Elements and Atoms
Chapter Two
Introduction
This chapter describes the origins of atomic theory, the
discovery and characterization of subatomic particles and
the locations of these particles in atoms, as well as the
method of designating the composition of individual
atoms. The existence of isotopes of elements are also
discussed. The arrangement of the elements in the
periodic table is then described followed by the
introduction of the mole as the quantity of matter that
connects the submicroscopic world of atoms with the
macroscopic world in which we live.
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Origins of the Atomic Theory
•
Ancient Greeks
The simple picture of atoms as tiny spheres in constant
motion goes back a long way in history - to the Greek
philosopher Leucippus and his student, Democritus
(460 - 370 BC). Democritus reasoned that if a bit of
matter were divided into smaller and smaller pieces,
one would ultimately arrive at a tiny particle that could
not be further divided. He described this particle with
the word "atom" which means uncuttable. This theory
was, however, untestable and unsupported and
remained so for more than 2000 years.
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Lavoisier

Antoine Lavoisier (1743-1794) - showed
that when a compound is combusted (reacted
with oxygen) in a closed system, there is no
loss in mass. This lead to the Law of
Conservation of Mass, which states "mass
can neither be created or destroyed in
chemical reactions."
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Proust


Joseph Proust (1754-1826) - formulated the
Law of Definite Proportions, which states
"different samples of a pure chemical
substance contain the same proportion of
elements by mass.“
Also called the Law of Constant Composition -
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John Dalton and his Atomic Theory


John Dalton (1766-1826) combined the
Law of Definite Proportions and the Law of
Conservation of Mass and formulated an
Atomic Theory.
1803 - Dalton linked the existence of
elements, which cannot be decomposed
chemically, to the idea of atoms, which are
indivisible.
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Postulates


All matter is made of atoms. These
indivisible and indestructible objects are
the ultimate chemical particles.
All atoms of a given element are identical,
both in mass and in properties. Atoms of
different elements have different masses
and different properties.
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Postulates cont.


Compounds are formed by combination of two
or more different kinds of atoms. Atoms
combine in the ratio of small whole numbers, fro
example, one atom of A with one atom of B, or
two atoms of A with one atom of B.
Atoms are the units of chemical change. A
chemical reaction involves only combination,
separation, or rearrangement of atoms, but
atoms are not created, destroyed, divided into
parts, or converted into other kinds of atoms
during a chemical reaction.
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Explained


Lavoisier's Law of Conservation of Matter Matter cannot be created or destroyed because
atoms cannot be created or destroyed.
Proust's Law of Constant Composition - If all
atoms of a given element are identical, and all
compounds are formed by atoms in small whole
number ratios - then the mass percent of an
element in a given compound will always be the
same.
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Proposed

Law of Multiple Proportions - When two
elements form two different compounds,
the mass ratio in one compound is a small
whole number times the mass ratio in the
other.
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Law of Multiple Proportions
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Atomic Structure
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Electricity




Electricity is involved in many of the experiments from
which the theory of atomic structure was derived.
Electrical charge was first observed and recorded by the
ancient Egyptians - amber attracted small objects when
rubbed with wool or silk.
Benjamin Franklin - positive and negative charges;
conservation of charge
Experiments with electroscope - unlike charges attract
one another, like charges repel one another
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Result
Electrical charges must be associated with
matter - perhaps atoms!
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Radioactivity


Henri Becquerel (1896) - Discovered that
uranium ore emitted rays when exposed to
photographic plates
Marie Curie (1898) - Isolated polonium and
radium, which also emitted the same kind of
rays. (1899) She suggested that atoms of
radioactive substances disintegrated when they
emit these rays - named the phenomenon
radioactivity.
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Types of Radiation
alpha, beta, gamma
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Result
The suggestion that atoms disintegrate
contradicts Dalton's idea that atoms are
indivisible, requires an extension of
Dalton's theory.
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Electrons

Michael Faraday (1833) - Same current
used in electrolysis caused different
quantities of different metals to deposit.
these quantities were related to the
relative masses of the atoms of those
elements. A fundamental particle of
electricity must exist - electron.
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Cathode Rays

Travel in straight lines, cause sharp
shadows, cause gases and fluorescent
materials to glow, can heat metal objects
red hot, can be deflected by a magnetic
field, and are attracted toward positively
charged plates.
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J.J. Thomson

(1897) - Calculated the charge to mass
ratio for particles in cathode ray beam by
applying magnetic and electrical fields and
using basic laws of electricity and
magnetism.
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J.J. Thomson’s Experiment
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Robert Millikan

