Chapter 3. Types of Models – Why do we use models?. Physical Model vs. Conceptual Model. Physical model. is a representation of a very large or small object shown at a convenient scale. . ID: 676898
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Atomic Theory & Structure
Chapter 3
Slide2Types of Models – Why do we use models?
Physical Model vs. Conceptual Model
Physical model
is a representation of a very large or small object shown at a convenient scale.
Conceptual model
is an explanation that treats what is being explained as a system. In a system, it is the interactions among component parts that matter most; knowing or seeing the parts themselves does not provide the understanding we are looking for.
Slide3Timeline
2000
1000
300
CE
American
Independence
(1776)
Issac Newton
(1642 - 1727)
400
BCE
Greeks
(Democritus ~450
BCE)
Discontinuous
theory of matter
Greeks
(Aristotle ~350
BCE)
Continuous
theory of matter
Slide4The Greeks
History of the Atom
Not the history of atom, but the idea of the atom
In 400 B.C.E. the Greeks tried to understand matter (chemicals) and broke them down into earth, air, fire,
and water.
Democritus and Leucippus Greek philosophers
~
~
Slide5The Hellenic Market
Fire Water Earth Air
~
~
Slide6Greek Model
Greek philosopher
Idea of ‘democracy’
Idea of ‘
atomos
’
Atomos
= ‘indivisible’
‘Atom’ is derived No experiments to support ideaContinuous vs. discontinuous theory of matter
Democritus’s model of atom
No protons, electrons, or neutrons
Solid and INDESTRUCTABLE
Democritus
“To understand the very large,
we must understand the very small.”
Slide7DEMOCRITUS
(400 BC) – First Atomic Hypothesis
Atomos
: Greek for “
uncuttable
”. Chop up a piece of matter until
you reach the
atomos.Properties of atoms: indestructible.
changeable, however, into different forms. an infinite number of kinds so there are an infinite number of elements. hard substances have rough, prickly atoms that stick together.
liquids have round, smooth atoms that slide over one another. smell is caused by atoms interacting with the nose – rough atoms hurt.
sleep is caused by atoms escaping the brain. death – too many escaped or didn’t return. the heart is the center of anger.
the brain is the center of thought. the liver is the seat of desire.
“Nothing exists but atoms and space, all else is
opinion.”
Democritus
Slide8Anaxagoras
(Greek, born 500 B.C.)
Suggested every substance had its own kind of “
seeds
” that clustered together to make the substance, much as our atoms cluster to make molecules.
Some Early Ideas on Matter
Empedocles
(Greek, born in Sicily, 490 B.C.)
Suggested there were only
four basic seeds – earth, air, fire, and water
. The elementary substances (atoms to us) combined in various ways to make everything.
Democritus
(Thracian, born 470 B.C.)Actually
proposed the word atom (indivisible) because he believed that all matter consisted of such tiny units with voids between, an idea quite similar to our own beliefs. It was rejected by Aristotle and thus lost for 2000 years.
Aristotle
(Greek, born 384 B.C.)Added the idea of “qualities”
– heat, cold, dryness, moisture – as basic elements which combined as shown in the diagram (previous page).
Hot + dry made fire; hot + wet made air, and so on.
Slide9Alchemy
After that chemistry was ruled by alchemy.
They believed that that could take any cheap metals and turn them into gold.
Alchemists were almost like magicians.
elixirs, physical immortality
Slide10Alchemy
. . . . .
. . . .
. . .
. .
.
GOLD
SILVER
COPPER
IRON
SAND
Alchemical symbols for substances…
transmutation
: changing one substance into another
In ordinary chemistry, we cannot transmute elements.
D
Slide11Slide12Contributions
of
Alchemists
:
Information about elements
- the elements mercury, sulfur, and antimony were discovered
- properties of some elements
Develop lab apparatus / procedures / experimental techniques
- alchemists learned how to prepare acids.
- developed several alloys
- new glassware
Slide13Dalton Model of the Atom
Late 1700’s - John Dalton- England
Teacher- summarized results of his experiments and those of others
Combined ideas of elements with that of atoms in
Dalton’s Atomic Theory
Slide14The Atomic Theory of Matter
In 1803, Dalton proposed that elements consist of individual particles called
atoms.
