Download Introductory Chemistry The Evolution of Atomic Theory

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Chapter 3 Lecture Presentation
Introductory
Chemistry
Fourth Edition
Chapter 3
The Evolution
of Atomic Theory
© 2011 Pearson Education, Inc.
3.1 Dalton’s Atomic Theory
• Atoms proposed around 400 B.C.
– Took 2000 years to be accepted
• Dalton’s atomic model (early 1800s)
formulated explanations for a variety of
laws.
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3.1 Dalton’s Atomic Theory (Continued)
• Law of conservation of mass
– When a reaction takes place, matter is neither
created or destroyed.
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3.1 Dalton’s Atomic Theory (Continued)
• Law of constant proportions
– Multiple samples of any pure compound
always contain the same percent by mass of
each element making up the compound.
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3.1 Dalton’s Atomic Theory (Continued)
• Percent by mass
– Consider that 50.0 g of water (H2O) is
decomposed into the component elements
yielding 5.6 g of H and 44.4g of O2. What is
the percent by mass of the components?
– It is always the same for pure water.
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3.1 Dalton’s Atomic Theory (Continued)
• Dalton’s atomic theory
1.
2.
3.
4.
5.
All matter is made up of atoms.
Atoms can neither be created nor destroyed.
Atoms of a particular element are alike.
Atoms of different elements are different.
A chemical reaction involves the union or
separation of individual atoms.
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3.1 Dalton’s Atomic Theory (Continued)
• Ball-and-hook model
– Different size balls represent different atom
types.
– Different types have different numbers of
hooks representing bonding.
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3.1 Dalton’s Atomic Theory (Continued)
• No point of the theory is entirely true, so
updates have occurred.
– Atoms are not the most fundamental unit.
• Protons, electrons, and neutrons
– Atoms can be created and destroyed in
nuclear reactions.
• Although updates were needed, it does
not diminish the theory or its usefulness.
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3.2 Development of a Model
for Atomic Structure
• What do atoms look like?
– Problem tackled by J.J. Thomson, James
Chadwick, and others
– Found atoms are comprised of even smaller
particles.
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3.2 Development of a Model
for Atomic Structure (Continued)
• J.J. Thomson
– In 1897, he discovered the electron.
• The first subatomic particle
– Very small and lightweight
• 1/1836 mass of a hydrogen atom
– Has a negative charge, which is referred to as
negative one
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3.2 Development of a Model
for Atomic Structure (Continued)
• J.J. Thomson and E. Goldstein
– In 1907, they discovered the proton.
– Much heavier than the electron
• Mass is roughly equal to a hydrogen atom.
– Exhibits a positive charge referred to as
positive one
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3.2 Development of a Model
for Atomic Structure (Continued)
• James Chadwick
– 25 years later, he discovered the neutron.
– Roughly same mass as a proton
– Did not exhibit a charge
• Was electrically neutral
– Very difficult to study due to the absence of
charge
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3.2 Development of a Model
for Atomic Structure (Continued)
Properties of Subatomic Properties
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3.2 Development of a Model
for Atomic Structure (Continued)
How could these subatomic particles lead to the
variety of the periodic table?
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3.2 Development of a Model
for Atomic Structure (Continued)
How could these subatomic particles lead to the
variety of the periodic table?
By containing differing numbers of each subatomic
particles, the elements of the periodic table exhibit
differing behaviors.
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3.2 Development of a Model
for Atomic Structure (Continued)
• The first model proposed by J.J. Thomson
is called the plum-pudding model.
– Knew two basic facts:
1. Atoms contain small, negatively charged
particles.
2. Atoms of an element behave as if they have no
charge.
– Reasoned that something must encapsulate
the negative charge of the electron
– Proposed a “cloud of positive electricity”
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3.2 Development of a Model
for Atomic Structure (Continued)
Plum-Pudding Model
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3.3 The Nucleus
• Rutherford’s gold foil experiment
– Fired alpha particles at metal foils
• Alpha particle has two protons and two neutrons
with a charge of +2.
• Gold foil is only a couple of thousand atoms thick.
– Used a glass substrate covered with zinc
sulfide to monitor the alpha particles
– Expected a straight flight path for the particles
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3.3 The Nucleus (Continued)
• Results were almost the expected.
• While most flew straight through, some
bent or were “bounced” backward.
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3.3 The Nucleus (Continued)
• This non-linear flight was very unexpected.
• Rutherford’s quote:
“It was … as if you fired a 15-inch shell at a
piece of tissue paper and it came back and hit
you.”
