Transcript nucleus - Gordon State College

Ch 2. Atoms and Elements
Nucleus (Rutherford’s experiment)
Atom
Electrons (Thomson’s experiment)
Atomic Nucleus
Nucleus carries almost all the mass of an atom
Nucleus carries positive charges
Each electron carries one negative charge
Protons
(each carries a positive charge)
Nucleus
Atom
Neutrons (neutral)
Electrons (each carries a negative charge)
Atoms are neutral
number of electrons = number of protons
= atomic number
Atoms that have the same number of protons belong to
one kind of element.
isotopes: same number of protons, different number of neutrons
mass of a proton ≈ mass of a neutron >> mass of an electron
number of protons + number of neutrons = mass number
Mass number
A
Z
X
Atomic number
or
X-A
Chemical symbol
Practice – Complete the table
27
13 Al
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Practice – Complete the table
13
6C
96
42Mo
27
13 Al
133
55 Cs
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No change occurs inside a nucleus in chemistry
Atoms can lose or gain electrons
Na − e−  Na+
positive ion = cation
Mg − 2e−  Mg2+
Cl + e−  Cl−
O + 2e−  O2−
negative ion = anion
Practice – Complete the table
Al3 
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Practice – Complete the table
S
2
Mg2
3
Al
Br 
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Symbol
Number of Number of
Number of
Protons in Neutrons
Net charge
Electrons
Nucleus in Nucleus
87Rb+
16
18
36
2−
28
1+
Symbol
Number of Number of
Number of
Protons in Neutrons
Net charge
Electrons
Nucleus in Nucleus
87Rb+
37
50
36
1+
32S2−
16
18
18
2−
65Cu+
29
36
28
1+
isotopes: same number of protons, different number of neutrons
mass of a proton ≈ mass of a neutron >> mass of an electron
number of protons + number of neutrons = mass number
1H
1.6735 x 10−24 g
16O
2.6560 x 10−23 g
One atomic mass unit (amu) is defined as 1/12 of
the mass of a 12C atom.
12C
atom: 6 protons, 6 neutrons, 6 electrons.
1 amu = 1.6605 x 10−24 g
1H
1.6735 x 10−24 g
16O
2.6560 x 10−23 g
1 amu = 1.6605 x 10−24 g
1.6735  10
24
1 amu
g
 1.0078 amu
24
1.6605  10 g
1 amu
2.6560  10 g 
 15.995 amu
24
1.6605  10 g
23
Mass Spectrometer
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Mass Spectrum of Natural Copper
69.09 %
63Cu
30.91 %
65Cu
natural abundance:
percent of an isotope
in nature
Mass Spectrum of Natural Copper
69.09 %
63Cu:
62.93 amu
30.91 %
65Cu:
64.93 amu
Average atomic mass of Cu = ?
How to find the average?
1, 1, 1, 1, 2
Average = (1 + 1 + 1 + 1 + 2) / 5 = 1.2
1  1  1  1  2 4 1  1 2

5
5
4
1
 1   2
5
5
 80 %  1  20 %  2
 0.8  0.4
 1.2
Mass Spectrum of Natural Copper
69.09 %
63Cu:
62.93 amu
30.91 %
65Cu:
64.93 amu
Average atomic mass of Cu = ?
63.55 amu
listed in the periodic table
35Cl
34.967 amu
75.78 %
37Cl
36.966 amu
24.22 %
Average atomic mass of Cl
= 34.967 amu x 75.78 % + 36.966 amu x 24.22 %
= 35.45 amu
listed in the periodic table
mass of a proton ≈ mass of a neutron ≈ 1 amu
mass number of an isotope
≈ atomic mass of the isotope in amu
35Cl
34.967 amu
75.78 %
37Cl
36.966 amu
24.22 %
One atomic mass unit (amu) is defined as 1/12 of
the mass of a 12C atom.