His
atomic theory of
matter
contains four hypotheses:
All matter is composed of tiny particles called atoms
, indivisible and indestructible.
All atoms of an element are identical in mass and fundamental chemical properties.
A chemical compound is a substance that always contains the same atoms in the same ratio.
In chemical reactions, atoms from one or more compounds or elements redistribute or rearrange in relation to other atoms to form one or more new compounds. Atoms themselves do not undergo a change of identity in chemical reactions
.
Slide15Foundations of Atomic Theory
Dalton’s Theories were supported by these scientific laws:
Law of Conservation of Mass
Mass is neither destroyed nor created during ordinary chemical reactions.
Law of Definite Proportions
The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.
Law of Multiple Proportions
If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first elements is always a ratio of small whole numbers.
Slide16Conservation of Mass (Atoms)
John Dalton
2 H
2
+ O
2
2 H
2
O
4 atoms hydrogen
2 atoms oxygen
4 atoms hydrogen
2 atoms oxygen
H
H
O
O
O
O
H
H
H
H
H
H
H
2
H
2
O
2
H
2
O
H
2
O
+
Slide17Law of Definite Proportions
Joseph Louis Proust (1754 – 1826)
Each compound has a specific ratio of elements
It is a ratio by mass
Water is always 8 grams of oxygen for every one gram of hydrogen
i.e. H
2
0, so two hydrogen for every one oxygen
Slide18The Law of Multiple Proportions
Dalton could not use his theory to determine the elemental compositions of chemical compounds because he had no reliable scale of atomic masses.
Dalton’s data led to a general statement known as the
law of multiple proportions
.
Law states that when two elements form a series of compounds, the ratios of the masses of the second element that are present per gram of the first element can almost always be expressed as the ratios of integers.
Slide19The Atomic Theory of Matter
Dalton’s atomic theory is essentially correct, with four minor modifications:
Not all atoms of an element must have precisely the same mass.
Atoms of one element can be transformed into another through nuclear reactions.
The composition of many solid compounds are somewhat variable.
Under certain circumstances, some atoms can be divided (split into smaller particles: i.e. nuclear fission).
Slide20
Dalton’s Symbols
John Dalton
1808
Slide21Structure of Atoms
Scientist began to wonder what an atom was like.
Was it solid throughout with no internal structure or was it made up of smaller, subatomic particles?
It was not until the late 1800’s that evidence became available that atoms were composed of smaller parts.
Slide22Thomson Model of the Atom
J. J. Thomson - English physicist 1897
Used a piece of equipment called a cathode ray tube.
It is a vacuum tube - all the air has been pumped out.
This was made available by prior scientists including William Crookes
Slide23Crookes Tube
William Crookes
Mask holder
Cathode
(-)
Anode
(+)
Crookes tube
(Cathode ray tube)
Mask holder
Glow
Slide24Thomson’s Experiment
+
-
voltage
source
OFF
ON
Passing an electric current makes a beam appear
to move from the negative to the positive end
Slide25Thomson’s Experiment
+
-
voltage
source
OFF
ON
+
-
By adding an electric field…
he found that the moving pieces were negative.
Slide26Cathode Ray Experiment
1897 Experimentation
Using a cathode ray tube, Thomson was able to deflect cathode rays with an electrical field.
The rays bent towards the positive pole, indicating that they are negatively charged.
Slide27Conclusions
He compared the value with the mass/ charge ratio for the lightest charged particle.
By comparison, Thomson estimated that the
cathode ray particle
weighed 1/1000 as much as hydrogen, the lightest atom.
He concluded that
atoms do contain subatomic particles
- atoms are divisible into smaller particles.
This conclusion contradicted Dalton’s postulate and was not widely accepted by fellow physicists and chemists of his day.
Since any electrode material produces an identical ray, cathode ray particles are present in all types of matter - a universal negatively charged subatomic particle later named the electron
Slide28J.J. Thomson
He proved that atoms of any element can be made to emit tiny negative particles.
From this he concluded that ALL atoms must contain these negative particles.
He knew that atoms did not have a net negative charge and so there must be balancing of the negative charge.
J.J. Thomson
Slide29William Thomson (Lord Kelvin)
& J.J. Thomson
Lord Kelvin with other scientists came up with several models in early 1900s but most writings say…
In 1910, J.J. Thomson proposed the Plum Pudding model
Negative electrons were embedded into a positively charged spherical cloud.
Slide30Thomson Model of the Atom
J.J. Thomson discovered the electron and knew that electrons could be emitted from matter (1897).
Thomson proposed that atoms consist of small, negative electrons embedded in a massive, positive sphere.
The electrons were like currants in a plum pudding.
This is called the ‘plum pudding’ model of the atom.
-
electrons
-
-
-
-
-
-
-
Slide31Millikan’s Oil
Drop Experiment
Slide32Actual Apparatus used by Millikan
in his Oil Drop Experiment
Slide33Millikan Oil Drop Experiment
A fine mist of oil droplets is introduced into the chamber as the gas molecules inside the chamber are ionized (split into electrons and positive ions) by a beam of x-rays (not represented)
The electrons adhere to the oil droplets, some droplets having one electron, some two electrons, and so forth
These negatively charged oil droplets fall under the force of gravity into the region between the electrically charged plates
If you carefully adjust the voltage on the plates, the force of gravity can be exactly counterbalanced by the attractive force between the negative oil drop and upper, positively charged plate
Analysis of these forces leads to a value for the charge on thee electron
Robert Andrews Millikan (1868 – 1953) won the Nobel physics prize in 1923 for his work in isolating and weighing the electron
Slide34Oil Drop Experiment
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oil droplets
oil droplet
under observation
Charged plate
Small hole
Charged plate
-
+
Telescope
oil atomizer
Robert Millikan
(1909)
Balancing electrical and gravitational forces allowed the electron charge to be determined.
Mass was calculated using charge to mass ratio (9.1093 x 10
-28
g).
Slide35Rutherford
Model of the Atom
Slide36Ernest Rutherford (1871-1937)
Learned physics in J.J. Thomson’s lab.
Noticed that ‘alpha’ particles were sometime deflected by something in the air.
Gold-foil experiment
Slide37Rutherford ‘Scattering’
In 1909 Rutherford undertook a series of experiments
He fired
a
(alpha)
particles at a very thin sample of gold foil
According to the Thomson model the
a
particles would only
be slightly deflectedRutherford discovered that they were deflected through large angles and could even be reflected straight back to the source
particle
source
Lead collimator
Gold foil
a
q
Slide38Rutherford’s Apparatus
beam of alpha particles
radioactive
substance
gold foil
circular ZnS - coated
fluorescent screen
Rutherford received the
1908 Nobel Prize in Chemistry
for his pioneering work in nuclear chemistry.
Slide39What he expected…
Slide40What he got…
richocheting
alpha particles
Slide41Density and the Atom
Since most of the particles went through, the atom was mostly empty.
Because the alpha rays were deflected so much, the positive pieces it was striking were heavy.
Small volume and big mass = big density
This small dense positive area is the
nucleus
Slide42Explanation of Alpha-Scattering Results
Plum-pudding atom
+
+
+
+
+
+
+
+
-
-
-
-
-
-
-
-
Alpha particles
Nuclear atom
Nucleus
Thomson’s model
Rutherford’s model
Slide43Rutherford Model
In
the early twentieth century, Rutherford showed that
most of an atom’s mass is concentrated in a small,
positively charged
region called the
nucleus
.
Nucleus
Electron
Slide44Niels Bohr
In the Bohr Model (1913) the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun.
They are not confined to a planar orbit like the planets are.
Slide45Bohr Model
After Rutherford’s discovery, Bohr proposed that
electrons travel in definite
orbits
around the nucleus.
Planetary
model
Niels
Bohr
Slide46Bohr Model
The Bohr Model will be continued and explained in much greater detail next unit with the electron cloud
Quantum Model:
Electron Cloud – Physicists contributions to layers of sub-orbitals
Standard Model:
Quarks, Leptons, Higgs-Boson, etc.
Slide47Particles in the Atom
Slide48Evidence for Particles
In 1886, Goldstein, using equipment similar to cathode ray tube, discovered particles with charge equal and opposite to that of electron, but much larger mass
.
