2-Chemical-Basis-of-Life.pptx

CHEMICAL BASIS OF LIFE
HUMAN ANATOMY AND PHYSIOLOGY
CHAPTER 2

CHAPTER SECTIONS
1. Basic chemistry
2. Chemical Reactions
3. