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CHEMISTRY REVIEW –
INTRODUCTORY CHEMISTRY
Why is it necessary to study chemistry? Living things are made of matter (anything that occupies space & has
mass) & matter follows the laws of chemistry. Even the characteristics we consider to be unique to living things
are the result of chemical reactions (ex. movement as a result of muscle contractions).
I. THE ATOM
A. Atom - smallest unit of matter unique to a particular element.
B. Element - A substance made up of only one kind of atom - ex. carbon contains only carbon atoms.
Elements can't be broken down into other substances by ordinary chemical means. Each element displays
unique properties (ex. some are gases, some are solids, some are metals, etc.).About 92 elements occur
naturally (there are also some man-made elements). Some of the elements important to our study of living
systems are carbon, oxygen, hydrogen, nitrogen, sodium, chlorine, phosphorus, potassium, calcium, sulfur,
sodium, magnesium, iodine, and iron. Carbon, oxygen, hydrogen, and nitrogen make up 96% of body
weight!
C. Some Important Things to Know about Atoms & Elements:
1. An atom consists of 2 basic parts:
a. nucleus - the nucleus contains protons & neutrons:
1.) protons - positively charged; all atoms have protons; protons give the nucleus a positive
charge.
2.) neutrons - neutral in charge; fxn.: "stabilizers"; all atoms except hydrogen have one or more
neutrons.
b. electrons – negatively charged; occupy orbit energy levels or shells around the nucleus; attracted
to the positive charge of the nucleus; in an atom, the number of electrons always equals the number
of protons, so the atom, as a whole, has no charge; electrons determine the chemical properties of
elements (ex. whether they are a liquid, solid, or gas, etc.).
2. atomic number = number of protons (or number of electrons); in an atom, the number of protons
always equals the number of electrons; this number differs for each element.
3. atomic mass number = number of protons + number of neutrons; the number of electrons is not
included in the mass number due to their insignificant mass.
Note: Nearly all elements have 2 or more variants called isotopes (atomic numbers are the same, but
mass numbers are different; ex. the mass number for carbon on the periodic table is Carbon 12, but you
may have heard of Carbon 14. The heavier isotopes of many elements are unstable and they
decompose into more stable forms. The process of decay is called radioactivity. As they decay
radioisotopes produce radiation and gradually lose their radioactivity at a certain rate. Functions:
Radioisotopes can be tagged to biological molecules that show up on special X-rays.
4. The 6 elements important for building organic molecules like carbohydrates, lipids, proteins, and
nucleotides are: carbon, hydrogen, oxygen, sulfur, nitrogen, & phosphorus.
Note: Be able to determine the number of electrons, protons, & neutrons in an atom, given the atomic
number & atomic mass. Ex. Sodium (Na) has an atomic number of 11 & an atomic mass of 23. A
Sodium atom then has 11 protons, 11 electrons, & 12 neutrons.
II. CHEMICAL REACTIONS
A. More about Electrons
1. Electrons orbit around the nucleus at different energy levels or shells the 1st shell (k shell) can hold no
more than 2 electrons; the next shell (l shell) can hold 8 electrons, 2 electrons in each of 4 orbitals; the
next shell (m shell) also holds 8 electrons. These will be the only shells that we will deal with in this
class.
2. An atom is the most stable when all of its shells are completely filled (the k shell fills first, then the l
shell, then the m shell, & so on).
3. Atoms with partially filled outer shells react with other atoms to completely fill their outer shells & in
doing so form chemical bonds; it is important to remember that it's the electrons of an atom that
participate in the chemical bonds that form between atoms. Molecules are formed when 2 or more
atoms are joined together by interactions between the electrons of their outer electron shells.
