Transcript History of the Atom Scientists and Their Discoveries 1

History of the Atom
Scientists and Their
Discoveries
1
Atoms
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+
+
+
+
-
-
• Atoms- smallest
possible unit into
which matter can be
divided, while still
maintaining
itsis the
For example, what
smallest possible unit
properties.
into which a long essay can be
divided and still have some meaning?
• Made
up of:
+
-
– protons
– neutrons
– electrons
• Electrically they are
2
NEUTRAL!!!!
Elements
• Element- made up of
one kind of atom
that can’t be broken
down into simpler
substances by
physical or chemical
means
• 90 occur naturally on
Earth
• 29 were synthesized
(made) by scientists
3
The Periodic Table of
Elements-Reference
• Periodic table tiles contain a lot of information
and to understand it, it’s necessary to know the
parts of the atom and some terminology
concerning them
8
O
Oxygen
15.999
Atomic Number (number
of protons)
Element Symbol (a capital
letter or a capital followed
by a lower case letter)
Element Name
Atomic Mass (weighted
average of all isotopes mass)
4
The Atom’s “Center”
• Nucleons- particles in the nucleus of atoms
– Protons
– Neutrons
Notice that the electrons are not a part of the nucleus
-
+
-
+ +
5
Neutrons
-
++ +
+ + +
+ +
-
-
-
-
-
• Neutrons- neutral
particles; have no
electric charge
– Help make up the
nucleus of the atom
– Contribute to the
atomic mass
1.67 X 10-24 g
6
Protons (+)
• Protons- positively
charged particles
– In nucleus
– They ID an atom—
atomic number
– Contribute to the mass
of the atom
1.67 X 10-24 g
– Charge = + 1.6 X 10-19
coulombs (same value as
electron but positive)
-
++ +
+ + +
+ +
-
-
-
-
-
+
7
Atomic Number
• Atomic number - the number of protons in
the nucleus of an atom…its ID #!!!!
• Represented by “Z”
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-
+
++
-
What would be the
atomic number of
the atom to the
left?
What element is it?
8
Electrons (-)
• Electrons- negatively
charged particles
– Outside the nucleus of the
atom in electron
orbits/levels
– Move rapidly and create an
electron cloud
– Mass is insignificant
– Valence electrons- the
outermost electrons involved
in the formation of chemical
bonds
-
++ +
+ + +
+ +
-
-
-
-
-
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Neutral Atoms
• Most atoms are neutral and the number of
protons = the number of electrons
• 1 +protons and 1 –electrons=0 (neutral)
• Atomic number = protons = electrons
10
Electromagnetic Force
• Electromagnetic force is the
force that results from the
+ +
repulsion of like charges and
the attraction of opposites
+
AND NEUTRALIZE
ONE ANOTHER
- • This the force that holds
Notice how the
the electrons around the
particles with the
nucleus
same charge move
• Bill Nye Atoms
apart and the
Why are neutrons not
pictured above?
particles with
different charges
move together.
