Download Atomic Structure Note Packet

ATOMIC STRUCTURE
Name:_________________________________________Period:______Date:_______
I. LAW vs. THEORY:
1) ____________________________ = a generalization of scientific observations
that ________________what happens (does ____ explain)
2) __________________ (model) = a set of assumptions used to
explain observations and predict new observations
a) Can __________ be truly proven 100% correct.
b) Inevitably change (and must sometimes be abandoned) as more
information becomes available.
c) Considered successful when they _______________ observations and, more
importantly, ________________new observations.
II. THE ATOM: FROM PHILOSOPHICAL IDEA TO
SCIENTIFIC THEORY:
1) EARLY ATOMIC THEORY:
a) 400 B.C- __________________(Greek philosopher)
• Coined the term “________” (meaning indivisible) to
describe the smallest particles of matter
• Did not experiment
2) JOHN DALTON (EARLY 1800’S):
a) “Father of the”_____________________________”
b) Dalton’s Atomic Theory helped explain measurable
observations and successfully predict new observations (while
stimulating additional research from Dalton and other
scientists).
c) Assumptions of Dalton’s Atomic Theory (proposed by Dalton
in 1803):
i. All matter is composed of ____________
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ii. All atoms of the same element are_____________ in size, mass, other
properties. Atoms of different elements differ in size, mass, and other
properties.
iii. Atoms cannot be subdivided, created, or destroyed
iv. ____________________________: Atoms of different elements combine in
simple whole-number ratios to form chemical compounds.
v. In______________________, atoms are combined, separated, or rearranged.
d) Dalton used his Atomic Theory to help correctly _____________new observations
leading to the “Law of Multiple Proportions”.
e) _________________________________ = The masses of one element that
combine with a constant mass of another element to form more than one compound
are in the ratio of small whole numbers
•
Carbon monoxide and carbon dioxide contain oxygen in a 1:2 ratio.
•
______ of carbon reacts with ______of oxygen to form carbon monoxide, CO.
•
_______ of carbon will react with _____ of oxygen to form carbon dioxide, CO2.
f) Dalton’s Atomic Theory explained many observations and correctly predicted many
additional observations. It ___________ correctly predict all new observations (no
theory ever does).
g) Dalton’s Atomic Theory has been __________to explain the new observations:
•
Atoms of one element can have ______________masses (_____________).
•
Atoms ______________subdivided (but not in a chemical reaction or
physical change)…Nuclear reactions
•
An atom of one element ___________ changed into an atom of another
element (but not in a chemical reaction or physical change)……Nuclear
reactions
h) As with all theories, Atomic Theory has been changed and expanded over time (and
will continue to change and expand) to explain new observations.
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3)
CATHODE-RAY TUBE EXPERIMENTS (late 1800’s):
•
_________________________= A glass tube containing a gas at a very low
pressure which contains a negative
electrode (cathode) and a positive
electrode (anode). When high voltage is
applied, a “cathode ray” passes from the
cathode to the anode causing the low
pressure gas in the tube to glow
(different gases glow different colors).
a) Early cathode-ray tube experiments:
•
Proposed that a cathode ray consists of tiny particles with mass and they are
negatively charged.
b) ____________________–discovered the ___________
• Used early cathode-ray tube experiments and some of his own
findings to support the hypothesis that electrons are
negatively charged particles
• Used a _________________- as voltage across the tube was
increased, a beam of light (cathode ray) became visible
• ______________= beam of electrons seen b/c of excited gas
• Found that the beam was _______________by both magnetic
and electrical fields
• OBSERVATION: Noticed that the “cathode rays” were attracted to the positive
electrode, called the anode.
• What conclusion did Thompson have?
The cathode rays were made up a very small ___________charged particle –
_________________
• OBSERVATION: Measured the bending of the path of the cathode rays and was
able to determining the ratio of an electron’s charge to its mass. Found that the
ratio was always the ________________, regardless of the metal used.
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• What conclusion did Thompson have?
Concluded that the negatively charged particles were much _____________
than the lightest know atom (hydrogen), which meant that atoms had a
_____________________________!
c) Later _____________were also discovered using a modified cathode-ray tube
d) Thomson’s Model of the Atom (______________________________________)
•
Thomson postulated that an atom consisted of a diffuse cloud of positive charge
with the negative ________________________ embedded randomly in it.
•
This represented a major __________________from Dalton’s model of atoms as
indivisible.
