Download The Structure of the Atom 1 Philosophers And Early Scientists

The Structure of the Atom
Section 4.1 Early Theories of Matter
Early Greeks believed in 4
fundamental substances:
Early Greeks believed
matter could
continuously be divided
into smaller and smaller
particles
Philosophers
And Early Scientists
fire, earth, water, and air.
Democritus (460-370 B.C.)





First person to propose the idea
that matter was not infinitely
divisible.
Matter is composed of empty
space that small particles move
through
Matter made up of tiny individual
particles – atomos (Greek for
atom)
o Atoms are solid; cannot be
created, destroyed, or
further divided
o Various sizes and shapes of
atoms
o Different properties of
matter are due to size,
shape, and movement of
atoms – different atom for
each substance
Changes in matter result from
changes in the groupings of
atoms, not changes in the atoms
themselves
General ideas, not science – not
supported experimentally
1
Aristotle (384-322 B.C.)



Rejected Democritus’ atomic
theory (didn’t agree with his own
ideas on nature)
Didn’t believe in the ―nothingness
of space‖
Because Democritus could not
experimentally support his theory
and prove what held atoms
together, Aristotle’s reputation
killed the idea of the atom for
1600 years!
Next 2000 years of chemistry history
is dominated by alchemy –
pseudoscience that was obsessed
with turning cheap metals into gold.
John Dalton (1766-1844)


Beginning of development of modern
atomic theory
Revised Democritus’ ideas based
upon scientific research
Dalton’s Atomic Theory
Isaac Newton (1642 – 1727) &
Robert Boyle (1627 – 1691)


Renewed support of Democritus
Still no experiments or predictions
Boyle



First ―chemist‖ to perform quantitative
experiments
Measured relationship between
pressure and volume of air
Substance is an element if it can’t be
broken down into 2+ simpler substances






All matter is composed of extremely
small particles called atoms
All atoms of a given element are
identical – same size, mass, chemical
properties
Atoms of one element are different
from atoms of another element
Atoms cannot be created, divided
into smaller particles or destroyed
Different atoms combine in wholenumber ratios to form compounds
In a chemical reaction, atoms are
separated, combined, or rearranged
The Structure of the Atom
Dalton and the Law of Conservation of Mass


Chemical reactions merely separate,
combine, or rearrange atoms
Experimental evidence and explanation of
the composition of compounds (with the
conservation of mass) led to the
acceptance of Dalton’s atomic theory
Atom (definition):
The smallest particle of an
element that retains all the
properties of that element
2
Problems with Dalton’s atomic theory


Atoms are indivisible – we now know
that there are subatomic particles that
make up atoms
All atoms of the same element are
identical – some atoms of an element
may have slightly different masses
(isotopes)
Atoms are incredibly small, but can be seen with a
scanning tunneling microscope (STM). Using the
STM, scientists can:


Move individual atoms around to form shapes,
patterns and simple machines
Nanotechnology – field of manipulating atoms – to
manufacture machines the size of molecules.
Section 4.2: Subatomic Particles and
the Nuclear Atom
History of discovery


Due to the presence of static electricity,
scientists in the 1800s began to look for a
connection between matter and electric
charge
o How would electricity behave in the
absence of matter (in a vacuum with
the air/matter removed)?
William Crookes’ experiment with a cathode
tube – he saw a flash of light produced by
some form of radiation strike a coating at the
opposite end of the tube.
o Because of the experiment, these
became known as cathode rays.
http://micro.magnet.fsu.edu/electroma
g/java/crookestube/

Impact of Thomson’s work


Particle smaller than an atom
o Atoms are divisible into smaller
subatomic particles
Dalton’s theory was wrong
o Not believed initially, but further
experiments proved Thomson right.
Lessons from the cathode ray tubes



Cathode rays are streams of
charged particles
Particles have a negative charge,
but the exact value of the charge
is unknown
Found in all forms of matter
o Now known as electrons
J.J. Thomson


Because experiments to measure
the particle’s mass directly had
failed, Thomson worked to
compare the ratio of the cathode
ray charge to its mass
Determined that the mass of the
charged particle was less than a
hydrogen atom (the lightest atom
known)
Thomson discovered the electron in
the late 1890s!
The Structure of the Atom
3
Next step…
Questions about the atom:
Robert Millikan determined the charge of an
electron (1909).
1. How can the atom be
electrically neutral, if it is
known to have a negative
charge?
2. Mass of an electron is
miniscule, what accounts for
the rest of the mass in a
typical atom?