11
Atomic Models
• Model – a familiar idea used to explain
unfamiliar facts observed in nature
• Theory- an explanation of observable
facts and phenomena
• To remain valid, models and theories
must:
• Explain all known facts
• Enable scientists to make correct
predictions
12
Democritus
(460 BC – 370 BC)
• Proposed the existence of
atoms from Greek word
“atomos” which means “not to
cut” or “indivisible”
• Thought you could cut
matter in half until you got
an indivisible (not dividable)
particle
Image taken from: https://reichchemistry.wikispaces.com/T.+Glenn+
Time+Line+Project
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Aristotle
(384 BC – 322BC)
• Rejected idea of the atom
• Said matter could be cut
continually
• Aristotle was more
influential than Democritus
so atoms were forgotten
about until late 1700’s
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Antoine Lavoisier
(1743 – 1794)




“Father of Modern Chemistry”
Generated a list of 33 elements
Devised the metric system
Discovered/proposed the Law
of Conservation of Mass
 - matter can’t be created or
destroyed, it just changes form
(beginning mass = end mass)
Image taken from:
www.ldeo.columbia.edu/.../v1001/geo
time2.html
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John Dalton
(1766 – 1844)
 In 1803 he proposed first
experimentally based
Atomic Theory that states
atoms:
o are building blocks of matter
o are indivisible
o of the same element are
identical
o of different elements are
different
o “Billiard Ball Model”
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John Dalton
(1766 – 1844)
o Dalton’s Atomic Theory also explained The
Law of Multiple Proportions -ratio of the
masses of combined elements are WHOLE
numbers which become subscripts for
chemical formulas
o Nitrogen and Oxygen combine to form NO or
NO2, but not NO1.5
17
Law of Definite Composition
• Law of Definite Composition (Proust’s Law)elements combine in a definite (constant)
ratio by atomic mass
– Water (H2O) is always 2 hydrogens for each
oxygen
– 16:2 oxygen:hydrogen mass ratio or 8:1 reduced
– For Carbon dioxide (CO2) there is always a 32:12
oxygen to carbon mass ratio, or 8:3 reduced.
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J.J. Thomson
(1856 – 1940)
Put electricity through a vacuum tube
and produced a beam that was negatively
charged
Cathode Ray Tube Experiment
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J.J. Thomson
(1856 – 1940)
• Credited with discovery of
electron; a blow to Dalton’s
indivisible atom idea
– “Plum Pudding Model”
– Also because atoms are neutral,
the negative electrons must be
embedded in a ball of positive
charge
20
Millikan Oil Drop
Experiment
• Calculated the mass and
quantified the charge
of electrons!
• Mass of electron =
9.1 X 10-28 grams
(0.000000000000000000
00000000091 g)
• Charge on the electron
= -1.6 X 10-19 coulombs
(unit of charge)
- Millikan Animation and
Interactive
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Millikan Oil Drop
Experiment-Reference
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JJ Thomson and E. Goldstein
• Realized if neutral atoms contain negative
electrons, they must contain positive
particles
• Used vacuum tube similar to discovery of
electron to discover protons
– Mass = 1.67 X 10-24 grams (much heavier than
electron!)
– Charge = + 1.6 X 10-19 coulombs (same value as
electron but positive)
Ernest Rutherford
(1871 – 1937)
• Gold-foil
experiment
 Positively
charged alpha
particles aimed
at thin gold foil,
but most passed
through
 A few were
deflected and
some even
bounced right
back
 Gold foil
experiment
24
Rutherford’s Gold Foil
Experiment-Reference
(1871 – 1937)
25
Ernest Rutherford
(1871 – 1937)
Conclusions:
Disproved Thomson because
showed most of atom is
empty space
Discovered dense, positively
charged core, or nucleus,
repels the + alpha particles
Protons are surrounded by
negatively charged electrons
“Planetary Model”
You’ll never see life the
same way again
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James Chadwick
(1891 – 1974)
 Discovered the atomic mass
of most elements was
double the number of
protons  discovery of the
neutron in 1932
 Worked on the Manhattan
Project
 Worked with Ernest
Rutherford
 Like many others before him,
he won a Nobel Prize
27
Review
Early Atomic Theory video 5 min
• Number your paper from 1-5 and answer
the following questions. Two will be
cumulative review!
– 1. Which of these has 3 significant
figures?
•
•
•
•
a. 3340
b. 3.340
c. 0.001334
d. 334.00
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Review
• A
• 2. Which of these is a homogenous
mixture?
•
•
•
•
a. salt
b. iced tea
c. pizza
d. your computer
29
Review
• B
• 3. Which of these is true about subatomic
particles?
