1. Historical Context
● Ancient Greece: Leucippus and Democritus proposed the concept of atomos (indivisible
particles). However, Aristotle’s view—that matter is composed of four elements (earth,
air, fire, water) and is infinitely divisible—dominated for 2,000 years.
● Scientific Revolution: In 1807, John Dalton moved beyond philosophy by using
experimental data to support the existence of atoms.
2. Dalton’s Atomic Theory (The 5 Postulates)
Dalton’s theory serves as the foundation of modern chemistry. Its core tenets are:
1. Smallest Units: Matter is made of tiny particles called atoms; an atom is the smallest
unit that can participate in chemical change.
2. Elemental Identity: An element contains only one type of atom. All atoms of a specific
element have the same mass and chemical properties.
3. Distinct Differences: Atoms of one element are fundamentally different from atoms of
any other element.
4. Compound Formation: Compounds consist of atoms of different elements combined in
small, whole-number ratios. These ratios are constant for any given compound.
5. Conservation of Matter: Atoms are never created or destroyed during a chemical
reaction—they are only rearranged to create new substances.
3. Fundamental Chemical Laws
Dalton’s theory provides a microscopic explanation for two macroscopic laws:
The Law of Definite Proportions (Constant Composition)
● Definition: All samples of a pure compound contain the same elements in the exact
same proportion by mass.
● Example: Every sample of isooctane has a carbon-to-hydrogen mass ratio of $5.33:1$.
The Law of Multiple Proportions
● Definition: When two elements form more than one compound, the masses of one
element that combine with a fixed mass of the other are in a ratio of small, whole
numbers.
● Example: Copper and Chlorine can form two different compounds. One has twice as
much chlorine as the other ($2:1$ ratio), suggesting the atoms combine in different ratios
like $CuCl$ vs $CuCl_2$.
4. Key Takeaways from Examples
, ● Chemical Changes: If a visual representation of a reaction shows atoms disappearing
or appearing out of nowhere, it violates Dalton’s postulate regarding the conservation of
mass.
● Identifying Compounds: By calculating mass ratios (e.g., $grams\ of\ O \div grams\ of\
C$), we can determine if two samples are the same substance (Definite Proportions) or
different substances related by whole numbers (Multiple Proportions).
This section details the transition from Dalton’s solid-sphere model to the modern nuclear model
of the atom. Through a series of landmark experiments, scientists discovered that atoms are not
indivisible but are composed of specific subatomic particles.
1. The Discovery of the Electron (J.J. Thomson)
In the late 1800s, J.J. Thomson used cathode ray tubes to investigate electrical discharges.
● The Experiment: He applied high voltage to electrodes in a vacuum tube, creating a
"cathode ray." This beam was deflected toward positive charges and away from negative
ones.
● The Discovery: Thomson concluded the ray was made of negatively charged particles,
now called electrons.
● Key Finding: He calculated the charge-to-mass ratio of the electron and realized it
was significantly lighter than a whole atom, proving atoms have internal structure.
● The Model: Thomson proposed the "Plum Pudding" model, where negative electrons
were embedded in a positively charged "soup."
,
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2. Quantifying Charge (Robert A. Millikan)
In 1909, Millikan's "Oil Drop" experiment provided the missing pieces of the electron puzzle.
● The Experiment: He sprayed tiny oil droplets into a chamber and used an electric field
to halt their fall.
● The Discovery: He found that the charge on any drop was always a multiple of a
specific base value: $1.6 \times 10^{-19}\text{ C}$.
● Ancient Greece: Leucippus and Democritus proposed the concept of atomos (indivisible
particles). However, Aristotle’s view—that matter is composed of four elements (earth,
air, fire, water) and is infinitely divisible—dominated for 2,000 years.
● Scientific Revolution: In 1807, John Dalton moved beyond philosophy by using
experimental data to support the existence of atoms.
2. Dalton’s Atomic Theory (The 5 Postulates)
Dalton’s theory serves as the foundation of modern chemistry. Its core tenets are:
1. Smallest Units: Matter is made of tiny particles called atoms; an atom is the smallest
unit that can participate in chemical change.
2. Elemental Identity: An element contains only one type of atom. All atoms of a specific
element have the same mass and chemical properties.
3. Distinct Differences: Atoms of one element are fundamentally different from atoms of
any other element.
4. Compound Formation: Compounds consist of atoms of different elements combined in
small, whole-number ratios. These ratios are constant for any given compound.
5. Conservation of Matter: Atoms are never created or destroyed during a chemical
reaction—they are only rearranged to create new substances.
3. Fundamental Chemical Laws
Dalton’s theory provides a microscopic explanation for two macroscopic laws:
The Law of Definite Proportions (Constant Composition)
● Definition: All samples of a pure compound contain the same elements in the exact
same proportion by mass.
● Example: Every sample of isooctane has a carbon-to-hydrogen mass ratio of $5.33:1$.
The Law of Multiple Proportions
● Definition: When two elements form more than one compound, the masses of one
element that combine with a fixed mass of the other are in a ratio of small, whole
numbers.
● Example: Copper and Chlorine can form two different compounds. One has twice as
much chlorine as the other ($2:1$ ratio), suggesting the atoms combine in different ratios
like $CuCl$ vs $CuCl_2$.
4. Key Takeaways from Examples
, ● Chemical Changes: If a visual representation of a reaction shows atoms disappearing
or appearing out of nowhere, it violates Dalton’s postulate regarding the conservation of
mass.
● Identifying Compounds: By calculating mass ratios (e.g., $grams\ of\ O \div grams\ of\
C$), we can determine if two samples are the same substance (Definite Proportions) or
different substances related by whole numbers (Multiple Proportions).
This section details the transition from Dalton’s solid-sphere model to the modern nuclear model
of the atom. Through a series of landmark experiments, scientists discovered that atoms are not
indivisible but are composed of specific subatomic particles.
1. The Discovery of the Electron (J.J. Thomson)
In the late 1800s, J.J. Thomson used cathode ray tubes to investigate electrical discharges.
● The Experiment: He applied high voltage to electrodes in a vacuum tube, creating a
"cathode ray." This beam was deflected toward positive charges and away from negative
ones.
● The Discovery: Thomson concluded the ray was made of negatively charged particles,
now called electrons.
● Key Finding: He calculated the charge-to-mass ratio of the electron and realized it
was significantly lighter than a whole atom, proving atoms have internal structure.
● The Model: Thomson proposed the "Plum Pudding" model, where negative electrons
were embedded in a positively charged "soup."
,
Shutterstock
Explore
2. Quantifying Charge (Robert A. Millikan)
In 1909, Millikan's "Oil Drop" experiment provided the missing pieces of the electron puzzle.
● The Experiment: He sprayed tiny oil droplets into a chamber and used an electric field
to halt their fall.
● The Discovery: He found that the charge on any drop was always a multiple of a
specific base value: $1.6 \times 10^{-19}\text{ C}$.
