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Section 4.1 Early Ideas About Matter (cont.)
Dalton's atomic theory
The ancient Greeks tried to explain
matter, but the scientific study of the
atom began with John Dalton in the
early 1800's.
The History of Atomic Theory
• Many ancient scholars believed matter was
composed of such things as earth, water,
air, and fire.
• Many believed
matter could be
endlessly divided
into smaller and
smaller pieces.
The History of Atomic Theory
• Democritus (460–370 B.C.) was the first
person to propose the idea that matter was
not infinitely divisible, but made up of
individual particles called atomos.
• Aristotle (484–322 B.C.) disagreed with
Democritus because he did not believe empty
space could exist.
• Aristotle’s views went unchallenged for 2,000
years until science developed methods to test
the validity of his ideas.
The History of Atomic Theory
The History of Atomic Theory
• Late 18th Century
• Lavoisier proposes the Law of
conservation of mass and Proust proposes
the Law of constant composition.
• Early 19th Century
• Using the previously unconnected ideas
above, John Dalton formulates his Atomic
Theory.
The History of Atomic Theory
• John Dalton revived the idea of the atom in
the early 1800s based on numerous
chemical reactions.
• Dalton’s atomic theory easily explained
conservation of mass in a reaction as the
result of the combination, separation, or
rearrangement of atoms.
Dalton -1803
Billard Ball Model of Atom
• Atoms are small, round, solid spheres
Atomic Theory (5 parts)
• Elements are made from tiny particles
called atoms.
• All atoms of a given element are identical.
(N.B. see isotopes below).
• The atoms of a given element are
different to those of any other element.
Dalton
Atoms of different elements combine to
form compounds. A given compound
always has the same relative numbers and
types of atoms. (Law of constant
composition).
Atoms cannot be created or destroyed in a
chemical reaction they are simply
rearranged to form new compounds. (Law
of conservation of mass).
The Atom
• The smallest particle of an element that
retains the properties of the element is
called an atom.
• To give you an idea of how small an atom is
we can look at a model. If we use an orange
to represent the size of an atom, on the same
scale, the orange would be the size of the
Earth.
• An instrument called the scanning tunneling
microscope (STM) allows individual atoms to
be seen
The Electron
• When an electric charge is applied, a ray of
radiation travels from the cathode to the
anode, called a cathode ray.
• Cathode rays are a stream of particles
carrying a negative charge.
• The particles carrying a negative charge are
known as electrons.
The Electron (cont.)
• This figure shows a typical cathode ray
tube.
The Electron (cont.)
• J.J. Thomson measured the effects of both
magnetic and electric fields on the cathode ray to
determine the charge-to-mass ratio of a charged
particle, then compared it to known values.
• The mass of the charged particle was much less
than a hydrogen atom, then the lightest known
atom.
• Thomson received the Nobel Prize in 1906 for
identifying the first subatomic particle—the
electron. This proved that Dalton was
incorrect…atoms were not the smallest
particles.
The Electron (cont.)
• In the early 1910s, Robert Millikan used the oildrop apparatus shown below to determine the
charge of an electron.
The Electron (cont.)
• Charges change in discrete amounts—
1.602  10–19 coulombs, the charge of one
electron (now equated to a single unit, 1– ).
• With the electron’s charge and charge-tomass ratio known, Millikan calculated the
mass of a single electron.
the mass of
a hydrogen
atom
The Electron (cont.)
• Matter is neutral, it has no electric charge.
• J.J. Thomson's plum pudding model of the
atom states that the atom is a uniform,
positively charged sphere containing
electrons.
The Nucleus
• In 1911, Ernest Rutherford studied how
positively charged alpha particles
interacted with solid matter.
• By aiming the particles at
a thin sheet of gold foil,
Rutherford expected the
paths of the alpha
particles to be only
slightly altered by a
collision with an electron.
The Nucleus (cont.)
• Although most of the alpha particles went
through the gold foil, a few of them
bounced back, some at large angles.
The Nucleus (cont.)
• Rutherford concluded that atoms are
mostly empty space.
• Almost all of the atom's positive charge and
almost all of its mass is contained in a dense
region in the center of the atom called the
nucleus.