• a. electrons are negatively charged and in
the nucleus
• b. protons are negatively charged and in the
nucleus
• c. protons are positively charged and fly
around the outside of the nucleus
• d. neutrons are neutral and are in the
nucleus
30
Review
• D
• 4. Which of these is true about the
discovery of Millikan’s oil drop experiment
– a. He discovered the electron
– b. He discovered the mass of the neutron
– c. He discovered the mass and the charge of
the electron
– d. He discovered the proton
31
Review
• C
• 5. Which of these is false?
–
–
–
–
a. Neutrons are neutral
b. Protons are positive and two will repel
c. Electrons are negative and two will attract
d. Protons are positive and they will attract
negative electrons
32
Review
• C
33
Mass of Sub-Atomic ParticlesReference
(protons, neutrons, electrons)
Neutron = 1.6749286 x10-24 g
Proton = 1.6726231 x10-24 g
Electron = 9.1093897 x10-28 g
- - - - - - - - - - - - - -
- - - - - - - - - - - - - -
1839 electrons = 1 neutron
+
1836 electrons = 1 proton
+
How do you think the mass of a neutron
compares to that of a proton?
1 neutron ≈ 1 proton ≈ 1.67 x10-24 g
34
Mass Number
• Mass number – number of particles of
significant mass in the atom
• Represented by “A”
protons + neutrons = mass number
• Electrons are NOT included, their mass is zero
• NOT found directly on the periodic table!
Particle
Charge
Mass
number
Proton
1
Neutron
1
Electron
0
Location
in atom
35
Let’s Do It!!!
What would be the mass
number of this atom?
+
-
 3
 4
+
++
3 protons + 4 neutrons =
a mass number of 7
Why did we not account for the
electrons when calculating the
mass number?
-
36
Calculating the Actual Mass
of 1 Atom
• Actual mass of an atom is determined by the
protons and neutrons (electrons have virtually
no mass)
• Each proton and neutron mass= ~1.67 x 10-24 g
• Ex: a hydrogen atom has 1 electron and 1
proton:
–
–
–
–
Proton =
1.67 x10-24 g
No neutrons
0g
Electron =
+
0g
Mass of the entire hydrogen atom = 1.67 x10-24 g
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Calculating the Actual Mass of
1 Atom-Reference
• The actual mass of an atom=
(#protons + neutrons)(mass of p’s and n’s)=
(#protons + neutrons) (1.67 x 10-24 g) =
don’t count sd’s
• Ex: What’s the actual mass of a Lithium
atom with 3 protons and 4 neutrons?
(#protons + neutrons)(mass of p’s and n’s)=
(7)(1.67 x 10-24g) =
1.17 x 10-23 g 38
Relate Actual Mass to Mass
Number
• We can say the actual mass of an atom=
(protons + neutrons)(mass of p’s and n’s)=
OR
(mass number) (1.67 x 10-24 g) =
don’t count sd’s
39
Let’s Do It!!!
• What’s the actual mass of a Carbon atom
with 6 protons and 8 neutrons?
40
Let’s Do It!!!
• (14)(1.67 x 10-24g) = 2.34 x 10-23 g
41
Calculating the Actual Mass of
1 Atom
• Actual mass of an atom based on the idea
that the whole atom is equal to the sum of
the parts
• Not exactly correct because binding
energy is needed to hold the parts of an
atom together
• Some mass converted to this binding energy
in a nuclear reaction so the calculation gives
a value that is a little larger than reality
• (E = mc2)
Isotopes
• What mass numbers do these atoms have?
• What elements are these?
• How do you know?
43
Isotopes
• There mass numbers are 1, 2, and
3 but they all have one proton and
therefore are all hydrogen!
• So what is going on? They are isotopes.
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Isotopes
• Isotopes – different versions of atoms of
an element that have same # of protons but
different # of neutrons
• This discovery disproved one of Dalton’s
idea that atoms of the same element are
exactly alike!
