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Transcript
Chapter 2
Atoms, Molecules,
and Ions
Chapter 2
Table of Contents
2.1
2.2
2.3
2.4
2.5
2.6
2.7
The Early History of Chemistry
Fundamental Chemical Laws
Dalton’s Atomic Theory
Early Experiments to Characterize the Atom
The Modern View of Atomic Structure: An Introduction
Molecules and Ions
An Introduction to the Periodic Table
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2
Section 2.1
The Early History of Chemistry
Objectives
1. To learn about the relative abundances of
the elements
2. To learn about early Greek thinking
3. To learn the names and symbols of some
elements
4. To learn how a formula describes a
compound’s composition
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3
Section 2.1
The Early History of Chemistry
The Elements
• All of the materials in the universe can be chemically
broken down into about 100 different elements.
• Compounds are made by combining atoms of the
elements just as words are constructed from the letters
in the alphabet.
Words
Compounds
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4
Section 2.1
The Early History of Chemistry
Greek Thinking circa 400 B.C.
• Empedocles (492-432 )
•
•
•
•
•
•
•
All things made up of 4 elements
Fire
Earth
Water
Air
Democritus (460 -370)
Was the first to propose the atom
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5
Section 2.1
The Early History of Chemistry
The Atom: Shelved for 2,000 Years
• 400 B.C. ~ 1600 A.D. (Alchemist Period)
• Robert Boyle (1627-1691) defined an
element as something that could not be
broken down into simpler substances. He
was one of the first to bring
experimentation and repetition thereby
bringing an end to the Greek view of the 4
elements.
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6
Section 2.1
The Early History of Chemistry
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7
Section 2.1
The Early History of Chemistry
Names and Symbols for the Elements
• Each element has a name and a symbol.
– The symbol usually consists of the first one or two
letters of the element’s name.
• Examples:
Oxygen
Krypton
O
Kr
– Sometimes the symbol is taken from the element’s
original Latin or Greek name.
• Examples:
gold
lead
Au
Pb
aurum
plumbum
Bromine Br
“stench” in Greek
Chuck Norris element, Deceased Chemist
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8
Section 2.1
The Early History of Chemistry
Formulas of Compounds
• A compound is represented by a chemical formula in
which the number and kind of atoms present is shown by
using the element symbols and subscripts.
Example: the simple sugar, glucose
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9
Section 2.1
The Early History of Chemistry
Formula Examples
•
•
•
•
•
•
•
•
•
•
•
H2O
Water
SO3
Acid rain
C6H12O6 Glucose
Fe2(CO3)3
How many of each atom type in the following formulas?
P4O10
UF6
AlCl3
(NH4)2C8H4O2
Al2(Cr2O7)
Al(NO3)3
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10
Section 2.1
The Early History of Chemistry
Periodic Table Name Game
• Someone starts by naming an element.
• They then call on another player who must
name an element that starts with the last
letter of the previous element name.
• They then call on the next player and on
and on…
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11
Section 2.1
The Early History of Chemistry
Objectives Review
1. To learn about the relative abundances of the
elements
2. To learn about early Greek thinking
3. To learn the names and symbols of some
elements
4. To learn how a formula describes a
compound’s composition
5. Work Session: pg 71 # 35,37
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12
Section 2.2
Fundamental Chemical Laws
Objectives 2.2 – 2.6
1. To learn about Dalton’s Theory of Atoms
2. To understand and illustrate three important Laws of
Chemistry
3. To understand Thomson’s, Millikan’s, and Rutherford’s
work
4. To understand the modern view of the atom
5. To learn about the terms isotope, atomic number, and
mass number
6. To understand the use of the symbol
to describe a
given atom
7. To describe the formation of ions from their parent
atoms
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13
Section 2.2
•
•
•
•
•
Fundamental Chemical Laws
John Dalton English Schoolteacher
(1766-1844) 100 years after Boyle
Dalton started his own school at the age of 12!
At age 15, went into business with uncle
because he was threatened by his older
students when he tried to discipline them!
Chemical analysis and synthesis can go no further than
to the separation of particles from one another, and to
their reunion. No new creation or destruction of matter
is within reach of chemical agency. We might as well try
to introduce a new planet into the solar system and to
annihilate one already in existence, as to create or
destroy a particle of hydrogen- J.D.
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14
Section 2.3
Dalton’s Atomic Theory
Dalton’s Atomic Theory (1808)
•
•
•
Each element is made up of tiny particles
called atoms.
The atoms of a given element are identical; the
atoms of different elements are different in
some fundamental way or ways.
Chemical compounds are formed when atoms
of different elements combine with each other.
A given compound always has the same
relative numbers and types of atoms.
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15
Section 2.3
Dalton’s Atomic Theory
Dalton’s Atomic Theory (continued)
•
•
Chemical reactions involve reorganization
of the atoms—changes in the way they are
bound together.
The atoms themselves are not changed in
a chemical reaction.
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16
Section 2.3
Dalton’s Atomic Theory
Concept Check
Which of the following statements regarding
Dalton’s atomic theory are still consistent with
modern atomic theory?
