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UNIT 3 – ATOMS:
THE
BUILDING
BLOCKS OF MATTER
- NUCLEAR CHEMISTRY
Learning Objectives and Learning Target Packet
-Text support Ch. 3 and 21
Atomic Theory Continues…
Six Easy Pieces Read Aloud
■ We have come so far in our understanding of atoms. Centuries of
researching and countless scientists devoting their lives to create
the understanding of the atom today (textbook concepts).
■
However it is NOT over. The more we understand about atoms
(how they work, their make up, etc…) the greater our ability to
advance science and technology in all aspects of our lives (i.e.
medicine).
Check for Understanding
■ What is the charge of a proton?
■ Where would you find a proton in an atom?
■ What is the charge of an electron?
■ Where would you find an electron in an atom?
■ What is the charge of a neutron?
■ Where would you find a neutron in an atom?
■ How big is an electron compared to a proton?
■ How big is a neutron compared to a proton?
The Atom
■ The smallest particle of an element that retains the chemical
properties of that element.
■ Consists of two regions:
– Nucleus
■
1. Small region located at the center of an atom
■
2. Made up of at least one positively charged particle (proton)
■
3. Made up of usually one or more neutral particles (neutrons)
– Region surrounding the nucleus – Electron Cloud
■
1. Very large compared to size of nucleus
■
2. Contains the negatively charged particles (electrons)
Refresh: What are Atoms?
■ Atoms are tiny particles that determine
properties of all matter.
■ Atoms are the building blocks for molecules.
■ Atoms form elements.
■ Element: A substance that cannot be broken
down into simpler substances by chemical
means.
Parts of an Atom
■ Proton: A subatomic particle that has a positive
charge and is found in the nucleus of the atom.
■ Neutron: A subatomic particle that has NO charge
and is found in the nucleus of the atom.
■ Electron: A subatomic particle that has a negative
charge and moves around the outside of the
nucleus.
Label the Atom
Subatomic Particles
Particle Charge
Proton
+1
Mass (kg)
Location
1.67 x 10-27
nucleus
(1 AMU)
Neutron
0
1.67 x 10-27
nucleus
(1 AMU)
Electron
-1
9.11 x 10-31
(~0 AMU)
Outside
nucleus
Electron Orbital
-Electrons orbit the nucleus in orbital clouds.
-Electrons with different amounts of energy exist in different energy
levels.
The Electron Cloud Model
Electrons in each energy level
• Each energy level can hold a limited number of
electrons.
• The lowest energy level is the smallest and the
closest to the nucleus.
Atomic Number
• The atomic number of
an element is the
number of protons in
the nucleus of an atom
of that element.
Charge of Atoms
■ Atoms are not charged even though they have particles that contain
charges.
■ Atoms are neutral because they have EQUAL numbers of protons and
electrons.
Ex: Helium Atom
Charge of 2 protons:
Charge of 2 neutrons:
Charge of 2 electrons:
Total charge of He atom:
+2
0
-2
0
Mass Number
■ The sum of the protons and neutrons in the nucleus is the mass
number of that particular atom. (mass # = p + n)
Atomic Number vs. Mass Number
■ Atomic Number: Equal to the number of protons in the nucleus of
the atom.
(number of electrons = the number of protons)
■ Mass Number: Equal to the number of protons AND neutrons in an
atom’s nucleus.
■ Average Atomic Weight (below symbol on PT):
– Weighted average of the atomic masses of the naturally occurring
isotopes of an element
– We will round this number to the nearest hundredth (TWO decimal
places)
■ Example Oxygen’s average atomic mass is 15.9994 = 16.00
Application- TOGETHER
Iodine (I)
Iron (Fe)
Atomic number
____
Atomic Number
____
Mass Number
126
Mass Number
____
Number of Protons ____
Number of Protons____
Number of Neutrons ____
Number of Neutrons 33
Number of electrons ____
Number of Electrons ____
Application- ON YOUR OWN
Nickel (Ni)
Atomic number
Radon (Rn)
____
Atomic Number____
Mass Number ____
Mass Number 222
Number of Protons ____
Number of Protons____
Number of Neutrons 31
Number of Neutrons____
Number of electrons ____
Number of Electrons ____
Isotopes
• Isotopes of an element have different mass numbers because they
have different numbers of neutrons, but they all have the same
atomic number.
