Transcript ic Structure - Phillips Scientific Methods

Unit 2
Atomic
Structure &
Nuclear
Chemistry
‘new’ book: Ch 2
Early Theories of Matter
 Democritus (460-370 B.C.) proposed &
believed that



Matter was not infinitely divisible
Matter is made up of tiny particles called
atomos (“uncuttable”)
Atoms could not be created, destroyed, or
further divided
Democritus’ model of atom
Solid and INDESTRUCTABLE
“Billiard Ball” Model
Development of
Chemistry
Robert Boyle
 Beginnings of modern chemistry were seen in
the 16th and 17th centuries, where great
advances were made in metallurgy, the
extraction of metals from ores.
 In the 17th century, Robert Boyle described the
relationship between the pressure and volume
of air and defined an element as a substance
that cannot be broken down into two or more
simpler substances by chemical means.
Development of Chemistry
 During the 18th century, Priestley
discovered oxygen gas and the process
of combustion where carbon-containing
materials burn vigorously in an oxygen
atmosphere.
Priestley
Development of Chemistry
Lavoisier
 In the late 18th century, Lavoisier wrote
the first modern chemistry text. His most
important contribution was the law of
conservation of mass, which states that
in any chemical reaction, the mass of the
substances that react equals the mass of
the products that are formed.

He is known as the father of modern
chemistry.
Development of Chemistry
 In the 19th century, John Dalton revised
Democritus's ideas based upon the
results of scientific research he
conducted

Led to Dalton’s atomic theory
Dalton
Dalton’s Atomic Theory
1. Elements are composed of tiny
indivisible particles called atoms
2. Atoms of the same element are
identical. The atoms of any one
element are different from those of any
other element.
Dalton’s Atomic Theory
3. Atoms of different elements can
physically mix together or can
chemically combine with one another in
simple whole-number ratios to form
compounds.
Dalton’s Atomic Theory
4. Chemical reactions occur when atoms
are separated, joined, or rearranged.

Atoms of one element, however, are
never changed into atoms of another
element as a result of a chemical
reaction.
Legos are Similar to Atoms
H
H2
H
H
O
+
H2
H
H
O2
H
O
H 2O
H
O
O
H
H 2O
Legos can be taken apart and built into many different things.
Atoms can be rearranged into different substances.
Foundations of Atomic Theory
Law of Conservation of Mass
Mass is neither destroyed nor created during ordinary chemical
reactions.
Law of Definite Proportions
The fact that a chemical compound contains the same elements
in exactly the same proportions by mass regardless of the size
of the sample or source of the compound.
Law of Definite
Proportions
Joseph Louis Proust
(1754 – 1826)
 Each compound has a specific ratio of
elements
 It is a ratio by mass
 EX: Water is always 8 grams of oxygen
for every one gram of hydrogen
Law of Definite Proportions
 Whether synthesized in the laboratory or
obtained from various natural sources, copper
(II) carbonate always has the same
composition.
 Analysis of this compound led Proust to
formulate the law of definite proportions.
+
103 g of copper (II) carbonate
53 g of copper
+
40 g of oxygen 10 g of carbon
Structure of Atoms
 Scientists began to wonder what an
atom was like.
 Was it solid throughout with no internal
structure or was it made up of smaller,
subatomic particles?
 It was not until the late 1800’s that
evidence became available that atoms
were composed of smaller parts.
Radioactivity
 One of the pieces of evidence that
atoms are made of smaller particles
came from the work of Marie Curie
(1876 - 1934).
 She discovered radioactivity, the
spontaneous disintegration of some
elements into smaller pieces.
Marie Curie
Discovery of the Electron
 J.J. Thomson (1856 – 1940)
performed experiments that involved
passing electric current through gases
at low pressure

