A&P – Chapter 2: Basic Chemistry
Concepts of Matter and Energy
• Matter is divided into classes and forms
• 3 Classes of Matter
– Elements – simplest form
– Compounds – 2 or more substances are combined chemically
– Mixtures – 2 or more substances combined
• 4 Forms of Matter
– Plasma – biology: liquid part of blood or lymph; physics: hot ionized material consisting of
nuclei and electrons (found in the sun, most stars and fusion reactors)
– Liquids – ex. Blood plasma and gastrointestinal fluid – have a definite volume, but conform
to the shape of their container
– Gas – air we breathe – has neither a definite shape or volume
– Solids – ex. Bones and teeth – have a definite shape and volume
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Physical Change – does not alter the nature of the substance (ex: water to ice, stretched rubber band)
Chemical Change – does alter the nature of the science (ex: paper to ashes, iron to rust)
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Energy – commonly defined as the ability to do work or to put matter into motion; massless and does
not take up space
Kinetic energy – moving energy
Potential energy – energy of position, stored or at rest
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4 Forms of Energy
– Chemical – the energy stored in the bonds of chemical substances; gasoline molecules are
broken down apart in your car engine so the energy released powers your car
– Electrical – energy that results from the movement of charged particles; electrical energy
flowing through wiring or electrical current produced by ions moving across cell membranes
– Mechanical – directly involved in moving matter; pedaling a bicycle
– Radiant – energy that travels in waves; x-ray, infrared, light, radio, and ultraviolet; light
energy
Dmitrii Mendeleev
• 1834-1907
• Created the 1st periodic chart
• He arranged the chart according to
the elements atomic weight
– Weights can change
Henry Moseley
• 1887-1915
• Created the 2nd period chart
• He arranged the chart according to
the atomic number
• This is the same chart we use today
John Dalton
• 1766-1844
• Formulated the Atomic Theory:
– He held that all the atoms of an element are of exactly the same size and
weight and are in these two respects unlike the atoms of any other
element. He stated that atoms of the elements unite chemically in simple
numerical ratios to form compounds. The best evidence for his theory
was the experimentally verified law of simple multiple proportions,
which gives a relation between the weights of two elements that
combine to form different compounds. (www.factmonster.com)
JJ Thomson
• 1856 – 1940
• Discovered the electron-part of an atom
– Smallest part of an atom
– Awarded the Nobel Prize in Physics in 1906
• Discovered the isotope
Niels Bohr
• 1885-1962
• Made one of the 1st atomic models
• Nobel Prize in Physics in 1922
• We use the Bohr model today
• Part of the team that worked on the
Manhattan Project
– project that created one of the 1st
atomic bombs
Composition of Matter
• 118 total elements
– 94 are natural occurring elements
– The remaining elements are synthetic (produced artificially)
• 4 elements make up approx. 96% of the body: carbon, oxygen, hydrogen, and nitrogen
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Atom – the building block (smallest unit) of an element; the smallest particle that still retains its
special properties
– Protons – p+ – a positive charge; found in the nucleus
– Neutrons – n0 – uncharged; found in the nucleus
– Electrons – e- – a negative charge; equal in strength to the protons; orbit outside the nucleus
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Planetary Model of an atom portrays the atom as a miniature solar system with the
electrons moving around the nucleus in a fixed orbit
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Orbital Model of an atom is more modern of an atomic structure and is more useful
in predicting the chemical behavior of atoms because electrons do not follow specific orbits
– The electron cloud shows regions by denser shading where electrons are most likely to be
found
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Identifying Elements:
– Atomic number = the # of protons of the particular atom
– Atomic mass number = the sum of the protons and neutrons contained in its nucleus
– Atomic weight = atomic mass number (as long as you are talking about a single type of
element)
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Isotopes
– Variations of an element; still have the same atomic number but different atomic masses.
• Almost all elements exhibit 2 or more structural variations
– The # of neutrons are different from the original element
– Their chemical properties remain the same
– As a general rule, the atomic weight of any element is approximately equal to the mass
number of its most abundant isotope.
