• Study Resource
• Explore

Survey

* Your assessment is very important for improving the work of artificial intelligence, which forms the content of this project

Document related concepts

History of molecular theory wikipedia, lookup

Unbinilium wikipedia, lookup

Ununennium wikipedia, lookup

Tennessine wikipedia, lookup

Transcript
```Chapter 5 - Electrons in Atoms
THE NUCLEAR ATOM AND
1. In the early 1900’s scientists were able to
determine that an element’s chemical
behavior is related to the arrangement
of electrons in its atoms
WAVE NATURE OF LIGHT
1. Electromagnetic radiation is a form of
energy that exhibits wavelike behavior
as it travels through space
Examples: a) visible light (ROYGBIV)
b) microwaves
c) x-rays
WAVE NATURE OF LIGHT
2. Primary characteristics of all waves
include wavelength, frequency,
amplitude, and speed
SEE Fig. 5-2, Page 119
3. Wavelength () is the shortest distance
between equivalent points on a continuous
wave is usually expressed in meters,
centimeters, or nanometers
WAVE NATURE OF LIGHT
4. Frequency (  ) is the number of waves
that pass a given point per second
5. The SI unit for frequency is the hertz (Hz)
which equals one wave per second
Example: 652 Hz = 652 waves/second =
652/s = 652s-1
WAVE NATURE OF LIGHT
6. The amplitude of a wave is the wave
height from the origin to a crest, or
from the origin to a trough
7. All electromagnetic waves travel at the
speed of 3.00 X 108 m/s in a vacuum
WAVE NATURE OF LIGHT
8. The speed of light ( c ) is the product of its
wavelength (  ) and its frequency (  )
c=
9. Although the speed of all electromagnetic
waves is the same, waves may have
different wavelengths and frequencies
SEE Fig. 5-3, Page 119
WAVE NATURE OF LIGHT
10. White light is a combination of all the
colors of the visible spectrum
R O Y G
e r e r
d a l e
n l e
g o n
e w
SEE Fig. 5-4, Page 1`20
B
l
u
e
I V
n i
d o
i l
g e
o t
WAVE NATURE OF LIGHT
11. The electromagnetic spectrum
encompasses all forms of
SEE Fig. 5-5, Page 120
Example Problem 5-1, Page 121
Practice Problems, Page 121
Practice Problems
1. A helium-neon laser emits light with a
wavelength of 633 nm. What is the
frequency of this light?
2. What is the wavelength of X-rays having
a frequency of 4.80 X 1017 Hz?
frequency of 98.5 MHz. What is the
signal?
PARTICLE NATURE OF LIGHT
1. The temperature of an object is a measure
of the average kinetic energy of its
particles
2. As an object gets hotter it possesses a
greater amount of energy, and emits
different colors of light.
SEE Fig. 5-6, Page 122
PARTICLE NATURE OF LIGHT
3. These different colors correspond to
different frequencies and wavelengths
4. Max Planck determined that matter can
gain or lose energy only in small,
specific amounts called quanta
PARTICLE NATURE OF LIGHT
5. A quantum is the minimum amount of
energy that can be gained or lost by an
atom
6. The energy of a quantum is related to the
frequency of the emitted radiation by the
equation
Equantum = h 
where E is energy, h is Planck’s constant,
and  is frequency
PARTICLE NATURE OF LIGHT
7. Planck’s constant has a value of
6.626 X 10-34 J•s
8. The photoelectric effect is a phenomenon in
which photoelectrons are emitted from a
metal’s surface when light of a certain
frequency shines on the surface
SEE Fig. 5-7, Page 123
Example: Calculator in Fig. 5-8, Page 123
PARTICLE NATURE OF LIGHT
9. While a beam of light has many wave
characteristics, it can also be thought of as
a stream of tiny particles, or bundles of
energy called photons
10. A photon is a particle of electromagnetic
radiation with no mass that carries a
quantum of energy
PARTICLE NATURE OF LIGHT
11. Albert Einstein calculated that a
photon’s energy depends on its
frequency
Ephoton = h 
Example Problem 5-2, Page 124
Practice Problems, Page 124
ATOMIC EMISSION SPECTRA
1. The atomic emission spectrum of an
element is the set of frequencies of the
electromagnetic waves emitted by
atoms of the element
2. The atomic emission spectrum is
characteristic of the element and
is used to identify the element.
