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Chapter 4
Atoms
Table of Contents
Section 1 Development of the Atomic Theory
Section 2 The Atom
Copyright © by Holt, Rinehart and Winston. All rights reserved.
Chapter 4
Section 1 Development of the
Atomic Theory
Objectives
• Describe some of the experiments that led to the
current atomic theory.
• Compare the different models of the atom.
• Explain how the atomic theory has changed as
scientists have discovered new information about the
atom.
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Chapter 4
Section 1 Development of the
Atomic Theory
The Beginning of the Atomic Theory
• Early Greeks proposed matter comprised of:
•Fire, earth, water, air
• Democritus, a Greek philosopher
•Around 440 BCE
•thought that you could have a particle that could
not be cut.
•Called it an atomos.
• Aristotle, another Greek philosopher, disagreed
•Basis: What holds the atoms together?
•Democritus couldn’t answer the question
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Chapter 4
Section 1 Development of the
Atomic Theory
The Beginning of the Atomic Theory,
continued
• Democritus was right:
•Matter is made of particles,
•called atoms.
•atom is the smallest particle into which an element
can be divided and still be the same substance.
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2000 Years of Psuedochemistry
• Alchemy
– Most alchemists were mystics or fakes
– Obsessed with turning cheap lead into
expensive gold
– But:
• Discovered many elements
• Learned how to prepare mineral acids
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First True Chemist
• Robert Boyle (1627-1691)
– The Skeptical Chemist (1661)
• First quantitative experiments
• Defined an element if it could not be broken down
into a simpler substance
• Based on experience in metallurgy
• Boyle still thought you could turn one metal into
another
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The Problem with Combustion
• 17th & 18th Centuries dealt with issues of
‘combustion’
– Georg Stall, Germany, 1660-1734
• Suggested “phlogiston” flowed out of burning
material
• Burning in a closed container stopped because the
jar became saturated with ‘phlogiston’.
– Joseph Priestly, English, 1733-1804
• Discovered oxygen
• Proposed that oxygen was ‘dephlogisticated air’
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Combustion & Fundamental
Laws
• By late 18th century, combustion had been
studied.
– Carbon dioxide, nitrogen, oxygen had been
discovered.
– Many other element discovered.
• Antoine Lavoisier, French (1743-1794)
– Based on experiments postulated “Mass is neither
created or destroyed”
– Law of Conservation of Mass
– Proved combustion involved oxygen, not ‘phlogiston’
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What is Proust’s Law?
• Law of Definite Proportion
– A given compound always contains exactly
the same proportion of elements by mass.
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What is the Law of Multiple
Proportions?
• When 2 elements form a series of
compounds, the ratios of the masses of
the second element that combine with 1
gram of the first element can always be
reduced to small whole numbers
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Example:
• Nitrogen and Oxygen. What is the ratio of
masses of Nitrogen that combines with
oxygen:
– Compound A: 1.750g
– Compound B: 0.8750g
– Compound C: 0.4375g
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Example Answers:
• Ratios of Element 2:
– A/B = 1.750/0.875 = 2/1
– B/C = 0.875/0.4375 = 2/1
– A/C = 1.750/0.4375 = 4/1
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Chapter 4
Section 1 Development of the
Atomic Theory
Dalton’s Atomic Theory Based on Experiments
• John Dalton published his atomic theory in 1803.
• His theory stated
•all substances are made of atoms.
•Atoms are small particles that cannot be
created, divided, or destroyed.
•Atoms of the same element are exactly alike,
and atoms of different elements are different.
•Atoms join with other atoms to make new
substances.
• Not Quite Correct
• Prepared first table of atomic masses
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Keys to Atomic Formulas
• Actually published by:
– Joseph Gay-Lussac, French, 1778-1850
• Measured volumes of gasses that reacted together
– Amadeo Avogadro, Italian, 1776-1856
• Interpreted Gay-Lussac results by proposing that
at the same temperature and pressure, equal
volumes of different gasses contain the same
number of particles
• Not accepted for 50 years.
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Definition of an Atom
• The smallest particle of an element that
retains the properties of the element
• How small?