• Electrons are held within the atom by their
attraction to the positively charged nucleus.
The Nucleus (cont.)
• The repulsive force between the positively
charged nucleus and positive alpha
particles caused the deflections.
The Nucleus (cont.)
• Rutherford refined the model to include
positively charged particles in the nucleus
called protons.
• James Chadwick received the Nobel Prize in
1935 for discovering the existence of
neutrons, neutral particles in the nucleus
which accounts for the remainder of an
atom’s mass.
The Nucleus (cont.)
• All atoms are made of three
fundamental subatomic
particles: the electron, the
proton, and the neutron.
• Atoms are spherically
shaped.
• Atoms are mostly empty
space, and electrons travel
around the nucleus held by
an attraction to the positively
charged nucleus.
Atomic Structure
Scientist
Experiment
J. J.
Thomson
Cathode Ray Mass/charge ratio
of the electron
determined
Millikan
Oil Drop
Experiment
Gold Foil
Experiment
Rutherford
Knowledge
Gained
Relating
to
Electron
Electron
Charge on the
electron
Nucleus present The nucleus
of an
in atom
atom and the
proton
Chadwick NA
Neutrons present
in nucleus
Neutrons
Atomic Structure
Particle
Charge
Position in
atom
+1
Mass in
atomic
mass units
1
Proton
Neutron
0
1
Nucleus
Electron
-1
1/1836
Outside of
Nucleus
Nucleus
Section 4.3 How Atoms Differ
• Explain the role of atomic number in determining the
identity of an atom.
• Define an isotope.
• Explain why atomic masses are not whole numbers.
• Calculate the number of electrons, protons, and
neutrons in an atom given its mass number and
atomic number.
Section 4.3 How Atoms Differ (cont.)
periodic table: a chart that organizes all known
elements into a grid of horizontal rows (periods)
and vertical columns (groups or families) arranged
by increasing atomic number
atomic number
atomic mass unit (amu)
isotopes
atomic mass
mass number
The number of protons and the mass
number define the type of atom.
ATOM
ATOM
NUCLEUS
NUCLEUS
ELECTRONS
ELECTRONS
PROTONS
PROTONS
NEUTRONS
NEUTRONS
POSITIVE
CHARGE
NEUTRAL
CHARGE
NEGATIVE
CHARGE
NEGATIVE CHARGE
equal in a
Atomic
Most Number
of the atom’s mass.
neutral atom
equals the # of...
QUARKS
Atomic Number
• Each element contains a unique positive
charge in their nucleus.
• The number of protons in the nucleus of an
atom identifies the element and is known as
the element’s atomic number.
Isotopes and Mass Number
• All atoms of a particular element have the
same number of protons and electrons but
the number of neutrons in the nucleus can
differ.
• Atoms with the same number of protons but
different numbers of neutrons are called
isotopes.
Isotopes and Mass Number (cont.)
• The relative abundance of each isotope is
usually constant.
• Isotopes containing more neutrons have a
greater mass.
• Isotopes have the same chemical behavior.
• The mass number is the sum of the protons
and neutrons in the nucleus.
Atomic & Mass Numbers
• Mass Number = protons + neutrons
• Atomic Number = protons only
• # of Neutrons = mass # - atomic #
• Mass Number
– always a whole
number
– NOT on the
Periodic Table!
© Addison-Wesley Publishing Company, Inc.
• Nuclear symbol:
Mass #
Atomic #
12
6
• Hyphen notation: carbon-12
C
4.3 Mass Number
Mass number is the number of protons and
neutrons in the nucleus of an isotope.
Mass # = p+ + n0
Nuclide
Oxygen - 16
Arsenic-
75
Phosphorus - 31
p+
n0
e-
8
8
8
33
42
33
16
75
15
16
15
31
Mass #
• Chlorine-37
– atomic #:
17
– mass #:
37
– # of protons:
17
– # of electrons:
17
– # of neutrons:
20
37
17
Cl
Mass of Atoms
• One atomic mass unit (amu) is defined as
1/12th the mass of a carbon-12 atom.
• One amu is nearly, but not exactly, equal to
one proton and one neutron.
•
12C
atom = 1.992 × 10-23 g
Mass of Atoms (cont.)