Isotope Examples
■ Carbon – 12
– 6 protons
– 6 electrons
– 6 neutrons
■ Carbon – 13
– 6 protons
– 6 electrons
– 7 neutrons
■ Carbon 14
– 6 protons
– 6 electrons
– 8 neutrons
Example
6,7
Symbol
Li
3
Element
Name
Mass
Numbers
Lithium
Atomic
Number
One More Example
107
Ag
47
Silver
Isotope Tables
■ Breaking it down…
– To find the symbol – determine the atomic number of the element. This is
the number of protons
– To find the protons- determine the atomic number of the element.
– To find the electrons – equal to the number of protons of a neutral atom
– To find neutrons: Mass Number – Atomic Number = Number of Neutrons
– To find Mass Number: Atomic Number + Number of Neutrons = Mass
Number
Isotopes
Practice
Molar Mass
■ Molar Mass: mass in grams of one mole of a pure substance.
– Units = g/mol
■ To calculate molar mass you must find the atomic mass (Units = g).
Ex: Find the atomic mass of Al.
Ex: Find the atomic mass of O2.
Ex: Find the atomic mass of CH4.
* To change atomic mass to molar mass just put the answer (g) over
moles.
Calculate the molar mass of the following …
show ALL your work!
■ 1)
NaBr
■ 2)
PbSO4
■ 3)
Ca(OH)2
■ 4)
Na3PO4
■ 5)
(NH4)2CO3
■ 6)
C6H12O6
■ 7)
Fe3(PO4)2
■ 8)
(NH4)2S
■ 9)
Zn(C2H3O2)2
■ 10) AgF
Warm-up
■
–
–
–
Identify the element:
65 neutrons
48 protons
48 electrons
■ How many protons, neutrons, electrons in the
following isotopes
– Mn-56
S-31
■ Determine the molar mass of the following
– H2O
– Na2CO3
AVERAGE ATOMIC
MASS
Why is the mass on the periodic table not a whole number?
Can we have parts of a proton or neutron?
Average Atomic Mass - POGIL
■ Group Roles and Members
What is the average atomic mass of
Copper?
■
There are two naturally occurring isotopes of copper.
Isotope
a.m.u. (atomic mass unit)
% Abundance
Copper-63
62.929601
69.17
Copper-65
64.927794
30.83
■
a.m.u. means atomic mass unit and is the mass of one atom of that
particular isotope.
■
–
Average atomic mass = Σ (a.m.u. x % abundance in decimal form)
Σ means the “sum of”
■
=(62.929601 x 0.6917) + (64.927794 x 0.3083)
■
=63.55 amu and also grams/mol….check the periodic table…..the mass
number is the average atomic mass for copper on your periodic table!
Partner Practice: What is the average atomic
mass of Oxygen?
Isotope
a.m.u.
% Abundance
Oxygen-16
15.994915
99.757
Oxygen-17
16.99132
0.038
Oxygen -18
17.999160
0.205
■ 15.999 a.m.u. or 15.999 grams/mol … check the periodic table!
Write it Out!!
■ Describe in your OWN words how you can calculate the
average atomic mass of an element.
– Be sure to describe what information you need to know
and how you use it.
– **Be prepared to share out your response!**
PERCENT
ABUNDANCE
What isotopes are most common?
Lithium has two isotopes… Li-6 and Li-7.
What is the % abundance of each?
Isotope
a.m.u.
% Abundance
Lithium-6
6.015122
?
Lithium-7
7.016004
?
■ Look up Lithium’s average mass on the periodic table: 6.94
■ Now it is time for some ALGEBRA!
– 6.94 = 6.015122(x) + 7.016004 (1-x). Solve for x …
– 6.94 = 6.015122x + 7.016004-7.016004x (apply distributive
property here!)
– Combine like terms
■
6.94 – 7.016004 = -1.000882x
■
-0.076004 = -1.000882 x and therefore x=0.0759 and 1-x = 0.9241
Continued…
Isotope
a.m.u.
% abundance
Lithium-6
6.015122
7.59%
Lithium-7
7.016004
92.41%
Partner Practice: Calculate the average
atomic mass of argon to two decimal places.