Oil Drop Experiment - Measured charge on
particle. (1.60 x 10-19 C)
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RESULT
electron mass = 9.109389 x 10-28 g;
charge = -1.60217733 x 10-19 C.
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So! Thompson has
discovered the
electron, and it must
live inside atoms.
It is much less massive
than the the atom itself,
so perhaps we have
little electrons stuffed
into the ‘rest’ of the
atom like raisins in the
oatmeal, or:
Plum Pudding...
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Protons

Canal rays - observed in a special cathode-ray
tube with perforated cathode. Ray opposite to
cathode ray is observed. Charge to mass ratio
for the positive particles differed. Hydrogen has
largest charge to mass ratio- suggests hydrogen
provides positive particles with the smallest
mass - these were considered to be fundamental
positively charged particles of atomic structure
and were later called protons by E. Rutherford.
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Result

proton mass = 1.672623 x 10-24 g; charge
is equal and opposite to electron.
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Neutrons


Chadwick (1932) discovers neutron
RESULT - neutron mass = 1.6749286 x
10-24 g; charge = 0
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Nucleus


J.J. Thomson - proposed that the atom was a
uniform sphere of positively charged matter
within which thousands of electrons circulated in
coplanar rings. Tested theory by directing a
beam of electrons at a very thin metal foil expected results - many deflections - did not
occur.
Ernest Rutherford (1910)- tested Thomson’s
model by directing a beam of alpha particles at
thin foil of metal.
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Rutherford’s Experiment
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Rutherford’s Experiment
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Top: Expected results: alpha particles passing through the plum
pudding model of the atom undisturbed.
Bottom: Observed results: a small portion of the particles were
deflected, indicating a small, concentrated positive charge.
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Results

All of the positive charge and most of the
mass of the atom is concentrated in a very
small volume - the nucleus. The electrons
occupy the rest of the space in the atom.
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Important Notes


To chemists the electrons are the most
important part of the atom, because they
are the first part of the atom that contacts
another atom.
Number of protons and electrons are
equal in a given atom
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Atomic Composition

electron
mass = 9.109389 x 10-28 g;
charge = -1.60217733 x 10-19 C

proton
mass = 1.672623 x 10-24 g;
charge is equal and opposite to
electron.

neutron
mass = 1.6749286 x 10-24 g;
charge = 0
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Nucleus

Nucleus is made up of protons and
neutrons; electrons are found in space
about the nucleus.
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Electrically Neutral Atoms

For an atom which has no net electrical
charge, the number of negatively charged
electrons around the nucleus equals the
number of positively charged protons in
the nucleus
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Atomic Number

All atoms of the same element have the
same number of protons in the nucleus.
this number is called the atomic number
and is given the symbol Z.
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Mass Number

The sum of the number of protons and
neutrons in an atom is called the mass
number and is represented with the
symbol A.
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Nuclear Symbol of an Atom
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Isotopes

Atoms having the same atomic number Z
but a different mass number A. Isotopes
are atoms of the same element, but they
have different masses because they have
different numbers of neutrons.
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Atomic Mass

The atomic mass of a given atom is based on a
scale that is relative to a standard. That
standard is the mass of a carbon atom that has
six protons and six neutrons in its nucleus. Such
an atom is defined to have a mass of exactly 12
atomic mass units. The actual masses of atoms
have been determine experimentally using mass
spectrometers.
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Calculation of Atomic Mass



The atomic mass given for an element on the
periodic table is a weighed average determined
from the masses and abundances of the various
isotopes of a given element.
atomic mass =
(%isotope 1/100) (mass of isotope 1) +
(%isotope 2/100) (mass of isotope 2) + etc
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Calculating atomic masses

for boron, 19.9% occurs as 10B and 80.1%
occurs as 11B. The isotopic mass of 10B is
10.013 and 11B is 11.009
0.199  10.013 amu = 1.99 amu
11B 0.801  11.009 amu = 8.82 amu
Atomic mass of boron = 10.81 amu
 10B


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Atomic Mass and Periodic Table

An element is a group of atoms all having
the same atomic number although the
mass number may differ. The mass of the
atoms is a relative mass based on using
carbon -12 as the standard reference.
The mass reported on the periodic table is
a weighted average based on the
percentages of all the naturally occurring
isotopes.
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Relative Mass

Each element has an average relative
mass based on C-12. It follows logically
that if you mass out the gram relative
mass there must be the same number of
atoms for all the elements.
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Element



