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When you crush a lump of sugar, you can see that it is made
up of many smaller pieces of sugar. You may grind these
particles into a very fine powder, but each tiny piece is still
sugar. Now suppose you dissolve the sugar in water. The
tiny particles seem to disappear completely. Even if you
look at the sugar-water solution through a powerful
microscope you cannot see any sugar particles. Yet if you
were to taste the solution, you’d know that the sugar is still
there. Observations like these led early philosophers to
ponder the fundamental nature of matter. Is it continuous
and infinitely divisible, or is it divisible only until a basic,
invisible particle that cannot be divided further is reached?
The Atom: From Philosophical Idea to
Scientific Theory
The particle theory of matter was supported as
early as 400 b.c. by certain Greek thinkers, such as
Democritus. He called nature’s basic particle an
atom, based on the Greek word meaning
“indivisible.” Aristotle was part of the generation
that succeeded Democritus. His ideas had a lasting
impact on Western civilization, and he did not
believe in atoms. He thought that all matter was
continuous, and his opinion was accepted for
nearly 2000 years. Neither the view of Aristotle
nor that of Democritus was supported by
experimental evidence, so each remained
speculation until the eighteenth century. Then
scientists began to gather evidence favoring the
atomic theory of matter.
Foundations of Atomic
Theory

Nearly all chemists in late 1700s accepted the definition of an
element as a substance that cannot be broken down further

Knew about chemical reactions

Great disagreement as to whether elements always combine in the
same ratio when forming a specific compound
Law of Conservation of Mass

With the help of improved balances, investigators could
accurately measure the masses of the elements and compounds
they were studying

This lead to discovery of several basic laws

Law of conservation of mass  states that mass is neither destroyed
nor created during ordinary chemical reactions or physical changes
Law of Definite Proportions

law of definite
proportions  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
Each of the salt crystals shown
here contains exactly
39.34% sodium and 60.66%
chlorine by mass.
Law of Multiple Proportions

Two elements sometimes combine to form more than one compound

For example, the elements carbon and oxygen form two compounds,
carbon dioxide and carbon monoxide

Consider samples of each of these compounds, each containing 1.0 g
of carbon

In carbon dioxide, 2.66 g of oxygen combine with 1.0 g of carbon

In carbon monoxide, 1.33 g of oxygen combine with 1.0 g of carbon

The ratio of the masses of oxygen in these two compounds is exactly
2.66 to 1.33, or 2 to 1
1808 John Dalton

Proposed an explanation for the law of conservation of mass,
the law of definite proportions, and the law of multiple
proportions

He reasoned that elements were composed of atoms and that
only whole numbers of atoms can combine to form compounds

His theory can be summed up by the following statements
Dalton’s Atomic Theory
1.
All matter is composed of extremely small particles called atoms.
2.
Atoms of a given element are identical in size, mass, and other
properties; atoms of different elements differ in size, mass, and
other properties.
3.
Atoms cannot be divided, created or destroyed.
4.
Atoms of different elements combine in simple whole-number ratios
to form chemical compounds.
5.
In chemical reactions, atoms are combined, separated, or
rearranged.
Modern Atomic Theory

Dalton turned Democritus’s idea into a scientific theory which was
testable

Not all parts of his theory have been proven correct

Ex. We know atoms are divisible into even smaller particles

We know an element can have atoms with different masses
Although john dalton thought atoms were indivisible,
investigators in the late 1800s proved otherwise. As
scientific advances allowed a deeper exploration of matter,
it became clear that atoms are actually composed of several
basic types of smaller particles and that the number and
arrangement of these particles within an atom determine
that atom’s chemical properties. Today we define an atom
as the smallest particle of an element that retains the chemical
properties of that element.
The Structure of the Atom
All atoms consist of two regions
Nucleus  very small region located near the center of an atom
1.
1.
2.
2.
In the nucleus there is at least one positively charged particle called
the proton
Usually at least one neutral particle called the neutron
Surrounding the nucleus is a region occupied by negatively
charged particles called electrons
Discovery of the Electron

Resulted from investigations into the
relationship between electricity and matter

Late 1800s, many experiments were performed
 electric current was passed through different
gases at low pressures

These experiments were carried out in glass
tubes known as cathode-ray tubes
Cathode Rays and
Electrons

Investigators noticed that
when current was passed
through a cathode-ray
tube, the opposite end of
the tube glowed

Hypothesized that the
glow was caused by a
stream of particles, which
they called a cathode ray

The ray traveled from the
cathode to the anode
when current was passed
through the tube
Observations

1. cathode rays deflected by magnetic field in same way as wire
carrying electric current (known to have negative charge)

2. rays deflected away from negatively charged object

Observations led to the hypothesis that the
particles that compose cathode rays are
negatively charged

Strongly supported by a series of experiments
carried out in 1897 by the English physicist
Joseph John Thomson

He was able to measure the ratio of the charge of cathode-ray
particles to their mass

He found that this ratio was always the same, regardless of the
metal used to make the cathode or the nature of the gas inside the
cathode-ray tube