– World population in 2000 was about 6 billion
– One penny contains 5 billion times more
atoms
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Can we See Atoms
• Only with ‘scanning tunneling microscopes
• Nanotechnolgy manipulates individual
atoms to make very small devices
– Next generation of computers will have wires
one or two atoms wide (smaller, less power,
less heat)
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Pictures
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Pictures
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Chapter 4, Section 2
Sub-Atomic Particles and Nuclear
Atoms
Accidental Discoveries?
• Does anything get discovered by
accident?
• Yes
– Vulcanized rubber
– Aspartame (Nutrasweet)
– Electrons
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Self Taught Class
•
•
•
•
Who it is?
When?
What did He do?
How was it important to understanding the
Atom?
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Discovering the Electron
• Sir William Crookes, early 1800’s
– What is the relationship between electricity
and matter?
• Static from combs
• Static from carpets
• Recent inventions:
– Vacuum pump
– Cathode Ray Tube (CRT)
• Cathode (+) at one end of vacuum tube
• Anode (-) at other end
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Discovering the Electron
• Crookes was in a darkened room.
– Noticed flashes of light within his tube (coated
inside with light producing chemicals)
– Further work: “rays” going from cathode end
to anode end (hence cathode ray tube)
– Cathode Ray Tube is basis for TV and
computer monitors
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Crookes CRT
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Chapter 4
Section 1 Development of the
Atomic Theory
Thompson’s Discovery of Electrons
• Thompson experimented with a cathoderay tube.
• He discovered negatively charged
particles known as electrons.
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Chapter 4
Section 1 Development of the
Atomic Theory
Thompson’s Cathode-Ray Tube Experiment
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Thompson’s Model
• Thompson proposed
a new model of the
atom.
– electrons are mixed
throughout an atom,
like plums in a pudding
(or raisins in raisin
bread).
– Called Plum Pudding
model
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Discovering the Electron
• By late 1800’s Further work led to
conclusion that:
– Cathode Rays were actually stream of
charged particles
– Particles carried a negative charge
– These particles were found in all matter
– Particles were called ‘electrons’
• CRISIS: Dalton was wrong, Atoms did
have smaller particles
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Discovering the Electron
• 1909 – Robert Millikin (US)
– Determined charge of an electron
– Determined mass of an electron
• 9.11 X 10-28g = 1/1840 mass of a hydrogen atom
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Chapter 4
Section 1 Development of the
Atomic Theory
Rutherford’s Atomic “Shooting Gallery”
• In 1909, Ernest Rutherford aimed a beam of small,
positively charged particles at a thin sheet of gold
foil. The next slide shows his experiment.
• Surprising Results Rutherford expected the
particles to pass right through the gold in a straight
line. To Rutherford’s great surprise, some of the
particles were deflected.
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Chapter 4
Section 1 Development of the
Atomic Theory
Rutherford’s Gold-Foil Experiment
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The Nuclear Atom
• Rutherford concluded Thompson was
wrong:
– There must be a tiny, very dense region of the
atom, called the ‘nucleus’
• Must be very dense (like all the mass of an atom)
• Must have a positive charge to keep the electrons
attracted
– Between atoms and nucleus must be a lot of
empty space
• How Much?
– Nucleus the size of a quarter has electrons over 1 mile
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away
The Nuclear Atom
• Rutherford Model Explains:
– Why alpha particles (electrons) bend on their
way through nucleus
– Why some alpha particles are deflected at
very sharp angles
• Did not explain all of the Atom’s Mass
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Chapter 4
Section 1 Development of the
Atomic Theory
Where Are the Electrons?
• Far from the Nucleus
Rutherford proposed that in
the center of the atom is a
tiny, positively charged part
called the nucleus.