I. Elements are made of tiny particles called atoms.
II. All atoms of a given element are identical.
III. A given compound always has the same relative
numbers and types of atoms.
IV. Atoms are indestructible.
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17
Section 2.3
Dalton’s Atomic Theory
Three Important Laws
•
Law of conservation of mass (Lavoisier):

•
Law of definite proportion (Proust):

•
Mass is neither created nor destroyed.
A given compound always contains exactly the
same proportion of elements by mass. CO2 is
always CO2!
Law of multiple proportions (Dalton):

When two 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. C2O4 = CO2
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18
Section 2.4
Early Experiments to Characterize the Atom
The Structure of the Atom
• Experiments by J.J. Thomson showed that atoms
contain electrons.
• Cathode ray tube
• Video Clip & CRT Monitor
Rigby, ID 1 2 3
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19
Section 2.4
Early Experiments to Characterize the Atom
The Structure of the Atom
The Plum Pudding Model (Tapioca)
JJ Thompson (1856-1940)
Nobel Prize in Physics 1908
Intellectual discovery of the + particle to keep neutrality
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20
Section 2.4
Early Experiments to Characterize the Atom
Robert Millikan (1909)
•
•
•
Performed experiments involving charged oil
drops.
Determined the magnitude of the charge on a
single electron.
Calculated the mass of the electron.
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21
Section 2.4
Early Experiments to Characterize the Atom
Millikan Oil Drop Experiment
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Section 2.4
Early Experiments to Characterize the Atom
Ernest Rutherford
• 1871-1937
• New Zealand
• 1908 Nobel Prize in Chemistry
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Section 2.4
Early Experiments to Characterize the Atom
The Structure of the Atom
Rutherford’s Experiment
Student of Thomson
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Section 2.4
Early Experiments to Characterize the Atom
The Structure of the Atom
• Results of the Rutherford experiment
(a) The results that the metal foil
experiment would have yielded if the
plum pudding model had been correct
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(b) Actual results
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25
Section 2.4
Early Experiments to Characterize the Atom
Rutherford’s Gold Foil Experiment pHet
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26
Section 2.4
Early Experiments to Characterize the Atom
Introduction to the Modern Concept of Atomic
Structure
• Ernest Rutherford showed that
atoms have internal structure.
– The nucleus, which is at the
center of the atom, contains
protons (positively charged)
and neutrons (uncharged
1932).
– Neutrons were another
intellectual discovery for
reasons of mass and repulsive
space between the protons
– Electrons move around the nucleus.
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27
Section 2.5
The Modern View of Atomic Structure: An Introduction
Introduction to the Modern Concept of Atomic
Structure Comparing the Parts of an Atom
If a grape is a nucleus, the nearest electron is 1 mile away! Neutron went to dinner...
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28
Section 2.5
The Modern View of Atomic Structure: An Introduction
•
The nucleus is:
 Small compared with
the overall size of the
atom.
 Extremely dense;
accounts for almost
all of the atom’s
mass.
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Section 2.5
The Modern View of Atomic Structure: An Introduction
Isotopes
• Isotopes are atoms with the same number of protons but
different numbers of neutrons.
Remember the bean!
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30
Section 2.5
The Modern View of Atomic Structure: An Introduction
Isotopes
• Show almost identical chemical properties;
chemistry of atom is due to its electrons.
• In nature most elements contain mixtures of
isotopes.
• A particular isotope is represented by the symbol
.
The Simpson’s Baseball.. Querque Topes
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Section 2.5
The Modern View of Atomic Structure: An Introduction
Examples
Carbon-12
• # proton
12
6C
Carbon-13 136C
Carbon-14 146C
• # neutron
• # electron
• The number of protons DEFINES the atom type.
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Section 2.5
The Modern View of Atomic Structure: An Introduction
Practice – List the number of Protons, Neutrons, and
Electrons in each neutral species
•
•
•
•
•
•
•
90
38Sr
201 Hg
80
Magnesium-24
Silver with 61 neutrons
Phosphorus with 17 neutrons
Ag-108
What did the chemist say when she found two isotopes
of Helium?
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33
Section 2.5
The Modern View of Atomic Structure: An Introduction
Exercise
A certain isotope X contains 23 protons and 28
neutrons.
• What is the mass number of this isotope?
• Identify the element.
Mass Number = 51
Vanadium
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Section 2.5
The Modern View of Atomic Structure: An Introduction
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Section 2.6
Molecules and Ions
Ions
• Atoms can form ions by gaining or losing electrons.
– Metals tend to lose one or more electrons to form
positive ions called cations.
– Cations are generally named by using the name of
the parent atom.
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36
Section 2.6
Molecules and Ions
Ions
• Nonmetals tend to gain one or more electrons to form
negative ions called anions.
• Anions are named by using the root of the atom name
followed by the suffix –ide.
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Section 2.5
The Modern View of Atomic Structure: An Introduction
Objectives Review 2.2 – 2.6
1.
2.
3.
4.
5.
6.
7.
8.