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3.3 The Nucleus (Continued)
• From the study, he knew two things:
1. Most of the alpha particles went straight
through the foil.
•
Most of the atom must be empty space.
2. A few of the alpha particles (about 1 in
20,000) were deflected from a straight-line
path.
•
•
There must be a small and massive something
inside the atom.
This massive unit must have a positive charge.
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3.3 The Nucleus (Continued)
• Rutherford’s model of the atom
– An atom is mostly empty space.
• Contains the electrons spread throughout the atom
– A nucleus is a tiny, massive, positively
charged unit in the atom.
• Placed in the center of the atom
• Contains the protons and neutrons
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3.3 The Nucleus (Continued)
Rutherford’s Model of the Atom
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3.3 The Nucleus (Continued)
• Why don’t opposite charges simply
collapse into each other?
• Rutherford knew more work was needed.
• Ended in a new form of physics
– Quantum physics was born.
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3.4 The Structure of the Atom
• What we know:
– The nucleus at the center of the atom
contains:
• Protons—relatively massive and positively (+1)
charged
• Neutrons—relatively massive and neutrally
charged
– Electrons orbit the nucleus with a -1 charge.
– Positive and negatively charged atoms are
drawn to one another.
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3.4 The Structure of the Atom (Continued)
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3.4 The Structure of the Atom (Continued)
• Atomic number (Z)
– The number of protons in the nucleus
– Determines the identity of the atom
– The atom with one proton is always hydrogen.
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3.4 The Structure of the Atom (Continued)
• Mass number
– Number of protons plus the number of
neutrons
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3.4 The Structure of the Atom (Continued)
Sketch a neutral carbon atom in the following three
ways, showing the correct numbers of protons,
neutrons, and electrons.
(a) Make the mass number equal to 12.
(b) Make the mass number equal to 13.
(c) Make the mass number equal to 14.
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3.4 The Structure of the Atom (Continued)
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3.4 The Structure of the Atom (Continued)
• Isotopes
– Atoms with the same number of protons
(same element), but different numbers of
neutrons
– Exhibit identical chemical properties
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3.4 The Structure of the Atom (Continued)
• Isotope symbol
– Shows both the mass number and the atomic
number along with the element symbol
12
6
C
– Often, the atomic number is omitted as it is
implied by the element symbol.
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3.4 The Structure of the Atom (Continued)
• Isotopes
– All isotopes are not present in the same
amounts.
•
12C
= 98.89%, 13C = 1.11%, and 14C = trace
amounts
– Nearly all elements have isotopes.
• Hydrogen has three isotopes with special names.
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3.4 The Structure of the Atom (Continued)
1. Identify an atom with a mass number of 16,
containing 9 neutrons.
2. Give the full atomic symbol for an atom with 16
neutrons and an atomic number of 15.
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3.4 The Structure of the Atom (Continued)
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3.4 The Structure of the Atom (Continued)
• Atomic mass
– The actual mass of any atom
– Have units of atomic mass units (amu) or
daltons (Da)
– Relative atomic mass
• Measures how massive an atom is in comparison
to a 12C atom
• 1H would be 1/12 the mass of 12C.
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3.4 The Structure of the Atom (Continued)
• Weighted average of atomic mass
– Values reported on the periodic table
– Weighted average of the isotope masses
•
14C
is ignored due to its only having trace
amounts present.
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3.5 The Law of Mendeleev—Chemical
Periodicity
• By 1860, nearly 70 elements had been
isolated and studied.
• As discovered, a means to organize the
elements was needed.
• Enter Mendeleev and his arrangements
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3.5 The Law of Mendeleev—Chemical
Periodicity (Continued)
• A major breakthrough occurred when the
elements were ordered by increasing atomic
mass
• He noticed a regular repeating of properties
– Every eighth element exhibited similar properties.
– K, Na, and Li for example
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3.5 The Law of Mendeleev—Chemical
Periodicity (Continued)
• K, Na, and Li all:
– React vigorously with water
– Form oxides (K2O) and hydroxides (KOH)
– Are conductive
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3.5 The Law of Mendeleev—Chemical
Periodicity (Continued)
• Law of octaves
– Elements that are eight elements apart by
mass react in similar manners.
– Called chemical periodicity or periodic
behavior
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3.5 The Law of Mendeleev—Chemical
Periodicity (Continued)
• Law of Mendeleev
– Properties of the elements recur in regular
cycles (periodically) when elements are
arranged in order of increasing atomic mass.