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Discovering Protons and Neutrons
• 1919 Rutherford Later Experiments
– Concluded nucleus must contain positive
particles called ‘protons’
– With co-worker James Chadwick showed
nucleus also contained a neutral particle
called ‘neutron’
• Mass of neutron almost same as proton
• No electrical charge
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Summary to Date
• Atoms are composed of:
– Protons (+ charge, 1 mass unit)
– Neutrons (no charge, 1 mass unit)
– Electrons (- charge, very little mass)
• Most of an atom’s size is electrons moving
through empty space
– Electrons are held to nucleus by +/- electrical
attraction
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Summary of Models
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Section 3
How Atoms Differ
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Chapter 4
Section 3 The Atom
Objectives
• Describe the size of an atom.
• Name the parts of an atom.
• State how atoms of different elements differ.
• State how isotopes differ.
• Calculate atomic masses.
• Describe the role of electrons in an atom.
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Chapter 4
Section 3 The Atom
How Small Is an Atom?
• Scientists know that aluminum is made of averagesized atoms. An aluminum atom has a diameter of
about 0.00000003 cm.
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Chapter 4
Section 3 The Atom
What Is an Atom Made Of?
• The Nucleus
•Protons are positively charged particles
•Neutrons have no electrical charge.
• Outside the Nucleus
•Electrons are the negatively charged particles in
electron clouds.
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Chapter 4
Section 3 The Atom
Parts of an Atom
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Chapter 4
Section 3 The Atom
What Is an Atom Made Of?
• The Nucleus
•positively charged particles called protons.
•Each proton has a mass of about 1 amu.
• The SI unit used to express the masses of particles
in atoms is the atomic mass unit (amu).
• Neutrons
•In nucleus that have no electrical charge.
•Neutrons have a mass of about 1 amu
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Chapter 4
Section 3 The Atom
What Is an Atom Made Of?, continued
• Outside the Nucleus
•Electrons are negatively charged particles in
atoms.
•Electrons are found around the nucleus within
electron clouds.
• The charges of protons and electrons are opposite
but equal, so their charges cancel out.
• Because an atom has no overall charge, it is
neutral.
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Chapter 4
Section 2 The Atom
How Do Atoms of Different Elements Differ?
• Starting Simply
•The hydrogen atom
has one proton and
one electron.
•The helium atom has
two protons, two
neutrons, and two
electrons.
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Chapter 4
Section 2 The Atom
How Do Atoms of Different Elements
Differ?, continued
• Building Bigger Atoms For bigger atoms, simply
add protons, neutrons, and electrons.
• Protons and Atomic Number
•atomic number = number of protons
•Atomic mass = number of protons + number of
neutrons
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Atoms
• Atoms are electrically neutral, so:
• Protons = Electrons = Atomic Number
• Neutrons does not have a specific
relationship to protons
• Atomic Mass = Protons + Neutrons
– Electrons have almost no mass
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Reading the Periodic Table
Name
Atomic Number
Symbol
Avg. Atomic Mass
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Periodic Table
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Reading the Periodic Table - Quiz
• How Many Protons are in
– Boron (B)
– Platinum (Pt)
5
78
• How many electrons are in:
– Radium (Ra) 88
– Magnesium (Mg)
12
• An element contains 66 electrons. What is
it?
Dysprosium
• An element contains 14 protons. What is it?
Silicon
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Calculating Protons & Neutrons
Element Atomic
Mass
B
Mg
O
K
Atomic
Number
Protons
Neutrons
5
5
6
5
12
12
12
12
8
8
8
20
19
11
24.305
16
39
8
19
19
Electrons
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Why are Atomic Masses not Even
Numbers?
• What is the atomic mass of Carbon (C)?
– 12.011
• What is the atomic mass of Chlorine (Cl)?
– 35.453
• If Protons = 1 and Neutrons = 1, where
does the .453 come from?
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Why not whole Numbers?
• Mass of both Neutron and Proton is
1.67x10-24
• Small units/hard to work with
• Scientists set standard based on Carbon
12
– 1 atomic mass unit = 1/12 of carbon atom
• So Silicon is 29.974 instead of 30
– This is only part of the reason…
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Chapter 4
Section 2 The Atom
Isotopes
• have the same number of protons but different
numbers of neutrons.