Section 5.1 Assessment, Page 126
BOHR MODEL OF THE ATOM
1. Niels Bohr proposed a quantum model
for atoms which included:
a) predicted the frequencies of the lines in
hydrogen’s atomic emissions spectrum
b) proposed that the hydrogen atom has
only certain allowable energy states
BOHR MODEL OF THE ATOM
c) related the hydrogen atom’s energy
states to the motion of the electron
within the atom
d) the single electron in a hydrogen atom
moves around the nucleus in only
certain allowed circular orbits
BOHR MODEL OF THE ATOM
e) the smaller an electron’s orbit, the lower
the atom’s energy state, or energy level
f) assigned a quantum number, (n), to each
orbit and calculated the orbits radius
SEE Table 5-1, Page 127
BOHR MODEL OF THE ATOM
2. The lowest allowable energy state of an
atom is called its ground state
3. In the ground state, an atom does not
4. When energy is added to an atom by an
outside force, electrons can move up to a
higher-energy orbit (called the excited
state)
BOHR MODEL OF THE ATOM
5. Electrons can also drop to a lower-energy
level and emit energy in the form of
photons
6. To calculate the energy (E) of a photon
we can use the following equation:
ΔE = Ehigher-energy orbit - Elower-energy orbit = Ephoton = h 
THE QUANTUM MECHANICAL
MODEL OF THE ATOM
1. Louis de Broglie proposed that
particles of matter can behave like
waves
2. de Broglie derived an equation for the
wavelength (  ) of a particle of mass (m)
moving at a velocity ( v )
h
 =
mv
3. The de Broglie equation predicts that all
moving particles have wave characteristics
Example: 910 kg traveling at 25 m/s
m = 910 kg
v = 25 m/s
h = 6.626 X 10-34 J
=?
 =
h
mv
= 6.626 X 10-34 J = 2.9 X 10-38 m
(910kg)(25m/s)
THE HEISENBERG UNCERTAINTY PRINCIPLE
1. Werner Heisenberg concluded that it is
impossible to make any measurement on an
object without disturbing the object
2. Consider the energy of a photon:
A high-energy photon of electromagnetic
electron.
The interaction between the two particles
changes both the wavelength of the photon
and the position and velocity of the electron
SEE Fig. 5-13, Page 132
3. The Heisenberg uncertainty principle states
that it is fundamentally impossible to know
precisely both the velocity and position of a
particle at the same time.
Example: The effect of a photon emitted by a
flashlight on a helium balloon is so
small that it is virtually impossible
to measure
4. Erwin Schrodinger derived an equation
that treated the hydrogen atom’s
electron as a wave
THE HEISENBERG UNCERTAINTY PRINCIPLE
5. The atomic model proposed by
Schrodinger is called the quantum
mechanical model of the atom
6. The region around the nucleus of an atom
where electrons are located is called an
atomic orbital
SEE Fig. 5-13, Page 132
HYDROGEN’S ATOMIC ORBITALS
1. Principle Quantum Numbers ( n ) indicate
the relative sizes and energies of atomic
orbitals
2. As n increases, the orbital becomes larger,
the electron spends more time farther
from the nucleus, and the atom’s
energy level increases
HYDROGEN’S ATOMIC ORBITALS
3. n specifies the atom’s major energy levels,
called principle energy levels
4. Up to seven energy levels have been
detected
5. Principle energy levels contain energy
sublevels
HYDROGEN’S ATOMIC ORBITALS
6. The quantum number assigned to an
energy level also indicates the number
of sublevels contained within the energy
level
Examples:
First energy level (n = 1) has one sublevel
Second energy level (n = 2) has two sublevels
Third energy level (n = 3) has three sublevels
HYDROGEN’S ATOMIC ORBITALS
7. Sublevels are labeled s, p, d, or f
according to the shape of the atoms
orbitals
Examples:
n = 1 has a 1s sublevel
n = 2 has a 2s and 2p sublevel
n = 3 has a 3s, 3p, and 3d sublevel
n = 4 has a 4s, 4p, 4d, and 4f sublevel
SEE Table 5-2, Page 134
Section 5.2 Assessment, Page 134
GROUND STATE ELECTRON CONFIGURATION
1. The arrangement of electrons in an atom is
called the atom’s electron configuration
2. Electrons in an atom tend to assume the
arrangement that gives the atom the lowest