Equivalent to a single unit of negative charge
Mass of an electron = 9.1 x 10-28 g (1/1840 of
the mass of a hydrogen atom)
Thomson’s atomic model

―Plum pudding model‖ (for us, it would
be known as the chocolate chip cookie
dough model).
o Spherical model
o Chocolate chips are the electrons
o Cookie dough is the uniformly
distributed positive charge
Plum pudding
model didn’t
last long!
Ernest Rutherford (1911)




Completing the atom


1920 Rutherford refined his concept of the
nucleus
o Nucleus contains positively charged
particles called protons
 Proton subatomic particle
with a charge equal to but
opposite of electron
 +1 charge
1932 Chadwick showed that the nucleus
contains another subatomic particle, a
neutral particle called the neutron
o Mass of neutron is nearly equal to
that of the proton
o No electrical charge

Gold foil experiment
Plum pudding model is incorrect
Atom is mostly empty space
Center of the atom is a dense
region (he called it the nucleus),
that contains most of the atom’s
mass and all of its positive charge
Beginning of the modern model of
the atom
Rutherford’s new nuclear atomic model

Electrons move rapidly through
empty space and are held by the
attraction of a positively-charged
nucleus
Atom:
Electrically
neutral
Spherical
Protons (+) =
Electrons (―)
Nucleus
(+ charged)
Neutrons = no
charge
Electrons (―)
surround nuc.
http://www.mhhe.com/physsci/chemist
ry/essentialchemistry/flash/ruther14.sw
f
The Structure of the Atom
4
Modern Atomic Theory:



Atoms are composed of a nucleus (containing positively charged protons and neutrally
charged neutrons) surrounded by negatively charged electrons)
Protons and electrons are identical in number in a neutral atom.
Nucleus of the atom:
o Positively charged
o Composed of protons and neutrons
o 99.97% of the mass of the atom
Section 4.3: How Atoms Differ
Henry Moseley (1887 – 1915)

Discovered that atoms of each element contain
a unique positive charge in their nuclei
o Number of protons in an atom identifies
that atom as being a particular element
o # proton = atomic number
o Atomic number determines the element’s
position on the periodic table
o Moseley reordered Mendeleev’s periodic
table
o Periodic table is organized left-to-right
and top-to-bottom by increasing atomic
number
o All atoms are neutral, so the # of protons
equals the #s of electrons
chemical name
1
Atomic number
H
Chemical symbol
1.008
Average atomic
mass
Isotopes and mass number
Dalton’s errors: atom is indivisible,
all atoms of an element are identical

All atoms of an element have the
same # of protons (and electrons)

# of neutrons in nuclei may differ
o Potassium (20 neutrons, 21
neutrons, 22 neutrons)

Solving for number of neutrons:
Atoms with the same number of
protons but different numbers of
neutrons are called isotopes
o Most elements in nature
exist as a mixture of
isotopes
Mass number is the sum of the number of protons
and neutrons in an element
# N0 (# of neutrons) = atomic mass # — atomic #
Example: K-39 = 39-19 = 20 N
Hydrogen
o All react the same
chemically and physically
0

Isotopes differ in mass: more
neutrons = more mass

Identify isotopes by their mass
number, ex. potassium-39,
potassium-40, potassium-41.
The Structure of the Atom
5
Naming and writing isotopes
Mass of individual atoms
Chemists often write out isotopes using a shortened
type of notation using only the chemical symbol,
atomic number, and mass number.
Mass of neutrons and protons are
~ 1.67 x 10-24g
40
41
K
K
19
19
Mass number is written as a superscript to the left
of the chemical symbol
Atomic number is written as a subscript to the left
of the chemical symbol
Atomic mass