The number of carbon atoms in exactly 12 g of 12C
is called Avogadro’s number: 6.022 x 1023 (exact number)
One Avogadro’s number of particles is called a mole.
1 mol = 6.022 x 1023 particles
1 pair = 2 particles
1 dozen = 12 particles
1 mol = 6.022 x 1023 particles
Example 2.6, page 67
Calculate the number of atoms in 2.45 mol of Cu.
1 mol = 6.022 x 1023 particles
Practice 2.6, page 67
A pure Ag ring contains 2.80 x 1022 Ag atoms.
How many moles of Ag does it contain?
For an element X, its atomic mass is x amu.
What is the mass of 1 mol of X in grams?
The mass of 1 mol of X is x g.
The molar mass of an element is the mass in
grams per mole of the element.
Unit: g/mol
Two ways to find molar mass
1) Read from the periodic table
2) Use the definition of molar mass:
mass in grams
molar mass 
moles
(recall d = m/V)
1 mol = 6.022 x 1023 particles
mass in grams
molar mass 
moles
Unit: g/mol
molar mass and Avogadro’s number are exact numbers
A piece of Cu has a mass of 200. g. How many copper
atoms are present?
1 mol = 6.022 x 1023 particles
molar mass 
mass in grams
moles
A silicon chip has a mass of 5.68 mg. How many silicon
atoms are present in the chip?
1 mol = 6.022 x 1023 particles
molar mass 
mass in grams
moles
Compute the mass in grams of a sample of six
Americium atoms.
1 mol = 6.022 x 1023 particles
mass in grams
molar mass 
moles
Calculate the number of moles in a sample of cobalt (Co)
containing 5.00 x 1020 atoms and the mass of the sample.
1 mol = 6.022 x 1023 particles
molar mass 
mass in grams
moles
Mendeleev
• Ordered elements by atomic mass
• Saw a repeating pattern of properties
• Periodic Law – when the elements are arranged
in order of increasing atomic mass, certain sets
of properties recur periodically
• Put elements with similar properties in the same
column
• Used pattern to predict properties of
undiscovered elements
• Where atomic mass order did not fit other
properties, he re-ordered by other properties
– Te & I
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Periodic Pattern
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Periodic Patterns
NM
H2O
a/b
H
1.0
H2
M
Li
6.9
M
Li2O
b
M
LiH
9.0
Be
Na2O M
b
NaH 24.3
M
K2O
b
39.1
KH
M
Mg
23.0
B
BeH2 10.8
MgO
b
Na
K
BeO NM
a/b
M
Al
MgH2 27.0
CaO
b
B2O3 NM
a
C
CO2 NM
a
N
N2O5 NM
a
F
NH3 16.0
H2O 19.0
HF
Al2O3 M/NM
a/b
SiO2 NM
a
P4O10 NM
a
SO3 NM
a
Cl2O7
a
AlH3 28.1
SiH4 31.0
Si
P
PH3 32.1
M = metal, NM = nonmetal, M/NM = metalloid
a = acidic oxide, b = basic oxide, a/b = amphoteric oxide
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O
CH4 14.0
CaH2
42
NM
BH3 12.0
Ca
40.1
O2
S
H2S 35.5
Cl
HCl
Most
About
A
fewofelements
¾
the
of remaining
the elements
are classified
elements
are classified
asare
metalloids.
classified
as metals.
Their
as nonmetals.
