isotopes of an
element.
• The natural abundance is used to weight
each isotope in the calculation.
𝐴𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 =
𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑛 × (𝑚𝑎𝑠𝑠 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑛)
𝑛