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RSPT 1060
MODULE B: Basic Chemistry
Lesson #1
Atomic & Subatomic Matter
Why Chemistry?
• Respiratory Therapists must have a
basic knowledge of the principles of
chemistry …
– To better understand the functioning of the
human body
– To better appreciate the clinical concepts of:
•
•
•
•
Arterial blood gas interpretation
Fluid and electrolyte physiology
Nutrition
Pharmacology
Objectives
•
At the end of this module, the student will:
– Define terms associated with atomic and
sub-atomic matter.
– Differentiate between the types of
matter.
– Describe what each item in an element’s
box on the periodic table represents.
– Compare the composition of the
elements of the universe, the earth’s
crust and the human body.
– Differentiate between an atom,
elements, molecules and compounds.
Objectives
•
At the end of this module, the student will:
– Describe the components of an atom
and the purpose of each.
– Differentiate between atomic number,
atomic mass and mass number.
– Explain what an isotope is.
– Explain what determines physical and
chemical properties of an element.
Matter
• What is Matter?
– Anything that
• Takes up space
• Has mass (weight), and
• Can be perceived by the senses.
– If it’s “something” it’s matter, if it’s “nothing”,
it’s not matter
• The primary states of matter are:
– Solid
– Liquid
– Gas
Divisions of Matter
Matter
Pure
Substance
(Homogeneous)
Mixture
(Homogeneous or
Heterogeneous)
Elements
Solution
Compounds
Suspension
Colloid
Matter - Pure
Substances
• Matter in it’s simplest form.
– Atom, Element, Molecule, Compound
• Always the same regardless of where it is
found.
– Oxygen (O), water (H2O), table salt (NaCl)
• It cannot be broken down any further without a
chemical or nuclear reaction.
– It will then become a different substance.
• Uranium in a nuclear bomb
• Pure substances are homogenous.
– Uniform in structure or composition throughout .
Matter
• So…what is an element or compound?
Elements
• An element is a pure form of matter.
• Other pure forms of matter include:
– Atoms
– Molecules
– Compounds
Elements
• Large collection of atoms of the same
type.
– Substance that cannot be broken down
further and still maintain its identity.
– All atoms have same atomic number.
– Not bonded together, only existing
together.
• A listing of all the elements known to
man is called the Periodic Table.
Introduction to the Periodic
Table
117 Elements
1
Atomic Number
Hydrogen
Element Name
H
Symbol or Abbrev.
1.01
Atomic Mass Unit
(AMU)
http://www.ceet.niu.edu/mrdl/software/Periodic%20Table.htm
Elements of the
Universe
91% of all atoms are Hydrogen
9% of all atoms are Helium
The other 115 elements are found in traces.
Elements of the Earth’s
Crust
60.1% =
21.1% =
6.1% =
2.9% =
2.6% =
2.4% =
2.2% =
2.1% =
oxygen
silicon
aluminum
hydrogen
calcium
magnesium
iron
sodium
Elements of the Human
Body
Major
Water
Macronutrients
Micronutrients
Elements
•Hydrogen
•Oxygen
•Carbon
•Nitrogen
•66% of
the body.
•Calcium
•Iron
•Arsenic
•Chlorine
•Fluorine
•Boron
•Potassium
•Cobalt
•Copper
•Phosphorus
•Iodine
•Manganese
•Magnesium
•Zinc
•Molybdenum
•Sulfur
•Selenium
•Chromium
•Sodium
•Silicon
•Nickel
Atoms
• Smallest “particle” of an element.
Molecule
• Smallest “particle” of a pure
substances (element or compound)
bonded together.
– Combination of similar atoms (O2 element)
– Combination of different atoms (H2O compound)
Compound
• Substance composed of a large
collection of molecules. Can be broken
down by chemical means into
molecules or elements.
• Often will have properties unlike those
of its constituent elements.
Pure Substances
A
B
Elements
– A large
collection
of atoms
of a given
type.
Atoms
of
element
A
Atoms of
element
B
Atoms of
element
A&B
existing
together
Molecules
made from
element A &
B through a
process
called
bonding.
Compound – a
large collection
of molecules
made from
atoms from
element A & B
The Atom
• Smallest particle of an element which
still maintains the chemical properties
of the element.
• Head of a pin could hold 100 trillion
atoms.
The Atom
• If broken down further by a
nuclear reaction, an atom would
become particles:
– Electrons
– Protons
– Neutrons.
The Atom
• If broken down further, protons and
neutrons are made of subatomic
particles:
– Positrons
– Mesons
– Neutrinos
The Helium Atom
Nucleus
Proton
2 Protons (+) and
2 Neutrons (No Charge)
2ENeutron
Electron
Smallest particle of an element.
Atom - Composition
• Nucleus
– Proton (+) nucleon
– Neutron (No charge) nucleon