Hydrogen (1H)
(1p+, 0n0, 1e-)
Deuterium (2H)
(1p+, 1n0, 1e-)
Tritium (3H)
(1p+, 2n0, 1e-)
Carbon-14 is well known as a tool for dating geological and archeological artifacts. It decays to form Nitrogen-14 with a half-life of
5730 yr. It gives reasonably accurate setting for the age of objects between 500 and 50,000 years old.
Carbon-13 is used in magnetic resonance studies of organic molecules.
Carbon-12 is more abundant than others because natural processes favor it over the others; the element shown on the periodic table.
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Radioisotopes
– Heavier isotopes of certain atoms are unstable and tend to decompose so that they become
more stable
– The “glue” that hold the atomic nuclei is weaker in the heavier isotopes; this process of
spontaneous atomic decay is called radioactivity (it can be compared to a tiny explosion)
– All types of radioactive decay involve the ejection of particles (alpha or beta) or
electromagnetic energy (gamma rays) from the atom’s nucleus and are damaging to living
cells.
• Alpha emission has the least penetrating power
• Gamma radiation has the most
– Radioisotopes are used in minute amount to tag biological molecules so that they can be
followed or traced through the body and are valuable tools for medical diagnosis and
treatment
• PET scans (positron-emission tomography) – Alzheimer’s
• Radioisotope of iodine – thyroid tumor screenings
• Radium, cobalt (and certain others) – cancer treatment
The Periodic Table of Elements
• Vertical columns - Groups (or families) – 18 columns; similar properties and configurations in their
outermost electron shells
• Horizontal rows – Periods – 8 rows; elements within the same period do not have the same properties
but have the same # of electron shells; the higher the # of protons, the less metallic an element
becomes
• Stair-steps – separates nonmetals and metals
– Metals - Elements that form cations (positively charged ions) when compounds of it are in
solution and oxides of the elements form hydroxides rather than acids in water. Most metals
are conductors of electricity, have crystalline solids with a metallic luster and have a high
chemical reactivity. Many of these elements are hard and have high physical strength. The
metal series includes all elements of the alkali, alkali-earth, inner-transition (lanthanides and
actinides series), transactinides and transition series as well as some elements of the metalloid
series.
• **The exception is Hydrogen…
• Alkali Metals - With the exception of francium, these metals are all soft and silvery.
They may be readily fused and volatilized with their melting and boiling points
becoming lower with increasing atomic mass. They are the strongest electropositive
metals. These elements react vigorously, even violently with water.
• Alkaline Earth Metals - These elements are in general white, differing by shades of
color or casts; they are malleable, extrudable and machinable. These elements may be
made into rods, wire or plate. Also, these elements are less reactive than the alkali
metals and have higher melting points and boiling points.
• Inner Transition Metals - The thirty elements of the Lanthanides and Actinides series,
which are sub-series of the Transition Metals.
• Lanthanides - (rare earth metals) - The fourteen elements of the upper row on
the inner-transition metals on the periodic table that follow the element
lanthanum (#57). Some reference sources include lanthanum in this series
others do not.
• Actinides - The fourteen elements in the bottom row of the inner-transition
elements of the periodic table that follow the element actinium (Ac #89). Some
reference sources include actinium in this series others do not.
• Transition Metals/Elements - This series include all elements in the sub-series
Lanthanides and Actinides of the inner-transition elements and at least part of the sub-
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series Transactinides, which are the elements following the Actinides series. In general
these elements are known for their hardness, high density, high melting point and
boiling point and heat conduction although there are exceptions.
Other Metals
Metalloids - refers to elements that exhibit some properties of metals and nonmetals. These
elements tend to be semiconductors.
• Example: Silicon
• Without the semiconductive properties silicon you would not be reading this
text right now as most microchips and microprocessors are made with silicon
and without these processors computers as we now know them would not exist.
Nonmetals - These elements form anions (negatively charged ions) and differ markedly
from metals in respect to electronegativity and thermal and electrical conductivity. These
elements, in general, are poor conductors and have a high electronegativity.
• Other nonmetals
• Halogens - The reactive nonmetals that are in Group 17 of the periodic table. All of
these elements are electronegative.