The existence of the electron raised new questions: if electrons are part of all
matter and they possess a _____________________charge, how can all matter be
__________________________? Also, if the mass of an electron is so small,
what accounts for the rest of the _________________ in a typical atom?
o More experiments to come…
•
4)
ERNEST RUTHERFORD & THE GOLD FOIL
EXPERIMENT
a) Rutherford discovered the _____________________using
alpha particles & gold foil
• __________________________________ = a relatively
large positively charged particle
b) Gold-Foil Experiment (around 1908-1909): Thin
___________________________was bombarded by alpha
particles and the path of the alpha particles was charted after
they passed through the gold foil.
c) Expected results: The massive ________________________
(positively charged) were expected to “crash” through the gold foil with little or
_________ deflections; there was nothing in Thomson’s model of the atom to cause
anything more than minor deflections of the alpha particles.
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(a) When a beam of alpha particles is directed at a thin gold foil, most particles pass
through the foil undeflected, but a small number are deflected at large angles and a few
bounce back toward the particle source. (b) A closeup view shows how most of an atom
is empty space and only the alpha particles that strike a nucleus are deflected.
d) OBSERVATION: Some alpha particles were deflected at large angles, and some
were redirected backward ( ________________________________________!)
• What conclusion did Rutherford draw from this evidence?
The POSITIVE particles of the atom must
_________________________________________________________,
but instead must be concentrated at the center of the atom-the
(__________________).
e) OBSERVATION: Most of the alpha particles passed through the gold foil with few
deflections.
• What conclusion did Rutherford draw from this evidence?
Most of the alpha particles did not hit anything and passed straight through
the gold atoms so therefore, most of the volume of an atom consists of
__________________________________.
f) Rutherford’s Nuclear Model
• Explained the ____________________________nature of matter: the
positive charge of the nucleus balances the negative charge of the electrons.
•
Suggested that the electrons travel
______________________ the positively
charged nucleus.
•
The early nuclear model did not account for all
the atom’s __________________
•
By 1920, Rutherford refined the concept of
the nucleus and concluded that the nucleus
contained positively charged particles called
______________________.
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5)
JAMES CHADWICK (1932)
a) Discovered the__________
III. THE MODERN MODEL OF THE ATOM:
1) WHAT IS THE DIFFERENCE BETWEEN AN ATOM AND AN
ELEMENT?
a). _________________= a substance that cannot be broken down to other
substances by a chemical reaction.
b). ___________: the smallest particle of an element that can exist either alone or in
combination with other atoms.
2) SUBATOMIC PARTICLES OF AN ATOM
a) __________________________________________= a particle smaller than an
atom
• Ex: proton, neutron, electron
b) An atom is composed of subatomic particles including protons, neutrons, and
electrons (plus scientists have determined that protons and neutrons have their
own structures and they are composed of quarks- these particles will not be
covered since scientists do not yet understand if or how they affect chemical
behavior).
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c) SUMMARY OF SUBATOMIC PARTICLES
Particle
Symbol
Relative
Relative
Electron
Mass
Charge
Location
Actual Mass
(g)
Electron
In the space surrounding
the ________________
1/1840
9.110 × 10−28 g
Proton
In the _____________
1
1.673 × 10−24 g
Neutron
In the _____________
1
1.675× 10−24 g
d) ______________________ give the nucleus the positive charge &
____________________ the identity of an atom
e) Protons and neutrons have about the ________________ mass and are over 1840
times more massive than electrons.
f) Most of the mass of an atom is located in the __________________ (protons &
neutrons). (Nuclear forces hold the particles of a nucleus together. These forces
only act over a very short range.)
g) _________________ are electrically ___________________ b/c of the number
of protons EQUALS the number of electrons.
h) ________________________= the positively charged, dense central portion of an
atoms
• Contains nearly all of the atom’s _______________, but takes up a
__________________small fraction of its ___________________
3) ATOMIC NUMBER & AVERAGE ATOMIC MASS
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K
Potassium
39.098
•
_______________
___
___________________
_________________: the number of protons in the nuclei
Also represents the number of _____________that an atom has since atoms
are electrically neutral
Ex: Atomic number of Na is 11; Na has 11 ___________ and 11 ___________
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•
__________________________: the weighted average of the atomic masses of
naturally occurring isotopes of an element.
__________________= atoms of the same element with different number of
NEUTRONS, therefore, different masses.
____________________lists average atomic mass
Weighted averages account for the percentages of each isotope of a given
element
Important b/c they indicate relative mass relationships in chemical reactions
Ex: How average atomic mass is calculated
4) ISOTOPES
__________________= atoms of the same element with different number of
NEUTRONS, therefore, different masses.
• Isotopes of an element all have the SAME NUMBER of _____________and
______________
• Named by their __________________
• ________________ = the total number of protons and neutrons in the
nucleus of an isotope.