To learn about Dalton’s Theory of Atoms
To understand and illustrate three important Laws of
Chemistry
To understand Thomson’s, Millikan’s, and Rutherford’s work
To understand the modern view of the atom
To learn about the terms isotope, atomic number, and mass
number
To understand the use of the symbol
to describe a given
atom
To describe the formation of ions from their parent atoms
Work Session: pg 68 #2,4,5 pg 69-73 # 5,9,17,18,19,21,
22 skip b, 39,51,49,73 Challenge! (Hint Na makes -1
charge)
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38
Section 2.7
An Introduction to the Periodic Table
Objectives
1. To learn the various features of the
periodic table
2. To learn some of the properties of metals,
nonmetals and metalloids
3. To learn the natures of the common
elements
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Section 2.7
An Introduction to the Periodic Table
• The periodic table is organized to group elements with
similar properties in vertical columns. (Families)
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Section 2.7
An Introduction to the Periodic Table
• Most elements are metals and occur on the left side.
• The nonmetals appear on the right side.
• Metalloids are elements that have some metallic and
some nonmetallic properties.
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Section 2.7
An Introduction to the Periodic Table
•
Physical Properties of Metals
1.
2.
3.
4.
Efficient conduction of heat and electricity
Malleability (can be hammered into thin sheets)
Ductility (can be pulled into wires)
A lustrous (shiny) appearance
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Section 2.7
An Introduction to the Periodic Table
Natural States of the Elements
• Most elements are very reactive.
• Elements are not generally found in uncombined form.
– Exceptions are:
• Noble metals – gold, platinum and silver
• Noble gases – Group 8 (Bumped the Table…)
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Section 2.7
An Introduction to the Periodic Table
Natural States of the Elements
• Diatomic Molecules
Nitrogen gas contains
N2 molecules.
Oxygen gas contains
O2 molecules.
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Section 2.7
An Introduction to the Periodic Table
Natural States of the Elements
• Diatomic Molecules (The Diatomic 7 Yeehaw!)
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Section 2.7
An Introduction to the Periodic Table
Natural States of the Elements
• Elemental Solids (Allotrope same element different form)
Carbon
atoms
Diamond
Graphite
Buckminsterfullerene
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Section 2.7
An Introduction to the Periodic Table
Objectives Review
1. To learn the various features of the
periodic table
2. To learn some of the properties of metals,
nonmetals and metalloids
3. To learn the natures of the common
elements
4. QUOTE NEXT SLIDE
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Section 2.7
An Introduction to the Periodic Table
Food For Thought
• The only objective reality in the universe
is that which constitutes a combination of
every point of experience.
• Absolute truth can be ascertained only
through the sum total of all relative
observations.
-EinsteinReturn to TOC
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Section 2.7
An Introduction to the Periodic Table
This is the end of the required material………….
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Section 2.7
An Introduction to the Periodic Table
Acids
•
•
Acids can be recognized by the
hydrogen that appears first in the
formula—HCl.
Molecule with one or more H+ ions
attached to an anion.
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Section 2.7
An Introduction to the Periodic Table
Acids
•
•
If the anion does not contain oxygen, the
acid is named with the prefix hydro– and the
suffix –ic.
Examples:
HCl
Hydrochloric acid
HCN
Hydrocyanic acid
H2S
Hydrosulfuric acid
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Section 2.7
An Introduction to the Periodic Table
Acids
•
If the anion does contain oxygen:
 The suffix –ic is added to the root name if
the anion name ends in –ate.
• Examples:
HNO3
Nitric acid
H2SO4
Sulfuric acid
HC2H3O2 Acetic acid
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Section 2.7
An Introduction to the Periodic Table
Acids
•
If the anion does contain oxygen:
 The suffix –ous is added to the root name
if the anion name ends in –ite.
• Examples:
HNO2
Nitrous acid
H2SO3
Sulfurous acid
HClO2
Chlorous acid
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Section 2.7
An Introduction to the Periodic Table
Flowchart for Naming Acids
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Section 2.6
Molecules and Ions
Chemical Bonds
•
Covalent Bonds
 Bonds form between atoms by sharing
electrons.
 Resulting collection of atoms is called a
molecule.
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Section 2.6
Molecules and Ions
Covalent Bonding
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Section 2.6
Molecules and Ions
Chemical Bonds
•
Ionic Bonds
 Bonds form due to force of attraction
between oppositely charged ions.
 Ion – atom or group of atoms that has a net
positive or negative charge.
 Cation – positive ion; lost electron(s).
 Anion – negative ion; gained electron(s).
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Section 2.6
Molecules and Ions
Molecular vs. Ionic Compounds
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Section 2.6
Molecules and Ions
Exercise
A certain isotope X+ contains 54 electrons and
78 neutrons.
• What is the mass number of this isotope?
133
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Section 2.7
An Introduction to the Periodic Table
The Periodic Table
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Section 2.7
An Introduction to the Periodic Table
Groups or Families
•
Table of common charges formed when
creating ionic compounds.
Group or Family
Charge
Alkali Metals (1A)
1+
Alkaline Earth Metals (2A)
2+
Halogens (7A)
1–
Noble Gases (8A)
0
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