Electron cloud or shell

Electron (-)
Atom - Nucleus
• The nucleus is the small, dense
positively charged center of the
atom
– It contains protons and neutrons
(nucleons)
– The nucleus only comprises 1/100,000
of the size of the atom even though it
is constitutes the vast majority of the
atom’s mass.
Atoms - Nucleus
• Nucleons
– Protons
• One Proton is 1836 times the size of an electron
• The number of protons determines the atomic
number.
• The number of protons is always equal to the
number of electrons
– This allows for a neutral charge of the atom.
– Neutrons
• The number of neutrons can vary
• The number of neutrons determines the number
of isotopes an element will have.
– Isotope: One of two or more atoms having the same
atomic number but different mass numbers.
Atom - Electrons
• 99.9% of the atom is open space
where the electrons travel (electron cloud
or shell)
– 99.99% of an atom is the negatively charged
electron cloud
– This cloud actually determines the size of the
atom
Atom - Electrons
•Electrons do not contribute to the
mass of the atom; only the size.
EXAMPLE: If the electron cloud was the size
of Ford Field, the nucleus would be smaller
than a pea at the center of the field.
•The nucleus determines the mass.
Atom – Size & mass
• Electrons determine size
• Nucleus (protons & neutrons) determine
mass
Atom – Electron Number and
Arrangement
• The number and arrangement of the
electrons determine the chemical
properties of an element.
– How it acts in relation to other elements
– How it acts in a chemical reaction
The Periodic Table
112 Elements
1
Atomic Number
Hydrogen
Element Name
H
Symbol or Abbrev.
1.01
Atomic Mass Unit
(AMU)
Atomic Number
• The number of protons in the atom of a
given element.
• All atoms of an element have the same
number of protons and electrons.
– This never changes.
• Because atoms are neutral, the atomic
number also indicates the number of
electrons.
EXAMPLE: Boron has an Atomic # 5
5 protons & _____
5 electrons
This means there are _____

Atomic Mass Unit
• Abbreviated as (amu).
• Reflects the mass of the most frequently
found form of an element in nature.
• The unit amu is a unit of measure made up
by scientists.
• It is used as a unit of measure for a particle
that is extremely small.
– 1 amu = 1.6606 x 10-24 grams
Mass of an Atom and the
amu
• The mass of an atom is too small to
express in grams
– Hydrogen atom = 1.7 x 10-24 gram.
• The relative scale of atomic mass
units is used instead of grams &
scientific notation.
Comparative Example
• 12 eggs = One dozen
• Dozen is a unit of measure made up by
farmers. (not really)
• Dozen is a simple unit of measure that
represents a larger number (12)
Mass of an Atom
• Mass is composed mainly of the
mass of protons & neutrons
– Proton = 1 amu
– Neutron = 1 amu
Carbon and the amu
• All elements are compared to the mass of
carbon.
– 1 amu = 1/12 the mass of a Carbon atom
– Carbon has 6 protons & 6 neutrons
– It’s atomic mass is 12.011 AMU
• Carbon is a point of reference for all other
elements
– Hydrogen is 1/12 the mass of carbon so it has a
mass of 1 amu
• 1 proton & 0 neutron
– Magnesium has twice the mass of carbon so it has a
mass of 24 amu
• 12 protons & 12 neutrons
Isotope
• There may be different forms of atoms of
the same element.
– This occurs when the number of neutrons
varies.
– Atoms of the same element with differing
numbers of neutrons are called isotopes
Isotopes and Physical
Properties
• Neutrons will determine the physical
properties which vary slightly between
isotopes.
– Result: The same element may “appear”
slightly different depending on which isotope
you look at. All isotopes should “act” the
same because the electron numbers don’t
change.
• Only 20 elements exist without isotopes.
Isotopes and Medicine
• We hear about isotopes most often in
nuclear medicine.
– Body scans use isotopes (Xenon)
• Ventilation & perfusion of the lungs
• Bone scans
– Radioactive material is injected in the
blood or inhaled into the lungs
– Image forms on radiology film showing
areas that isotope has been exposed to
Key Facts about Isotopes
• Isotopes:
– Atoms of the same element
– BUT, Have different numbers of neutrons
• Atomic number on periodic table does not
change (Same # of Protons)
• Atomic Mass (amu) on periodic table does not
change (Average of most common isotopes)
• Mass number changes (Actual number of
protons and neutrons)
Example of an Isotope:
Chlorine
Example:
Chlorine #17 Atomic mass 35.45
• Most common form (76% of the time)
– Cl-35 with a mass of 34.97 amu
• Less common form (24% of the time)
– Cl-37 with a mass of 36.97 amu
Calculation:
– (0.76)(34.97) = 26.5772 amu
– (0.24)(36.97) = 8.8727 amu
35.45 amu (average on periodic table)
Isotope
Chlorine:
• Atomic mass – Atomic # = the average #neutrons
• 35 – 17 = 18
neutrons in most common form
Isotopes and Mass
Number
• “Mass Number”
– Each isotope has its own mass number.
– Not on the periodic table
– Is the actual total number of protons + neutrons
– The number of neutrons can change so the mass
number can change.
• Protons + Neutrons = Mass number
Isotope
Example:
Chlorine #17 Atomic mass 35.45
• Most common form 76% of the time
–
–
–
–
Cl-35 with a mass of 34.97 amu
Mass # 35
Mass # - Atomic # = Neutrons
35 – 17 = 18
• Less common form 24% of the time
–
–
–
–
Cl-37 with a mass of 36.97 amu
Mass # 37
Mass # - Atomic # = Neutrons
37 – 17 = 20
Example:
(K)
Potassium
• Atomic Number = 19
• Mass Number = 39
• Mass Number = 40
• Protons = _______
• Electrons = _________
• Neutrons = Mass # – Atomic #
= ___________ or __________
Oxygen Isotopes
(AMU
15.9994)
Mass number
14
O
15
O
16
O
Atomic number
8
8
8
Number of
neutrons =
(Mass # Atomic #)
?
?
?
Number of
protons &
electrons
?
?
?
Symbol
ASSIGNMENTS
• Read: Chemistry Book to assist in
completing the objectives.
• Self-Assessment