Download Chapter 3 Atoms and Moles

Chapter 4 & 5
Atomic Models and Structure
1. Atomic Theory
A. General
i. as early as 400BC, a few people believed in an
atomic theory
a. the idea that matter is made of tiny particles
called atoms
b. Democritus was one of the first supporters
of this idea
ii. until recently scientists had no evidence of atoms
2. Dalton's Atomic Theory
A. General
i. in 1808, John Dalton developed an atomic theory
a. he believed that a few kinds of atoms made
up all matter, that elements are composed
of one kind of atom and compounds are made
from 2 or more kinds of atoms
b. he also reasoned that only whole numbers
of atoms could combine to form compounds
B. Dalton's Atomic Theory contains 5 principles
i. all matter is composed of extremely small particles
called atoms which cannot be subdivided,
created or destroyed
ii. atoms of a given element are identical
iii. atoms of different elements are different
iv. atoms of different elements combine in simple,
whole-number ratios to form compounds
v. in chemical reactions, atoms are combined,
separated, or rearranged but never created or
destroyed
-today, we have revised his theory
a. we know now that atoms can be divided into
smaller particles, many atoms of the same element
combine to form molecules, and all atoms of the same
element have the same number of protons but they
can have different numbers of neutrons
-Subatomic Particles
1. General
A. in the mid 1800's scientists discovered that
atoms were made of smaller particles, called
subatomic particles
i. the three subatomic particles are the
electron, proton, and the neutron
2. Electrons were discovered by using cathode rays
A. J.J. Thomson (1897) pumped most of the air out of
a glass tube and applied an electrical current to it
i. the positive end was called the anode and the
negative end was called the cathode
ii. he observed a glowing beam in the tube that
came from the cathode and struck the anode,
so he called them cathode rays and the tube
was called a cathode ray tube
3. An electron has a negative charge
A. Thomson knew the rays must have come from the
cathode which had a negative charge, so he reasoned
that the rays were negatively charged
B. To test this hypothesis, he placed a magnet near the
tube
i. the beam deflected (see fig. 6, p.80)
C. He also placed a paddle wheel in the path of the
beam, and it turned
D. These 2 experiments showed that electrons are indeed
particles and they have a negative charge
E. An electron has a negative charge, but an atom is neutral.
Also, electrons have much less mass than that of an atom.
That means there must be other particles in an atom and
some must have positive charges that balance out the
negatives
F. we can use the symbol e- to represent an electron
*Millikan discovers the charge of
an individual electron and
calculates the mass of an
electron using an oil drop
experiment
-relative charge = -1
-mass – 1/1840 the mass of
a hydrogen atom
Goldstein discovers protons (1886)
-he found rays in a cathode ray tube that
traveled in the opposite direction then the
cathode rays so they must be positively
charged
Chadwick discovered neutrons (1932)
-subatomic particles with no charge (see next
few slides)
Physicists think protons and neutrons are made of
even smaller particles called quarks
4. Rutherford Discovered the Nucleus
A. Thomson proposed that the electrons of an atom
were embedded in a positively charged ball of matter
i. this is called the plum pudding model
B. In 1909, Ernest Rutherford (one of Thomson's
former students) performed experiments that
disproved the plum pudding model
i. the experiment is referred to as "the Gold
Foil" experiment because; Rutherford directed
a beam of small, positively charged particles
(called alpha particles) at a thin sheet of gold
foil
ii. the results of the experiment showed that
most of the alpha particles went right through
the foil, with a few alpha particles being
deflected (some even went backwards)
C. The fact that most of the alpha particles went straight
through the foil shows that atoms are made mostly of empty
space
D. The fact that a few of the alpha particles were deflected
backwards shows that there is a very tiny positively charged
core to the atom
i. this small positively charged core is called the
nucleus
ii. the nucleus contains almost all of the atoms mass
and all of its positive charge, but only a very small
fraction of its volume
iii. the nucleus is made up of proton and neutrons
5. Protons and Neutrons Compose the Nucleus
A. The nucleus of an atom is very small compared to
the overall size of the atom
i. if you think of the nucleus as the size of a
marble, than the entire atom would be the
size of a football field
B. The positively charged particles are called protons
i. they have a charge opposite but equal to the
charge of an electron
ii. their mass is about 2000 times greater that
the mass of an electron
C. Since protons and electrons have opposite but
equal charges and the overall charge of the atom is
neutral, atoms must contain equal numbers of protons
and electrons
D. The mass of an atom is still greater than the mass of the
protons and electrons combined
i. this means that there is another subatomic particle
E. About 30 years after the discovery of the electron, Irene
Joliot-Curie discovered that when alpha particles hit a
sample of beryllium, a beam that could penetrate almost
anything was produced. James Chadwick found that this
beam was not deflected by electric or magnetic fields.
i. that means the beam had no charge, it was neutral
ii. these neutral particles are called neutrons and they
are found in the nucleus of all atoms except most
hydrogen
F. The number of protons in an atom is called the atomic
number and it determines the identity of the element
i. each element has a specific number of protons in its
nucleus
6. Protons and Neutrons form a stable nucleus
-Atomic Number and Mass Number
1. General
A. all atoms are made of protons and electrons and
most have neutrons
B. protons and neutrons make up the nucleus
C. the electrons occupy the space surrounding the
nucleus
D. elements differ from each other in the number of
the protons their atoms contain
2. Atomic Number is the Number of Protons of the Nucleus
A. the atomic number of an atom is its number of
protons
B. the atomic number is always a whole number,
atoms do not contain partial protons
C. no 2 elements have the same number of protons or
atomic number
D. we can use the number of protons (the atomic
number) to identify an element
E. To date, scientists have identified 118 elements
F. Since the number of protons in a neutral atom is
equal to the number of electrons, the atomic number
also tells us the number of electrons in a neutral atom
3. Mass Number is the Number of Particles in the Nucleus
A. mass number is equal to the number of protons
plus the number of neutrons
i. therefore, to calculate the number of neutrons
in an atom, subtract the atomic number from the
mass number
*mass number - atomic number = number of neutrons
B. unlike the atomic number, mass number can vary
among atoms of the same element
i. all atoms of the same element have the same
number of protons, but they can have different
numbers of neutrons
ii. atoms of the same element with different
numbers of neutrons (different mass numbers)
are called isotopes
C. Now we know how to find the number of protons,
electrons and neutrons in an atom
5. Atomic Structure Can Be Represented by Symbols
A. Each element has a name and the same name is
given to all of the atoms of that element
B. Each element also has a symbol
i. the symbol is taken from the name of the
element
a. it is usually the first letter of the
elements name (ex. Oxygen - O)
b. if more than one element begins with
the same letter, the symbol can be
the first and second letter or the first
and some other letter in the elements
name (Helium He)
c. in some cases the symbols come from
the latin or greek form of the elements
name (ex. Iron - Fe)
-these symbols do not use any letters
from the elements name
C. The elements atomic number and mass number are
sometimes written with an elements symbol
i. the atomic number is the smaller whole number
ii. the mass number is the larger of the two and it
may have a decimal in it
iii. both numbers can be written with the symbol, or
one, or neither
Isotopes of an element have the same atomic number
A. all atoms of the same element have the same number of
protons, but they can have different numbers of neutrons and
electrons
B. Remember, atoms of the same element with different
numbers of neutrons (different mass numbers) are called
isotopes
-Atomic Mass
A. Masses of Atoms are Expressed in Atomic Mass units
1. atoms are so small that they do not have much
mass, this means that the gram is not a very
convenient unit for measuring their masses.
2. therefore we use a special mass unit to express
atomic mass (the mass of an atom)
i. this unit has 2 names: the atomic mass unit
(amu) and the Dalton (Da)
ii. we will use amu
3. you can find the atomic mass for an element by checking the
periodic table
i. the mass number on the periodic table is an average
atomic mass of all the isotopic forms of the element
a. remember, isotopes are atoms of the same
element with different masses due to different
numbers of neutrons
b. the mass listed on the periodic table takes all
of the isotopic masses into account
c. to calculate the average atomic mass for an
element given the isotopic forms and their
percentages:
-change all percentages to decimals
-multiply the masses by their relative
percentages
-add all of the answers together
d. this is the reason that the mass numbers given on
the periodic table have decimals in them
Atomic Models
1. General
A. models help scientists imagine what may be happening
at the microscopic level
B. models have limitations and may need to be revised or
replaced as new information arises
2. Rutherford's Model Proposed Electron Orbits
A. Rutherford's goldfoil experiment led to the
replacement of Thomson's plum-pudding model of the
atom
B. Rutherford suggested that electrons are around
the nucleus
C. 2 years later, Niels Bohr revised Rutherford's
model by explaining that the electrons orbit the
nucleus in specific energy levels (planetary atomic
model)
3. Bohr's Model Confines Electrons to Energy Levels
A. Bohr reasoned that electrons can be only certain distances
from the nucleus
i. each distance corresponds to a certain quantity of energy that
electrons can have
ii. An electron in its lowest energy level is closest to the nucleus (this is
called the ground state)
iii. The higher the energy level the farther from the nucleus an electron
is
iv. The difference between 2 energy levels is known as a quantum of
energy
B. the energy levels in Bohr's model can be compared to the rungs
on a ladder
C. an electron can be in only one level or another, not between
levels
D. electrons can move between levels by gaining or giving off
energy, but while remaining in the same level electrons do not
change energy
4. Electrons Act Like Both Particles and Waves
A. Thomson's experiments demonstrated that
electrons act like particles that have mass
B. Louis de Broglie pointed out that the electrons in
Bohr's model behave as waves with certain
frequencies that correspond to the energy levels in
which electrons are found
C. This is called the wave-particle duality of nature
(electrons have properties of both waves (energy) and
particles)
D. the present day model of the atom takes into account both
the wave and particle properties of electrons
i. in this model, electrons are located in orbitals
a. regions around the nucleus that correspond
to specific energy levels
b. a region in an atom where there is a high
probability of finding electrons
ii. we call this model of orbitals an electron cloud
because they do not have sharp boundaries
E. when an orbital is drawn it shows the area an electron is
likely to be, but electrons can be found elsewhere
F. ex. the blades of a fan, when the blades are spinning, you
know that each blade is within the spinning image but you
cannot tell exactly where one blade is
-Electrons and Light
1.General
A. Albert Einstein proposed that light has the
properties of both waves and particles
i. wave properties - light has frequencies,
wavelengths and speeds
-wavelength is the distant between wave
peaks & frequency is the number of
wave that pass a certain point in a given
time period
-the wave length of light can vary , the
broad range of light wavelengths makes
up the electromagnetic spectrum
ii. particle properties
-when light strikes a metal, electrons are
released (this is called the photoelectric effect)
a. to remove an electron, a particle has
to have at least a minimum energy and
therefore a minimum frequency
B. According to Einstein, light can be described as a stream
as particles, the energy of which is determined by the lights
frequency
2.Light is an Electromagnetic Wave
A. the electromagnetic spectrum is made of waves
with many different wavelengths and frequencies
B. the visible spectrum is only a small section of the
electromagnetic spectrum
C. some of the other waves in the spectrum include Xrays, Ultraviolet, Infrared, Microwaves, and Radio
waves
D. All of these waves are referred to as light, even
though we can't see them
E. The frequency and wavelength of a wave vary
inversely
F. The visible spectrum consists of ROYGBIV
i. red has the longest wavelength in the visible
spectrum and violet has the shortest
3.