Isotope
a.m.u.
% Abundance
Argon-36
35.97
0.337%
Argon-38
37.96
0.063%
Argon-40
39.96
99.600%
39.95 g/mol
Application… Skittlium!
■ Lab Reflection
■ Question #5:
– Determine the average atomic mass for the
following sample of skittlium: (Please show your
work for full credit!)
Type of Skittlium
Skittlium-58
Number of Atoms
3
Skittlium-59
Skittlium-60
Skittlium-61
9
1
6
Skittlium-62
3
THE MOLE
Mole Music Video
■ http://www.youtube.com/watch?v=oIkC7SRqXP0
THE MOLE &
AVOGADRO’S
NUMBER
NA =
23
6.02x10
http://ed.ted.com/lessons/daniel-dulek-how-big-is-a-mole-not-the-animal-the-other-one
Counting Particles
Scientists Counting Unit
■ Since atoms are so small, it is hard to count individual atoms.
■ To solve this problem, chemists count by moles.
– Moles: SI unit for measuring the amount of a
substance.
One mole of
Carbon
Mole Continued
– One mole of anything contains
6.02 x 1023 “particles”
– Particles can be atoms, ions,
molecules, electrons, formula units, etc.
Avogadro’s Number: 6.02x1023 = 1 mole
602,200,000,000,000,000,000,000
How Big is Avogadro’s Number?
■ One mole of sheets of
paper stacked one on top
of the other would reach
beyond the solar system.
■ One mole of basketballs
could create a new planet
the size of the earth.
■ One mole of rice grains
would cover the land
masses of Earth in a depth
of 75 meters.
■ If you had a mole of
pennies and gave away 1
million dollars of it
(100,000,000 pennies) a
day to everyone in the world
it would take you more than
3000 years to distribute all
your money.
Conversions with Avogadro’s
Number
1.
Start with what you know.
2.
Use the following conversion factor:
6.02 x 1023 particles
1 mole
3.
Cancel units.
4.
Solve (preform the math).
“Mole Map”
Practice Conversions:
Moles to Particles
1.
Determine how many particles of sucrose are in 3.50 moles of
sucrose.
2.
Determine the number of atoms in 2.50 mol of Zn.
Practice Conversions:
Moles to Particles
3. Given 3.25 mol AgNO3, determine the number of formula units.
4. Calculate the number of molecules in 1.15 mol of water.
Warm-Up
Calculate the number of molecules in 1.15 mol of water.
“Mole Map”
Practice Conversions:
Particles to Moles
1.
How many moles are in 5.75x1024 atoms Al?
Practice Conversions:
Particles to Moles
2.
How many ATOMS are in 3.75x1024 molecules of CO2?
3.
How many moles are in 2.50x1020 atoms of Fe?
– Mole Conversion Worksheet
Practice, Practice, Practice
Check for Understanding
■ 1.5 moles Cu = ________ atoms Cu
■ 8.51 x 1023 SO3 = ________ moles SO3
Solutions:
– 9.0 x 1023atoms Cu
– 28.3g SO3
MOLAR MASS AND
CONVERSIONS
Adding Mass-Mole and Mole-Mass
Finally, put it all together and what do you get?!
Molar Mass
■ Molar Mass: mass in grams of one mole of a pure substance.
– Units = g/mol
■ To calculate molar mass you must find the atomic mass (Units = g).
Ex: Find the atomic mass of Al.
Ex: Find the atomic mass of O2.
Ex: Find the atomic mass of CH4.
* To change atomic mass to molar mass just put the answer (g) over
moles.
Mole-Mass Conversions
Conversion Factor:
Mass (g)
Moles
P. Table
1
Ex: While working in a Chemistry lab, Gary needs 3.00
moles of Mn for a chemical reaction. How much Mn
does Gary need to mass?
Practice with Mole-Mass
Conversions
■ Determine the moles in
each of the following:
■
Determine the mass in
grams of each of the
following:
4. 25.5 g Ag
1.
3.57 mol Al
5. 125 g Zn
1.
42.6 mol Si
6. 1.45 kg Fe
1.
3.45 mol Co
Mass-Atoms Conversions
■ This involves two conversions.