J.J. Thomson (~1897) discovered the electron
Rutherford later (1911) found
“Goldstein’s particles”
to be identical to hydrogen atoms minus one
electronnamed these particles
protonsChadwick (1932) discovered particles with similar mass to proton but
zero charge.discovered neutrons
Slide49Discovery of the Neutron
James Chadwick bombarded beryllium-9 with alpha particles,
carbon-12 atoms were formed, and neutrons were emitted.
+
+
Slide50
Electrons
(-) charge no mass located outside the nucleus
Protons
(+) charge 1 amu located inside the nucleus
Neutrons
no charge 1 amu located inside the nucleus
Particles in the Atom
Slide51
Atoms
consist of
electrons
,
protons
, and
neutrons.
1. Electrons and protons have electrical charges that are identical in magnitude but opposite in sign. Relative charges of
1 and +1 are assigned to the electron and proton, respectively.
2. Neutrons have approximately the same mass as protons but no charge—they are electrically neutral.
3. The mass of a proton or a neutron is about 1836 times greater than the mass of an electron. Protons and neutrons constitute the bulk of the mass of the atom.
Particles in the Atom
Slide52Subatomic particles
Electron
Proton
Neutron
Name
Symbol
Charge
Relative
mass
Actual
mass (g)
e
-
p
+
n
o
-1
+1
0
1/1840
1
1
9.11 x 10
-28
1.67 x 10
-24
1.67 x 10
-24
Slide53Structure of the Atom
There are two regions
The nucleus
With protons and neutrons
Positive charge
Almost all the mass
Electron cloud
Most of the volume of an atom
The region where the electron can be found
Slide54Counting the Pieces
Atomic Number
= number of protons
# of protons determines kind of atom
Atomic Number
= number of electrons in a neutral atom
Mass Number
= the number of protons + neutrons
12
6
C
14
6
C
12
6
C
Mass Number = A
Atomic Number = Z
Slide55Mass Number
Mass Number = Protons + Neutrons
Always a whole number
NOT on the
Periodic Table!
+
+
+
+
+
+
Nucleus
Electrons
Nucleus
Neutron
Proton
Carbon-12
Neutrons 6
Protons 6
Electrons 6
e-
e-
e-
e-
e-
e-
Slide56Symbols
The symbol of the element, the mass number and the atomic number
X
Mass
number
Atomic
number
# protons
# protons
+ # neutrons
mass number
Slide57Symbols
Find the
number of protons
number of neutrons
number of electrons
Atomic number
Mass number
F
19
9
= 9
= 10
= 9
= 9
= 19
+
Slide58Electron Number (atom’s reactivity)
If the symbol or atom is not specified as positive or negative, then it has a no overall charge and is thus neutral
This means it has equal number of protons and electrons
If they are not equal, then the atom will have either a net positive or negative charge
These are IONS
Positive ions are called
Cations
Negative ions are call Anions
Slide59Symbols
Find the
number of protons
number of neutrons
number of electrons
Atomic number
Mass number
Br
-1
80
35
= 35
= 45
=
36
= 35
= 80
What kind of atom is this?
What kind of ion is it?
Slide60Isotopes
Atoms of the same element with different mass numbers.