possible energy
3. The most stable, lowest-energy
arrangement of the electrons in an atom
of each element is called the element’s
ground-state electron configuration
GROUND STATE ELECTRON CONFIGURATION
4. The aufbau principle states that each
electron occupies the lowest energy orbital
available
5. The sequence of atomic orbitals from
lowest to highest energy (aufbau
diagram) is shown in Figure 5-17, Page
135
Aufbau Diagram
7p
GROUND STATE ELECTRON CONFIGURATION
6. Key features of the aufbau principle
include:
a) all orbitals related to an energy sublevel
are of equal energy
b) in a multi-electron atom, energy
sublevels within a principal energy
level have different energies
GROUND STATE ELECTRON CONFIGURATION
c) in order of increasing energy, the
sequence of energy sublevels within a
principal energy level is s, p, d, and
f
d) orbitals related to energy sublevels
within one principal energy level can
overlap orbitals related to energy
sublevels within another principal
level
The Pauli Exclusion Principle
7. The Pauli Exclusion principle states that
a maximum of two electrons may
occupy a single orbital, and the
electrons must have opposite spins
• Each type of subshell contains a different
number of orbitals.
• Each orbital can hold at most 2 electrons.
www.wwnorton.com/chemistry/tutorials/ch3.html
The Pauli Exclusion Principle
The following table shows how many
electrons each type of subshell can hold.
Subshell
# of Maximum # of
Type Orbitals
Electrons
s
1
2
p
3
6
d
5
10
f
7
14
Table 3-6b Orbitals and Electron Capacity of the First Four Principle Energy
Levels
Principle
energy level
(n)
Type of
sublevel
Number of
orbitals per
type
Number of
orbitals per
level(n2)
Maximum
number of
electrons (2n2)
1
s
1
1
2
s
1
p
3
4
8
s
1
p
3
9
18
d
5
s
1
p
3
d
5
16
32
f
7
2
3
4
GROUND STATE ELECTRON CONFIGURATION
7. Electrons can spin in one of two
directions. () represents one direction
and () represents the opposite direction
8. The Pauli Exclusion principle states that a
maximum of two electrons may occupy a
single orbital, and the electrons must have
opposite spins
GROUND STATE ELECTRON CONFIGURATION
9. An atomic orbital containing paired
electrons with opposite spin are written
as 
10. Hund’s Rule states that single electrons
with the same spin must occupy each
electrons with opposite spins can
occupy the same orbitals
ORBITAL DIAGRAMS AND ELECTRON
CONFIGURATION NOTATIONS
1. Two methods of representing an atom’s
electron configuration are:
a) orbital diagram
b) electron configuration notation
SEE Table 5-3, Page 137
2. Using Sodium (Na) as an example:
Na 1s22s22p63s1
 
1s 2s
  
2p

3s
3. Electron configurations for noble gases use
bracketed symbols
Example: a) Helium = [ He]
b) Neon = [ Ne ]
ORBITAL DIAGRAMS AND ELECTRON
CONFIGURATION NOTATIONS
4. The electron configuration for an element
can be represented using the noble-gas
notation for the noble gas in the
previous period and the electron
configuration for the energy level
being filled
Example: Sodium (Na) = [Ne]3s1
Topic for 11.03.05
Ch. 5 Electrons in Atoms
• Discuss Valence electrons and
Electron-dot structure
ORBITAL DIAGRAMS AND ELECTRON
CONFIGURATION NOTATIONS
5. SEE Table 5-4 for electron
configurations for elements in period
three
6. A memory aid called a sublevel diagram is
shown in Fig. 5-19, Page 138
Example Problem 5-3, Page 139
Sample Problems, Page 139
VALENCE ELECTRONS
1. Valence electrons are defined as electrons
in the atom’s outermost orbital
2. Number of valence electrons is the same
as an elements group number
(A-families)
VALENCE ELECTRONS
3. Number of energy levels is the same as an
elements period number
4. An atom’s electron-dot structure consists
of the elements symbol surrounded by
dots representing each of the atom’s
valence electrons
SEE Table 5-5, Page 140
Ch 5 test preparation:
Example Problem 5-4, Page 141
Practice Problems, Page 141
Section 5.3 Assessment, Page 141
Chapter 5 Assessment, Pages 146 - 148
Standardized Test Practice - Ch. 5, Page 149
```
Related documents