■ Conversion #1: Mass to moles
■ Conversion #2: Moles to particles
Mass
Moles
“Particles”
(g)
1
6.02x1023
■ Ex: How many atoms of gold are in a pure gold nugget having a
mass of 25.0 g?
Practice Mass-Atom Conversions
How many atoms are in
the following samples?
1.
55.2 g Li
2.
0.230 g Pb
3.
45.6 g Si
Reverse It…
Atoms-Mass Conversions
■ This involves two conversions.
■ Conversion #1: particles to moles
■ Conversion #2: moles to grams.
Mass
Moles
“Particles”
(g)
1
6.02x1023
REVERSE IT….
Practice Atom-Mass Conversions
How many grams are in the following samples?
4. 6.02 x 1024 atoms Bi
5. 3.40 x 1022 atoms He
6. 1.50 x 1015 atoms U
Tie up your shoes its…
PRACTICE TIME!!!
Mole Conversion Worksheet
Mole Music Video
■ http://www.youtube.com/watch?v=oIkC7SRqXP0
REVIEW
Unit 3 – Atoms: The Building Blocks of Matter &
Nuclear Chemistry
Review Questions
■
What is the charge of a proton?
■
Where would you find a proton in an atom?
■
What is the charge of an electron?
■
Where would you find an electron in an atom?
■
What is the charge of a neutron?
■
Where would you find a neutron in an atom?
■
How big is an electron compared to a proton?
■
How big is a neutron compared to a proton?
■
What part of the atom takes up the most space?
■
What part of the atom contains the most mass?
Complete the following table
Particle Charge
+1
Neutron
-1
Mass
(amu)
1
Location
Review Questions
■ Complete the following table.
Isotope
Atomic
Number
Mass
Number
Number
of
Protons
Number Number
of
of
Neutrons Electrons
Zn-64
9
10
24
20
Individual Review Time
■
–
–
–
–
–
–
Review Ch. 3 Worksheets, Practice Packet, and Notes!
Atomic Structure
Isotopes
Mole Conversions
Molar Mass
Average Atomic Mass
% Abundance
NUCLEAR
CHEMISTRY
Explore Pages 649-671
10 minutes
Unit Packet Questions and Practice
Nuclear Chemistry – Pages 649-671
– Nuclear Reactions = Changes in the # of protons or neutrons in
an atom’s nucleus
– Nuclide - The nucleus of an atom
– Stability of nuclide depends on the ratio of neutrons to protons
■
Too few or too many neutrons make the nuclide unstable and nuclide
decays to form more stable nuclides.
■
The most stable nuclei cluster over a range of neutron-proton ratios
called the band of stability (page 683 in book)
■
For lower atomic numbers, typically 1:1 ratios of protons to neutrons, for
higher atomic numbers typically 1:1.5
■
Nuclides are most stable when they reach a magic numbers (natural
occurrences that are more stable).
Types of Radiation
■
alpha = rapidly moving helium nuclei (2 protons and 2 neutrons)
■
beta negative (-) = extremely fast electrons, little mass
■
gamma – electromagnetic radiation – no mass, no charge, energy only
■
positron or + = particle with same mass as an electron, but a positive
charge
■
Other types of radiation do exist – microwaves, radar, light, etc.
■
Radiation is classified as ionizing or non-ionizing
– Ionizing radiation = enough energy to change atoms and
molecules into ions (alpha, beta, X-rays, gamma)/Can cause
changes in living cells
– Non-ionizing radiation = cannot ionize matter (radio, light)
What is nuclear decay?
■ Nuclear decay is an exothermic process discovered by Henri
Becquerel
– Each decay results in a nucleus that contains less energy
– Stability of nuclide depends on the number of neutrons to
protons
■
Too few or too many neutrons makes the nuclide unstable
■
Nuclide decays to form more stable nuclides
■ Describing Nuclear Decay = Equations are used to represent
– Initial nucleus = parent nucleus, new nucleus = daughter
nucleus
– After any decay, the daughter nuclide is less energetic than the
parent
– The daughter is also more stable – better neutron/proton ratio
Alpha Emission
– Occurs with heavy nuclei – decreases neutrons and protons to reach
stability.
