Transcript Chemistry: Matter and Change

Section 4.1 Early Ideas About Matter
• Compare and contrast the atomic models of
Democritus, Aristotle, and Dalton.
• Understand how Dalton's theory explains the
conservation of mass.
theory: an explanation supported by many
experiments; is still subject to new experimental
data, can be modified, and is considered
successful if it can be used to make predictions
that are true
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.
• How do we know atoms even exist?
– Atoms are too small to see
– Indirect Evidence
– For example, does wind exist?
Greek Philosophers (cont.)
• 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.
Greek Philosophers (cont.)
• 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.
Moving past the Greek Philosophers
• John Dalton revived the idea of the atom in
the early 1800s based on numerous
chemical reactions.
• Dalton’s _____________________easily
explained conservation of mass in a reaction
as the result of the combination, separation,
or rearrangement of atoms.
• The invention of the chemical balance
allowed for Dalton’s theories.
Dalton’s Atomic Theory
1. Matter is composed of extremely small particles
called atoms.
2. Atoms are indivisible and indestructible
3. Atoms of a given element are identical in size,
mass and chemical properties.
4. Atoms of a specific element are different
from those of another element.
5. Different atoms combine in simple wholenumber ratios to form compounds
6. In a chemical reaction, atoms are
separated, combined, or rearranged.
Daltons Theory needed small revisions over time
• Atoms can be
divisible (rule 2)
– Protons, neutrons,
and electrons
• Atoms of the same
element can have
different masses (rule3)
– Isotopes – different forms
of the same elements
Carbon-13,
Uranium -235
Section 4.2 Defining the Atom
• Define atom.
• Distinguish between the subatomic particles in
terms of relative charge and mass.
• Describe the structure of the atom, including the
locations of the subatomic particles.
model: a visual, verbal, and/or mathematical
explanation of data collected from many
experiments
Section 4.2 Defining the Atom (cont.)
atom
cathode ray
electron
nucleus
proton
neutron
An atom is made of a nucleus containing
protons and neutrons; electrons move
around the nucleus.
The Atom
• The smallest particle of an element that
retains the properties of the element is
called an ____________.
• An instrument called the scanning tunneling
microscope (STM) allows individual atoms to
be seen.
The Atom
Use of the cathode ray tube
Different gases at very low pressure, with an
electric current running through
Results
1)different colors produced by different gases
2)paddle wheel moves in ray = Mass
3)deflected away from (-) magnet = (-)
charge
The Cathode Ray Tube
• When an electric charge is applied, a ray of
radiation travels from the cathode to the
anode, called a ___________________.
• Cathode rays are a stream of particles
carrying a negative charge.
• The particles carrying a negative charge are
known as _______________.
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
Could calculate charge / mass ratio of
electron
Ratio says electron smaller than
Hydrogen, the smallest element
• Thomson received the Nobel Prize in 1906
for identifying the first subatomic particle—the
electron
J.J. Thomson
• Experiment
– Cathode Ray Tube
• Outcomes
– Proof of subatomic
particles
– Evidence of (-) charge
– Charge / mass ratio
– Ratio was same for all
metals in electrode
and all gases in tube
Actual Charge and Mass of Electron
• American physicist Robert Millikan, 1909
• Millikan’s Oil drop experiment
• Very small mass: 9.109 x 10-28 g
• 1/1840 the mass of a hydrogen atom
Robert Millikan
• Experiment
– Oil Drop Experiment
• Outcome
– Electric charge from
an electron
Millikan’s Oil Drop Experiment
The Electron (cont.)
• 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
Questions left unanswered
• Atoms are neutral, so there must have a positive to
balance electrons
• Because electrons are so much smaller in mass than
atoms, some additional particles must be present to
account for mass
• How are particles arranged in atom
– J.J. Tomson proposed a plum pudding model
JJ. Thomson’s Plum Pudding
(choc. chip) model of the atom
Ernest Rutherford
• Experiment
– Gold Foil Experiment
• Outcome
– Nucleus is discovered
– Evidence shows atoms
are mostly empty space
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.
Rutherford’s quote
• "It was as if you fired a 15-inch shell at a sheet of tissue paper
and it came back to hit you."
Finding’s of foil experiment
• Rutherford waited 2 years before proposing his
answer
• ______________________ is positively charged,
very dense, central portion of atom that contains
nearly all of the mass of the atom
• 2 types of particles in the nucleus
– Protons, 1.673 x 10-24 g
– Neutrons, 1.675 x 10-24 g, discovered by James Chadwick in
1932
– 1,836 times larger than the mass of an electron
– 99.95% of the mass of hydrogen-1.
JJ Thomson’s Plum Pudding Model vs.
Rutherford’s nucleus
The Nucleus (cont.)
• James Chadwick received the Nobel Prize in
1935 for discovering _______________,
which are neutral particles in the nucleus
which accounts for the remainder of an
atom’s mass.
• Why were they discovered last?
• Subatomic particles:
Electron, e-
Proton, p+
Neutron, n0
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.
The Nucleus (cont.)
• Scientists have determined that protons
and neutrons are composed of subatomic
particles called quarks.
• Read about the Higgs Boson particle
The Nucleus (cont.)
Chemical behavior of atoms can be
explained by considering only an atom's
___________________.
The vast majority of chemistry deals with the
interactions of electrons.
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.
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, mass number, and
atomic mass
• Atomic number
– # of protons
– If neutral, it is also the #
of electrons
• Mass number
– # of protons and
neutrons
• Mass number is used to
calculate the atomic mass of
the isotope
• Since electrons are so much
smaller than protons and
neutrons, they are not a major
factor in atomic mass
Isotopes and Mass Number
• All atoms of the same element have the
same number of protons 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
• 2 ways of writing isotopes
• Carbon -13
•
13
6
C
Isotopes of Hydrogen
• Protium, Hydrogen-1
– 99.985% of all natural hydrogen
• Deuterium, Hydrogen-2
– 0.015% of all natural hydrogen
• Tritium, Hydrogen-3
– Radioactive
– Spiderman 2
Isotopes and Mass Number (cont.)
The relative abundance of each isotope is
usually constant.
Relative abundance = How much of
each isotope is present
• Isotopes containing more neutrons have a
greater mass.
• Isotopes have the same chemical behavior.
• The ____________________is the sum of
the protons and neutrons in the nucleus.
Isotopes and Mass Number (cont.)
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 or one neutron.
Relative Atomic Mass
• Atomic mass is the mass of an atom expressed in
a.m.u
• Hydrogen-1 has atomic mass of 1.007825 u
– Very close to mass number
• Oxygen-16 has atomic mass of 15.994915 u
– Very close to mass number
Mass of Atoms (cont.)
• The _________________________of an element
is the weighted average mass of the isotopes of
that element.
Weighted average of the isotopes of an element
Copper-63, 69.17%, 62.939 a.m.u
Copper-65, 30.83%, 64.927 a.m.u
• Atomic Mass vs Molar Mass
– Both are the weighted average of the isotopes of that
element.
– 1 atom of carbon weighs 12.011 a.m.u - on
weighted average
– 1 mole of carbon atoms weigh 12.011 g - on
weighted average