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Transcript
Chemistry Unit 3
Atomic Structure (Ch.3)
3-1 Early Models of Atoms
Democritus (450 BC)
Proposed that all matter
was made of tiny
indivisible particles.
He called these particles
atomos (meaning
indivisible).
We call them atoms.
Good looking guy!
Atom
An atom is the
smallest particle of an
element that retains
the identity of that
element.
If we repeatedly cut a
piece of Al, the
smallest possible
piece is an atom of
Al.
Classic model of an atom
Aristotle
Didn’t agree with
Democritus.
Believed matter
was continuous
and made up of
only one substance
called “hyle”
It wasn’t until the
1700’s when his
ideas were
reexamined.
Newton and Boyle (1600s)
Published articles
stating their belief in
the atomic nature of
elements
They had no proof
Antoine Lavoisier (1770’s)
French, The “Father
of Modern Chemistry”
Discovered the law of
conservation of
matter.
Matter is neither
created nor
destroyed.
Joseph Proust (1799)
French Chemist
Developed The Law
of Definite
Proportions
Compounds always
contain elements in
the same proportion
by mass.
Law of Definite
Proportions
H20 (by mass is
always)
88.9% Oxygen,
11.1% Hydrogen
If we had an 80g
sample of H20 how
much is O?
.889 x 80 = 71g
How much is H?
.111 x 80 = 9g
John Dalton (1803)
Proposed the atomic
theory of matter,
which is the basis for
present atomic theory
John Dalton,
English schoolteacher
Atomic Theory of Matter
Each element is
composed of
extremely small
particles called
atoms.
All atoms of a given
element are
identical, but differ
from those of any
other element.
Which element is this?
Atomic theory of matter
When elements unite
to form compounds, they
do so in a ratio of small
whole numbers. This is
called the Law of
Multiple Proportions.
Ex: C and O can
combine to form CO or
CO2, but not CO1.8.
Dalton’s Model of an Atom
All matter is composed of tiny particles
J.L. Gay-Lussac (early 1800s)
 Observed that working
with gas reactions at
constant volume,
temperature and
pressure are directly
related.
 He named the discovery
of this relationship
Charles Law, which is
represented by
P1/T1=P2/T2.
Amadeo Avogadro (early
1800s) – Italian Physicist
Explained GayLussac’s work using
Dalton’s theory: Equal
volumes of gases at
the same temp/pres
have the same
number of gas
molecules.
Michael Faraday (1839)
Suggested that
atomic structure was
related to
electricity.
Atoms contain
particles that have
electrical charges.
Positive (+)
Negative (-)
Opposite charges
attract
Like charges repel
William Crookes (1870’s)
English Physicist
Developed the
cathode ray tube to
find evidence for the
existence of particles
within the atom.
J.J. Thomson (1896)
Used a cathode ray
tube (CRT) to
identify negatively
charged particles,
called electrons.
Determined the ratio
of an electron’s
charge to its mass.
Developed the
“plum pudding”
model of an atom.
Cathode ray bending
toward a positive charge
Plum Pudding Model
-
- +
+
+ +
+
+
- + Atoms are composed of randomly arranged charged particles
Robert Millikan (early 1900s)
US Physicist
Used the oil drop
experiment to prove
the electron has a
negative charge
Was able to
determine the charge
of the electron
Millikan’s Oil Drop Experiment
Bothe/Chadwick (early 1930s)
English
Found high energy
particles with no
charge with the same
mass as the proton
called neutrons.
Ernest Rutherford (1909)
Used the gold foil
experiment to prove the
atom is mostly empty
space.
Rutherford concluded that
all of an atom’s positive
charge, and most of its
mass is located in the
center, called the
nucleus. Analogy: thumb
nail and the 50 yard line.
98% of the particles passes straight through
2% of the particles deflected off at varying angles
0.01% of the particles bounced straight back
Rutherford’s Planetary
Model of an atom
Positive charge and
majority of mass
located in the
nucleus.
Negatively charged
electrons orbit the
nucleus like planets.
-
++++
++
+
-
-
-
Most of an atom is empty space!
Problem
 He thought a moving electrical charge (-)
in a curved path should lose energy (give
off light).
 If it did, it would fall into the (+) nucleus.
Why don’t the (-) electrons fall into
the (+) nucleus?
Atom:The smallest particle of an
element that has the properties of that
element.
 Make up of nucleus
consists of protons
and neutrons
 Surrounded by an
electron cloud
Electron cloud
Sub-Atomic Particles
 Protons
Positively (+) charged
The number of protons in
an atom refers to the
atomic number (Z)
Composed of 3 quarks (2
up, 1 down)
Mass= 1.6726 x 10-27kg
Atomic mass  1 amu
(µ)
Sub-Atomic Particles
Neutrons
Found in nucleus
Neutral (no) charge
composed of 3 quarks
(1 up, 2 down)
Atomic mass 
1 amu (µ)
Isotopes- atoms of
the same element that
have a different
number of neutrons.