Thomson

He sealed the gases in glass tubes
fitted at both ends with metal disks
called electrodes
Electrodes were connected to a source
of high-voltage electricity
Cathode-Ray Experiment
 One electrode, the anode, became
positively charged
 The other electrode, the cathode,
became negatively charged
 A glowing beam formed between the 2
electrodes (called a cathode ray)
Cathode-Ray Experiment
-
voltage
source
+
vacuum tube
metal disks
Cathode-Ray Experiment
ON
-
OFF
voltage
source
+
Passing an electric current makes a beam appear
to move from the negative to the positive end
Cathode-Ray Experiment
ON
-
OFF
voltage
source
+
+
By adding an electric field…
he found that the moving particles were negative.
Thomson’s Findings
 Cathode rays are attracted to positively
charged metal plates and repelled by
negatively-charged plates
 He proved that atoms contain tiny
negative particles (electrons) and
concluded that ALL atoms must contain
these negative particles.
 He knew that atoms were neutral in
charge and deduced that there must be
a positive charge within the atom.
Plum-Pudding Model
Discovery of the Proton
 Goldstein discovered the proton using
the cathode ray tube in a similar way as
did Thomson and the electron
Discovery of the Nucleus
 Ernest Rutherford (1871-
Rutherford
1937) learned physics in
J.J. Thomson’ lab.
 Noticed that ‘alpha’ particles
were sometimes deflected
by something in the air.


Alpha particles are helium nuclei
Alpha particles are positivelycharged
 Gold-foil experiment
Rutherford’s Gold Foil
Experiment
 Alpha particles were fired at a thin sheet of
gold foil
 Particle-hits on the detecting screen (film)
are recorded
Lead
block
Uranium
Florescent
Screen
Gold Foil
What he expected…
Because he thought the mass was
evenly distributed in the atom.
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What he got…
Expected and Actual Results of
Rutherford’s Experiment
Alpha particles
Nucleus
+
+
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+
+
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+
+
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+
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+
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Plum-pudding atom
Nuclear atom
Thomson’s model
Rutherford’s model
Try it Yourself!
 In the following pictures, there is a target hidden
by a cloud. To figure out the shape of the
target, we shot some beams into the cloud and
recorded where the beams came out. Can you
figure out the shape of the target?
The Answers
Target #1
Target #2
Rutherford’s Findings
 Most of the particles passed right through
 A few particles were deflected
 VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The atom is mostly empty space
The nucleus is small
 The nucleus is dense
 The nucleus is positively charged
Size of the Nucleus
 If an atom is as
large as a stadium,
then the nucleus is
about the size of a
fly in the center of
the stadium!!!
Nuclear Model
n+
Discovery of the Neutron
James Chadwick
9
4
Be
+
4
2
He
12
6
C
+
1
0
n
In 1932 James Chadwick bombarded beryllium-9 with alpha
particles, carbon-12 atoms were formed, and neutrons were
emitted.
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 764
The Modern View of Atomic
Structure
 The atom contains:



electrons
protons: found in the nucleus,
they have a positive charge
equal in magnitude to the
electron’s negative charge.
neutrons: found in the nucleus,
virtually same mass as a proton
but no charge.
The Mass and Charge of the
Electron, Proton, and Neutron
Particle
Mass (kg)
Charge
Electron
9.11  1031
1
Proton
1.67  1027
1+
Neutron
1.67  1027
0
The Chemists’ Shorthand:
Nuclear Symbols
Mass number 
Atomic number 
39
K
19
 Element Symbol
Atomic Number
 Equal to the number of protons
 Equal to the number of electrons in an
atom
 Determines the element!
Mass Number
 mass # = protons + neutrons
 always a whole number
 NOT on the Periodic Table!
Ions
 Atoms that have lost or gained electrons

Cation – positive ion (lost electrons)
 Example:
How many electrons does Na1+
have?