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Isotopes
• Another thing Dalton got wrong
• Isotopes occur as a mixture in nature
– Example Potassium:
•
•
•
•
93.25% have 20 neutrons
6.7302% have 22 neutrons
0.117% have 21 neutrons
ALL have 19 Protons and 19 Electrons
• Isotopes have the Same Atomic Number
but a Different Atomic Mass
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Chapter 4
Section 2 The Atom
Isotopes, continued
• Telling Isotopes Apart
•by its mass number.
• How Many Neutrons?
•Calculate Neutrons = Atomic Mass – Atomic number
• Properties of Isotopes
•An unstable atom has a nucleus that will change
over time.
•This type is radioactive.
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Chapter 4
Section 2 The Atom
Isotopes, continued
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Chapter 4
Section 2 The Atom
Isotopes, continued
• Naming Isotopes
•Write the name of the element followed by a
hyphen and the mass number.
•Example: C-14 is Carbon 14 (8 neutrons)
instead of 6 (C-12 is normal)
• Calculating the Mass of an Element
•The atomic mass of an element is the weighted
average of the masses of the isotopes of that
element.
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Chapter 4
Section 3 The Atom
Math Focus
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Unstable Nuclei and
Radioactive Decay
Section 4.4
Radioactivity
• Some substances spontaneously emit radiation
– radioactivity
• The rays and particles emitted are called
“radiation”.
• Radioactive elements change their identity –
they can change into another element
– They do this because the nuclei are unstable
– Changes to atom’s nucleus is called a nuclear
reaction
Something else Dalton got wrong.
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Radioactive Decay
• Radioactive elements emit energy as
radiation
• This is called ‘radioactive decay’
• Unstable atoms undergo radioactive decay
until they form stable atoms
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Types of Radiation
• Alpha Radiation – emitting a positively
charged particle - α
• Beta Radiation – emitting a negatively
charged particle - β
• Gamma Radiation – High energy radiation
with no mass - γ
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Alpha Radiation
• Alpha radiation
– Made up of
Alpha particles
– Alpha particles
have 2 protons,
2 neutrons
– Alpha particles
have a 2+
charge
α
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Nuclear Equations
Nuclear Equations
transmutation - an element is transformed into a new
element. This can occur by natural or artificial means.
loss of an a particle
radium is emitting an a particle
+
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Beta Radiation
• Beta Radiation is fast moving electrons
– Attracted to positive charged items
– Charge is 1-
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Gamma Radiation
• High Energy radiation with no mass
– No charge
– Usually accompany alpha and beta radiation
– Account for most of the energy lost during
radioactive decay
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Chapter 4
Section 2 The Atom
The Important Role of Electrons
• The electrons at the
outer layer of the
atom are important to
the atom’s
interactions with its
environment.
• Energy Levels Each
electron cloud exists
at a certain energy
level.
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Chapter 4
Section 2 The Atom
The Important Role of Electrons, continued
• Valence Electrons
•Outer layer (or cloud or energy level) of the atom
are Valence Electrons
•Most likely to be lost if the atom loses electrons.
•The outermost energy level is also where the
atom is most likely to gain electrons.
• Valence Electrons Are The BONDING Agents
•Gaining/losing/sharing electrons results in
chemical bonds
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• Chemical Bonds
– Chemical bonds are made by gaining,
losing, or sharing electrons to fill up the
outer electron shells
• Ions: Electron-Proton Imbalance
• Ions are formed when an atom loses or
gains electrons, leaving an unequal
number of protons and electrons.
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What Makes Nuclear Stability?
• Biggest factor is neutron to proton ratio
• Will be discussed in Chapter 25
• Atoms with either too many or too few
neutrons will be radioactive
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Review
• Explain how unstable atoms gain
stability?
– Atoms gain stability by losing energy as
emitted radiation.