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3.5 The Law of Mendeleev—Chemical
Periodicity (Continued)
• Gaps in Medeleev’s table
– He left gaps for undiscovered elements in his
table.
– He predicted the properties of the missing
elements.
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3.5 The Law of Mendeleev—Chemical
Periodicity (Continued)
• Predicted values matched those in the
later found elements
• Giving further support to his arrangement
of elements
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3.6 The Modern Periodic Table
• Current organization of the 113 elements
• Elegant and simplistic in showing information
• Relays both subatomic structures and chemical
properties
• Shows the atomic number and mass for each
element
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3.6 The Modern Periodic Table (Continued)
• Periods of the periodic table
– The horizontal rows of the table
– Numbered 1 through 7 as you descend
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3.6 The Modern Periodic Table (Continued)
• Groups
– The vertical columns of the table
– Also called families
– Numbered in a few different manners:
• Using roman numerals
• Using Arabic numbers
– Groups with varying names are shown in
violet.
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3.6 The Modern Periodic Table (Continued)
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3.6 The Modern Periodic Table (Continued)
• Sections of the table
– Main-group elements—groups 1, 2, and
13−18
• Much early chemistry based here
• Shows strongest periodic nature
– Transition metals—groups 3−12
– Lanthanides (rare earths) and actinides
• In the lower section of the table, not numbered
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3.6 The Modern Periodic Table (Continued)
• Transition-metal numbering
– Numbered with roman numeral and B
– I and II are at the high end of the B numbers.
– This awkwardness leads to the adaption of
the 1-through-18 numbering scheme.
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3.6 The Modern Periodic Table (Continued)
• Phase of matter at room
temperature
– Solid, liquid, and gas
– Group 18 elements are the
noble gases.
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3.6 The Modern Periodic Table (Continued)
• Metallic nature
– Metal, nonmetal, or metalloid
– Separated by stair-step starting at boron
– Roughly 75% are metals.
– All life is based on the nonmetal carbon.
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3.6 The Modern Periodic Table (Continued)
• Metals
– Shiny solids
– Bendable and malleable
• Nonmetals
– Brittle
– Do not conduct electricity or heat well
• Semimetals
– Can act as a metal or nonmetal
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3.6 The Modern Periodic Table (Continued)
1. How many groups constitute main-group
elements?
2. What do the groups of the periodic table have in
common with one another?
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3.6 The Modern Periodic Table (Continued)
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3.6 The Modern Periodic Table (Continued)
• Group names
– Five common names used
– The first two and last three columns
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3.6 The Modern Periodic Table (Continued)
• Law of octaves adjustments
– Works well when only considering the maingroup elements
– After row 4, repeats after 18 spaces
• Due to inclusion of transition metals
– After row 6, increases again to 32 elements
• Due to inclusion of lanthanides and actinides
• We will return to this topic in later
chapters.
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3.7 An Introduction to Ions
and the First Ionization Energy
• Atoms are neutral due to balanced
numbers of protons and electrons.
• Ions are when this balance is not present.
– Either an electron is added or removed.
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3.7 An Introduction to Ions
and the First Ionization Energy (Continued)
• Ions
– Negatively charged ions are anions.
– Positively charged ions are cations.
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3.7 An Introduction to Ions
and the First Ionization Energy (Continued)
• Ion notations
– Charges are shown as superscripts after
symbol.
– + and – are used to show positive and
negative charges
– For +1 and -1, the one is generally implied.
– Any element written without charge is neutral.
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3.7 An Introduction to Ions
and the First Ionization Energy (Continued)
• First ionization energy
– Minimum amount of energy it takes to
completely remove an electron forming +1 ion
– Always requires energy to overcome coulombic
attractions
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3.7 An Introduction to Ions
and the First Ionization Energy (Continued)
• First ionization energy
– Exhibits periodicity similar to chemical properties
– Metals exhibit lower values than nonmetals.
• Metals like to give up electrons.
• Nonmetals like to gain electrons.
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3.7 An Introduction to Ions
and the First Ionization Energy (Continued)
• Looking ahead
– Rutherford’s model cannot explain the
periodicity seen.
– An improved model is needed.
• Hint: Comes from light spectra
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3.7 An Introduction to Ions
and the First Ionization Energy (Continued)
• Sunlight
– Exhibits a continuous spectra
– All colors are seen and blend together.
• Light from elements
– Exhibits a line spectra
– Not all colors are seen
and there are distinct
breaks in spectra.
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