Mass Number = # Protons + # Neutrons
• Ex:
H-1
H-2
H-3
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• Example 1: ALL carbon atoms have how many protons? _______________
Most carbon atoms have 6 neutrons. What is their mass number? _____
Some carbon atoms have 8 neutrons. What is their mass number? _____
C-12 and C-14 are isotopes of carbon
o Both isotopes have ___electrons & ___ protons, but differ in
the # of ___________________!!
• Example 2: How many neutrons are in a sodium-23 atom? ________
• Example 3: How many protons, electrons and neutrons are there in an atom of
chlorine-37?
__________________________________________________
__________________________________________________
Hyphen Notation:
• Write the hyphen notation for a hydrogen isotope with a mass number of 3.
_________ or _________________________
Nuclear Symbol Notation:
• Write the nuclear symbol for a hydrogen isotope with a mass number of 3.
_________
Complete the following table.
• Remember:
o atomic number = number of protons = number of electrons
o mass number = number of protons + number of neutrons
Nuclear
Symbol
Atomic
Number
Mass
Number
30
70
87
Number of
Protons
Number of
Neutrons
Number of
Electrons
108
74
50
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_____________ = SI unit for amount of substance
A mole simply represents a______________________, much in the
same way that a dozen represents a set of twelve.
Ex: Just as you can have a dozen cans of soda or a dozen donuts, you
can have a mole of stars or a mole of water molecules.
So a dozen eggs represents 12 eggs, a mole of carbon represents
___________________carbon atoms
__________________is also called “___________________________” = the number
of particles in exactly one mole of a pure substance.
There are _____________________atomic mass units (amu) in______ gram!!
• This conversion allows us to use the masses listed on the periodic table for both
atomic mass and molar mass.
• _____________= is the mass of one atom, is measured in atomic mass units
(amu)
Mass of C-12 is arbitrary assigned a mass of 12 amu
1/12 the mass of a carbon-12 atom is __________
Mass of any atom is expressed relative to the mass of one atom of
carbon-12
• _________________= is the mass of one mole of atoms or molecules, is
measured in grams.
Numerically equal to the atomic mass
1 mole of a substance is 1 molar mass of that substance
Ex: 1 mole of nitrogen is equal to 14.007g of nitrogen!
Atomic & Molar Masses
Element & Symbol
Atomic Mass—Mass of 1 Atom
Carbon (C)
12.011 amu
Helium (He)
4.0026 amu
Copper (Cu)
63.546 amu
Potassium (K)
39.1098 amu
Molar Mass—Mass of 6.02 X
1023 atoms
REMEMBER: one mole of atoms contains 6.02 X 1023 atoms. One mole of molecules contains
6.02 X 1023 molecules. One mole of formula units contains 6.02 X 1023 formula units. One mole
of ions contains 6.02 X 1023 ions.
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Conversions: Using Charts!
Ex: Conversion factors:
• 12 inches = 1 foot
• 3 feet = 1 yard
Set up conversion charts so your units will cancel out!
You multiple across the top of the chart and divide by what is on the bottom of the
chart.
1) How many feet are in 60.72 yards?
2) How many yards are in 678.35 feet?
3) How many inches are in 964.8 feet?
4) How many feet are in 1,743.2 inches?
Mrs. H’s Rounding:
•
We will use 6.02 X 1023 atoms
•
We will round the molar mass values from the Periodic Table (0.5+ round to the whole #)
o Ex: Phosphours 30.974g would be rounded to 31g
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MOLES MASS (grams)
o KEY IDEA: It is important to know the following conversion:
The __________________of an element is equal to the mass of _____________ of
atoms of that element.
Example: 1 mole of zinc is equal to 65.39 grams of zinc!
Use the factor label method (units will cancel out!)
Ex: What is the mass in grams of 3.5 mol of the element copper (Cu)?
3.50 mol Cu 64g Cu
1 mol Cu
= 224g Cu
1. What is the mass in grams of 2.25 mol of iron (Fe)?
2. What is the mass in grams of 0.375 mol of K?
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MASS (grams)
MOLES
1 mol of an element = the molar mass of that element
Ex: A chemist produced 11.9g of Al. How many moles of Al has been produced?
11.9 g Al
1 mol Al
27g Al
= 0.441 mol Al
1. How many moles of Ca are contained in 5.0g of Ca?
2. How many moles of Au are contained in 3.6 x 10-10g of Au?
ATOMS
MOLES
____________of an element = ____________________of that element
Ex: How many moles of Ag are in 3.01 x 1023 atoms?