• Complete the following table:
Particle
α
β
γ
Symbol
4
2
0
-1
He
β
γ
0
0
Mass (amu) Charge
4
1/1840
0
+2
-1
0
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Review
Particle
α
β
γ
Symbol
4
2
He
β
γ
0
-1
0
0
Mass (amu) Charge
4
+2
1/1840
-1
0
0
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Review
• Classify each as a chemical reaction, nuclear
reaction, or neither
• Thorium emits a beta particle:
– Nuclear
• Two atoms share electrons to form a bond:
– Chemical
• A sample of pure sulfur emits heat as it slowly
cools:
– Neither
• A piece of iron rusts:
– chemical
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Chapter 4
Atoms
Concept Mapping
Use the terms below to complete the concept map
on the next slide.
nucleus
mass number
isotopes
protons
atoms
electrons
atomic number
neutrons
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Chapter 4
Atoms
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Chapter 4
Atoms
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End of Chapter 4 Show
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Chapter 4
Standardized Test Preparation
FCAT
For the following questions, write your answers on a
separate sheet of paper.
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Chapter 4
Standardized Test Preparation
1. British chemist and schoolteacher John Dalton
published a theory that defined atoms in 1803.
Included in his theory was the idea that atoms are
small particles which cannot be divided. One of
the first major challenges to this theory came
nearly 100 years later. Another British scientist,
J.J. Thomson, created an experiment using a
cathode-ray tube and discovered the existence of
negatively charged subatomic particles. What
was the effect of this new information?
Continued on next slide
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Chapter 4
Standardized Test Preparation
Question 1, continued
A. Dalton’s theory was not changed and is still
believed to be true.
B. Dalton’s theory had to be modified in response to
the new information.
C. Two different theories of atoms were developed
and used by different scientists.
D. Thomson’s experiment was changed so that its
results matched Dalton’s theory.
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Chapter 4
Standardized Test Preparation
Question 1, continued
A. Dalton’s theory was not changed and is still
believed to be true.
B. Dalton’s theory had to be modified in response to
the new information.
C. Two different theories of atoms were developed
and used by different scientists.
D. Thomson’s experiment was changed so that its
results matched Dalton’s theory.
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Chapter 4
Standardized Test Preparation
2. The illustration below shows a model of an isotope
of boron. What is the mass of the isotope shown?
F. 5
G. 10
H. 11
I. 16
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Chapter 4
Standardized Test Preparation
2. The illustration below shows a model of an isotope
of boron. What is the mass of the isotope shown?
F. 5
G. 10
H. 11
I. 16
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Chapter 4
Standardized Test Preparation
3. What is the difference between an isotope and an
ion?
A.
An isotope is an atom that has a different number of electrons than
other atoms of the same element have. An ion is a particle that has
an equal number of protons and neutrons.
B.
An isotope is an atom that has a different number of protons than
other atoms of the same element have. An ion is a particle that has
an equal number of protons and electrons.
C.
An isotope is an atom that has a different number of neutrons than
other atoms of the same element have. An ion is a particle that has
an unequal number of protons and electrons.
D.
An isotope is an atom that has a different number of protons than
other atoms of the same element have. An ion is a particle that has
an unequal number of protons and electrons.
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Chapter 4
Standardized Test Preparation
3. What is the difference between an isotope and an
ion?
A.
An isotope is an atom that has a different number of electrons than
other atoms of the same element have. An ion is a particle that has
an equal number of protons and neutrons.
B.
An isotope is an atom that has a different number of protons than
other atoms of the same element have. An ion is a particle that has
an equal number of protons and electrons.
C.
An isotope is an atom that has a different number of neutrons than
other atoms of the same element have. An ion is a particle that has
an unequal number of protons and electrons.
D.
An isotope is an atom that has a different number of protons than
other atoms of the same element have. An ion is a particle that has
an unequal number of protons and electrons.
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Chapter 4
Standardized Test Preparation
4. British scientist Ernest Rutherford proposed a new
model of the atom in 1911. The diagram below
shows his model of the atom.