• Noble gases - These elements are very unreactive, however, they are not nonreactive
as compounds containing these elements have been synthesized. There are no naturally
occurring compounds that are made up of these elements.
Left side stair-step separates other metals
Right side stair-step separates Nonmetals
Molecules and Compounds
• Molecules
– When 2 or more of the same atoms combine chemically
– The smallest unit of a compound and retains the properties of the compound
• Compound
– When 2 or more different atoms combine, they are referred to as a compound
• Chemical reactions occur whenever atoms combine with, or dissociate from, other atoms
• When atoms unite chemically, chemical bonds are formed
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Role of electrons
– Electrons closest to the nucleus are those most strongly attracted to its positive charge
– Electrons farther away are less securely held and easier to interact with other atoms
The only electrons that are important when considering bonding behavior are those in the outermost
shell, called the valence shell
– Its electrons determine the chemical behavior of a atom
– The “rule of 8s” – atoms interact in such a way that they will have 8 electrons in their valence
shell, this is the key to chemical reactivity
• Valence shell complete – chemically inactive – Inert element
• Valence shell incomplete – chemically active – Reactive element
Valence number - the number of electrons that are gained or lost during bonding
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Chemical Bond – the transferring or sharing of electrons by a chemical reaction
2 Types of Bonds
– Ionic Bonds – transfers electrons completely from one atom to another; combines metals and
nonmetals
– Covalent Bonds – shares electrons; combines nonmetals to nonmetals OR have hydrogen with
it (the exception)
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Ionic Bonds
– Ions – charged particles
• Bonding cause the positive and negative charges to become unbalanced
• Gaining electrons – atom acquires a negative charge (more electrons (-) than protons
(+)); specifically called anions
• Losing electrons – atoms become positively charged (more protons (+) than electrons
(-)); called cations
– Both anions and cations result when an ionic bond is formed
– Occurs between metals and nonmetals
Sodium atom (Na)
(11p+; 12n0; 11e-)
Chlorine atom (Cl)
(17p+; 18n0; 17e-)
Sodium (Na+)
Chlorine atom (Cl-)
Sodium Chloride
Sodium gains stability by losing one electron.
Chlorine becomes stable by gaining one electron.
After electron transfer, sodium becomes a sodium ion (Na +) and chlorine becomes a chloride ion (Cl-)
The oppositely charged ions attract each other.
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Covalent Bonds
– Atoms can become stable by sharing electrons in such a way that each atom is able to fill its
valence shell at least part of the time
– Occurs between nonmetals and nonmetals
• With the exception of Hydrogen….
• Any time Hydrogen is being bonded with another atom, it is a covalent bond!!
Oxygen atom (O)
(8p+; 8n0; 8e-)
Hydrogen atoms (H)
(1p+; 0n0; 1e-) each
Hydrogen (H2+) Oxygen (O-2)
Water
Oxygen gains stability by gaining one electron through the sharing of 2 Hydrogen electrons.
The compound is held together by Hydrogen bonding.
Although as a whole the compound is neutral, Oxygen now carries a slight negative charge and Hydrogen carries a slight
positive charge – this is a polar covalent bond.
The oppositely charged ions attract each other.