– In alpha emission, the parent nuclei decays into the daughter nuclide
by emitting alpha particles (helium nucleus, 42He)
– Example – Uranium-238
– Nuclide loses 2 protons and 2 neutrons
Beta Emission
■
1. Occurs if too many neutrons to protons
■
2. Either electron or positron is released from the parent nuclide
■
3. -, an electron is given off, but is it not one of the original electrons
a. The electron is formed from the change of a neutron into 1 proton
and 1 electron
b. The daughter nuclide thus has a new atomic number because the
number of protons increases by
1; mass stays the same
c. New atomic number = new element symbol
 0-1 +
■
4. Example: Thorium-234
234 Pa
91
■
5. Positron or + emission occurs when the parent nuclide has too few
neutrons for the number of protons
234 Th
90
a. Results in conversion of a proton into a neutron – Releases positron
b. Decreases the atomic number by one, mass stays the same
c. New atomic number = new element symbol
■
6. Example: Potassium-38
38 Ar
18
38 K
19
 0+1 +
Electron Capture
■ 1. Occurs with too few neutrons to protons, like positron
■ 2. An inner orbital electron is captured by the nucleus of its own
atom.
■ 3. This electron combines with a proton and a neutron is formed.
■ 4. Decreases the number of protons and increases neutrons.
■ 5. The atomic number is decreased by one, but the mass stays
the same.
■ 6. Electron capture is shown on the left side of the equation.
106
47Ag
+ 0-1e  10646Pd
Neutron and Gamma Decay
■ Gamma Decay
– Equation results in no change in parent nuclide.
■ Neutrons can also be released from nuclear reactions
– 25399Es + 42He  10n + 256101Md
Types of Radiation Compared
Type of
Radiation
Symbol
Change in
Mass #
Change in
Atomic #
Charge
Alpha
Decreases by
4
Decreases by
2
2+
Beta
No Change
Increases by 1
1-
Gamma
No Change
No Change
0
Nuclear Fission
■ Nuclear Fission = a very heavy nucleus splits into more stable
nuclei or intermediate mass and releases large amounts of
energy
– Can occur spontaneously or when nuclei are bombarded with
particles.
– A chain reaction = a reaction in which the material that starts
the reaction is also one of the products and can start another
reaction.
■
■
■
■
■
■
For example, when fission of uranium=235 occurs, the 2 or 3 neutrons
given off cause the fission of other uranium-235 nuclei.
Continues until all atoms are split or neutrons fail to hit uranium-235
nuclei. (mass hits a certain level.
The minimum amount of nuclide that provides the number of neutrons
to sustain a chain reaction = critical mass.
Uncontrolled chain reactions = nuclear bombs
Controlled chain reactions = nuclear reactors (used to produce energy or
radioactive nuclides.)
Nuclear Power Plants use heat from reactors to produce electrical
energy.
Nuclear Fusion
– Light mass nuclei combine to form heavier, more stable nucleus.
■
Like Nuclear Fission, releases high amounts of energy - Releases even
more energy per gram of fuel than nuclear fission.
■
Occurs on the sun and in other stars.
– Like Nuclear Fission, uncontrolled fusion reactions are used to
make bombs – uncontrolled fusion reactions of hydrogen are the
source of energy for the hydrogen bomb.
– Fission reactions are used to provide heat and pressure to
trigger the fusion.
– We are investigating ways to use this power in a controlled way.
Calculating Half-lives
■ A. Half-life (t1/2) = time required for half of the atoms of a
radioactive nuclide to decay.
■ B. Solve half-live problems, first determine the rate constant (k)
k = ln(0.5)/half life years
■ C. Use the k value to determine how long it takes for decay to
occur
t1/2 = ln(percent left as decimal)/k
Half-life Examples
■ The radioactive isotope cobalt-60 has a half-life of 5.27 years.
What is the rate constant?
K = ln(0.5)/5.27 = -0.132
■ How long will it take 20% of the cobalt-60 to change to nickel60?
t1/2 = ln(.80)/-0.132 = 1.69 years
■ How long will it take 70% of the cobalt-60 to decay to nickel-60?
t1/2 = ln(.30)/-0.132 = 9.12 years
BE PREPARED FOR
THE SPIRAL
ASSESSMENT!