45
Isotopes
– Hydrogen-1 (Protium): 1 proton, no neutrons and
is most common
– Hydrogen-2 (Deuterium): 1 proton and 1 neutron
– Hydrogen-3 (Tritium): 1 proton and 2 neutrons
– All versions of hydrogen!
46
Hyphen Notation-Reference
• We use mass numbers to distinguish between
isotopes because they differ in their number
of neutrons
• Hydrogen-1 =1 proton + 0 neutron=mass # 1
• Hydrogen-2 = 1 proton + 1 neutron =mass # 2
• Hydrogen-3 = 1 proton + 2 neutrons=mass # 3
• This is written in hyphen notation
47
Nuclear Symbols-Reference
• Nuclear symbols are a way to write atoms
using the mass number and atomic number
• Format:
• Hydrogen-1
1
Hydrogen-2
2
H
1
Hydrogen-3
3
H
1
H
1
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Let’s Do It!!!
• Naturally occurring carbon consists of three
isotopes, Carbon-12, Carbon-13, and Carbon14. State the number of protons, neutrons,
and electrons in each of these carbon atoms
• Think-Pair-Share
12C
13C
14C
6
6
6
#p _______
_______
_______
#n _______
_______
_______
#e _______
_______
_______
49
Review
Atomic Number video 9 min
• Number your paper from 1-5 and answer
the following questions. Two will be
cumulative review!
– 1. Which of these is proper scientific
notation?
• a. 4.56 X 106
• b. 4.5 X 32
• c. 45.6 X 1010
• d. 456 X 106
50
Review
• A
• 2. Which of these would be the correct
answer with the proper number of
significant digits if you multiplied 0.05 X
3.01
–
–
–
–
a. 0.1505
b. 0.150
c. 0.15
d. 0.2
51
Review
• D Remember beginning zeros are never
significant, so 0.05 has only 1!
• 3. What does this mean? 14C
6
– a. this atom has 6 neutrons and 20 electrons
– b. this atom has 6 protons and 8 neutrons for
a combined mass number of 14
– c. this atom has 6 protons and 14 electons
– d. this atoms has 6 electrons and 14 neutrons
52
Review
• B
• 4. Which of these is a correct definition
of an isotope?
– a. different versions of an element that have a
different number of neutrons
– b. atoms of the same element with the same
atomic number but different mass number
– c. different version of an element that have
the different number of electrons
– d. A and B
53
– e. B and C
Review
• D
• 5. How do you calculate the actual mass of
an atom
– a. add up all the protons, electrons, and
neutrons
– b. add up all the protons and neutrons and
divide by two
– c. add up all the protons and neutrons and
multiply by 1.67 x 10-24 g
– d. add up all the protons and neutrons
54
Review
• C
55
Actual Mass of 1 AtomReference
• The actual mass of an atom is a super small
number and is cumbersome in calculations, so
scientists assigned a relative scale to the mass
of these particles and created a new unit
called “atomic mass unit”, or amu ()
• amu () = atomic mass unit
56
Finding Atomic Mass of a
Single Atom
– To convert from actual mass to this new
amu, scientists set Carbon as the standard
and the value of the amu unit is defined by
the actual mass of Carbon-12:
Actual mass of C-12= (12)(1.67x10-24g) = 2.00x10-23g
and we use this as a standard to create a
conversion factor:
2.00x10-23g = 12 (amu)
57
Reference-Atomic Mass of
a Single Atom
• Find the atomic mass of an Oxygen-18 atom
• Step 1: calculate the actual mass
of an Oxygen-18 atom
18 x 1.67x10-24 g = 3.01 x10-23 g
• Step 2: use dimensional analysis to convert
to amu with our Carbon standard conversion
factor
• 3.01 x10-23 g x
12  = 18.1 amu
2.00 x10-23 g
• **Don’t use 12 when figuring SDs because it58
is a standard, not a measurement
Let’s Do It!!!
• What’s the atomic mass of a
7
Li atom?
