Transcript Atoms: The Building Blocks of Matter

Atoms: The Building
Blocks of Matter
Chapter 3
Development of
Atomic Theory
In the Beginning…
Democritus—matter is composed
of small, indivisible particles
 Atoms – “indivisible”
 Greeks—four “elements” make
up all substances: earth, water,
fire, and air

Slow Progress
No experimental evidence for
atoms—even though Democritus
was right, everyone believed
Aristotle.
 Next 2 000 years--alchemy
 Discoveries: elements and
preparation of mineral acids

“Modern” Chemistry
17th century—chemists worked
with metals and used elements &
compounds for medicine.
 Robert Boyle—first quantitative
experiments—pressure/volume
relationships in gases
 Evidence led to support for atomic
theory of matter

New Ideas
Boyle—An element is any
substance that cannot be
reduced to a simpler substance.
 Greek idea of elements died.
 Joseph Priestly—discovered
oxygen, “dephlogisticated air”—
led to study of combustion

Fundamental Laws
Lavoisier
Law of Conservation of
Mass—mass is neither
created nor destroyed in a
chemical reaction
Proust
Law of Definite Proportions—a
given compound always
contains the same proportion
of elements by mass
H20 NaCl NaOH AlCl3
Example: H2O is always 2
parts hydrogen and 16 parts
oxygen by mass.
 Example: CO2 is always 12
parts carbon and 32 parts
oxygen by mass.

Dalton
Law of Multiple Proportions—When
two elements form a series of
compounds, the ratios of the
masses of he second element that
combines with 1 g of the first
element can always be reduced to
small whole numbers.
Law of Multiple
Proportions (Cont’d)
Examples:
Always a 1:2 ratio in the
CO
amount of oxygen used
CO2
to make CO & CO2
Dalton’s Atomic
Theory
1.
2.
3.
Each element is made up of tiny
particles called atoms.
Atoms of a given element are
identical in size, mass,and other
properties. Atoms of different
elements are different.
Atoms cannot be subdivided,
created, or destroyed.
Dalton’s Atomic
Theory
4.
5.
Atoms of different elements
combine in simple, wholenumber ratios to form
compounds.
Chemical reactions involve the
reorganization of atoms, but the
atoms themselves are not
changed.
Dalton’s Atomic
Theory
By relating atoms to mass
(something that could be
measured) Dalton developed
Democritus’ idea into a theory.
 Notice, not all of Dalton’s points
are correct!

Modern Definition of
an Atom
An atom is the smallest
particle of an element that
retains the chemical
properties of that element.
Structure of the Atom
Discovery of Electrons

J.J. Thomson—experiments
with cathode ray tubes (CRT’s)
Cathode Ray Experiment
Applying voltage causes a glow
to travel from the negative to
positive end glass tube in which
a partial vacuum exists.
What can we learn from a ray??
Observations and
Conclusions

Ray is deflected by a magnetic
field.


Has electromagnetic properties
Ray travels away from negative
and toward positive

probably is negatively charged
Deflection of cathode rays by an
applied electric field.
Observations and
Conclusions

Any metal will produce a ray.


All atoms must contain these
particles.
Atoms are electrically neutral.
 Some positive particle must be
present to balance the negative
charge.
Charge vs. Mass
Thomson was able to calculate
the charge to mass ratio
 1.7 x 1011 Coulombs/kg
 VERY LARGE charge for a very
small mass

New Model
Plum pudding—atoms are
spherical masses of positive
charge with electrons scattered
throughout
 Since electrons have a small
mass, the mass must come
from something else.

Charge of an Electron
Robert Millikan’s oil drop
experiment
 Oil drops can be suspended
in an electric field by
adjusting voltage of charged
plates

Millikan’s Experiment
Millikan’s Conclusion

The charge produced on the
oil drop was always a whole
number multiple of an
electron’s charge.
Gold Foil Experiment
Question: How will positive alpha
particles behave when passing
through gold foil?
Background: “Plum Pudding” model
of the atom
Hypothesis: All particles should
crash through the foil—slight
deflection.
Expected Results
Experiment: -particle
bombardment of metal foil.
Actual Results
Gold Foil Experiment
Result: Most particles went through,
but some were deflected and some
were reflected.
Conclusion: “Plum Pudding” model
cannot be correct. Atoms must have
a dense, positively charged nucleus.
Nuclear Structure
Nucleus—protons (+) and
neutrons (0)
 Atomic number (Z)--# of
protons (element)
 Atomic mass (A)—protons +
neutrons (different isotopes)

