Transcript Chapter 5: Electrons in Atoms • John Dalton (1766- 1844), an English

Chapter 5: Electrons in Atoms
Development of the Modern Atomic
Theory
Dalton’s Atomic Theory
• John Dalton (17661844), an English
schoolteacher and
chemist.
Dalton’s Atomic Theory
• The following statements are the main points
of Dalton’s atomic theory.
1. All matter is made up of atoms.
2. Atoms are indestructible and cannot be
divided into smaller particles. (Atoms are
indivisible.)
3. All atoms of one element are exactly
alike, but are different from atoms of
other elements.
The Electron
• Dalton’s atomic theory, led most scientists in
the 1800s believed that the atom was like a
tiny solid ball that could not be broken up
into parts.
• In 1897, a British physicist, J.J. Thomson,
discovered that this solid-ball model was not
accurate.
• Thomson’s experiments used a vacuum tube.
The Electron
• A vacuum tube has
had all gases pumped
out of it.
• At each end of the tube is a metal piece called
an electrode, which is connected through the
glass to a metal terminal outside the tube.
• These electrodes become electrically charged
when they are connected to a high-voltage
electrical source.
Cathode-Ray Tube
• When the electrodes are charged, rays travel
in the tube from the negative electrode, which
is the cathode, to the positive electrode, the
anode.
• These rays originate at the cathode, they are
called cathode rays.
• Thomson found that the rays bent toward a
positively charged plate and away from a
negatively charged plate.
Cathode-Ray Tube
• He knew that objects with like charges repel
each other, and objects with unlike charges
attract each other.
• Thomson concluded that cathode rays are
made up of invisible, negatively charged
particles referred to as electrons.
• These electrons had to come from the matter
(atoms) of the negative electrode.
Cathode-Ray Tube
• From Thomson’s experiments, scientists had
to conclude that atoms were not just neutral
spheres, but somehow were composed of
electrically charged particles.
• Reason should tell you that there must be a
lot more to the atom than electrons.
• Matter is not negatively charged, so atoms
can’t be negatively charged either.
• If atoms contained extremely light, negatively
charged particles, then they must also contain
positively charged particles—probably with a
much greater mass than electrons.
Rutherford’s Gold Foil Experiment
• In 1909, a team of
scientists led by Ernest
Rutherford in England
carried out the first of
several important
experiments that
revealed an arrangement
far different from the
cookie-dough model of
the atom.
Protons
• In 1886, scientists discovered that a cathoderay tube emitted rays not only from the
cathode but also from the positively charged
anode.
• These rays travel in a direction opposite to
that of cathode rays.
Protons
• Like cathode rays, they are deflected by
electrical and magnetic fields, but in
directions opposite to the way cathode rays
are deflected.
• Thomson was able to show that these rays
had a positive electrical charge.
• Years later, scientists determined that the rays
were composed of positively charged
subatomic particles called protons.
Protons
• At this point, it seemed that atoms were made
up of equal numbers of electrons and protons.
• However, in 1910, Thomson discovered that
neon consisted
of atoms of
two different
masses. This
led to isotopes.
• He didn’t know where they were in the atom.
protons.
Protons
• Atoms of an element that are chemically
alike but differ in mass are called isotopes
of the element.
• Today, chemists know that neon consists of
three naturally occurring isotopes.
• The third was too scarce for Thomson to
detect.
Rutherford’s Gold Foil Experiment
• The experimenters set up a lead-shielded box
containing radioactive polonium, which
emitted a beam of positively charged
subatomic particles through a small hole.
• Today, we know that the particles of the
beam consisted of clusters containing two
protons and two neutrons and are called alpha
particles.
• The sheet of gold foil was surrounded by a
screen coated with zinc sulfide, which glows
when struck by the positively charged
particles of the beam.
The Gold Foil Experiment
The Nuclear Model of the Atom
• The new model of the atom as pictured by
Rutherford’s group in 1911 is shown below.
The Nuclear Model of the Atom
• To explain the results of the experiment,
Rutherford’s team proposed a new model
of the atom.
• Most of the particles passed through the foil,
they concluded that the atom is nearly all
empty space.
• So few particles were deflected, they
proposed that the atom has a small, dense,
positively charged central core, called a
nucleus.
Neutrons
• The discovery of isotopes, scientists
hypothesized that atoms contained still a
third type of particle that explained these
differences in mass.
• Calculations showed that such a particle
should have a mass equal to that of a proton
but no electrical charge.
• The existence of this neutral particle, called a
neutron, was confirmed in the early 1930s.
The Nuclear Model of the Atom
• Because so few particles were deflected, they
proposed that the atom has a small, dense,
positively charged central core, called a
nucleus.