59
Let’s Do It!!!
• What’s the atomic mass of a 7 Li atom?
• Step 1: calculate the actual mass of this
lithium atom
7 x 1.67x10-24 g = 1.17 x10-23 g
• Step 2: use dimensional analysis to convert
to amu with our Carbon standard conversion
• 1.17 x10-23 g x
12  = 7.02 amu
2.00 x10-23 g
60
Isotopes
• What elements are these?
• What are their mass numbers?
++
+
+ +
61
Isotopes
• How would we write the hyphen notation of
these isotopes? Nuclear notation?
• Students on board
++
+
+ +
62
Isotopes
• If we pick up one of the trillions of boron
atoms in the world, it could be either of these
2 types because they are both present
++
+
+ +
•
Boron-10
Boron-11
63
Isotopes
• As it turns out, any mass of Boron, and all
Boron in the world, is ~20% Boron-10, and
~80% Boron-11—their relative abundances
++
+
+ +
•
Boron-10
20%
Boron-11
80%
64
Isotopes
• Many elements are like this
• All chlorine in the world is 75% Chlorine35 and 25% Chlorine-37
• The majority are Chlorine-35
65
Isotopes-Reference
• In nature there are always mixtures of
isotopes and this can pose difficulties when
we do calculations. Why?
• Pencil lead (carbon) has some carbon-12,
carbon-13 and carbon-14 mixed
• The mass of one Carbon-12 atom with 12
protons and neutrons would be this:
(12)(mass of protons and neutrons) =
(12)(1.67 x10-24 g) = 12 amus ()
but that is just for Carbon-12!!
66
Isotopes and Atomic Mass
of an Element
• STOP! “Mass” confusion and Reference Chart
• In calculations we need a mass value that
represents the whole element mixture
(carbon-12, carbon-13 and carbon-14 mixed),
not just one isotope, like Carbon-12 How?
67
Atomic Mass of an ElementReference
• Atomic mass- WEIGHTED average of the
atomic masses of all the element’s isotopes as
they are found in nature (don’t confuse it with
mass number which is just p + n)
(abundance isotope #1) (atomic mass isotope #1 )
+
(abundance isotope #2) (atomic mass isotope #2 )
+
continue for all isotopes
68
Isotopes and Atomic Mass
of an Element
• The atomic masses of each element (the
weighted average of all its isotopes) is
found in the periodic table
++
+
+ +
•
Boron-10
20%
Boron-11
80%
69
Atomic Mass of an Element –
Reference
Use the following data to calculate the atomic mass
for the element Magnesium
Isotope Atomic Mass of Isotope Abundance
Mg - 24
23.982628 
78.600 %
Mg - 25
24.963745 
10.11 %
Mg - 26
25.960802 
11.29 %
(.78600) (23.982628  ) +
(.1011) (24.963745  )
+
(.1129) (25.960802  ) =
18.850  + 2.52  + 2.931  = 24.305 
You do SD’s for every individual calculation
70
Let’s Do It!!!!
• The element copper has naturally
occurring isotopes with mass numbers of
63 and 65
• The relative abundance and atomic masses
are 69.2% for a mass of 62.93amu and
30.8% for a mass of 64.93amu. Calculate
the atomic mass of the element copper
71
Let’s Do It!!!!
• Divide the percentages by 100 to convert
to decimals
(.692) (62.93amu) + (.308)(64.93amu) =
43.5 amu +
20.0 amu
=
63.5 amu
This shows that the majority of the
isotopes found in nature are 62.93 amu
(closer to 63 than 65)
72
Let’s Do It!!!
• There are two isotopes of silicon. The
atomic mass of the element silicon found
on the periodic table is 28.086amu.
• Which of these is not possible to be one
of the atomic masses of the individual
isotopes?