B. Chemical Bonds
1. Ionic bonds
a. definition - in ionic bonds electrons are transferred to other atoms to completely fill outer shells;
atoms are electrically neutral, but when they gain or lose electrons in combining with other atoms,
they are called ions (charged atoms) & they take on a positive or negative charge; in other words, the
transfer of electrons upsets the balance of protons & electrons in an atom; atoms that lose electrons
are positively charged, atoms that gain electrons are negatively charged; ionic bonds involve the
attractions between these oppositely charged ions. So, before you can have an ionic bond, you have
got to have oppositely charged ions, & to create ions you have got to transfer electrons.
b. example: NaCl (sodium chloride- table salt); Na (at. # 11) has one electron in its outer m shell - it
needs 7 electrons to fill this shell - it is easier for Na just to give this electron away, & eliminate the m
shell entirely. Cl (at. # 17) has 7 electrons in its outer m shell - it only needs 1 electron to fill its outer
shell. Therefore, when Na & Cl atoms react, Na gives up its outer electron to Cl. Because Na gives
up an electron, it now has 11 protons & 10 electrons, resulting in a positively charged atom. Cl now
has 17 protons & 18 electrons, resulting in a negatively charged atom. Na+ & Cl- ions are attracted to
each other because of their opposite charges & an ionic bond is formed. The attraction is the ionic
bond! Only the electron # changes when ions are formed!
2. Covalent bonds - more common in the human body & are more stable.
a. definition - Electrons are not transferred, but are shared. The shared electrons spend part of their
time around the nucleus of one atom & part of their time around the other. Each pair of electrons
shared equals one covalent bond (if 2 pair of electrons are shared between 2 atoms, a double
covalent bond is formed, a triple covalent bond occurs when 3 pr. of electrons are shared). We will
limit out discussion to single covalent bonds.
b. example: methane (CH4); a carbon atom can form four covalent bonds - it has 4 electrons in its
outer shell, therefore it needs 4 more electrons to fill its outer shell; hydrogen has one electron in its
outer shell, therefore it needs one electron to fill its outer shell. Rather than give 4 electrons away or
accept 4 electrons, carbon shares its 4 electrons in its outer shell with 4 Hydrogen atoms.
c. polar vs. nonpolar - covalent bonds can be nonpolar (ex. in methane) or polar (ex. in water); if both
atoms exert the same pull on the shared electrons (equal sharing), the bond is nonpolar (example:
methane); if there is unequal sharing of electrons, the bond is polar; in molecules with polar covalent
bonds, there is an atom that has a much larger nucleus (more protons) than the other atoms in the
molecule; the atom with the most protons is more attractive to the shared electrons, & so the
electrons spend most of their time around this atom's nucleus; all of these electrons spending most of
their time around a particular nucleus gives this part of the molecule a partial negative charge; the
other atom(s) in the molecule acquire a partial positive charge, because the shared electrons are not
spending much time around them. It’s important to realize that the shared electrons in polar covalent
bonds are not spending all of their time around a particular nucleus - if this were the case then we
would be talking about electrons being transferred (ionic bonds) not shared.
3. Hydrogen bonds - These bonds result from polar covalent bonds; they form between molecules &
occur between the slightly negative atom of one molecule & the slightly positive atom of another
molecule. These bonds can occur between hydrogen & oxygen & between hydrogen & nitrogen.
Hydrogen bonds are weaker than ionic & covalent bonds, because the charges on the molecules are
"partial" or weak charges.
C. Chemical Reactions - A chemical reaction occurs when atoms or molecules (called reactants) collide and
are transformed into different combinations of the same atoms or molecules (called products). In this
process, chemical bonds break and new ones form. In living systems special proteins called enzymes
catalyze these chemical reactions (that is they increase the rate of the chemical reaction). We'll talk more
about enzymes later.
Metabolism = Anabolism + Catabolism
Anabolic Reactions (constructive or building up)
Catabolic Reactions (destructive or breaking down)
INORGANIC COMPOUNDS
Inorganic compounds include water, salts, and many acids and bases.
I. WATER – most abundant and important compound in living material; makes 60-80% of the volume of
most living cells.
A. STRUCTURE (already discussed) – Water is a polar molecule (a molecule that has partially positive &
negative regions). Its polarity allows water molecules to react with one another, forming hydrogen bonds.