They
solidshave
havea
Their
reflective
some
solids
characteristics
surface,
have a conduct
non-reflective
of metals
heatand
and
surface,
some
electricity
of
dononmetals.
not
better
conduct
thanheat
other
and
elements,
electricity
andwell,
are malleable
and are brittle.
and ductile
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Metals
• Solids at room temperature, except Hg
• Reflective surface
– shiny
• Conduct heat
• Conduct electricity
• Malleable
– can be shaped
• Ductile
– can be drawn or pulled into wires
• Lose electrons and form cations in
reactions
• About 75% of the elements are metals
•44 Lower left on the table
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Sulfur, S(s)
Nonmetals
• Found in all three states
• Poor conductors of heat
• Poor conductors of
electricity
• Solids are brittle
• Gain electrons in reactions
to become anions
• Upper right on the table
– except H
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Bromine, Br2(l)
Chlorine, Cl2(g)
Metalloids
• Show some
properties of
metals and some
of nonmetals
• Also known as
semiconductors
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Properties of Silicon
shiny
conducts electricity
does not conduct heat well
brittle
The Modern Periodic Table
• Elements with similar chemical and
physical properties are in the same
column
• Columns are called Groups or Families
– designated by a number and letter at top
• Rows are called Periods
• Each period shows the pattern of
properties repeated in the next period
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The Modern Periodic Table
• Main group = representative elements = “A”
groups
• Transition elements = “B” groups
– all metals
• Bottom rows = inner transition elements =
rare earth elements
– metals
– really belong in Period 6 & 7
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= Alkali metals
= Halogens
= Alkali earth metals
= Lanthanides
= Noble gases
= Actinides
= Transition metals
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Important Groups - Hydrogen
• Nonmetal
• Colorless, diatomic gas
– very low melting point and density
• Reacts with nonmetals to form molecular
compounds
– HCl is acidic gas
– H2O is a liquid
• Reacts with metals to form hydrides
– metal hydrides react with water to form H2
• HX dissolves in water to form acids
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Important Groups – Alkali Metals
• Group IA = Alkali Metals
• Hydrogen usually placed here,
though it doesn’t really belong
• Soft, low melting points, low
density
• Flame tests  Li = red, Na =
yellow, K = violet
• Very reactive, never find
uncombined in nature
• Tend to form water-soluble
compounds, therefore salt is
crystallized from seawater then
molten salt is electrolyzed
• colorless solutions
• React with water to form basic
(alkaline) solutions and H2
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2 Na + 2 H2O  2 NaOH + H2
• releases a lot of heat
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lithium
sodium
potassium
rubidium
cesium
Important Groups – Alkali Earth
Metals
• Group IIA = Alkali earth metals
• Harder, higher melting, and denser than alkali
metals
– Mg alloys used as structural materials
• Flame tests  Ca = red, Sr = red, Ba = green
• Reactive, but less than corresponding alkali metal
• Form stable, insoluble oxides from which they are
normally extracted
• Oxides are basic = alkaline earth
• Reactivity with water to form H2
– Be = none; Mg = steam; Ca, Sr, Ba = cold water
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Important Groups – Halogens
• Group VIIA = halogens
• Nonmetals
• F2 and Cl2 gases; Br2 liquid; I2
solid
• All diatomic
• Very reactive
• Cl2, Br2 react slowly with water
Br2 + H2O  HBr + HOBr
• React with metals to form ionic
compounds
• HX all acids
– HF weak < HCl < HBr < HI
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fluorine
chlorine
bromine
iodine
astatine
Important Groups – Noble
Gases
• Group VIIIA = Noble Gases
• All gases at room temperature
– very low melting and boiling points
• Very unreactive, practically inert
• Very hard to remove electron from or give
electron to
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Ion Charge and the Periodic Table
• The charge on an ion can often be
determined from an element’s position
on the Periodic Table
• Metals always form positively charged
cations
• For many main group metals, the charge
= the group number
• Nonmetals form negatively charged
anions
• For nonmetals, the charge = the group
number − 8
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Practice – What is the charge
on each of the following ions?
