Download The Development of the Atomic Theory

Unit 3
Chapter 3 & Chapter 11
The Structure of the Atom &
Nuclear Chemistry
Early Theories of Matter
 Democritus (460-370 B.C.) proposed &
believed that
1. Matter was not infinitely divisible
2. Made up of tiny particles called atomos
3. Atoms could not be created, destroyed, or
further divided
- This sounds like the beginning of…..
John Dalton
19th century
Dalton revised Democritus's ideas based
upon the results of scientific research he
conducted
Dalton’s atomic theory
Dalton’s Atomic Theory
 Elements are made of extremely small particles called
atoms
 All atoms of a given element are identical
Same size, mass and chemical properties
 Atoms are indivisible in chemical processes.
They can neither be created nor destroyed in a
chemical reaction
A chemical reaction simply changes the way the atom
is grouped together
 A compound has a constant composition of its elements
Different atoms combine in simple whole number
ratios to form compounds
Dalton’s Atomic Theory
Dalton’s theory was supported by 2
Laws:
The Law of Conservation of Matter (Lavoisier)
The Law of Definite or Constant Composition
(Proust)
Not all of his theory is correct!!
What is an atom?
An atom is the smallest particle of an
element that retains the properties of the
element
The Nuclear Atom
 Ernest Rutherford developed a model of the
atom.
 His model consisted of the following ideas:
An atom consists mostly of empty space through which
electrons move
Electrons are held within the atom by their attraction to
the positively charged nucleus
Tiny, dense region called the nucleus, which is located
in the center of the atom
 Nucleus contains all of an atom’s positive charge
and virtually all of its mass
Three Subatomic Particles
Electron
Proton
Neutron
Subatomic Particles in the Nucleus
Positively charged particles called
protons
Proton’s charge: equal to but opposite that of an
electron
Why don’t the protons repel each other in the
nucleus?
Neutral particles called neutrons
Neutron mass nearly equal to that of proton
Neutron has no charge
Outside the nucleus
Negatively charged particles are called
electrons
move through empty space in atom
Negatively charged
Atomic Number
The number of Protons in the nucleus
Examples:
Carbon (C) has 6 protons
Atomic number is 6
Copper has 29 protons
Atomic number is 29
Atomic Number
In uncharged atom, atomic number is
also the number of electrons
Why?
If an atom is charged, then it is an ion
Uncharged atom:
Atomic number = # of protons = # of electrons
 Take out your periodic table
Mass Number
To find the Mass number
# protons + # neutrons
To find # neutrons
mass number – proton (or atomic number)
Mass numbers are always WHOLE #’s!!
Symbols for Atoms
X= symbol of element
A= mass number
Z= number of protons
A
Z
X
A
or
X
Isotopes and Mass Number
12
6
C
carbon-12
13
6
C
carbon-13
Isotopes are atoms with the same number
of protons but different number of
neutrons
Isotopes and Mass Number
Example:
3 types of Potassium
All 3 types contain 19 protons and __
electrons
# of Protons
19
19
19
# of Neutrons
20
21
22
Mass Number
What’s the difference between mass
number and average atomic mass?
12
6
C
carbon-12
But if you look on the periodic table, the
number states 12.01…
Mass of Individual Atoms
Mass of protons and neutrons:
1.67 x 10-24 g
Mass of electron is 1840 times smaller
than that of protons and neutrons
9.11 x 10-28 g
Atomic Mass Unit (amu)
Small mass #’s are not easy to work with,
so the atomic mass unit (amu) was
developed
One atomic mass unit (amu) is equal to
1/12 the mass of a carbon-12 atom
The mass of 1 amu is nearly equal to the
mass of one proton or neutron
Mass of Individual Atoms
Atomic mass:
The weighted average mass of the isotopes of
an element
Example: Chlorine
Mixture of 75% chlorine-35 and 25% chlorine-37
Atomic mass = (0.75)*35 + (0.25)*37 = 35.5 amu
What’s the difference between mass
number and average atomic mass?
Mass number- specifically about one
isotope
Average atomic mass- includes the
masses of all the different isotopes for that
atom
Nuclear Reactions
Chemical reactions
What can NOT change in a chemical reaction?
Nuclear Reactions: changes that occur in
the nucleus of an atom
The nucleus is unstable!!
Most atoms have unstable nuclei.
Why are some stable, while others are
not?
Primary Reason:
ratio of the neutrons to the protons (n/p)
 An atom is most stable when the ratio is 1:1
 The maximum ratio of stability is around 1.5 : 1
Radioactivity
An unstable nucleus emits rays and
particles, called radiation, to become
stable
The process is called radioactivity
Gain stability by LOSING energy
Radiation
Three types of radiation produced:
Alpha
Beta
Gamma
3 Radioactive Particles
1. Alpha Particle (+)
4
2
He
o It travels about 1/10 the speed of light (slowest)
o It is the largest, most massive particle
o It is the most dangerous if ingested
o It has the least penetrating ability - paper can
stop this particle
3 Radioactive Particles
2. Beta Particle (-)
0
1