The same type of interaction is possible between water & many other polar substances. Polar substances
are hydrophilic (water loving) & nonpolar substances are hydrophobic (water-fearing).
B. PROPERTIES OF WATER - Hydrophilic & hydrophobic interactions underlie several properties of
water that are biologically important.
1. Temperature-Stabilizing Effects:
Note: The temperature of a substance is a measure of how fast its molecules are moving; the higher
the temperature of a substance, the faster its molecules are moving.
Heating water - it takes considerable heat to raise the temp. of water because the hydrogen bonds
between the water molecules restrict the movement of the molecules; in order for the temp. of water to
rise, a number of H bonds must be broken - this takes a lot of energy. This resistance to temp. change
helps living organisms to maintain a relative constant internal body temp.; this is important because
biochemical reactions take place within a narrow temperature range (this has to do with the action of
enzymes).
2. High Heat of Vaporization – when water evaporates it changes from a liquid to a gas. This
transformation requires that large amounts of heat be absorbed to break the hydrogen bonds; this is
beneficial to us when we sweat (when perspiration evaporates from out skin it take s heat with it).
3. Water As a Solvent - the polarity of water is also responsible for water's capacity as a solvent
(something that dissolves something else); water is an excellent solvent for ions & other polar
molecules (solutes) in living systems - it transports nutrients to and wastes away from body cells.
4. Reactivity – water is an important reactant in many chemical reactions; example: when some foods are
broken down into smaller building blocks it takes water - called hydrolysis reactions; when large
molecules are synthesized in the body (ex. proteins), water is removed and the reaction is called
dehydration synthesis.
II. ACIDS AND BASES
When molecules of inorganic acids, bases, or salts dissolve in water of body cells, they undergo ionization or
dissociation (they break apart into their individual ions). These molecules are called electrolytes.
A. Acids & Bases Generally Defined
Acid - a solute that releases or donates H+ ions in a solution [ex. HCl - hydrochloric acid dissociates into H+
ions & Cl- ions].
Base - a solute that removes H+ ions from a solution; many release OH- ions in this process [ex. Mg(OH)2 magnesium hydroxide dissociates into OH- ions & Mg++ ions; the OH- then combines with H+ and forms H2O].
B. pH Scale - Fluids are assigned a pH value (0 -14), which refers to the hydrogen ion concentration present
in the fluid. The hydrogen ion concentration is abbreviated as [H+].
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acid - pH below 7.0; base - pH above 7.0; neutral - pH = 7.0
pH = - log [ H+ ] (formula for calculating pH)
It is a common misconception to think that as the [H+] increases, the pH also increases! As [H+]
increases, pH decreases! This can be seen from the following example:
solution A: [H +] = 0.01
solution B: [H+ ] = 0.00000001
pH = -log[0.01] = 2
pH = -log[0.00000001] = 8
Can you figure out the quick way to find the pH of these solutions?
Solution A is more acidic than Solution B - Solution has a higher [H+] than Solution B (solution A has
more H+ than solution B Æ 0.01 > 0.0001).
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When you think about a pH value, think that this number is the number of decimal places in the
hydrogen ion concentration. The larger the pH value, the (more or less) decimal places there are in the
hydrogen ion concentration. As the pH increases it becomes more (acidic or basic).
C. Buffers - help maintain a constant pH by removing or adding H+ ions; the pH inside living systems is
generally between 7.35-7.45 (exception: the hydrochloric acid in the digestive system makes the pH here 23); this pH range is important, as many biochemical reactions take place only within this range; buffers can
combine with hydrogen ions &/or release them, & so help stabilize the pH. An important buffer system in
the blood is the bicarbonate buffer system.
H2CO3
ÆÆÆÆÆÆÆÆÆÆÆÆ
ÅÅÅÅÅÅÅÅÅÅÅÅ
H+ donor
(carbonic acid)
HCO3-
+
H+
H+ acceptor
(bicarbonate ion - base)
The above reaction is reversible – it responds to changes in pH. If pH is too high (too basic – not
enough H+), then the above reaction proceeds to the right. If pH is too low (too acidic – too many H+),
then the reaction proceeds to the left.