Anion – negative ion (gained electrons)
 Example:
have?
How many electrons does S2-
PRACTICE WITH IONS
Find the
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number of protons = 55
number of neutrons = 78
number of electrons = 54
Atomic number = 55
Mass number = 133
133
55
1+
Cs
Cesium ion
Isotopes
 Atoms of the same element with different mass
numbers.
12
6
C
carbon-12
14
6
C
carbon-14
 Isotopes are atoms with the same number of
protons but different number of neutrons
PRACTICE
 Atomic Number & Mass Number WS
Calculating Relative
Atomic Mass
https://www.youtube.com/watch?v
=SdhLTfma_Eg
Relative Atomic Mass
https://www.youtube.com/watch?v=vqTg4cYwHXY
 12C atom = 1.992 × 10-23 g (*watch above video!)
 atomic mass unit (amu)
 1 amu
1 p
1n
1 e-
= 1/12 the mass of a 12C atom
= 1.007276 amu
= 1.008665 amu
= 0.0005486 amu
Average Atomic Mass
For help:
http://www.docbrown.info/page04/4_73calcs01ram.htm
 weighted average of all isotopes
 on the Periodic Table
 round to 2 decimal places!
Avg.
= (mass x relative abundance) + (mass
Atomic
x relative abundance) +….
Mass
% divided by 100 or part over whole
Average Atomic Mass
Example #1
Magnesium has 3
isotopes. 78.99% Mg-24
with a mass of
23.985042 amu, 10.00%
Mg-25 with a mass of
24.985837 amu, and the
rest Mg-26 with a mass
of 25.982593 amu. What
is the average atomic
mass of magnesium?
12
Mg
24.31
Isotope
Percent
Abundance
Mass
Relative
Mass
Mg-24
78.99
23.985042
18.94578
Mg-25
10.00
24.985837
2.49858
Mg-26
11.01
25.982593
2.86068
If not told otherwise, the mass of the
isotope is the mass number in amu.
24.31 amu
Average Atomic Mass
Example #2
 EX: Calculate the avg. atomic mass of oxygen if its
abundance in nature is 99.76% 16O, 0.04% 17O, and
0.20% 18O.
Avg.
Atomic  (16)(.9976 )  (17)(0.000 4)  (18)(0.002 0)
Mass
= 16.00 amu
What’s the difference between mass
number and average atomic mass?
 Mass number- specifically about one
isotope
 Average atomic mass- includes the
masses of all the different isotopes for
that atom
Atomic Structure
 ATOMS

Differ by number of protons
 IONS
 Differ by number of electrons
 ISOTOPES

Differ by number of neutrons
Intro to Periodic Table (Book:
Section 2.3)
 Distinguish between:

Groups and periods Group
# of ‘A’ elements gives valence e(available for bonding); Can use this to
determine the ‘charge’ (oxidation #) the atom
needs to form a perfect ‘octet’
 Period # gives the # of energy levels where
electrons ‘reside’.

Metals and Nonmetals (and metalloids)
Where are they on PT?
Intro to Bonding (Book: Section 2.4)
 Distinguish between: Molecule and Ion
 Molecule: Covalent bond between atoms. Share e-.
Usually between nonmetals and H. Represent with
molecular formula which shows ratios using
subscripts. (Ex: CH4)
 Molecular Structural formula: Each bond between
atom represents a shared pair of e-.
 Molecules are weak electrolytes (don’t ionize in
solution). Ex- sugar and water
 What is a binary compound or molecule?
Intro to Bonding- con’t
 Ionic: Usually between metal and nonmetal. Give
and take of e- to form cations and anions.
Compounds held together by strong intermolecular
forces. Most are strong electrolytes in water. Ex
of ionic binary compound: NaCl
 Some ionic compounds are ternary (more than 2
atoms): Ex-NaOH. Na is 1+ and the (OH) is 1-. The
OH is actually covalently bonded and carries
overall 1- charge, BUT the Na+ is ionically bonded
to the (OH)-.
Note:
 You don’t have to know charges/
oxidation #’s for this test, except for the 5
polyatomic ions given on next slide
Polyatomic Ions
 Group of covalently bonded (share e-)
atoms carrying a charge. Will learn how
this is done in the next unit
 Start learning some common polyatomic
ions and their charges. For this test- see
table 2.2 in your book and know the
formula name and charge of these 5:
ammonium, hydroxide, carbonate, sulfate,
phosphate. You will have to know them all
for the next test.
HW- ‘New’ Book
 Go over Summary Problem p. 47 (omit h
and i)
 From p. 47- 50 (end of Ch 2): #3, 6,
8,10,14, 16, 18, 20, 22, 26, 44, 49, 50,
52
 Make sure you do fill in notes from PPt
(especially on the individuals and their
contributions to chemistry)
 Go over all worksheets and PPt’s