Thompson
1906
Rutherford
1913
Bohr
1924
Atomic Numbers
• The atomic number of an element is the
number of protons in the nucleus of an atom
of that element.
• It is the number of
protons that
determines the
identity of an element,
as well as many of its
chemical and physical
properties.
Atomic Numbers
• Atoms have no overall electrical charge, an
atom must have as many electrons as there
are protons in its nucleus.
• Therefore, the atomic
number of an element
also tells the number
of electrons in a
neutral atom of that
element.
Masses
• The mass of a neutron is almost the same as
the mass of a proton.
• The sum of the protons and neutrons in the
nucleus is the mass number of that particular
atom.
Masses
• Isotopes of an element have different mass
numbers because they have different numbers
of neutrons, but they all have the same
atomic number.
Atomic Mass
• In order to have a simpler way of comparing
the masses of individual atoms, chemists
have devised a different unit of mass called
an atomic mass unit, which is given the
symbol u.
• An atom of the carbon-12 isotope contains
six protons and six neutrons and has a mass
number of 12.
Atomic Mass
• Chemists have defined the carbon-12 atom as
having a mass of 12 atomic mass units.
• Therefore, 1 u = 1/12 the mass of a carbon-12
atom.
• 1 u is approximately the mass of a single
proton or neutron.
Calculating Atomic Mass
Calculating Atomic Mass
• Copper exists as a mixture of two isotopes.
• The lighter isotope (Cu-63), with 29 protons
and 34 neutrons, makes up 69.17% of copper
atoms.
• The heavier isotope (Cu-65), with 29 protons
and 36 neutrons, constitutes the remaining
30.83% of copper atoms.
Calculating Atomic Mass
• The atomic mass of Cu-63 is 62.930 amu,
and the atomic mass of Cu-65 is 64.928 amu.
• Use the data above to compute the atomic
mass of copper.
Calculating Atomic Mass
• First, calculate the contribution of each
isotope to the average atomic mass, being
sure to convert each percent to a fractional
abundance.
Calculating Atomic Mass
• The average atomic mass of the element is
the sum of the mass contributions of each
isotope.
Question 1
How does the atomic number of an element
differ from the element’s mass number?
Answer
The atomic number of an element is the number
of protons in the nucleus. The mass number is
the sum of the number of protons and neutrons.
Assessment Questions
Question 2
The table on the next slide shows the five
isotopes of germanium found in nature, the
abundance of each isotope, and the atomic
mass of each isotope.
Assessment Questions
Question 3
Calculate the atomic mass of germanium.
Isotope
Abundance
Atomic Mass (amu)
Germanium-70
21.23% or 0.2123
69.924
Germanium-72
27.66% or 0.2766
71.922
Germanium-73
7.73% or 0.0773
72.923
Germanium-74
35.94% or 0.3594
73.921
Germanium-76
7.44% or 0.0744
75.921
Answer 72.59 amu
Information in the Periodic Table
• The number
at the bottom
of each box is
the average
atomic mass
of that
element.
• This number is the weighted average mass
of all the naturally occurring isotopes of
that element.
Electrons in Motion
• Niels Bohr (1885-1962), a Danish scientist
who worked with Rutherford, proposed that
electrons must have enough energy to keep
them in constant motion around the nucleus.
• Electrons have energy of motion that enables
them to overcome the attraction of the
positive nucleus.
Electrons in Motion
• This energy keeps
the electrons
moving around the
nucleus.
• Bohr’s view of
the atom, which
he proposed in
1913, was called
the planetary
model.
The Electromagnetic Spectrum
• To boost a satellite into a higher orbit
requires energy from a rocket motor.
• One way to increase the energy of an electron
is to supply energy in the form of highvoltage electricity.
• Another way is to supply electromagnetic
radiation, also called radiant energy.
The Electron Cloud Model
• As a result of continuing research throughout
the 20th century, scientists today realize that
energy levels are not neat, planetlike orbits
around the nucleus of an atom.
• Instead, they are spherical regions of space
around the nucleus in which electrons are
most likely to be found.
The Electron Cloud Model
• Electrons themselves take up little space but
travel rapidly through the space surrounding
the nucleus.
• These spherical regions where electrons
travel may be depicted as clouds around the
nucleus.
• The space around the nucleus of an atom
where the atom’s electrons are found is called
the electron cloud.
The Electron Cloud Model
Electrons in Energy Level
• How are electrons arranged in energy levels?
• Each energy level can hold a limited number
of electrons.
• The lowest energy level is the smallest and
the closest to the nucleus.
Electrons in Energy Level
• This first energy level holds a maximum of
two electrons.
• The second energy level is larger because it is
farther away from the nucleus. It holds a
maximum of eight electrons.
• The third energy level is larger still and holds
a maximum of 18 electrons.
Energy Levels
• A hydrogen atom has only one electron. It’s in
the first energy level.