o Fast accelerated electron
o Ejected when a neutron is converted to a proton
in the nucleus
o Travels 1/4 the speed of light.
o It is lighter and faster than the alpha particle.
o Average penetrating ability - can be stopped by
heavy clothing
3 Radioactive Particles
3. Gamma Ray
0
0

o Not really a particle; it is a form of energy from
the electromagnetic spectrum (EM)
o Has no mass and no charge.
o Always accompanies either beta or alpha
radiation.
o Very high powered
o Travels at the speed of light.
o Highest penetrating ability - can be stopped by
heavy shielding such as lead.
Alpha Decay
4
2
 An alpha particle is released He
 A new atom is formed in the product in which:
a) the atomic number is lowered by 2
b) the mass number is lowered by 4
 Example:
226
88
Ra   Rn  He
222
86
4
2
Beta Decay
 A neutron is converted in the nucleus of an
atom resulting in:
0

a) beta particle released 1
b) a new atom in the product whose atomic number
increases by 1
c) mass number does not change
 Example:
14
6
C   N  B
14
7
0
1
Gamma Radiation
0
0

 It is the “Energy” of the reaction
 Always accompanies alpha and beta
emissions.
 Regardless of radiation or decay,
The Law of Conservation of Matter must be
observed!!
The nuclear equation must be BALANCED!!
226
88
Ra  
Rn  He
222
86
4
2
RADIOACTIVE PARTICLES
Summary
Alpha Particle
Beta particle
Gamma ray
Positive
Charge
Negative
Charge
No charge
Mass of a
Helium nucleus
Mass of an
electron
No Mass
1/10 the speed
of light
1/4 the speed
of light
Speed of Light, c
Half Life
The half life of a radioisotope is the time
it takes for half of the sample to decay.
Half-life is represented by “t1/2”
The number of half-lives that have
passed is represented by “n”
Amount Remaining
 You can use two equations, depending on
what information you have
If you have the # of half-lives that
have passed..
If you DON”T have the # of halflives that have passed..
initial amount x (1/2) n
initial amount x (1/2) P/ t1/2
n= # of half-lives that have
passed
p= time that has passed
t1/2= time of half-live
Half-life problem
If the half life element A is 3 hours and you
have 90 grams of it, how many grams
would be left after 9 hours?
What percent of it is remaining?
Handout #5 Equations to Add
To find the half-life
t1/2= time passed/ # of half-lives passed
To find the # of half lives that have
passed…
n= time passed/ one half-life
Total time passed= n times t1/2
Problem 3
Given information
I= 64g
r= 2.0g
p=12.5 hours
n=?
One half-life=??
FUSION and FISSION
 Fission is the splitting of nuclei
resulting in a tremendous release of energy
heavy atoms split so they can become more stable
 Fusion is the combining of nuclei
resulting in more energy being released than fission
The sun produces energy as a result of nuclear
fusion
2 H atoms combine to form Helium