Sub-Atomic Particles
Electrons
Found in electron
clouds surrounding
the nucleus.
Negative (-) charge
Mass =
.00091 x 10-27 kg
1800 times smaller
than protons &
neutrons
Mass  0 amu (µ)
Sub-atomic particles
Electrons
Orbit the nucleus at
very high speed in
energy levels
(electron clouds).
Negatively (-) charged
Have no mass (when
compared to protons
and neutrons)
Atomic Number = Protons
 The atomic number of
an element is the number
of protons an element
has.
 Located above the
symbol of the element
 The number of protons
determines the identity
of the element.
 Each element has a
different atomic number
Neils Bohr (1913)
 Improved Rutherford’s
work by saying electrons
do not lose energy in the
atoms so they will stay in
orbit
 Stated there are definite
levels in which the
electrons follow set paths
without gaining or losing
energy (Planetary Model)
 Each level has a certain
amount of energy
associated with it and the
electrons can only jump
levels if they gain or lose
energy
 Could not explain why (-)
electrons don’t fall into
the (+) nucleus.
Energy Levels
In the ground state
for an atom, electrons
are at their lowest,
most stable energy
levels.
In the excited state,
atoms require energy
and electrons move
to a higher energy
level.
How many electrons are in
an atom?
For an atom to have
an overall neutral
charge the number of
electrons must equal
the number of
protons.
#Protons=#electrons
What element is this?
Mass number
 The Mass number of an
atom is the sum of the
mass of protons and
neutrons
 Located below the
symbol of the element
 Atomic mass is measured
in amu’s, (atomic mass
units)
 Based on Carbon having
a mass of 12
Mass = Protons + Neutrons
How many neutrons are in
an atom?
 Mass=Protons+Neutrons
 195= 78 + Neutrons
 195-78= Neutrons
 Platinum has 117
Neutrons
 Find the number of
neutrons in:
Hydrogen
Carbon
Helium
Potassium
Boron
Gold
Mass =Protons + Neutrons
 Hydrogen (H)
1 =1 + Neutrons
Hydrogen has 0 neutrons
 Helium (He)
4 = 2 + Neutrons
Helium has 2 neutrons
 Boron (B)
11 = 5 + Neutrons
Boron has 6 neutrons
 Carbon (C)
12 = 6 + Neutrons
Carbon has 6 neutrons
 Potassium (K)
39 = 19 + N
Potassium has 20 neutrons
 Gold (Au)
197 = 79 + N
Gold has 118 neutrons
Atomic Mass
The average
mass of all of the
isotopes of an
element.
Aka: average
atomic mass
number, or
atomic weight.
Isotopes:Atoms
of the same
element with
different masses.
Average Atomic Mass
Ne-20 has a mass of 19.992 amu (u), and
Ne-22 has a mass of 21.991 amu (u). In
any sample of 100 Ne atoms, 90 will be
Ne-20. Calculate the average atomic
mass of Ne.
.90 x 19.992 =
17.9928
.10 x 21.991 =
2.1991
avg mass = 20.1919 amu
Ions
An atom that has
gained or lost an
electron.
It acquires a net
electrical charge.
If an atom loses an
electron (oxidation)
it has more protons
than electrons and
has a net positive
charge. (cation)
Na
11 P
11 e-
11 P
10 e-
Na+
Ions
If an atom gains an
electron (reduction) it
has more electrons
than protons and has
a net negative
charge.(anion)
7 valence e-
Full octet
Ionic Charges
Charge of ion = # protons - #
electrons
What is the charge of a magnesium
atom that loses 2 electrons?
Number of protons
12
-Number of electrons 10
charge of ion
+2
Mg2+ or Mg+2
Charge is written to the upper right of
the symbol.
Representations of atoms
General form:
(Elemental Notation)
X = Element Symbol
A = Atomic Mass
(P + N)
Z = Atomic Number
(P)
Ionic Charge
A
Z
Charge
X
What is the atomic
structure?
Determine the
number of:
P =
N =
e =
23
+
Na
11
What is the atomic
structure?
Determine the
number of:
P = 11
N = 12
e = 10
23
+
Na
11
What is the atomic
structure?
Determine the
number of:
P =
N =
e =
127
I
53
-
What is the atomic
structure?
Determine the
number of:
P = 53
N = 74
e = 54
127
I
53
-
Put into elemental notation
Atomic # = 29
Atomic Mass = 64
Ionic charge = +2
?
How many electrons?
Atomic # = 29
Atomic Mass = 64
Ionic charge = +2
# of electrons =
64
2+
29
Cu
Put into elemental notation
37 Protons
48 Neutrons
36 Electrons
?
Put into elemental notation
37 Protons
48 Neutrons
36 Electrons
85
37
+
Rb
Max Planck (early 1900s)
 Proposed Planck’s Theory
which says that energy is given
off in little packets or particles
called quanta which is based
on the particle nature of light
 Each quantum of energy
corresponds to the different
energy levels for electrons.
