Chapter 3
Atoms and Moles
Early Atomic theory
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4 element theory
Matter is made of atomos
Alchemy
Robert Boyle – proved water could be broken
down; disproved 4 element theory
Laws
• As many new elements were being
discovered, scientists realized that there were
certain laws that governed how they were
made into compounds
Law of definite proportions
• States that a chemical compound always
contains the same elements in exactly the
same proportions by mass
• In other words, a compound always has the
same chemical formula – change the formula,
you change the compound
Law of definite proportions (cont)
For example, water is always 88.8% oxygen and
11.2% hydrogen
So the formula for water must be H2O and
nothing else
Law of Conservation of Mass
• The mass of the reactants in a reaction equals
the mass of the products
• Mass cannot be created or destroyed in
ordinary chemical and physical changes.
Dalton’s Atomic Theory
•In 1808, John Dalton developed an atomic
theory.
•Dalton believed that a few kinds of atoms
made up all matter.
•According to Dalton, elements are composed
of only one kind of atom and compounds are
made from two or more kinds of atoms.
• Dalton’s model of the atom was a solid sphere
Dalton’s Atomic Theory (cont.)
Dalton’s Theory Contains Five Principles
1. All matter is composed of extremely small
particles called atoms, which cannot be
subdivided, created, or destroyed.
2. Atoms of a given element are identical in
their physical and chemical properties.
3. Atoms of different elements differ in their
physical and chemical properties.
Dalton’s Atomic Theory (cont.)
Dalton’s Theory Contains Five Principles (cont.)
4. Atoms of different elements combine in simple,
whole-number ratios to form compounds.
5. In chemical reactions, atoms are combined,
separated, or rearranged but never created,
destroyed, or changed.
• Data gathered since Dalton’s time shows that the
first two principles are not true in all cases.
Subatomic Particles
• Experiments by several scientists in the mid-1800s
led to the first change to Dalton’s atomic theory.
Scientists discovered that atoms can be broken
into pieces after all.
• The smaller parts that make up atoms are called
subatomic particles.
• The three subatomic particles that are most
important for chemistry are the electron, the
proton, and the neutron.
Subatomic Particles, continued
Electrons Were Discovered Using Cathode Rays
• To study current, J. J. Thomson pumped most of
the air out of a glass tube. He applied a voltage to
two metal plates, called electrodes, which were
placed at either end of the tube.
• One electrode, called the anode, was attached to
the positive terminal of the voltage source, so it
had a positive charge.
• The other electrode, called a cathode, had a
negative charge because it was attached to the
negative terminal of the voltage source.
CRT
Subatomic Particles, continued
Electrons Were Discovered Using Cathode Rays
•Thomson observed a glowing beam that came
out of the cathode and struck the anode and
the nearby glass walls of the tube.
• He called these rays cathode rays.
• The glass tube Thomson used is known as a
cathode-ray tube (CRT).
• CRTs are used in television sets, computer monitors,
and radar displays.
Subatomic Particles, continued
An Electron Has a Negative Charge
•Because the cathode ray came from the
negatively charged cathode, Thomson
reasoned that the ray was negatively charged.
• Thomson confirmed this prediction by seeing how electric and
magnetic fields affected the cathode ray.
•Thomson also observed that when a small
paddle wheel was placed in the path of the
rays, the wheel would turn.
• This suggested that the cathode rays consisted of tiny particles that
were hitting the paddles of the wheel.
Subatomic Particles, continued
An Electron Has a Negative Charge
• Thomson’s experiments showed that a cathode
ray consists of particles that have mass and a
negative charge.
• These particles are called electrons.
• An electron is a subatomic particle that has a
negative electric charge.
• Electrons are negatively charged, but atoms have
no charge.
• Atoms contain some positive charges that
balance the negative charges of the electrons.
• Thomson’s model of the atom
Rutherford
• Ernest Rutherford performed the gold foil
experiment, which disproved the plum-pudding
model of the atom.
• A beam of small, positively charged particles, called alpha
particles, was directed at a thin gold foil.
• Rutherford’s team measured the angles at which the
particles were deflected from their former straight-line
paths as they came out of the foil.
• Rutherford found that most of the alpha particles
shot at the foil passed straight through the foil.
But very few were deflected, in some cases
backward.