•
•
•
•
•
potassium cation
sulfide anion
calcium cation
bromide anion
aluminum cation
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K+
S2−
Ca2+
Br−
Al3+
Law of Conservation of Mass
• In a chemical reaction,
•
matter is neither created
nor destroyed
Total mass of the materials
you have before the
reaction must equal the
total mass of the materials
you have at the end
total mass of reactants =
total mass of products
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Antoine Lavoisier
1743-1794
Reaction of Sodium with
Chlorine to Make Sodium
Chloride
The mass of sodium and chlorine used is determined by the number of atoms
•
that combine
Because only whole atoms combine and atoms are not changed or destroyed
in the process, the mass of sodium chloride made must equal the total mass of
sodium and chlorine atoms that combine together
•
7.7 g Na
+ 11.9 g Cl2
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 19.6 g NaCl
Law of Definite Proportions
Joseph Proust
1754-1826
• All samples of a given compound,
regardless of their source or how they were
prepared, have the same proportions of their
constituent elements
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Proportions in Sodium Chloride
A 100.0 g sample of sodium
chloride contains 39.3 g of
sodium and 60.7 g of
chlorine
A 200.0 g sample of sodium
chloride contains 78.6 g of
sodium and 121.4 g of
chlorine
A 58.44 g sample of sodium
chloride contains 22.99 g of
sodium and 35.44 g of
chlorine
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Law of Multiple Proportions
John Dalton
1766-1844
• When two elements (call them A and B) form
two different compounds, the masses of B
that combine with 1 g of A can be expressed
as a ratio of small, whole numbers
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Oxides of Carbon
• Carbon combines with oxygen to form two
different compounds, carbon monoxide
and carbon dioxide
• Carbon monoxide contains 1.33 g of
oxygen for every 1.00 g of carbon
• Carbon dioxide contains 2.67 g of
oxygen for every 1.00 g of carbon
• Because there are twice as many
oxygen atoms per carbon atom in
carbon dioxide of in carbon monoxide,
the oxygen mass ratio should be 2
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Dalton’s Atomic Theory
•
1.
2.
3.
4.
Dalton proposed a theory of matter based on it
having ultimate, indivisible particles to explain
these laws
Each element is composed of tiny, indestructible
particles called atoms
All atoms of a given element have the same
mass and other properties that distinguish them
from atoms of other elements
Atoms combine in simple, whole-number ratios
to form molecules of compounds
In a chemical reaction, atoms of one element
cannot change into atoms of another element

they simply rearrange the way they are attached
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Practice – Decide if each statement is
correct according to Dalton’s model of the
atom
• Copper atoms can combine with zinc
atoms to make gold atoms
• Water is composed of many identical
molecules that have one oxygen atom and
two hydrogen atoms
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Practice – Decide if each statement is correct
according to Dalton’s model of the atom
• Copper atoms can combine with zinc atoms to make gold
atoms – incorrect; according to Dalton, atoms of one
element cannot turn into atoms of another element by a
chemical reaction. He knew this because if atoms could
change it would change the total mass and violate the Law
of Conservation of Mass.
• Water is composed of many identical molecules that have
one oxygen atom and two hydrogen atoms – correct;
according to Dalton, atoms combine together in
compounds in small whole-number ratios, so that you
could describe a compound by describing the number of
atoms of each element in a molecule. He used this idea to
explain why compounds obey the Law of Definite
67 Proportions.
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Practice – Decide if each statement is
correct according to Dalton’s Model of the
Atom
• Some carbon atoms weigh more than
other carbon atoms
• Because the mass ratio of Fe:O in wüsite
is 1.5 times larger than the Fe:O ratio in
hematite, there must be 1.5 Fe atoms in a
unit of wüsite and 1 Fe atom in a unit of
hematite
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Practice – Decide if each statement is
correct according to Dalton’s model of the
atom
• Some carbon atoms weigh more than other carbon
atoms – incorrect; according to Dalton, all atoms of an
element are identical.
• Because the mass ratio of Fe:O in wüsite is 1.5 times
larger than the Fe:O ratio in hematite, there must be 1.5
Fe atoms in a unit of wüsite and 1 Fe atom in a unit of
hematite – incorrect; according to Dalton, atoms must
combine in small whole-number ratios. If you could
combine fractions of atoms, that would mean the atom is
breakable and Dalton’s first premise would be incorrect.