Light Emission
A. Every element contains a certain # of electrons,
those electrons can move from 1 energy level to the next
B. they can move to higher levels by absorbing energy
and back to lower levels by releasing energy
C. This energy is released as light that has a specific
wavelength
D. movement between different levels releases light of
a different wavelength and these different lights make
what is called and line-emission spectrum
E. each element has its own specific line-emission
spectrum and we can use them to identify elements
4.Light Provides Information About Electrons
A. the lowest energy that an electron can occupy is
called the ground state
B. when an electron absorbs energy it moves to a
higher energy and it is in an excited state
-Quantum Numbers
1.the present day model of the atom where electrons are
located in orbitals is also known as the quantum model
2.to describe the region in which the electrons can be found
around the nucleus, scientist have assigned 4 quantum
numbers to each electron
A. these numbers describe the properties of an
electron and no 2 electrons in the same atom can
have the same set of numbers because they cannot be
in the exact same energy state
i. this is known as the Pauli Exclusion Principle
B. the four quantum numbers are as follows:
i. the principal quantum number, symbolized by n, indicates the
main energy level an electron occupies
a. as n increases, the electron's distance from the nucleus
increases and so does its energy
b. this can indicate the size of the electron cloud
ii. the angular momentum quantum number, symbolized by l,
represents the sublevel an electron occupies
a. whatever the level number you are in tells you how
many sublevels are in that level (until past the 4th
level)
-ex. If you are in level 3, there are 3 sublevels
b. there are 4 possible sublevels, they are referred to as
the s, p, d, f sublevels
c. this can indicate the shape or type of the orbital
-see page 131
iii. the magnetic quantum number, symbolized by m,
indicates the orbital (subset of the sublevel)
a. indicates the number and orientations of the
orbitals around the nucleus
b. each sublevel can hold a specific number of orbitals
-s=1, p=3, d=5, f=7
c. each orbital can hold 2 electrons
iv. the spin quantum number, symbolized by an up or down
arrow
a. indicates the orientation of an electrons magnetic
field
b. in order for 2 electrons to occupy the same orbital
they must have opposite spins
-Electron Configuration
1.General
A. Remember, no 2 electrons in an atom can have the
same set of 4 quantum numbers
B. The arrangement of electrons in an atom is called
its electron configuration
C. electrons tend to assume arrangements that have
the lowest possible energy
D. different sublevels have different numbers of
orbitals with different shapes
i. see page 131
2.An Electron Occupies the Lowest Energy Level Available
A. The aufbau principle states that the structure of
each successive elements is obtained by adding one
proton to the nucleus and one electron to the lowest
energy orbital that is available
i. electrons fill orbitals that have the lowest
energy first
ii. they start in the lowest level, the lowest
sublevel etc.
B. we use arrow diagrams to show electron configurations for
elements
i. in these diagrams, we begin by drawing the levels,
sublevels and orbitals for the electrons
ii. the orbitals are represented by circles, and each circle
can hold 2 electrons that are represented by arrows
a. the arrows must go in the circles with one
pointing up and one going down, this shows that
they have different spins and therefore different
quantum numbers (which agrees with the Pauli
Exclusion Principle)
iii. the arrows begin filling in the first level in its only
sublevel and orbital, then they go to the next level and fill
in the first sublevel in its only orbital before going to the
next sublevel with its 3 orbitals. These three orbitals are
filled in with one electron at a time and then the electrons
go back and double up in the orbitals. Etc.
C. the order for filling in the 1st 18 electrons is straight
forward
i. the electron fill in the 1s, 2s, 2p, 3s, & 3p sublevels
D. the order after the first 18, electrons begins to jump
around, but there is a pattern that is usually followed
i. see arrow diagram
E. there are elements that do not follow the pattern, or can
have more than one configuration
C. remember the p, d, and f sublevels have more than 1
orbital. Electrons in these sublevels fill in the orbitals one at
a time before backtracking to double up in the orbitals
i. this is known as Hund's rule which states: for an
atom in the ground state, the number of unpaired
electrons is the maximum number possible and these
unpaired electrons have the same spin
ii. think of these orbitals as seats on a bus. Students
fill in the bus seats singly until there are no empty
seats left. Then they go back and double up
3. An electron configuration is a shorthand notation
A. we can take our arrow diagram and use it to name
the electron configuration of an element
i. this "name" shows which levels and
sublevels contain electrons and how many are
in each
ii.1s2 2s2 2p4
B. each elements configuration builds on the previous
elements configuration, so to save space when writing
these configurations we can write the symbol of the
previous noble gas (last row of elements on the
periodic table, they have a full outer level) and then
write the additional electrons in the configuration out
(ex. [Ne]3s2)