–
–
–
–
A. 26.065
B. 29.543
C. 28.086
D. 27.439
73
Atomic Mass of an Element
• If an element has only 1 isotope, then the
atomic mass of that isotope IS the atomic
mass of the element
74
Review
• Number your paper from 1-5 and answer
the following questions. Two will be
cumulative review!
• 1. Which of these is the symbol for the
metric unit micro (and also for amu)?
– A. M
– B. m
– C. ρ
– D. μ
75
Review
• D
• 2. Which of these is the amount of
protons in an atom and therefore, the ID.
– a. atomic mass
– b. atomic number
– c. mass number
– d. isotope number
76
Review
• B
• 3. What is the bottom number on this
periodic table tile?
– a. atomic mass of the element
Boron
– b. atomic number
– c. mass number
– d. atomic mass of a Boron atom
77
Review
• A
• 4. The atomic mass of the element sulfur
is 32.1 g. There are four common isotopes
of sulfur, S-32 (32.2344 μ), S-33
(33.45676 μ), S-34 (34.15643 μ), S-36
(36.44321 μ) Which of these isotopes is
the most abundant?
–
–
–
–
a. S-32
b. S-33
c. S-34
d. S-35
78
Review
• A
• 5. How do you calculate the atomic mass of
an element?
– a. l x w x h
– b. D = m/V
– c. (abundance isotope #1)(atomic mass
#1) + (abundance isotope #2) (atomic
mass #2) etc.
– d. mass number – atomic number
79
Review
• C
80
The Mole
• The word “mole” in Chemistry is a term used
to describe a certain amount of something
• For example:
– 1 dozen = 12
– 1 baker’s dozen = 13
– 1 gross = 144
– 1 mole = 6.02 X 10 23 which is
602,000,000,000,000,000,000,000 or 602
hextrillion ………of whatever
81
The Mole-Reference
• A mole is the amount of a substance that
contains 6.02 x 1023 of something
– 1 mole of pencils = 6.02 x 1023 pencils
– 1 mole of eggs = 6.02 x 1023 eggs
– 1 mole of carbon = 6.02 x 1023 carbon atoms
• 6.02 x 1023 is called Avogadro’s number
• It helps us count atoms, which we can’t see,
using the measurement of mass
82
Let’s Do It!!!
• How many atoms of oxygen are in a mole of
oxygen?
• How many atoms of magnesium are in a
mole of magnesium?
• Would they have the same mass?
83
Atomic Mass of an Element =
Mole-Reference
• Each element has a unique atomic mass and
this is a standard for each element
• Atomic mass = mass of 1 mole of an element
• The atomic masses are from the periodic
table and we use grams
1 mole O = 6.02 x 1023 atoms O = 15.999 g O
1 mole Mg= 6.02 x 1023 atoms Mg= 24.305 g Mg
84
Molar Mass
• Molar mass –mass of 1 mole of a pure
substance
• Molar mass of element (g/mol) – the atomic
mass of an element
• Molar mass of a compound (g/mol)- the sum
of the molar masses of the elements making
up the compound
• What’s the molar mass of Cl? Don’t forget
units! We are going to round to 3 SDs for the
atomic mass
85
• What’s the molar mass of H2O?
Moles to Atoms-Reference
• 1 mole = 6.02x1023 atoms = atomic mass (g)
• These can be used as a conversion factors
6.02 x 1023 atoms
1 mole
atomic mass
atomic mass
6.02 x 1023 atoms
1 mole
• If you have 1.5 moles of potassium, how many
atoms is this?
• Question mark format for DA
? atoms of K=
6.02 x 1023 atoms K = 9.03 x 1023
1.50 mole K * --------------------atoms K
86
1 mole K
Moles to Grams-Reference
• 1 mole = 6.02x1023 atoms = Atomic Mass (g)
• If you have 1.25 moles of potassium, how
many grams is this?