Electrons in Energy Level
• The electrons in the outermost energy level
are called valence electrons.
• You can also use the periodic table as a tool
to predict the number of valence electrons in
any atom in Groups 1, 2, 13, 14, 15, 16, 17,
and 18.
• All atoms in Group 1, like hydrogen, have
one valence electron. Likewise, atoms in
Group 2 have two valence electrons.
Electrons in Energy Level
• An oxygen atom has eight electrons. Two of
these fill the first energy level, and the
remaining six are in the second energy level.
Lewis Dot Diagrams
• Valence electrons are so important to the
behavior of an atom, it is useful to represent
them with symbols.
Lewis Dot Diagrams
• A Lewis dot diagram illustrates valence
electrons as dots (or other small symbols)
around the chemical symbol of an element.
Lewis Dot Diagrams
• Each dot represents one valence electron.
• In the dot diagram, the element’s symbol
represents the core of the atom—the nucleus
plus all the inner electrons.
Question 2
Write a Lewis dot diagram for each of the
following.
A. Chlorine
B. Calcium
C. Potassium
Answer
A. Chlorine
B. Calcium
C. Potassium
Question 3
Give an example for each type of
electromagnetic energy listed below.
A. Ultraviolet light
B. Infrared light
C. Visible light
Energy Levels and Sublevels
• Each energy level has a specific number of
sublevels, which is the same as the number
of the energy level.
• For example, the first energy level has one
sublevel. It’s called the 1s sublevel.
• The second energy level has two sublevels,
the 2s and 2p sublevels
Energy Levels and Sublevels
• The third energy level has three sublevels:
the 3s, 3p, and 3d sublevels; and the fourth
energy level has four sublevels: the 4s, 4p,
4d, and 4f sublevels.
• Within a given energy level, the energies of
the sublevels, from lowest to highest, are s, p,
d, and f.
The Distribution of Electrons in Energy Levels
• A specific
number of
electrons can
go into each
sublevel.
The Distribution of Electrons in Energy Levels
• An s sublevel
can have a
maximum of
two electrons,
a p sublevel
can have six
electrons,
The Distribution of Electrons in Energy Levels
• A d sublevel
can have ten
electrons, and
an f sublevel
can have 14
electrons.
Orbitals
• In the 1920s, Werner Heisenberg reached the
conclusion that it’s impossible to measure
accurately both the position and energy of an
electron at the same time.
• This principle is known as the Heisenberg
uncertainty principle. In 1932, Heisenberg
was awarded the Nobel Prize in Physics for
this discovery, which led to the development
of the electron cloud model to describe
electrons in atoms.
Orbitals
• The electron cloud model is based on the
probability of finding an electron in a certain
region of space at any given instant.
• In any atom, electrons are distributed into
sublevels and orbitals in the way that creates
the most stable arrangement; that is, the one
with lowest energy.
Electron Configurations
• This most stable arrangement of electrons in
sublevels and orbitals is called an electron
configuration.
• Electrons fill orbitals and sublevels in an
orderly fashion beginning with the innermost
sublevels and continuing to the outermost.
Orbitals and the Periodic Table
• The shape of the modern periodic table is a
direct result of the order in which electrons
fill energy sublevels and orbitals.
• The periodic table is divided into blocks that
show the sublevels and orbitals occupied by
the electrons of the atoms.
Orbitals and the Periodic Table
• Notice that Groups 1
and 2 (the active
metals) have valence
electrons in s
orbitals, and Groups
13 to 18 (metals,
metalloids, and
nonmetals) have
valence electrons in
both s and p orbitals.
Building Electron Configurations
• Chemical properties repeat when elements
are arranged by atomic number because
electron configurations repeat in a certain
pattern.
• As you move through the table, you’ll notice
how an element’s position is related to its
electron configuration.
Building Electron Configurations
• Hydrogen has a single electron in the first
energy level. Its electron configuration is 1s1.
• This is standard notation for electron
configurations.
• The number 1 refers to the energy level, the
letter s refers to the sublevel, and the
superscript refers to the number of electrons
in the sublevel.
Building Electron Configurations
• Helium has two electrons in the 1s orbital.
Its electron configuration is 1s2.
• Helium has a completely filled first energy
level.
• When the first energy level is filled,
additional electrons must go into the second
energy level.
Building Electron Configurations
• Lithium begins the second period. Its first
two electrons fill the first energy level, so
the third electron occupies the second level.
• Lithium’s electron configuration is 1s22s1.
• Beryllium has two electrons in the 2s orbital,
so its electron configuration is 1s22s2 .
•
•
•
•
Building Electron Configurations
As you continue to move across the second
period, electrons begin to enter the p orbitals.
Each successive element has one more
electron in the 2p orbitals.