How can the positive
protons stay together in
the tiny nucleus?
Nuclear forces—strong forces
that only act when particles are
VERY close together
 Much stronger than electrostatic
forces that would cause
repulsion

Size
Atomic radii: 40-270 picometers
(billionths of a meter)
 Nuclear radii: 0.001 picometer
 Nucleus is very small compared to
the whole atom, but most of the
mass is there—VERY dense

Counting Atoms
Atomic Number
The number of protons in the
nucleus of an atom
 Sometimes written as Z
 Makes the element what it is
 Equal to number of electrons if
atom is neutral.

Isotopes
Atoms with the same number of
protons but different numbers of
neutrons and thus different
masses
 Examples: protium, deuterium,
tritium (isotopes of hydrogen
with masses of 1, 2, and 3)

Mass Number
Total number of protons &
neutrons in the nucleus of an
atom
 # Neutrons = mass # - atomic #

Writing Isotopes

Hyphen notation—symbol
followed by mass number


H-1, H-2, H-3, C-14, U-235
Nuclear symbol—symbol with
atomic number as subscript and
mass number as superscript

1
1
H
2
1
H
3
1
H
12
6
C
235
92
U
In Your Scientist’s
Notebook:
1.
2.
3.
How many protons and
neutrons are in Pb-208?
How many protons and
132
neutrons are in 53 I?
Write a nuclear symbol for
Ca-42.
Relative Atomic Mass
Because measuring atomic mass
in grams would be cumbersome,
relative mass is measured in
atomic mass units (amu)
 1 amu = 1/12 mass of carbon-12
 Mass of proton & neutron is very
close to 1 amu

Examples
Find atomic masses for
copper, chlorine, and helium.
 Find formula masses for
sulfur dioxide, magnesium
chloride, and ammonium
sulfate

In Your Scientist’s Notebook:
What is the atomic mass of:
4.
Sodium (Na)
 Iron (Fe)
 Nitrogen (N)

What is the formula mass of:
5.
Sodium chloride—NaCl
 Sulfuric acid—H2SO4
 Aluminum hydroxide Al(OH)3

Average Atomic Mass
Weighted average of naturally
occurring isotopes
 To calculate:


S (relative abundance) x (mass)
Examples:

Calculate the relative atomic
mass of copper if 69.17% of
the isotopes are Cu-63 (62.930
amu) and 30.83% are Cu-65
(64.928 amu).
69.17% Cu-63 (62.930 amu) &
30.83% Cu-65 (64.928 amu)
Examples:

Calculate the average atomic
mass of carbon if 98.90% is C12 (12.000 amu) and the
remainder is C-14 (14.003 amu).
98.90% C-12 (12.000 amu) &
C-14 (14.003 amu)
In your Scientist’s
Notebook:
6.
Calculate the average
atomic mass of oxygen if
99.76% is O-16 (15.995
amu), 0.04% is O-17
(16.999 amu), and 0.20%
is O-18 (17.999 amu)
Counting Demo
The Mole
SI unit for amount of a
substance
 The amount of a substance that
contains the same number of
particles as exactly 12 g of C-12
 6.022 x 1023 of anything

Avogadro’s Number
6.022 x
23
10
Molar Mass



Related to atomic mass because
1 mole of carbon weighs 12 g.
12 amu per carbon atom :
12 grams per mole of carbon
1 amu
: 1 gram per mole
Molar Mass
One mole of any element has a
mass in grams equal to its
atomic mass.
 The mass in grams of one mole
of any compound is the sum of
the atomic masses of its
elements.

Conversions

Use factor label (dimensional
analysis) to convert between
grams, moles & number of
particles
Examples:

How many grams of carbon
are in 3.2 moles?
Examples:

How many atoms of carbon
are in 3.2 moles?
Examples:

How many moles are there
in 60.0 grams of Ca?
Examples:

How many grams of water
are in 1 mole of H2O?
In Your Scientist’s
Notebook:
7.
8.
9.
How many moles of Al in
13.50 g?
How many atoms of Xe in
.0087 mol?
How many atoms of Sn in
43.23 g?