• Question mark format for DA
? grams of K =
39.1 g K
1.25 moles K * ----------- = 48.9 g K
1 mole K
87
Grams to Atoms-Reference
• 1 mole = 6.02x1023 atoms = Atomic Mass (g)
• If you have 12.5 grams of phosphorus, how
many atoms is this?
• Question mark format for DA
? atoms of P=
23 atoms P
6.02
x
10
12.5 g P
=
x --------------------31.0 g P
2.43 x 1023
atoms P
88
Let’s Do It!!
• 1 mole = 6.02x1023 atoms = Atomic Mass (g)
• If you have 2.93 x 1024 atoms of sodium,
how many moles is this?
• Question mark format for DA
? moles of Na=
2.93 x 1024
atoms Na
= 4.87 moles Na
89
Let’s Do It!!!
• 1 mole = 6.02x1023 atoms = Atomic Mass (g)
• If you have 1.25 grams of sodium, how
many moles is this?
• Question mark format for DA
? moles of Na =
1 mole Na
1.25 g Na
= 0.0543 mole Na
* -------------23.0 g Na
90
Let’s Do It!!!
• 1 mole = 6.02x1023 atoms = Atomic Mass (g)
• If you have 3.52 x 1022 atoms of carbon,
how many grams is this?
? grams of C =
1022
3.52 x
atoms C
12.0 g C
= .702 g C
x ---------------6.02 x 1023 atoms C
91
Mole vs. Molecule
• Don’t confuse mole and molecule
– A mole is an amount
– A molecule is a thing
– We can have a mole of molecules
• A mole of water molecules =
6.02 x 1023 molecules of water
• If I have an amino acid molecule
tryptophan, C11H12N2O2, how many
molecules of tryptophan are in a mole?
92
Review
• Number your paper from 1-5 and answer
the following questions. Two will be
cumulative review!
– 1. Which of these is proper scientific
notation?
•
•
•
•
a. 5.56 X 106
b. 3.5 X 32
c. 95.6 X 1010
d. 256 X 106
93
Review
• A
• 2. Which of the numbers in the nitrogen
periodic table tile is the atomic number?
•
•
•
•
a. N
b. 14
c. 14.007
d. 7
94
Review
• D
• 3. Which of these is NOT a correct
conversion factor for Carbon?
• a.
1 mole C
6.02 x 1023 g C
• b.
1 mole C
6.02 x 1023 atoms C
• c.
1 mole C
12.011 g C
• d.
12.011 g C
6.02 x 1023 atoms C
95
Review
• A
• 4. If I have 14.007 grams of nitrogen, how
many atoms do I have?
•
•
•
•
a. 1.67 x10-24
b. 1.17 x10-23
c. 2.24 x10-23
d. 6.02x1023
96
Review
• D
• 5. How many atoms are in 10 moles of
carbon?
•
•
•
•
a. 10 atoms
b. 6 atoms
c. 6.02x1023 atoms
d. 6.02x1024 atoms
97
Review
• D
• 1 mole of ANYTHING = 6.02x1023 atoms
• ? Atoms = 10 moles C x 6.02x1023 atoms C =
1 mole C
10 x 6.02x1023 atoms C= 6.02x1024 atoms
98
Radius of Nucleus of an
Atom-Reference
• Radius = 3 A (1.4 x 10-13 cm)
• Where A = mass number of the atom
3
A = # of protons + neutrons
and = the cube root of the mass number
and 1.4 x 10-13 cm = a constant that
describes the
effective range of force
99
Radius of Nucleus of an
Atom-Reference
• Calculate the radius of the nucleus of a
56Fe atom
3
• Radius = A (1.4 x 10-13 cm)
• Radius = 3
56
(1.4 X 10-13 cm)
• Radius = 5.4 X 10-13 cm
only round ONCE ! ! !