Carbon, for example, has four electrons in the
second energy level. Two of these are in the
2s orbital and two are in the 2p orbitals.
The electron configuration for carbon is
1s22s22p2.
Building Electron Configurations
Building Electron Configurations
• At element number 10, neon, the p sublevel
is filled with six electrons.
• The electron configuration for neon is
1s22s22p6.
• Neon has eight valence electrons; two are in
an s orbital and six are in p orbitals.
Building Electron Configurations
• Notice that neon’s configuration has an inner
core of electrons that is identical to the
electron configuration in helium (1s2).
• This insight simplifies the way electron
configurations are written.
• Neon’s electron configuration can be
abbreviated: [He]2s22p6.
Building Electron Configurations
• Notice that elements in the same group have
similar configurations.
• This is important because it shows that the
periodic trends in properties, observed in the
periodic table, are really the result of
repeating patterns of electron configuration.
Building Electron Configurations
Noble Gases
• Each period
ends with a
noble gas, so
all the noble
gases have
filled energy
levels and,
therefore,
stable electron
configurations.
Noble Gases
• These stable
electron
configurations
explain the lack
of reactivity of
the noble gases.
Noble gases
don’t need to
form chemical
bonds to acquire stability.
Transition Elements
• Notice in the periodic table that calcium is
followed by a group of ten elements beginning
with scandium and ending with zinc.
• These are transition elements. Now the 3d
sublevel begins to fill, producing atoms with
the lowest possible energy.
Transition Elements
• Like most metals, the transition elements lose
electrons to attain a more stable
configuration.
Inner Transition Elements
• The two rows beneath the main body of the
periodic table are the lanthanides (atomic
numbers 58 to 71) and the actinides (atomic
numbers 90 to 103).
• These two series are called inner transition
elements because their last electron occupies
inner-level 4f orbitals in the sixth period and
the 5f orbitals in the seventh period.
Assessment Questions
Question 1
Write electron configurations and abbreviated
electron configurations of the following
elements.
A. Boron
B. Fluorine
C. Phosphorus
Assessment Questions
Answer
A. Boron
B. Fluorine
C. Phosphorus
The Electromagnetic Spectrum
• Radiant energy travels in
the form of waves that
have both electrical and
magnetic properties.
• These electromagnetic
waves can travel through
empty space, as you know
from the fact that radiant
energy from the sun travels
to Earth every day.
The Electromagnetic Spectrum
• As you may already have guessed,
electromagnetic waves travel through space
at the speed of light, which is approximately
300 million meters per second.
The Electromagnetic Spectrum
• Electromagnetic radiation includes radio
waves that carry broadcasts to your radio and
TV, microwave radiation used to heat food in
a microwave oven, radiant heat used to toast
bread, and the most familiar form, visible
light.
• All of these forms of radiant energy are parts
of a whole range of electromagnetic radiation
called the electromagnetic spectrum.
The Electromagnetic Spectrum
Electrons and Light
• The spectrum of light released from excited
atoms of an element is called the emission
spectrum of that element.
Evidence for Energy Levels
• Bohr theorized that
electrons absorbed
energy and moved to
higher energy states.
• Then, these excited
electrons gave off
that energy as light
waves when they fell
back to a lower
energy state.
Evidence for Energy Levels
• Electrons can have only certain amounts of
energy, Bohr reasoned, they can move around
the nucleus only at distances that correspond
to those amounts of energy.
• These regions of space in which electrons can
move about the nucleus of an atom are called
energy levels.
Energy Levels and Sublevels
• The emission spectrum for each element has
a characteristic set of spectral lines.
• This means that the energy levels within the
atom must also be characteristic of each
element.
• But when scientists investigated multielectron atoms, they found that their spectra
were far more complex than would be
anticipated by the simple set of energy levels
predicted for hydrogen.
Energy Levels and Sublevels
• Notice that
these
spectra
have many
more lines
than the
spectrum
of
hydrogen.
Energy Levels and Sublevels
• Some lines
are grouped
close
together,
and there
are big gaps
between
these
groups of
lines.
Energy Levels and Sublevels
• The big gaps
correspond
to the energy
released
when an
electron
jumps from
one energy
level to
another.
Energy Levels and Sublevels
• The interpretation of the closely spaced lines
is that they represent the movement of
electrons from levels that are not very
different in energy.
• This suggests that sublevels—divisions
within a level—exist within a given energy
level.
Answer
Sample answers:
A. ultraviolet light:
part of sunlight
B. infrared light:
radiant heat
C. visible light:
the spectrum of light
we see as color
Energy Levels and Sublevels
• If electrons are distributed over one or more
sublevels within an energy level, then these
electrons would have only slightly different
energies.
• The energy sublevels are designated as s, p,
d, or f.