You can get the Fe:Fe mass ratio to be 1.5 if the formula
for wüsite is FeO and the formula for hematite is Fe2O3.
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Cathode Ray Tube
•
•
Glass tube containing metal electrodes from
which almost all the air has been evacuated
When connected to a high
voltage power supply,
a glowing area is
seen emanating from
the cathode
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J.J. Thomson
• Believed that the cathode ray was
composed of tiny particles with an
electrical charge
• Designed an experiment to demonstrate
that there were particles by measuring the
amount of force it takes to deflect their
path a given amount
– like measuring the amount of force it takes to
make a car turn
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Thomson’s Experiment
Investigate the effect of placing an electric field around tube
1. charged matter is attracted to an electric field
2. light’s path is not deflected by an electric field
+++++++++++
Cathode
Anode
(+)
(-)
-------------
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+
Power Supply
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Thomson’s Results
• The cathode rays are made of tiny particles
• These particles have a negative charge
– because the beam always deflected toward the + plate
• The amount of deflection was related to two
factors, the charge and mass of the particles
• Every material tested contained these same
particles
• The charge:mass ratio of these particles was
−1.76 x 108 C/g
– the charge/mass of the hydrogen ion is +9.58 x 104 C/g
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Thomson’s Conclusions
• If the particle has the same amount of
charge as a hydrogen ion, then it must
have a mass almost 2000x smaller than
hydrogen atoms!
– later experiments by Millikan showed that the
particle did have the same amount of charge
as the hydrogen ion
• The only way for this to be true is if these
particles were pieces of atoms
– apparently, the atom is not unbreakable
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Millikan’s Oil Drop Experiment
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Thomson’s Conclusions, cont’d
• Thomson believed that these particles were
therefore the ultimate building blocks of
matter
– “We have in the cathode rays matter in a new
state, a state in which the subdivision of matter
is carried very much further . . . a state in which
all matter . . . is of one and the same kind; this
matter being the substance from which all the
chemical elements are built up.”
• These cathode ray particles became known
as electrons
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Electrons
• Electrons are tiny, negatively charged
particles found in all atoms
• Cathode rays are made of streams of
electrons
• The electron has a charge of −1.60 x
1019 C
• The electron has a mass of 9.1 x 10−28 g
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A New Theory of the Atom
• Because the atom is no longer indivisible,
Thomson must propose a new model of
the atom to replace the first statement in
Dalton’s Atomic Theory
– rest of Dalton’s theory still valid at this point
• Thomson proposes that instead of being a
hard, marble-like unbreakable sphere, the
way Dalton described it, the atom actually
had an inner structure
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Thomson’s Plum Pudding Atom
• The structure of the atom contains
many negatively charged electrons
• These electrons are held in the
atom by their attraction for a
positively charged electric field
within the atom
– there had to be a source of positive
charge because the atom is neutral
– Thomson assumed there were no
positively charged pieces because
none showed up in the cathode ray
experiment
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Predictions of the
Plum Pudding Atom
• The mass of the atom is due to the mass
of the electrons within it
– electrons are the only particles in Plum
Pudding atoms, therefore the only source of
mass
• The atom is mostly empty space
– should not have a bunch of negatively
charged particles near each other as they
would repel
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Radioactivity
• In the late 1800s, Henri Becquerel and
Marie Curie discovered that certain
elements would constantly emit small,
energetic particles and rays
• These energetic particles could penetrate
matter
• Ernest Rutherford discovered that there
were three different kinds of emissions
– alpha, a, rays made of particles with a mass
4x H atom and + charge
– beta, b, rays made of particles with a mass
~1/2000th H atom and – charge
– gamma, g, rays that are energy rays, not
particles
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Marie Curie
1867-1934
Rutherford’s Experiment
• How can you prove something is empty
space?