100
Volume of a NucleusReference
• Calculate the volume of the nucleus of a
56Fe atom (we model as a sphere)
• Volume of a sphere = 4 r 3
3
• Volume = 4  (5.4 * 10 13 cm) 3
3
• Volume = 6.6 * 10-37 cm3
101
Size of the Nucleus
Atom Nuclear Radius
Nuclear Volume
4
A (1.4 x 10-13 cm)
r 3
Nuclear Mass
3
3
23
11
D
m
V
Na
197
79
4
2
A(1.67*10-24g )
Nuclear
Density
Au
He
64
29
Cu
80
35
Br
102
Size of the Nucleus
Atom Nuclear Radius
Nuclear Volume
4
A (1.4 x 10-13 cm)
r 3
Nuclear Mass
3
3
23
11
Na
197
79
4
2
Au
He
m
V
8.1 x 10-13 cm
2.2 x 10-13 cm
5.6 x 10-13 cm
80
35
6.0 x 10-13 cm
Br
D
4.0 x 10-13 cm
64
29
Cu
A(1.67*10-24g )
Nuclear
Density
103
Size of the Nucleus
Atom Nuclear Radius
Nuclear Volume
4
A (1.4 x 10-13 cm)
r 3
Nuclear Mass
3
4.0 x 10-13 cm
3
2.7 x 10-37 cm3
8.1 x 10-13 cm
2.2 x 10-36 cm3
2.2 x 10-13 cm
4.5 x 10-38 cm3
64
29
5.6 x 10-13 cm
7.4 x 10-37 cm3
80
35
6.0 x 10-13 cm
9.0 x 10-37 cm3
23
11
Na
197
79
4
2
Au
He
Cu
Br
A(1.67*10-24g )
Nuclear
Density
D
m
V
104
Size of the Nucleus
Atom Nuclear Radius
Nuclear Volume
4
A (1.4 x 10-13 cm)
r 3
Nuclear Mass
3
3
A(1.67*10-24g )
4.0 x 10-13 cm
2.7 x 10-37 cm3 3.84 x 10-23g
8.1 x 10-13 cm
2.2 x 10-36 cm3 3.29 x 10-22g
2.2 x 10-13 cm
4.5 x 10-38 cm3 6.68 x 10-24g
64
29
5.6 x 10-13 cm
7.4 x 10-37 cm3 1.07 x 10-22g
80
35
6.0 x 10-13 cm
9.0 x 10-37 cm3 1.34 x 10-22g
23
11
Na
197
79
4
2
Au
He
Cu
Br
Nuclear
Density
D
m
V
105
Size of the Nucleus-Reference
Atom Nuclear Radius
Nuclear Volume
4
A (1.4 x 10-13 cm)
r 3
Nuclear Mass
3
3
A(1.67*10-24g )
Nuclear
Density
D
m
V
4.0 x 10-13 cm
2.7 x 10-37 cm3 3.84 x 10-23g
1.4 x 1014
g/cm3
8.1 x 10-13 cm
2.2 x 10-36 cm3 3.29 x 10-22g
1.5 x 1014
g/cm3
2.2 x 10-13 cm
4.5 x 10-38 cm3 6.68 x 10-24g
1.5 x 1014
g/cm3
64
29
5.6 x 10-13 cm
7.4 x 10-37 cm3 1.07 x 10-22g
1.5 x 1014
g/cm3
80
35
6.0 x 10-13 cm
9.0 x 10-37 cm3 1.34 x 10-22g
1.5 x 1014
g/cm3106
23
11
Na
197
79
4
2
Au
He
Cu
Br
Densities of Atoms
• Why are these density values nearly the
same?
• Because all nuclei are made up of the same
material (protons & neutrons). The SAME
material always has the SAME density!
107
Densities of Atoms
• If we look up the density of the element
sodium (it’s on some periodic tables) we
see that its density is .971 g/cm3.
• Why is this different from the value we
calculated?
• We calculated values for the nucleus only
and the periodic table’s value is for the
whole atom – including the space that the
electrons occupy
108