2. Learning Objectives
• You should be able to:
– Discuss the development of the atom from its
earliest model to the modern day atom.
– Identify the correct number of subatomic
particles for atoms, ions, and isotopes.
– Calculate the average atomic mass and mass
number of an atom from isotopic data.
3. Atoms & Elements
• An atom is the smallest particle of an element
that retains it identity in a chemical reaction.
• Elements are the simplest form of matter that
has a unique set of properties; an element
cannot be broken down into simpler substances
by chemical means.
5. Ancient Philosophy
• Who: Democritus (460 BC – 370 BC)
• When: 400 BC
• Where: Greece
• What: Democritus believed that matter was made of
small particles he named “atomos”. Democritus
believed that atoms were indivisible and
indestructible, and that different types of atoms had
different sizes and shapes.
• How: Democritus used observation and inference to
explain the existence of everything; examined
broken sea shells to generate his concept of atoms.
6. Atomic Theory
• Who: John Dalton (1766-1844)
• When: 1803
• Where: England
• Why: Building on the ideas of Democritus in ancient
Greece; applied scientific reasoning.
• What: ”Billiard Ball Model” -Described atoms as tiny
invisible particles that could not be created,
destroyed, or divided. Thought each element was
made of its own kind of atom.
• How: Studied gases, pressure and temperature
changes.
7. 4 Hypotheses of Dalton’s Atomic Theory
1. All matter is composed of tiny 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.
3. Atoms of different elements can physically mix
together or can chemically combine in simple whole-
number ratios to form compounds.
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.
Dalton’s Atomic Theory
9. Foundations of Atomic Theory
Law of Definite Proportions (Proust)
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 Multiple Proportions (Dalton)
If two or more different compounds are composed of the
same two elements, then the ratio of the masses of the
second element combined with a certain mass of the first
elements is always a ratio of small whole numbers.
Law of Conservation of Mass (Lavoisier)
Mass is neither destroyed nor created during ordinary chemical
reactions.
10. Discovery of Electrons
• Who: J. J. Thomson (1856-1940)
• When: 1897
• Where: England
• What: ”Plum Pudding Model” –Described atoms as
a solid sphere of positively charged material having
negatively charged electrons spread throughout.
Thomson discovered that electrons were smaller
particles of an atom and were negatively charged.
• Why: Thomson knew atoms were neutrally
charged, but couldn’t find the positive particle.
• How: Cathode Ray Experiment
12. Atomic Structure I
• Who: Ernest Rutherford (1871-1937)
• When: 1911 & 1920
• Where: England
• What: “Planetary Model” -Decided that atoms were
mostly empty space, but had a dense central core
(nucleus-1911) consisting of positively charged
protons (1920); electrons orbit the nucleus.
• Why: He knew that atoms had positive and negative
particles, but could not decide how they were
arranged.
• How: Gold Foil Experiment – isolated the protons.
14. Atomic Structure II
• Who: Niels Bohr (1885-1962)
• When: 1913
• Where: England
• What: “Bohr Model” -Proposed that electrons traveled
in fixed paths (energy levels) around the nucleus;
every atom has a defined number of energy levels.
Scientists still use the Bohr model to show the number of electrons in each
orbit around the nucleus.
• Why: Bohr was trying to show why the negative
electrons were not sucked into the nucleus of the atom.
• How: Worked with frequency and wavelength of
radiation; applied math.
16. Modern Concept of the Atom
• Electrons travel around the nucleus in random orbits.
• Scientists cannot predict where they will be at any
given moment.
• Electrons travel so fast, they appear to form a “cloud”
around the nucleus.
• The analogy is that of a “beehive” where the bees
are the electrons moving around the nucleus in a
“cloud” of energy levels.
• Major contributions from many scientists, including
Schrödinger (1926) and Chadwick (1932).
17. Modern Concept of the Atom
• Who: Erwin Schrödinger (1887-1961)
• When: 1926
• Where: Austria
• What: “Electron Cloud Model” -Proposed that
electrons are not confined to fixed energy levels;
rather they occupy volumes of space outside the
nucleus; electron energy is based on its location
(increases away from nucleus)
• How & Why: Used math to explain the probable
location of electrons; the denser the electron cloud,
the more likely the electron will be there.
19. Modern Concept of the Atom
• Who: James Chadwick (1891-1974)
• When: 1932
• Where: England
• What: “Nuclear Model” –Discovered the neutron;
proposed that neutrons are neutral particles located in
the nucleus of an atom and have a mass about that of
a proton.
• Why: Expanded upon Rutherford’s experiments; If
nuclei contained only protons their charge would be
much higher than measurements suggested.
• How: Used a neutron chamber in his experiments.