Put the mass number after the name of the element
carbon- 12
carbon -14
uranium-235
Mass #
Atomic #
Nuclear symbol:
Hyphen notation:
carbon-12
12
6
C
Slide61Isotopes
+
+
+
+
+
+
Nucleus
Electrons
Nucleus
Neutron
Proton
Carbon-12
Neutrons 6
Protons 6
Electrons 6
Nucleus
Electrons
Carbon-14
Neutrons 8
Protons 6
Electrons 6
+
+
+
+
+
+
Nucleus
Neutron
Proton
Slide62
Isotopes
Chlorine-37
atomic #:
mass #:
# of protons:
# of electrons:
# of neutrons:
17
37
17
17
20
Cl
37
17
37
17
Cl
Slide63Using a periodic table and what you know about atomic number, mass, isotopes, and electrons, fill in the chart:
Element
Symbol
Atomic
Number
Atomic
Mass
# of protons
# of neutron
# of electron
charge
8
8
8
Potassium
39
+1
Br
45
-1
30
65
30
Atomic Number = Number of Protons
Number of Protons + Number of Neutrons = Atomic Mass
Atom (no charge) : Protons = Electrons
Ion (cation) : Protons
>
Electrons
Ion (anion) : Electrons > Protons
Slide64Using a periodic table and what you know about atomic number, mass, isotopes, and electrons, fill in the chart:
Element
Symbol
Atomic
Number
Atomic
Mass
# of protons
# of neutron
# of electron
charge
8
8
8
Potassium
39
+1
Br
45
-1
30
65
30
Oxygen
Bromine
Zinc
O
K
Zn
8
19
35
16
80
19
35
30
20
3
5
18
36
0
0
Atomic Number = Number of Protons
Number of Protons + Number of Neutrons = Atomic Mass
Atom (no charge) : Protons = Electrons
Ion (cation) : Protons
>
Electrons
Ion (anion) : Electrons > Protons
A N S W E R K E Y
Slide65Relative Atomic Mass
12
C atom = 1.992 × 10
-23
g
1 p =
1.007276
amu
1 n =
1.008665
amu
1 e- =
0.0005486 amu
atomic mass unit (
amu)
1 amu =
1/12 the mass of a
12C atom
+
+
+
+
+
+
Nucleus
Electrons
Nucleus
Neutron
Proton
Carbon-12
Neutrons 6
Protons 6
Electrons 6
Slide66
Isotopes
-> Weighted Atomic Mass
Because of the existence of isotopes, the mass of a collection of atoms has an average value.
Average mass =
ATOMIC WEIGHT
Boron is 20% B-10 and 80% B-11.
That is, B-11 is 80 percent abundant on earth.
For boron atomic weight
= 0.20 (10
amu
) + 0.80 (11
amu) = 10.8 amu
These values can be seen on the periodic table, which is the weighted average of all isotopes for each element.
Slide67Atomic Mass
Calculate the atomic mass of copper if copper has two isotopes. 69.1% has a mass of 62.93
amu
and the rest has a mass of 64.93
amu
.
Cu
29
63.548
Isotope
Percent
Abundance
Mass
A.A.M.
Cu-63
69.1
62.93
Cu-65
64.93
43.48463
20.06337
30.9
63.548
Slide68Challenge Problem
Assume you have only two atoms of chlorine.
One atom has a mass of 35 amu (Cl-35)
The other atom has a mass of 36 amu (Cl-36)
What is the average mass of these two isotopes?
35.5 amu
Looking at the average atomic mass printed on the periodic table...
approximately what percentage is Cl-35 and Cl-36?
55% Cl-35 and 45% Cl-36 is a good approximation
Cl
35.453
17
Slide69Mass Spectrophotometer
electron
beam
magnetic field
gas
stream
of ions of
different
masses
lightest
ions
heaviest
ions
Slide70Mass Spectrometry
-
+
Photographic plate
196
199
201
204
198 200
202
Mass spectrum of mercury vapor
Stream of positive ions
The first mass spectrograph was built in 1919 by
F. W. Aston
, who received the
1922 Nobel Prize
for this accomplishment
Slide71100
90
80
70
60
50
40
30
20
10
0
34
35
36
37
Abundance
Mass
Mass spectrum of chlorine. Elemental chlorine (Cl
2
) contains
only two isotopes: 34.97 amu (75.53%) and 36.97 (24.47%)
AAM = (34.97
amu
)(0.7553) + (36.97
amu
)(0.2447)
AAM = (26.412841
amu
) + (9.046559
amu
)
AAM = 35.4594
amu
Cl-35
Cl-37
Cl
35.4594
17
Slide72Mass spectrums reflect the abundance of naturally occurring
isotopes
Hydrogen
Carbon
Nitrogen
Oxygen
Sulfur
Chlorine
Bromine
1
H = 99.985%
2
H = 0.015%
12
C = 98.90%
13
C = 1.10%
14
N = 99.63%
15N = 0.37%
16O = 99.762% 17O = 0.038%
18O = 0.200%
32S = 95.02% 33S = 0.75%
34
S = 4.21% 36S = 0.02%
35Cl = 75.77% 37Cl = 24.23%
79Br = 50.69% 81Br = 49.31%
Natural Abundance of Common Elements
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