III. SALTS & ELECTROLYTES – A salt is an ionic compound containing cations other than hydrogen
ions and anions other than the hydroxide ion (ex. NaCl, KCl, calcium carbonate, calcium phosphate). When
salts are dissolved in water they dissociate into their component ions. All ions are electrolytes, substances that
conduct an electrical current in solution.
ORGANIC COMPOUNDS
4 Groups of Organic Compounds Important in Living Organisms:
Carbohydrates
Lipids
Proteins
Nucleotides
Organic compound defined – A compound containing carbon (with exception of carbon dioxide); found in all
living things.
THE CENTRAL ROLE OF CARBON
The Carbon Backbone - The processes of life are primarily the result of the chemistry compounds of carbon.
Because of carbon's tendency to form four covalent bonds in four different directions, carbon can form an
unbelievably large number of different compounds of high complexity; organic compounds derive their basic
shapes from the carbon atoms; this shape helps determine the compound's function in living systems.
Functional groups - The structure & behavior of organic compounds also depends on the properties of their
functional groups; functional groups are groups of atoms (ex. hydrogen, oxygen, nitrogen, phosphorus, and
sulfur) attached to the carbon backbone.
I. CARBOHYDRATES:
A. Structure: generally made up of only three elements: carbon, hydrogen, & oxygen
B. Three Principle Classes of Carbohydrates:
1. Monosaccharide
a. Structure - composed of single sugar molecule; the atoms in a sugar molecule can form straight
chain or a ring (rings are more common in the body).
b. Examples - Glucose, Ribose, Fructose, Galactose (notice the endings!)
c. Functions - monosaccharides are important energy molecules in living things; glucose is the
primary energy source for humans & many other animals; also important as building blocks of larger
sugars.
2. Oligosaccharides - composed of short chains of monosaccharides; Examples:
a. Sucrose (table sugar) is a disaccharide composed of glucose & fructose; sucrose is the form in
which sugars are transported in plants.
b. Lactose (milk sugar) is a disaccharide composed of glucose & galactose.
3. Polysaccharides - straight or branched chains of many monosaccharide units
a. Storage Polysaccharides
1.) Starch - sugar storage in plants.
2.) Glycogen – “animal starch;” principle storage form for glucose in higher animals; this energy
storage is short term; lipids are used for long term energy storage; glycogen is stored in liver &
muscle tissue.
b. Structural Polysaccharides
1.) Cellulose – forms the cell walls around cells in plants and some fungi; no cell walls around
animal cells! Our digestive system does not produce the enzymes to break down this type of
polysaccharide. Animals that have a mostly vegetarian diet (ex. cattle) require lots of bacteria
in their digestive systems. These bacteria do produce enzymes to break down theses
polysaccharides so that they can be used as an energy source by the animal.
II. PROTEINS
A. Protein Structure:
1. Proteins are composed of subunits called amino acids (there are 20 different amino acids that make up
proteins). Amino acids contain carbon, hydrogen, oxygen, and nitrogen. Some also contain sulfur.
Amino acids have a structure similar to the one below. The R stands for some other atom or group of
atoms bonded to the central carbon atom in the molecule. The sequence of amino acids in a chain
helps determine the structure & shape of the protein & therefore the function of the protein;
there are many possible combinations of amino acids that produce the many different kinds of proteins.
Diagram the general structure for an amino acid below:
2. Peptide bonds - linkage formed between one amino acid & another amino acid; the name of these
bonds is why chains of amino acids are called polypeptides.
3. Producing the three dimensional structure of a protein: We have been discussing proteins as
"chains of amino acids." However, the final structure of proteins is not a straight chain of amino acids.
Proteins are very complex, three-dimensional molecules, with numerous twists & folds. The amino acid
chain of every kind of protein is folded in a very specific way [the chain will twist & fold based on the
linkage of amino acids in a specific sequence & the environmental conditions (i.e., temperature & pH)].
There are several bonds & forces which give a protein its specific 3-D structure (i.e., hydrogen bonds,
ionic bonds, disulfide bridges, etc.); these bonds link distant parts of the molecule, forming loops, twists,
etc. Destruction of a protein's 3-D structure by extreme heat or pH is called denaturation (see
"enzymes" on the next page for how this occurs). Analogy for protein structure: Think of a phone cord.
Pull it straight (like a straight chain of a.a.), then let it twist, then roll the twisted cord into a ball. Every
type of protein folds and twists in a very specific way. This helps determine its function.
B.
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Some Specific Functions of Proteins:
cytoskeleton in cells
cell membrane transport proteins – transport substances across the cell membrane
cell movement – ex. flagellum on a sperm cell
the protein hemoglobin transports oxygen in the red blood cells
enzymes as catalysts (enzymes speed up the rate of chemical reactions); all life processes are primarily the
result of chemical reactions; molecules in living things require enzymes in order to react; without enzymes,
chemical reactions in living things can’t occur; see below for more information on enzymes.
antibodies in the immune response
many hormones are derived from proteins
clotting proteins in the blood (ex. albumin)
the proteins actin and myosin are involved in muscle contraction
the protein keratin waterproofs the outer layers of the skin
the protein collagen gives connective tissue elasticity
spindle fiber apparatus (involved in moving chromosomes around during cell division)
B. Enzymes – a large, globular protein molecule that accelerates a specific chemical reaction. Virtually all
chemical reactions that take place in cells involve enzymes!!! Most of a cell’s proteins are enzymes.
1. Why are enzymes needed? In order for particular molecules to react with one another, they must be in
close proximity & must collide with sufficient force to overcome the mutual repulsion of their negatively
charged electron clouds & to break existing chemical bonds within the molecules. The force with which
they collide depends on their kinetic energy (energy of motion). Most chemical reactions require an
initial input of energy to get started, which increases the kinetic energy of the molecules, enabling a
greater number of them to collide with sufficient force. In the chemistry lab, we can supply this energy
with heat. In a cell, many different reactions are going on at the same time, therefore heat cannot be
used as it would be nondiscriminatory (it would affect many reactions at the same time). Cells get
around this problem by using enzymes that serve as catalysts (they get the chemical reactions going).
The enzymes form a temporary association with the molecules that are to react, bringing them close to
one another & weakening the existing chemical bonds, making it easier for new ones to form.
2. Enzyme Structure & Function - Enzymes are large, complex, globular proteins consisting of one or
more polypeptide chains. The molecules that enzymes acts on are known as the substrates.
Enzymes are folded to form a groove or pocket (called an active site) on their surface into which the
substrate fits & where the chemical reactions take place. Diagram below:
3. Effects of Temperature & pH on Enzyme Function
a. Temperature: As the temperature increases, so does the rate of enzyme catalyzed reactions, but
only up to a certain point. At high temperatures, the enzymes are denatured (due to the vibration of
molecules at high temperatures, the bonds that maintain the enzyme's structure are broken & the
protein unfolds). If denaturation is severe, the damage to the enzyme is irreversible.
b. pH: The shape of the enzyme depends partly on attraction between positively & negatively charged
amino acids. As the pH changes (acidic - more H+, basic - fewer H+), these charges change,
changing the shape of the enzyme & its function. Remember: the optimum pH for most enzymes is
6-8. (stomach has a pH of 2)
III. LIPIDS
A.
General Structure - all lipids are mostly nonpolar (hydrophobic) & are insoluble in polar solvents such as
water; lipid structure varies greatly & is discussed below for each type.
B. Some General Functions:
1. long term energy storage (example: glycerides); energy is stored in the chemical bonds; excess
carbohydrates, proteins or fats are converted to triglycerides & are stored in adipose (fat) tissue.