3.01 x 1023 Ag atoms
1 mol Ag
= 0.5 mol Ag
23
6.02 x 10 Ag atoms
1. How many moles of Pb are equivalent to 1.5 x 1012 atoms?
2. How many moles of tin are equivalent to 2500 atoms?
3. How many atoms of Al are contained in 2.75 mol?
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The Mole , Molar Mass, # of Atoms Practice Problems
What is the conversion factor between moles and atoms?
What is the conversion factor between moles and mass?
1. What is the mass in grams of 4.25 mol of lithium?
2. How many moles of Mg are equivalent to 3.011 x 1023 atoms?
3. How many moles of Ca are contained in 150g of Ca?
4. How many atoms of K are equivalent to 5.83 mol?
5. How many moles of Ca are contained in 40.1g of Ca?
6. How many moles of Pb are contained in 3.25 x 105g of Pb?
7. How many moles of Zn are equivalent to 2.25 x 1025 atoms?
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8. What is the mass in grams of 6.52 mol of the C?
9. How many moles of Fe are contained in 2.65g of Fe?
10. How many atoms of Na are equivalent to 1.50 mol?
11. What is the mass in grams of 1.38 mol of the N?
12. How many moles of O are contained in 4.50 x 10-12g of O?
13. What is the mass in grams of 8.075 mol of the Au?
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History of the Atom Worksheet
Name:____________________________________________________________Period:______
1) What is the difference between a scientific law and theory?
2) Who was the first person to come up with the idea of an atom? _______________________________
3) Who is known as the Father of the Modern atomic theory? __________________________________
4) List 2 of the 5 statements that sum up Dalton’s atomic theory.
5) Is modern atomic theory the same as Dalton’s atomic theory? ___________
6) What 3 modifications to Dalton’s atomic theory have been necessary?
7) Who discovered electrons? _____________________________What did he use to discover
electrons?_____________________________________
8) a) How did cathode rays behave near positively charged objects?
b) What did these observations suggest about cathode rays?
9) What are cathode rays?_______________________________________________________
10) What did Ernest Rutherford discover? ________________________________
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11) During the gold foil experiment some alpha particles (positively charged) passed through the gold foil.
What did this suggest?
12) During the gold foil experiment some alpha particles (positively charged) bounced back toward the
main source. What did this suggest?
13) What is the name of Thomson’s model of the atom? ______________________________________
14) Describe Thomson’s model of the atom?
15) Rutherford’s model of the atom is called the “Nuclear Model”. Describe how this model is different
from Thomson’s.
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Worksheet: Parts of An Atom
1) What does the atomic number represent?
2) What does the mass number represent?
3) Complete the following table. Also, properly label each nuclear symbol with the correct
atomic number and mass number.
Nuclear
Symbol
Atomic
Number
Mass
Number
19
39
Number of
Protons
78
Number of
Neutrons
80
34
16
10
5
137
81
40
18
5
9
37
4
10
56
30
Number of
Electrons
26
17
35
34
29
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4) What is the difference between an atom and an element?
5) Compare and contrast the three types of subatomic particles in terms of location in the
atom, mass, and relative charge.
6) Atoms are electrically neutral. What does this mean in terms of the number of protons and
electrons in an atom?
7) The following are isotopes of oxygen: O-17 & O-18. In the table below, fill in how many
electrons, protons and neutrons each isotope has.
Isotope
# ELECTRONS
# PROTONS
# NEUTRONS
O-17
O-18
8) Write the nuclear symbol AND the hyphen notation for each isotope described below:
a) atomic number = 2 and mass number = 4
b) atomic number = 4 and mass number = 9
c) atomic number = 30 and mass number = 65
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Atomic Structure: Unit Review
Review Worksheet
Name:___________________________________Period:_____Date:________________
1. ___________________________________ was the first person to come up with the idea of an atom.
2. _____________________________is known as the “Father of the Modern Atomic Theory”.
3. _______________________________________discovered the neutron.
4. _____________________________________________discovered electrons.
5. ______________________________________________ discovered the nucleus.
6. _____________________________________are a beam of electrons
7. Cathode rays have a _________________________charge
8. Describe how electrons were discovered?
9. Describe how the nucleus was discovered?
10. Experiments & Findings:
a) Observation: During the cathode ray experiment, cathode rays were attracted to the positive
charge of a magnet.
Finding: What did this suggest?
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b) Observation: During the gold foil experiment some alpha particles (positively charged)
passed through the gold foil.
Finding: What did this suggest?
c) During the gold foil experiment some alpha particles (positively charged) bounced back
toward the main radiation source.
Finding: What did this suggest?