Atomic mass is the weighted average of the
mass of all the isotopes of an element

Most of the mass is in the nucleus
Mass of electron: ~ 9.1 x 10-28g
Small numbers are difficult to work
with.
A new measuring unit: mass of an
atom is relative to the mass of a
atomic standard, which is carbon12.
Carbon-12 atoms have an atomic
mass of exactly 12 atomic mass
units.
Calculating atomic mass
1. Mass of each isotope multiplied by the
percentage of the time it occurs naturally

1 atomic mass unit (amu) =
1/12 the mass of a carbon-12
atom

1 amu is approximately
equal to the mass of a single
proton or neutron
2. Add the products derived for each isotope
Analyzing atomic mass:
Example:
Looking at the atomic mass of an
element will help you determine
what the most abundant isotope
may be.
Cl-35 = 76% (.76 x 35) = 26.6
Cl-37 = 24% (.24 x 37) = 8.9
Atomic mass of Cl
35.5 amu
Ex.
Exception to this logic:
Bromine has an atomic mass of 79.904 amu,
leading you to expect that the most abundant
isotope is bromine-80. Not so. Bromine has 2
isotopes (bromine-79 (50.69%) and bromine-81
(49.31%), giving a weighted atomic mass of
approximately 80 amu.

Fluorine (F) has an atomic mass
close to 19 (18.998) amu.

Fluorine has an atomic number
of 9, so it has 10 neutrons.

If fluorine had several fairly
abundant isotopes, it would be
unlikely that its atomic mass
would be so close to a whole
number.
The Structure of the Atom
6
Section 4.4: Unstable Nuclei and Radioactive Decay
Radioactivity
Types of radiation
Chemical reaction: involves the change of
one or more substances into new
substances.
Alpha (α) radiation


Atoms may be rearranged, but their
identities do not change
Chemical reactions only involve an
atom’s electrons – the nucleus is not
changed




Each alpha particle contains 2
protons and 2 neutrons
(2 P+, 2 N0)
Has a 2+ charge
Equivalent to a helium-4 nucleus
Represented by
Reactions that change an atom of one
element into another element
Nuclear reactions: reactions that involve a
change in an atom’s nucleus



These reactions spontaneously emit
radiation – (radioactivity)
Changes the identity of the element
o Atoms of one element change
into atoms of a different element
Spontaneously emit radiation because
their nuclei are unstable
o Unstable systems gain stability by
losing energy
o Unstable nuclei lose energy by
emitting radiation – called
radioactive decay
o Unstable nuclei will undergo
radioactive decay until they form
a stable nonradioactive atom,
usually of a different element


Beta (β) radiation









Nuclear Equation

Equations that account for the changing
of atomic number and atomic masses
during radioactive decay.
During alpha decay, mass number
decreases by 4 and atomic number
decreases by 2
Since opposites attract, the positively
charged alpha particle is attracted to
the negatively charged plate shown
below
Radiation is deflected toward the
positive plate (see below)
Fast moving electrons, so has a 1—
charge
Represented by
Gamma (γ) radiation or gamma rays
High-energy radiation that possesses
no mass
Usually accompanies the other 2
types of radiation
Accounts for the most energy lost
during the radioactive decay process
Does not result in the formation of a
new atom
Represented by
The Structure of the Atom
7
Nuclear Equations

Representation of radioactive decay
Sample problems (pages 106-107)
1.
2.
3.
Nuclear Stability

Primary factor in determining nuclear
stability is the ratio of neutrons to
protons

Nuclei with too many or too few neutrons
are unstable

Unstable nuclei lose energy through
radioactive decay and become stable with
the appropriate ratio of neutrons to
protons

Few radioactive atoms are found in
nature – most have already decayed into
stable atoms.
Strong force

Force that holds N0 and P+ together
in the nucleus

Only effective over very short
distances (≈10-3 pm)

100,000 times stronger than
gravity