H
O
Na
Relative Mass of
one atom in "u"
Number of atoms in the
gram relative mass
1.0079
15.9994
22.98977
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N
N
N





One of the ways would be like this: take the actual
mass of the element and see how many atoms it would
take to get the relative mass in grams.
electron mass = 9.1093897 x 10-28
proton mass = 1.6726231 x 10-24
neutron mass = 1.6749286 x 10-24
The average value - when this is done over and over and when the mass defect of the nucleus is taken into
account is called Avogadro's number and it is equal to
6.022 x 1023. We call this a mole.
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The Mole - Definition

A mole is the amount of substance that
contains as many elementary entities
(atoms, molecules, or other particles) as
there are atoms in exactly 12 g of the
carbon - 12 isotope.
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The Mole – cont.

One mole of an element implies



The atomic mass expressed in grams,
different for each element
Avogadro’s number of atoms, which is the
same for all elements
A conversion factor between mass and
numbers of things (allows us to count atoms
by weighing)
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Molar Mass

When dealing with the elements in the
laboratory we work with gram masses and
when we use the same gram mass as the
relative atomic mass we call it a molar
mass - it represents the mass in grams of
one mole of the element.
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Mole of JuJuBes

The area of the 48 contiguous United
States is 3.02 x 106 square miles. If you
have a mole of JuJuBes and you covered
the United States, how thick of a layer
would you have?
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Examples

One mole of hydrogen atoms contains
Avogadro's number of hydrogen atoms
and has a mass of 1.0079 grams. We
usually write this as 1.0079 grams/mole
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
One mole of water molecules contains
Avogadro's number of water molecules
and has a mass of 2(1.0079) + 15.9994 =
18.0152 grams. Again - 18.0152
grams/mole.
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
One mole of sodium chloride formula units
contains Avogadro's number of formula
units and has a mass of 22.98977 +
35.453 = 58.443 grams - 58.443
grams/mole.
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Conversions
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Elemental Distribution



The Earth has 88 elements present in
measurable amounts, but 10 constitute
<99% of the crust
There are 114 know elements (26
spontaneously decay into other elements)
The human body is 93% carbon, hydrogen
and oxygen
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Names and Symbols of the Elements

Symbol - usually the first or second letters of the
name



for two letter symbols, the first letter is capitalized but
the second is not
two letters are used for most elements
When the name and symbol use different letters,
that means the element was known in antiquity,
and the Latin name was used for the basis of
the symbol. The only exception is Tungsten,
which uses the German name (Wolfram)
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Grouping the elements


The elements are grouped in several
categories, according to properties and
therefore according to position in the
periodic table.
Two major groupings are metals and
nonmetals
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The Periodic Table



Elements were originally arranged in
columns by chemical properties, then
atomic weight (Mendeleev)
The modern periodic table has the
elements by atomic number (Moseley)
The periodic law - the properties of the
elements are periodic (cyclically repeating)
functions of their atomic number
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Metals




to the left on the periodic table
ductile (can be drawn into wires)
malleable (can be pounded into sheets)
readily conduct electricity and heat
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Nonmetals


to the right on the periodic table
generally soft solids or gases
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Classes of metals

active metals



very reactive to air and water
include lithium, potassium and sodium
noble metals


very unreactive
gold, silver and copper
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Metalloids



elements that lie along the metalnonmetal border
have properties intermediate between
metals and nonmetals
most important current use is as
semiconductors
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Periodic divisions

Period




horizontal rows in the periodic table
properties of elements across a period change
dramatically
each period ends with a member of the noble gas
family
Group


vertical columns in the periodic table
properties of elements in the same group are similar
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Main group elements





Group
Group
Group
Group
Group
IA - alkali metals
IIA - alkaline earths
VIA - chalcogens
VIIA - halogens
VIIIA - noble gases
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Transition elements

Group B elements




metals with multiple oxidation states
includes noble or coinage metals (Cu, Ag, Au)
platinum group metals (Ru, Rh, Pd, Os, Ir, Pt)
structural metals such as Fe and Cr
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Inner transition elements

Lanthanides



often called rare earths due to the difficulty in
isolating pure samples
chemical properties almost identical
Actinides

radioactive elements, many synthetic
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Physical States of Elements




Reference condition is 1 atmosphere of
pressure and 25 °C
Gaseous elements - H2, N2, O2, F2, Cl2,
noble gases
Liquid elements - Hg, Br2
All other elements are solids
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