2. structural (example: phospholipids make up the cell membrane of cells)
C. Types of Lipids:
1. Lipids with Fatty Acids:
a. Glycerides
1.) Structure: classified as mono-, di-, & triglycerides, depending on the number of fatty acids
attached a single glycerol molecule; glycerol has 3 carbon atoms & 3 hydroxyl (OH) groups; fatty
acids are long, nonpolar chains composed of hydrogen & 4 to 24 carbon atoms, with a carboxyl
(COOH) group at one end.
2.) Saturated fatty acids - all carbons in the fatty acid tails are joined together by single carbon to
carbon bonds & as many hydrogen atoms as possible are linked to the carbons (the carbons are
said to be "saturated" with hydrogens); triglycerides with many saturated fatty acids are solid at
room temperature; occur mostly in animal tissues, but also in a few plant products; examples:
butter, lard, cocoa butter, palm oil, coconut oil; the liver produces cholesterol from some
breakdown products of saturated fats.
3.) Unsaturated fatty acids - one or more double bonds occur between carbon atoms in the fatty
acid tails; this cuts down on the number of hydrogen atoms that can bond to the carbons; liquids
at room temperature (the double bonds create a kink in the shape of the molecule, preventing
the fatty acids from packing close together & becoming solidified); unsaturated fatty acids are
more common in plants
4.) Functions of Glycerides – Energy! For most organisms and cellular microorganisms, sugars in
excess of what can be stored as glycogen are converted into fats for more permanent storage.
b. Phospholipids
1.)
Structure - 2 fatty acids & 1 phosphate group are linked to a glycerol molecule; this molecule
has a dual nature - the molecule has nonpolar, hydrophobic tails (composed of fatty acid tails) &
a polar, hydrophilic head (composed of glycerol and a phosphate ion). Phospholipid molecules
orient themselves so that the heads (water loving) are attracted to water and the tails (water
fearing) repel water. Look at the tails. Notice that only one tail has a kink in it created by a
double bond. Why is this important?
2.)
Function: Structural. Phospholipids form the structural basis of cellular membranes. If there
is water inside the cell and outside the cell, how do they do this? What is this structure called?
Diagram/label a phospholipid molecule below.
2. Lipids without Fatty Acids: Steroids or Sterols
a. Structure - different from other lipids; they consist of 4 interlocking carbon rings with numerous
hydrogens attached; while they have no fatty acids, they are still nonpolar & hydrophobic, so they are
classified as lipids.
b. Some Examples:
1.) cholesterol - important component in cell membranes & serves as the starting material for the
synthesis of other steroids.
2.) sex hormones – testosterone, estrogen, progesterone
3.) bile salts - produced in the liver (stored in the gall bladder) & released into the small intestine
where it is involved in the emulsification of fats.
4.) hormones produced by the cortex of the adrenal gland – ex. cortisol - involved in
metabolism, resistance to stress, & is an anti-inflammatory.
5.) vitamin D (activated form is called calcitriol) – causes you to absorb more calcium from what
you are eating, so that blood calcium levels increase.
How is Cholesterol is obtained by the body?
IV. NUCLEOTIDES
A. Structure: nucleotide = phosphate(s) + monosaccharide + a nitrogen-containing compound (called a
base); it’s the bases that spell out the genetic message in DNA & RNA).
B. Functions:
1. Nucleotides are the basic subunits of nucleic acids such as
a. DNA (deoxyribonucleic nucleic acid) - carrier of the genetic message - makes up chromosomes
in the nucleus of the cell. The structure of the DNA molecule is a twisted helix (it looks like a
twisted ladder). The sides of the ladder are made up of sugars and phosphates; the rungs of the
ladder are made up of the nitrogenous bases. The bases spell out the genetic code.
b. RNA (ribonucleic acid) - transcribes genetic message present in DNA & produces proteins from it
.
2.
Nucleotides also make up the adenosine phosphates (ex. ATP - adenosine triphosphate – the energy
molecule in the cell).