Thomson concluded that all cathode rays are composed of identical
negatively charged particles, which were later named electrons
Charge and Mass of the Electron

Confirmed that the electron carries a negative
electric charge

Because cathode rays have identical properties
regardless of the element used to produce them, it
was concluded that electrons are present in atoms
of all elements

Cathode-ray experiments provided evidence that
atoms are divisible and that one of the atom’s basic
constituents is the negatively charged electron

Thomson’s experiment revealed that the electron
has a very large charge for its tiny mass

Mass of the electron is about one two-thousandth
the mass of the simplest type of hydrogen atom
(the smallest atom known)

Since then found that the electron has a mass of
9.109 × 10−31 kg, or 1/1837 the mass of the
hydrogen atom

Based on information about electrons, two
inferences made about atomic structure
1.
b/c atoms are neutral they must have positive
charge to balance negative electrons
2.
b/c electrons have very little mass, atoms must
have some other particles that make up most of the
mass
Thomson’s Atom

Plum pudding model (based on English dessert)

Negative electrons spread evenly through positive charge of
the rest of the atom

Like seeds in a watermelon
Discovery of Atomic
Nucleus

1911 by New Zealander Ernest Rutherford and
his associates Hans Geiger and Ernest Marsden

Bombarded a thin, gold foil with fast-moving
alpha particles (positively charged particles
with about four times the mass of a hydrogen
atom)

Assume mass and charge were uniformly
distributed throughout atoms of gold foil (from
Thomson’s model of the atom)

Expected alpha particles to pass through with
only slight deflection
What Really Happened…

Most particles passed with only slight deflection

However, 1/8,000 were found to have a wide
deflection

Rutherford explained later it was “as if you have
fired a 15-inch artillery shell at a piece of tissue
paper and it came back and hit you.”
Explanation

After 2 years, Rutherford finally came up with an
explanation

The rebounded alpha particles must have experienced some
powerful force within the atom

The source of this force must occupy a very small amount of
space because so few of the total number of alpha particles
had been affected by it

The force must be caused by a very densely packed bundle
of matter with a positive electric charge

Rutherford called this positive bundle of matter the nucleus

Rutherford had discovered that the volume of a
nucleus was very small compared with the total
volume of an atom

If the nucleus were the size of a marble, then the size
of the atom would be about the size of a football
field

But where were the electrons?

Rutherford suggested that the electrons surrounded
the positively charged nucleus like planets around
the sun

He could not explain, however, what kept the
electrons in motion around the nucleus
Rutherford’s Atom
Composition of Atomic
Nucleus

Except hydrogen, all atomic nuclei made of two kinds of particles


Protons
Neutrons

Protons = positive

Neutrons = neutral

Electrons = negative

Atoms are electrically neutral, so number of protons and electrons IS
ALWAYS THE SAME

The nuclei of atoms of different elements differ in
the number of protons they contain and therefore in
the amount of positive charge they possess

So the number of protons in an atom’s nucleus
determines that atom’s identity
Forces in the Nucleus

Usually, particles that have the same electric charge
repel one another

Would expect a nucleus with more than one proton
to be unstable

When two protons are extremely close to each other,
there is a strong attraction between them

Nuclear forces  short-range proton-neutron,
proton-proton, and neutron-neutron forces that hold
the nuclear particles together
The Sizes of atoms

Area occupied by electrons is electron cloud – cloud
of negative charge

Radius of atom is distance from center of nucleus to
outer portion of cloud

Unit – picometer (10-12 m)

Atomic radii range from 40-270 pm

Very high densities – 2 x 108 tons/cm3
Counting Atoms
Consider Neon, Ne, the gas used in many illuminated signs. Neon is a
minor part of the atmosphere. In fact, dry air contains only about 0.002%
Ne. And yet there are about 5 x 1017 atoms of neon present in each breath
you inhale. In most experiments, atoms are too small to be measured
individually. Chemists can analyze atoms quantitatively, however, by
knowing fundamental properties of the atoms of each element. In this
section, you will be introduced to some of the basic properties of atoms.
You will then discover how to use this information to count the number
of atoms of an element in a sample with a known mass. You will also
become familiar the the mole, a special unit used by chemists to express
amounts of particles, such as atoms and molecules.
Atomic Number
3
Li
Lithium
6.941
[He]2s1

Atoms of different elements
have different numbers of
protons

Atomic number (Z)  number
of protons in the nucleus of
each atom of that element

Shown on periodic table

Atomic number identifies an
element
Isotopes

Hydrogen and other atoms can
contain different numbers of
neutrons

Isotope  atoms of same
element that have different
masses
Isotopes of Hydrogen
Mass Number

Mass number  total
number of protons and
neutrons in the nucleus
of an isotope
Naming Isotopes

Isotopes are usually identified by specifying their mass number

There are two methods for specifying isotopes
1.
Mass number is written with a hyphen after the name of the
element
1.
2.
Shows the composition of a nucleus as the isotope’s nuclear
symbol
2.
1.
2.

