Download Ch 3: Atoms

Ch 4: Atoms
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Scientists (key experiments and
contributions)
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Counting P,N,E
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Democritus
Lavosier
Dalton
Crookes
Thomson
Miliken
Rutherford
Atoms, Ions, Isotope
Cation vs anion
Define Ion and Isotope
Average Atomic Mass Problems
3.1 The Atom: From Idea to Theory
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Historical Background- In
approximately 400 BC, Democritus
(Greek) coins the term
“atom” (means indivisible). Before that
matter was thought to be one continuous
piece - called the continuous theory of
matter. Democritus creates the
discontinuous theory of matter. His
theory gets buried for thousands of years
18th century - experimental evidence
appears to support the idea of atoms.
Law of Conservation of Mass –
Antoine Lavosier (French) -1700’s
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The number of each kind of atoms
on the reactant side must equal the
number of each kind of atoms on
the product side
A +
2B +
C
—> AB2C
Law of Multiple Proportions – John
Dalton (English) - 1803
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The mass of one element combines
with masses of other elements
simple in whole number ratios.
Water (H2O) is always: 11.2% H;
88.8% O
Sugar (C6H1206) is always: 42.1%
C; 6.5% H; 51.4% O
Dalton’s Atomic Theory
1. Everything is made of atoms
2. Atoms of the same element are
identical (NOT TRUE)
3. Atoms can not be broken down,
created or destroyed. (NOT TRUE)
4. Atoms combine in simple whole
number ratios to form chemical
compounds
5. A chemical reaction is the
combining, separation, or
rearrangement of atoms.
3.2 The Structure of the Atom
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Updating Atomic Theory
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1870’s - English physicist William
Crookes - studied the behavior of gases
in vacuum tubes(Crookes tubes forerunner of picture tubes in TVs).
Crookes’ theory was that some kind of
radiation or particles were traveling
from the cathode across the tube. He
named them cathode rays .
3.2 The Structure of the Atom
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(video)
20 years later, J.J. Thomson (English)
repeated those experiments and
devised new ones.
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Thomson used a variety of materials,
so he figured cathode ray particles
must be fundamental to all atoms.
1897 - discovery of the electron.
Plum Pudding Model
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https://www.youtube.com/watch?v=O9Goyscbazk
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3.2 The Structure of the Atom
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Charge and Mass of the electron Thomson and Milliken (oil drop
experiment) worked together to
discover the charge and mass of the
electron
charge = 1.602 x 10-19 coulomb this is
the smallest charge ever detected
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mass = 9.11 x 10-28 g
pretty insignificant
this weight is
3.2 The Structure of the Atom
https://www.youtube.com/watch?v=wzALbzTdnc8
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1909 - Gold Foil Experiment
(Rutherford - New Zealand)
Nuclei are composed of ‘nucleons’:
protons and neutrons
A) Oil Drop
B) Found Nucleus
C) Law of conservation of mass
D) Plum Pudding Model
1) Lavosier
E) Coined the term cathode ray
2) Democritus
F) Term atom
3) Milliken
G) Law of multiple proportions
4) Rutherford
H) Cathode Ray Tube experiments
5) Dalton
I) Gold Foil Experiment
6) Thomson
J) Charge and Mass of electron
7) Crookes
Table: Subatomic particles
important in chemistry.
3.3 Weighing and Counting Atoms
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We look to the periodic table to give us
information about the number of particles
are in atoms and also to help us count atoms
in a sample.
Atomic Number (Z)
-Number of protons in the nucleus
- Uniquely labels each element
Mass Number (M)
- Number of protons + neutrons in the
nucleus
Counting electrons
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Atoms
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Same number of electrons and protons
Ions
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Ionic charge (q) = #protons #electrons
Positive ions are cations
Negative ions are anions
Review of formulas
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atomic # (Z) - (always a whole number, smaller
number on the periodic table) = # of protons in the
nucleus - also indicates the # of electrons if the
element is not charged
atomic mass – the average mass of all of the
isotopes of an element – is a number with a decimal
– is always the larger number on the periodic table.
mass number (A) - sum of the protons and
neutrons in a nucleus
this number is rounded from atomic mass due to the
fact that there are isotopes
# neutrons = A - Z
example - # of neutrons in
Li = 6.941-3 = 3.941 rounds to 4
Ion – a charged atom. Atoms become charged by
gaining electrons (become a negative charge) or
losing electrons (become a positive charge)
Practice- How many Protons?
23
16
7
Practice- How many Electrons?
23
16
7
Practice- How many Neutrons?
23
16
7
Lets try a few together
Lots of Practice!!!
p+
C
Ca
U
Cl
Mg
14C
S-2
Na+1
e-
n°
Atomic # =
(# of p+)
Mass # =
(p+ + n0)
Protons Neutrons Electrons
1) Sr
2) Sr
3) O
+2
-2
4) H+
5) B +3
Protons Neutrons Electrons
38
50
38
38
50
36
8
8
10
4) H+
1
0
0
5) B +3
5
6
2
1) Sr
2) Sr
3) O
+2
-2
Complete Practice Problems
Isotopes
Isotopes
Two atoms of the same element
# of p+) but with different weights
(different # of n0)
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H-1, H-2,H-3
C-12, C-13, C-14
(same
U-235, U-238
How do you determine your grades
in this class?
60% tests: 90
40% other: 100
Average Atomic Mass (“weighted
average”)
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Definition - The average weight of the
natural isotopes of an element in their
natural abundance.
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Carbon consists of two isotopes:
98.90% is C-12 (12.0000 amu).
The rest is C-13 (13.0034 amu).
Calculate the average atomic
mass of carbon to 5 significant
figures.
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(.9890)(12.0000)+(.0110)(13.0034)=x
11.8680+.1430=12.011
Ex1: Chlorine consists of two natural isotopes, 35Cl (34.96885)
at 75.53% abundance and 37Cl (36.96590) at 24.47%
abundance. Calculate the average atomic mass of Chlorine.
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(.7553)(34.96885)+(.2447)(36.96590)=x
26.41+9.045=35.46amu
Ex2: Antimony consists of two natural isotopes
57.25% is 121Sb (120.9038). Calculate the % and
mass of the other isotope if the average atomic
mass is 121.8.
(.5725)( 120.9038)+(.4275)(x) =121.8
69.2174255+.4275x=121.8
.4275x= 52.5825745
=123.0amu
Practice Worksheet Problems
Beanium Lab
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Scientists (key experiments and
contributions)
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Counting P,N,E
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Democritus
Lavosier
Dalton
Crookes
Thomson
Miliken
Rutherford
Atoms, Ions, Isotope
Cation vs anion
Define Ion and Isotope
Average Atomic Mass Problems