Continued on next slide
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Chapter 4
Standardized Test Preparation
Question 4, continued
4. What did this model add to atomic theory?
F.
the idea that an atom has a dense, negatively charged
nucleus with electrons surrounding the nucleus at a
distance
G. the idea that an atom has a dense, positively charged
nucleus with electrons surrounding the nucleus at a
distance
H. the idea that an atom has a dense, neutrally charged
nucleus with electrons surrounding the nucleus in an
electron cloud
I.
the idea that an atom has a dense, positively charged
nucleus with electrons surrounding the nucleus in an
electron cloud
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Chapter 4
Standardized Test Preparation
Question 4, continued
4. What did this model add to atomic theory?
F.
the idea that an atom has a dense, negatively charged
nucleus with electrons surrounding the nucleus at a
distance
G. the idea that an atom has a dense, positively charged
nucleus with electrons surrounding the nucleus at a
distance
H. the idea that an atom has a dense, neutrally charged
nucleus with electrons surrounding the nucleus in an
electron cloud
I.
the idea that an atom has a dense, positively charged
nucleus with electrons surrounding the nucleus in an
electron cloud
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Chapter 4
Standardized Test Preparation
5. Which one of the following is true of a neutron?
A. A neutron has half the mass of a proton.
B. A neutron has the same mass as an electron.
C. A neutron is a little more massive than a proton.
D. A neutron is a little more massive than an
electron.
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Chapter 4
Standardized Test Preparation
5. Which one of the following is true of a neutron?
A. A neutron has half the mass of a proton.
B. A neutron has the same mass as an electron.
C. A neutron is a little more massive than a proton.
D. A neutron is a little more massive than an
electron.
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Chapter 4
Standardized Test Preparation
6. Austrian physicist Erwin Schrödinger and German
physicist Werner Heisenberg expanded atomic
theory in the 20th century. They accepted some of
the work of earlier scientists, but they added to
atomic theory with new ideas about electrons.
They did not agree with Neils Bohr’s model that
had electrons moving in definite paths around the
nucleus of an atom. Schrödinger and Heisenberg
concluded that one cannot know exactly where
electrons are in an atom. One can only predict
where electrons are likely to be found.
Continued on next slide
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Chapter 4
Standardized Test Preparation
Question 6, continued
6. What was one of the main contributions of
Schrödinger and Heisenberg to atomic theory?
Current theory identifies regions where electrons
are likely to be found. What are these regions
called?
Copyright © by Holt, Rinehart and Winston. All rights reserved.
Chapter 4
Standardized Test Preparation
Question 6, continued
6. What was one of the main contributions of
Schrödinger and Heisenberg to atomic theory?
Current theory identifies regions where electrons
are likely to be found. What are these regions
called?
Full credit answers should include the following points: One
of the main contributions of Schrödinger and Heisenberg
was the idea that electrons do not travel in definite paths
around the nucleus. Electron clouds are the regions where
electrons are likely to be found, according to current atomic
theory.
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Chapter 4
Section 1 Development of the
Atomic Theory
Thompson’s Cathode-Ray Tube Experiment
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Chapter 4
Section 1 Development of the
Atomic Theory
Rutherford’s Gold-Foil Experiment
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Chapter 4
Section 1 Development of the
Atomic Theory
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Chapter 4
Section 1 Development of the
Atomic Theory
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Chapter 4
Section 2 The Atom
Parts of an Atom
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Chapter 4
Section 2 The Atom
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Chapter 4
Section 2 The Atom
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Chapter 4
Section 2 The Atom
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Chapter 4
Section 2 The Atom
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Chapter 4
Standardized Test Preparation
Copyright © by Holt, Rinehart and Winston. All rights reserved.
Chapter 4
Standardized Test Preparation
Copyright © by Holt, Rinehart and Winston. All rights reserved.
Chapter 4
Section 2 The Atom
Math Focus
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Chapter 4
Section 1 Development of the
Atomic Theory
Comparing Models of the Atom
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Chapter 4
Section 1 Development of the
Atomic Theory
Where Are the Electrons?, continued
• The Modern Atomic
Theory According to
the current theory, there
are regions inside the
atom where electrons
are likely to found.
These regions are
called electron clouds.
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Review
• Atoms are made of protons, neutrons, and
electrons
• Protons: Mass = 1 amu, charge = +1
– Location: nucleus
• Neutrons: Mass = 1 amu, charge = 0
– Location: nucleus
• Electrons: Mass = almost 0, charge = -1
– Location: Outside nucleus, far away
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Review
Copyright © by Holt, Rinehart and Winston. All rights reserved.