 Proposed the equation: E=hf,
where E is energy, f is
frequency, and h is Planck’s
constant (6.63 x 10^-34 J/Hz)
De Broglie (1923)
 Suggested that if waves can have a particle nature,
particles can have a wave nature, known as the “waveparticle duality” principle
 Wondered why the positive nucleus and negative
electrons do not attract. Proposed that electrons moved
so fast (speed of light) that they had properties of
waves instead of particles.
The Study of Waves
Wave: a progressive disturbance propagated
from point to point in a medium or space
without progress or advance by the points
themselves
Types of Waves
 Mechanical: a wave
that requires an energy
source and an elastic
material medium to
travel.
 Electromagnetic: a
wave that does not
require a material
medium to travel; it
propagates by electric
and magnetic fields
Wave Travel
 Transverse: displacement
of the medium is
perpendicular to the
direction of propagation
of the wave.
 Longitudinal:
displacement of the
medium is parallel to
the direction of
propagation of the wave
Properties of Waves
 Wavelength (‫)ג‬: The
distance between any part of
the wave (peak) and the
nearest part that is in phase
with it (another peak).
Standard unit is meters (m).
 Frequency (f ): The number
of peaks which pass a given
point each second. Standard
unit is cycles per second which
is a hertz (Hz).
 Amplitude (A): The
maximum displacement of a
vibrating particle from its
equilibrium position. Standard
unit is meters (m).
 Velocity (v): the distance a
wave (peak) travels in a given
time. Standard unit is meters
per second (m/s).
 Energy (E): The energy of a
single photon of radiation of a
given frequency. Standard unit
is the joule (J).
Some relationships between the properties of waves
are represented by the equations:
V=f*‫ ג‬and
E=h*f , where h=6.63x10^-34 J/Hz
Werner Heisenberg (1927)
 Proposed his “Heisenberg
uncertainty principle”,
which says that the position
and momentum of an
electron cannot
simultaneously be
measured and known
exactly.
 The arrangement of
electrons is discussed in
terms of the probability of
finding an electron in a
certain location.
Erwin Schrodinger (1926)
 Studied the wave nature
of the electron and
developed mathematical
equations to describe
their wave-like behavior.
 The most probable
location of the electrons
can be found and the plot
of this probability is
called the charge cloud
model.
The four quantum numbers
 Principal Quantum Number (n)
 Refers to the energy levels
in the atom which is the
distance from the nucleus and
designated with a positive
whole number (n=1,2,3,etc)
 Wavelength of emitted photon
is determined by the “energy
jump” between energy levels
 Energy levels (or shells)
means electrons are
contained in an area where
the probability of finding the
electron is 90%
 Angular Momentum Quantum
Number (l )
 Refers to the sublevel (within
an energy level) which is one
or more “partitions” each with
a slightly different energy.
 The number of sublevels in a
particular energy level is equal
to the principal quantum
number (n).
 The types of sublevels
include: s, p, d, f, etc.
The four quantum numbers
(continued)
 Magnetic Quantum Number (m)
Refers to the orientation in space of the electrons in a sublevel
Can have any whole number value from -1 to +1 which will tell
how many orbitals are in a sublevel.
A maximum of 2 electrons per orbital.
Sublevel
s
p
d
f
# of Orbitals
Total # of electrons
Four Quantum Numbers (continued)
Spin Quantum Number (s)
Refers to the spin of the electron.
Pauli Exclusion Principle : if two electrons
occupy the same orbital, they must have
opposite spin.
Half-filled orbital:
_____
Filled orbital:
_____
Permissible Values of Quantum Numbers for
Atomic Orbitals
n
1
l
0
m
0
Orbital
1s
# of Subshells
1
#of Orbitals
1
Max # of Electrons
2
2
0
1
0
-1,0,1
2s
2p
2
1
3
2
6
3
0
1
2
0
-1,0,1
-2,-1,0,1,2
3s
3p
3d
3
1
3
5
2
6
10
4
0
0
4s
1
-1,0,1
4p
2
-2,-1,0,1,2
4d
3 -3,-2,-1,0,1,2,3 4f
4
1
3
5
7
2
6
10
14
Distribution of Electrons for Different
Elements (Electron Configuration)
 Electrons will occupy the lowest energy levels and
sublevels first.
 Notation:
Principal Quantum
Number, n (energy
level)
2p
Type of Orbital
(sublevel)
Number of
electrons
2
y
Orientation of
Orbital
Long Notation: Pyramid Filling
“Rule of thumb” for
filling electrons at the
lowest energy level
possible.
Give the long notation
electron configuration for:
O
Ca
Ag
Give the short notation
O
Ca
Ag
Orbital Diagrams
Usually only done for the outer shell
electrons, which always includes the s and
p orbitals.
Electron Dot Diagrams
Shows the outer shell electrons for
elements.