• Put something through it!
– use large target atoms
• use very thin sheets of target so it will not absorb “bullet”
– use very small particle as bullet with very high
energy
• but not so small that electrons will affect it
• Bullet = alpha particles, target atoms = gold
foil
83
– a particles have a mass of 4 amu & charge of +2
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Rutherford’s Results
• Over 98% of the a particles went straight
through
• About 2% of the a particles went through
but were deflected by large angles
• About 0.005% of the a particles bounced
off the gold foil
– “...as if you fired a 15” cannon shell at a piece
of tissue paper and it came back and hit you.”
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Rutherford’s Conclusions
• Atom mostly empty space
– because almost all the particles went straight
through
• Atom contains a dense particle that is
small in volume compared to the atom but
large in mass
– because of the few particles that bounced
back
• This dense particle is positively charged
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– because of the large deflections of some of
the particles
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Plum Pudding
Atom
•
•
•
•
•
•
•
•
•
•
A few of the
a particles
do not go through
•
•
•
•
•
•
•
•
•
•
•
If atom was like
a plum pudding,
all the a particles
should go
straight through
•
Nuclear Atom
.
.
.
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Almost all a particles
go straight through
Some a particles
go through, but are deflected due to
+:+ repulsion from the nucleus
Rutherford’s Interpretation –
the Nuclear Model
1. The atom contains a tiny dense center called
the nucleus
–
the amount of space taken by the nucleus is only
about 1/10 trillionth the volume of the atom
2. The nucleus has essentially the entire mass of
the atom
–
the electrons weigh so little they give practically no
mass to the atom
3. The nucleus is positively charged
–
the amount of positive charge balances the negative
charge of the electrons
4. The electrons are dispersed in the empty space
of the atom surrounding the nucleus
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Structure of the Nucleus
• Rutherford proposed that the nucleus
had a particle that had the same
amount of charge as an electron but
opposite sign – these particles are
called protons
– based on measurements of the nuclear
charge of the elements
• protons are subatomic particles found
in the nucleus with a charge = +1.60 x
1019 C and a mass = 1.67262 x 10−24 g
• Because protons and electrons have
the same amount of charge, for the atom
to be neutral there must be equal
numbers of protons and electrons
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Relative Mass and Charge
• It is sometimes easier to compare things to each other
rather than to an outside standard
• When you do this, the scale of comparison is called a
relative scale
• We generally talk about the size of charge on atoms by
comparing it to the amount of charge on an electron, which
we call −1 charge units
– proton has a charge of +1 cu
– protons and electrons have equal amounts of charge, but opposite
signs
• We generally talk about the mass of atoms by comparing it
to 1/12th the mass of a carbon atom with 6 protons and 6
neutrons, which we call 1 atomic mass unit
– protons have a mass of 1 amu
– electrons have a mass of 0.00055 amu, which is generally too small
to be relevant
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Some Problems
• How could beryllium have four protons stuck
together in the nucleus?
– shouldn’t they repel each other?
• If a beryllium atom has four protons, then it should
weigh 4 amu; but it actually weighs 9.01 amu!
Where is the extra mass coming from?
– each proton weighs 1 amu
– remember, the electron’s mass is only about 0.00055
amu and Be has only four electrons – it can’t account
for the extra 5 amu of mass
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Tro: Chemistry: A Molecular Approach, 2/e
There Must Be Something Else!
• To answer these questions, Rutherford
and Chadwick proposed that there was
another particle in the nucleus – it is called
a neutron
• Neutrons are subatomic particles with a
mass = 1.67493 x 10−24 g and no charge, and
are found in the nucleus
1 amu
slightly heavier than a proton
no charge
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Tro: Chemistry: A Molecular Approach, 2/e
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Tro: Chemistry: A Molecular Approach, 2/e
Problems for Chapter 2
• 51-72, 75-78, 79-90