Valences of Common Elements
+1
+2
Silver
Ag+1
Iron
Sodium
Na+1
Calcium
Potassium K+1
Barium
Hydrogen H+1
Copper (II)
Lithium
Li+1
Magnesium
Fe+2
Ca+2
Ba+2
Cu+2
Mg+2
+3
Iron (III)
Aluminum
Fe+3
Al+3
-1
Chlorine
Iodine
Bromine
Fluorine
-2
Cl-1
I-1
Br-1
F-1
Sulfur
Oxygen
S-2
O-2
Magnesium Oxide
(MgO) = Mg+2O-2
Aluminum Chloride
(AlCl3) = Al+3Cl3-1
Hydrogen Sulfide
(H2S) = H2+1S-2
Biochemistry
• All chemical found in the body fall into one of two major classes of molecules:
– Organic compounds – carbon-containing compounds
• Carbohydrates
• Lipids
• Proteins
• Nucleic acids
– Inorganic compounds – lack carbon and tend to be simpler, smaller molecules
• Water
• Salts
• Acids and bases (not all)
• Organic and inorganic compounds are equally essential for life
Inorganic Compounds
• Water – the most abundant compound in the body; accounts for about 2/3 of body weight
– High-heat capacity – absorbs and releases large amounts of heat; prevents sudden body
temperature changes that might result from environmental factors
– Polarity/solvent properties – water is the “universal solvent”. Solvent is a substance in which
another substance can be dissolved or suspended. A solute is the substance being dissolved
(can be a solid, liquid, or gas). The resulting mixture is called a solution when the solute
particles are exceeding minute and equally distributed. A suspension is a mixture where the
solute particles are fairly large and settle over time. Translucent mixtures with solute particles
of intermediate size are called colloids
– Chemical reactivity – water is an important reactant in some types of chemical reactions such
as hydrolysis reactions where water is added to the bonds of larger molecules (i.e., digesting
foods or breaking down biological molecules)
– Cushioning – protective function; serves as a lubricant and cushion around certain organs
• Salts – the most plentiful salts found in the body are those containing calcium and phosphorus –
found in bones and teeth.
• Vital to body functioning:
– Sodium and Potassium essential for nerve impulses
– Iron forms part of the hemoglobin molecule
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Salts dissolved in the body easily separate into their ions through the process of dissociation
All salts are electrolytes – substances that conduct an electrical current in solution (compounds that
ionize when dissolved)
Acids and Bases – Acids and bases are electrolytes – ionize and dissociate in water and can then
conduct an electrical current
Acids have a sour taste and can dissolve many metals or “burn” a hole in your rug
– A substance that can release hydrogen ions (H+) in detectable amounts
– Acids are also defined as proton donors – when acids are dissolved in water, it is the release
of protons that determine an acid’s effects on the environment
• Hydrochloric acid – digestion – strong acid because it ionizes completely
• Acetic acid – weak acid because of incomplete ionization
• Carbonic acid – weak acid
Bases have a bitter taste, feel slippery, and are proton acceptors
– The release of hydroxyl ions (OH-) are proton seekers and is considered a strong base
– Bicarbonate ion (HCO3-) is a weak base (important in blood)
When acids and bases are mixed they react with each other to form water and a salt; a process called
neutralization reaction.
pH: Acid-Base Concentrations
pH units – concentration of hydrogen and hydroxyl ions
The idea for the pH scale was devised in 1909 by a Danish biochemist name Sorensen and is based on
the number of protons in solution expressed in terms of moles per liter
The pH scale runs from 0 to 14 – each successive change of 1 pH unit represents a 10-fold change in
hydrogen-ion concentration.
Below 7: acidic
7: neutral
Above 7: base (or alkaline)
Buffers
• Buffer – a substance or substances that help to stabilize the pH of a solution
• Present in body fluids (lessens/softens the shock of levels to body systems)
• Homeostasis of acid-base balance is carefully regulated by the kidneys, lungs, and blood.
• Normal blood pH varies in a narrow range from 7.35 to 7.45. when blood pH changes more than a
few tenths of a pH unit from these limits, death becomes a distinct possibility.
• When blood pH begins to dip into the acid range, the amount of life-sustaining oxygen that the
hemoglobin in blood can carry to body cells begins to decline rapidly to dangerously low levels.
Organic Compounds - Carbohydrates
• Carbohydrates, include sugars and starches, contain carbon, hydrogen, and oxygen – ratio is usually
1:2:1.