• The atomic mass of an element is the
weighted average mass of the isotopes of
that element.
Average Atomic Mass
• All periodic tables have atomic mass
numbers that are not integers.
• For example, the atomic mass of Cl is
often quoted on periodic tables as 35.5,
and may be represented thus, 35.5Cl17. This
does not mean that there are 17 protons,
17 electrons and 18.5 neutrons in an atom
of chlorine. It is not possible to have a
fraction of a neutron in an atom.
Average Atomic Mass
• The non-integer values mean that there is
more than one isotope of chlorine that
exists in nature, in the case of chlorine,
35Cl and 37Cl. A quick calculation shows
that these two species have the same
number of protons and electrons, but
different, whole numbers of neutrons (18
and 20 respectively). That is, they are
isotopes of one another. These isotopes
happen to exist naturally in the approx.
abundance, 35Cl, 75 % and 37Cl, 25 %.
Average Atomic Mass
• EX 1: Calculate the avg. atomic mass of
oxygen if its abundance in nature is
99.76% 16O, 0.04% 17O, and 0.20% 18O.
Avg.
Atomic  (16x99.76)  (17x0.04)  (18x0.20) 
100
Mass
16.00
amu
• EX 2: Find chlorine’s average atomic mass if
approximately 8 of every 10 atoms are
chlorine-35 and 2 are chlorine-37.
Avg.
Atomic 
Mass
(35x8)  (37x2)
 35.40 amu
10
Example 3:
• A naturally occurring sample of an
element consists of two isotopes, one
of mass 85 and one of mass 87. The
abundance of these isotopes is 71%
and 29%. Calculate the average atomic
mass of an atom of this element.
Molecules
• Molecules are formed when a definite
number of atoms are joined together by
chemical bonds.
• A molecule can consist of the atoms of
only one element, or the atoms of many
different elements but always in a fixed
proportion.
Molecules
• Molecules are usually formed between
non-metal elements. Formula show the
number of each type of atom present
written as subscripts. The lack of a
subscript means only one of that type of
atom is present.
Examples
•
•
•
•
H2O = Water = Compound
NH3 = Ammonia = Compound
N2 = Nitrogen = Element/Compound
H2 = Hydrogen = Element/Compound
Ions
• Atoms have equal numbers of protons
and electrons and consequently have no
overall charge.
• When atoms lose or gain electrons, the
proton:electron numbers are unbalanced
causing the particles to become charged.
• These charged particles are called ions.
Ions
• Positive ions (where the number of protons is
greater than the number of electrons) are called
cations,
• Negative ions (where the number of electrons is
greater than the number of protons) are called
anions.
• An ion made up of only one type of atom is
called a monatomic ion;
• An ion made up from more than one type of
atom is called a polyatomic ion.
Examples
Substance
Formula
Sodium Ion
Na+
Chloride Ion
ClCarbonate Ion
CO32Ammonium Ion
NH4+
Cation or
Anion
Description
Cation
Monatomic
Anion
Monatomic
Anion
Polyatomic
Cation
Polyatomic
Ions
• Metals tend to form cations. The metals in
groups 1, 2 and 13 tend to lose 1, 2 and 3
electrons respectively to form positive ions
with 1+, 2+ and 3+ charges respectively.
• The transition metals form cations with
various positive charges and are much
less easy to predict.
• Non-metals tend to form anions. The nonmetals in groups 17, 16 and 15 tend to
gain 1, 2 and 3 electrons respectively to
form negative ions with 1-, 2- and 3charge respectively.
Examples
• Consider sodium
(i) What is the charge on an ion of sodium?
(ii) What is the charge on a sodium nucleus?
(iii) What is the charge on a sodium atom?
(iv) What is the atomic number of sodium?
(v) How many protons in the sodium nucleus?
(vi) How many protons in the sodium atom?
Examples
• For the potassium ion, 39K with a
charge of +1;
(i) How many electrons, protons and
neutrons are present?
(ii) Which of these particles has the
smallest mass?
• An ion has a net charge of -1. It has 18
electrons and 20 neutrons.
(i) What is its isotopic symbol?
(ii) What is its atomic number?
(iii) What is its mass number?
(iv) What is the charge on the nucleus?
(v) How many protons are present?