Tritium is written as hydrogen-3
Refer to this method as hyphen notation
Uranium-235 is written as 23592U
The superscript indicates the mass number and the subscript
indicates the atomic number
Nuclide – general term for specific isotope of an element
Sample Problem

How many protons, electrons, and neutrons are there in an atom of
chlorine-37?

atomic number = number of protons = number of electrons

mass number = number of neutrons + number of protons

mass number of chlorine-37 − atomic number of chlorine = number of
neutrons in chlorine-37

mass number − atomic number = 37 (protons plus neutrons) − 17 protons =
20 neutrons
Practice Problems
1.
How many protons, electrons, and neutrons are in an atom of bromine80?

Answer 35 protons, 35 electrons, 45 neutrons
2. Write the nuclear symbol for carbon-13.

Answer 136C
3. Write the hyphen notation for the element that contains 15 electrons and
15 neutrons.

Answer phosphorus-30
Relative Atomic Masses

Because atoms are so small, scientists use a standard to control the
units of atomic mass  carbon-12

Randomly assigned a mass of exactly 12 atomic mass units (amu)

1 amu is exactly 1/12 the mass of a carbon-12 atom

Atomic mass of any atom determined by comparing it with mass
of C-12

Ex. H  1/12 C-12, so 1 amu
Average Atomic Masses

Average atomic mass  the weighted average of the atomic masses of
the naturally occurring isotopes of an element

Ex. You have a box containing 2 size of marbles


25% are 2.00g each
75% are 3.00g each
25 x 2.00g = 50g
75 x 3.00g = 225g
50g + 225g = 275g
Divide by 100 – average marble mass = 2.75g
Method 1
25 x 2.00g = 50g
75 x 3.00g = 225g
50g + 225g = 275g
Divide by 100 – average marble mass = 2.75g
Method 2
25% = 0.25
75% = 0.75
(2.00g x 0.25) + (3.00g x 0.75) = 2.75g
Calculating Average Atomic Mass
Copper
69.17% Cu-63 – 62.929599 amu
30.83% Cu-65 – 64.927793 amu
(0.6917 x 62.929599) + (0.3083 x 64.927793)
= 63.55 amu
Relating Mass to Numbers of Atoms

The relative atomic mass scale makes it possible to know how
many atoms of an element are present in a sample of the element
with a measurable mass

Three very important concepts provide the basis for relating
masses in grams to numbers of atoms
1.
The mole
2.
Avogadro’s number
3.
Molar mass
The Mole

SI unit for an amount of substance (like 1 dozen = 12)

Mole (mol)  amount of a substance that contains as many
particles as there are atoms in exactly 12g of C-12
Avogadro’s Number

Avogadro’s number  the number of particles in exactly one mole of a
pure substance

6.022 x 1023

How big is that?

If 5 billion people worked to count the atoms in one mole of an element,
and if each person counted continuously at a rate of one atom per
second, it would take about 4 million years for all the atoms to be
counted
Molar Mass

Molar mass  the mass of one mole of a pure substance

Written in unit g/mol

Found on periodic table (atomic mass)

Ex. Molar mass of H = 1.008 g/mol
Example
How many grams of helium are there in 2 moles of helium?
2.00 mol He x
= ? g He
2.00 mol He x 4.00 g He
1 mole He
= 8.00 g He
Practice Problems
What is the mass in grams of 3.50 mol of the element copper, Cu?
222 g Cu
What is the mass in grams of 2.25 mol of the element iron, Fe?
126 g Fe
What is the mass in grams of 0.375 mol of the element potassium, K?
14.7 g K
What is the mass in grams of 0.0135 mol of the element sodium, Na?
0.310 g Na
What is the mass in grams of 16.3 mol of the element nickel, Ni?
957 g Ni
1. A chemist produced 11.9 g of aluminum, Al. How many moles of
aluminum were produced?
0.441 mol Al
2. How many moles of calcium, Ca, are in 5.00 g of calcium?
0.125 mol Ca
3. How many moles of gold,Au, are in 3.60 × 10−10 g of gold?
1.83 × 10−12 mol Au
Conversions with Avogadro’s
Number
How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver?
Given: 3.01 × 1023 atoms of Ag
Unknown: amount of Ag in moles
Ag atoms ×
3.01x1023atomsAg x
= moles Ag
=0.500 mol Ag
Practice problems
1. How many moles of lead, Pb, are in 1.50 × 1012 atoms of lead?
2.49 × 10−12 mol Pb
2. How many moles of tin, Sn, are in 2500 atoms of tin?
4.2 × 10−21 mol Sn
3. How many atoms of aluminum, Al, are in 2.75 mol of aluminum?
1.66 × 1024 atoms Al
1. What is the mass in grams of 7.5 × 1015 atoms of nickel, Ni?
7.3 × 10−7 g Ni
2. How many atoms of sulfur, S, are in 4.00 g of sulfur?
7.51 × 1022 atoms S
3. What mass of gold,Au, contains the same number of atoms as 9.0 g
of aluminum,Al?
66 g Au