11. Compare and contrast the three types of subatomic particles in terms of location in the atom and
relative charge.
a) ______________________________- found in the nucleus; positive charge
b) ______________________________ – found in the nucleus; neutral charge
c) ______________________________– found in the electron clouds; negative charge
12. How are atoms electrically neutral?
13. Atoms of the same element with different number of NEUTRONS are called ________________
14. What are isotopes named by? _______________________________
15. What is the hyphen notation for the following: atomic number = 6; mass number = 14 __________
16. What is the nuclear symbol notation for the following: atomic number = 26; mass number = 56 ____
17. The sum of protons and neutrons in an atom equals the ____________
18. The number 34 in the name sulfur-34 represents______
19. The weighted average of the atomic masses of naturally occurring isotopes of an element is the
____________________________
20. The si unit for an amount of a substance is the _____
21. __________________ is the mass of one mole of atoms or molecules, is measured in grams
22. 1/12 the mass of a carbon-12 atom is _________________________________
23. What is the conversion factor between moles and grams?
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24. What is the conversion factor between moles and atoms?
25. What is the mass in grams of 4.89 mol of the O?
26. How many moles of Au are equivalent to 9.62 x 1018 atoms?
27. How many moles of B are contained in 236.9g of B?
28. How many atoms of Rb are equivalent to 8.51 mol?
29. Complete the following table. Also, properly label each nuclear symbol with the correct atomic
number and mass number.
Nuclear
Symbol
Atomic
Number
Mass
Number
35
80
Number of
Protons
15
Number of
Neutrons
Number of
Electrons
16
137
56
30. What is the difference between a scientific law and a theory?
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Unit Learning Map (6 days):
Atomic Structure
Mrs. Hostetter
Class: Academic Chemistry A - Grade 11
Unit Essential Question(s):
Optional
Instructional Tools:
How do scientists view
the atom?
Concept
Theory
Lesson Essential Questions:
How did scientists
contribute to the
modern atomic theory?
Vocabulary:
Scientific Law
Theory
Democritus
John Dalton
J.J. Thompson
Ernst Rutherford
James Chadwick
Atom
Element
Law of Multiple Proportions
Cathode ray
Nucleus
Protons
Neutrons
Electrons
Concept
Model of an atom
Lesson Essential Questions:
How has the model of
the atom evolved?
Vocabulary:
Isotopes
Atomic number
Average atomic mass
Mass number
Nuclear symbol
notation
Hyphen notation
Guided Notes
Lab Materials: Al Atom
Concept
Concept
The mole
Lesson Essential Questions:
Lesson Essential Questions:
What is a mole?
Vocabulary:
Vocabulary:
Mole
Molar mass
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Atomic Structure Vocabulary:
3) Scientific law- a generalization of scientific observations that describes what happens (does not
explain)
4) Theory - a set of assumptions used to explain observations and predict new observations
5) Democritus- Greek philosopher that coined the term “atom” (meaning invisible) to describe the
smallest particles of matter
6) John Dalton- “Father of the Modern Atomic Theory”
-All matter is composed of atoms
-All atoms of the same element are identical in size, mass and other properties.
-Atoms cannot be subdivided, created or destroyed
-Law of Multiple Proportions
-In chemical reactions, atoms are combined, separated or destroyed
7) J.J. Thomson- discovered the electron using a cathode ray tube
8) Ernst Rutherford- discovered the nucleus through the gold foil experiment
9) James Chadwick- discovered the neutron
10) Atom- the smallest particle of an element that can exist either alone or in combination with other
atoms
11) Element- a substance that cannot be broken down to other substances by a chemical reaction
12) Law of Multiple Proportions- Atoms of different elements combine in simple whole-number ratios
to form chemical compounds
13) Cathode ray- beam of electrons seen because of excited gas
14) Nucleus- positively charged, dense central portion on an atom
15) Protons- gives the nucleus a positive charge, determines the identity of the atom, and positive charge
is equal to the negative charge of the e16) Electrons (e-)- move nearly at the speed of light and form an electron cloud around the nucleus
17) Isotopes- atoms of the same element with different number of neutrons, therefore, different masses.
18) Atomic number- number of protons (and electrons) in the nucleus of an atom
19) Average atomic mass- the weighted average of the atomic masses of naturally occurring isotopes of
an element
20) Mass number- total number of protons and neutrons in the nucleus of an isotope
21) Hyphen notation – element symbol followed by the mass number; Ex: He-4
22) Nuclear symbol notation- mass number is on top; atomic number is on the bottom; Ex:
23) Mole- SI unit for amount of substance
24) Molar Mass- mass of one mole of atoms or molecules measured in grams
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