• Glucose – C6H12O6
• Ribose – C5H10O5
• Classified according to size:
– Monosaccharides – simple/single sugars
• Glucose – blood sugar – universal cellular fuel
• Fructose – fruit sugar – converted to glucose for use by body cells
• Galactose – milk sugar – converted to glucose for use by body cells
• Ribose – part of structure of nucleic acids (RNA)
• Deoxyribose – part of structure of nucleic acids (DNA)
– Disaccharides – double sugars – when two simple sugars combine through dehydration
synthesis (removal of water)
• Sucrose – glucose + fructose – cane sugar
• Lactose – glucose + galactose – milk
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• Maltose – glucose + glucose – malt sugar
– Polysaccharides – many sugars – polymers of many monosaccharides
• Starch – plants – storage polysaccharide – found in grains and root vegetables
• Glycogen – animals – slightly smaller than starch – storage polysaccharide (fat) –
found largely in muscles and the liver
Provides a ready to use source of food energy for cells – when broken down releases ATP molecules
Organic Compounds - Lipids
• Large, diverse group of insoluble compounds – contain large amounts of carbon hydrogen, and
smaller amounts of oxygen atoms (i.e., tristearin: C57H110O6)
• Enter the body as fat-marbled meats, egg yolks, milk products, and oils
• Found in the body as neutral fats, phospholipids, and steroids
– Neutral fats: triglycerides – found in fat deposits – protect and insulate body organs – major
source of stored energy in the body
– Phospholipids: cephalin and others – found in cell membranes – participate in transport of
lipids in plasma – abundant in brain and nervous tissue (form white matter)
– Steroids:
• Cholesterol – the basis of all body steroids
• Bile salts – released by the liver into the digestive tract to aid in fat digestion and
absorption
• Vitamin D – produced in the skin; necessary for normal bone growth and function
• Sex hormones – estrogen and progesterone (female hormones) and testosterone (male
sex hormone); necessary for reproductive function
• Adrenal hormones – coritsol – a long-term antistress hormone that is necessary for life;
aldosterone – helps regulate salt and water balance in body fluids by targeting the
kidneys
Organic Compounds - Proteins
• Account for over 50% of the organic matter in the body and have the most varied functions of the
organic molecules.
• The building blocks of proteins are amino acids.
– 20 common varieties of amino acids are found in proteins
• All amino acids are composed of:
– Amine group – gives the basic properties
– Acid group – allows them to act as acids
– R-group – makes each amino acid chemically unique
• Amino acids are joined together in chains to form large, complex protein molecules that contain from
50 to thousands of amino acids.
– The sequence in which amino acids are bound together produce proteins that vary widely both
in structure and function
• Structural proteins (fibrous proteins) – strand like – appear most often in body structures;
important in binding structures together and providing strength
– Collagen – found in bones, cartilage, and tendons – the most abundant in the body
– Keratin – hair and nails – makes skin tough
• Globular proteins – mobile, general spherical – play crucial roles in all biological processes; also
called Functional Proteins
– Antibodies – provide immunity
– Hormones – help regulate growth and development
– Transport proteins – transport various substances through the body
• Hemoglobin – transports oxygen in the blood
• Other blood proteins transport iron, cholesterol, and other substances
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Enzymes – biological catalysts that regulate essentially every chemical reaction that goes on
within the body
• Catalyst – a substance that increases the rate of a chemical reaction without becoming
part of the product or being changed itself
Organic Compounds - Nucleic Acids
• Fundamental role: make up the genes (basic blueprint of life)
– Not only do they determine what type of organism you will be, but also direct your growth
and development
• Done entirely by dictating protein structure
• Nucleic acids are composed of carbon, oxygen, hydrogen, nitrogen, and phosphorus atoms
– The largest biological molecules in the body
• The building blocks are nucleotides.
– Three basic parts:
• A nitrogen-containing base
• A pentose (5-carbon) sugar
• A phosphate group
• Five varieties of bases:
– 2-ring bases:
• Adenine
• Guanine
– Single-ring bases:
• Cytosine
• Thymine
• Uracil
• 2 major kinds of nucleic acids:
– Deoxyribonucleic acid (DNA) – genetic material found within the cell nucleus
• 2 fundamental roles:
• Replicates itself exactly before a cell divides; ensuring genetic information in
every body cell is identical
• Provides instruction for building every protein in the body
– Ribonucleic acid (RNA) – located outside the nucleus; the “molecular slave” of DNA –
carries out the orders for protein synthesis by DNA
• Although made up of nucleotides, RNA and DNA structures are completely different
– DNA – a long double chain coiled up into a spiral staircase like structure (also called a double
helix); bases are A, T,C, G
– RNA – a single nucleotide strand; bases are A, U, C, G