Gold Foil Experiment
Gold Foil Experiment (cont)
• Rutherford reasoned that only a very concentrated
positive charge in a tiny space within the gold atom
could possibly repel the fast-moving, alpha particles
enough to reverse the alpha particles’ direction.
• Rutherford also hypothesized that the mass of this
positive-charge containing region, called the nucleus,
must be larger than the mass of the alpha particle.
• Rutherford argued that the reason most of the alpha
particles were undeflected, was that most parts of
the atoms in the gold foil were empty space.
Gold Foil Experiment (cont)
Rutherford
•The nucleus is the dense, central portion of the
atom.
•The nucleus is made up of protons and
neutrons.
•The nucleus has all of the positive charge,
nearly all of the mass, but only a very small
fraction of the volume of the atom.
Rutherford Model
• Rutherford added a nucleus with protons to
the atom
Bohr
• Rutherford did not know why electrons did
not get attracted to nucleus
• Was solved by Bohr
– Put electrons in orbitals
– Said electrons were spinning very fast
– Planetary model
Exciting electrons
• Adding energy excites electrons
• When an electron becomes excited, it
jumps to a higher energy level
• the electron cannot stay in that excited
state forever and so it must fall back to its
original energy level
• The energy is then released as light
– Law of conservation of energy
Exciting electrons
Atomic Number and Mass Number
• Atomic # = # of protons
– Is the identifying characteristic of an atom
– In a neutral atom it also = # of electrons
• Mass # = # protons + # neutrons
Isotopes
• Isotopes
– An atom that has the same # of protons, but a
different # of neutrons
– Example:
• Carbon-12 has 6 protons and 6 neutrons
• Carbon-14 has 6 protons and 8 neutrons
Quantum Numbers
• Schrondinger – father of quantum
mechanics
• Devised quantum numbers to define the
region in which an electron can be found
• There are 4 quantum numbers
• No 2 electrons can have the same 4
quantum numbers
Quantum numbers (cont)
1. 1st quantum number
= principal quantum
number
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Symbol = n
Indicates the energy
level the electron
occupies
n = 1, 2, 3, …
Quantum numbers (cont)
2. 2nd quantum
number = angular
quantum number
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Symbol = l
Indicates the shape
of the orbital
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s = spherical
p = dumbbell
d = double dumbbell
f = star/asterik
Quantum numbers (cont)
3. 3rd quantum number = magnetic
quantum number
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Symbol = m
Indicates the orbital’s orientation around the
nucleus
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S has 1
P has 3
D has 5
F has 7
Quantum numbers (cont)
4. 4th quantum number = spin quantum
number
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Symbol = ms
Indicates the spin direction of the electron
Since each orbital orientation can only hold
2 electrons, ms can only = +1/2 or -1/2
Each electron must have opposite spins
Electron configuration
• Pauli exclusion principal
– Each orbital can only hold a maximum of 2
electrons
Electron configuration (cont)
• The Aufbau Principle
– Electrons fill orbitals that have the lowest
energy first
– The smaller the principal quantum number,
the lower the energy
Electron configuration (cont)
• Electron configuration tells the location of
each electron
– Its energy level
– The shape of its orbital
– Which orbital it is in
– The direction it is spinning
Electron configuration (cont)
• Placing the electrons of
an element in their proper
location:
– Start with the lowest
energy level first
(closest to the
nucleus) and work
your way out
From:
http://www.ccs.k12.in.us/chsBS/kons/kons/w
hat_is_organic_chemistry.htm
Electron configuration (cont)
• Fill in electrons in order
• Remember that:
– s can only hold 2 (1 orbital
x 2 electrons)
– p can hold 6 (3 orbitals x 2
electrons each)
– d can hold 10 (5 orbitals x
2 electrons)
– f can hold 14 (7 orbitals x
2 electrons)
Order of energy levels
Counting Atoms
• Masses of atoms are expressed as atomic
mass units (amu)
Moles
• Mole is the number of atoms in exactly 12
grams of carbon-12
– This is equivalent to 6.02 x 1023 atoms
• Avogadro’s number
– It is the SI unit for amount of a substance
– Molar mass = mass in grams of 1 mole of an
element or compound
• This equals the atomic mass of an element
Converting
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Grams/molar mass = moles
Moles x molar mass = grams
Moles x 6.02 x 1023 = atoms
Atoms/ 6.02 x 1023 = moles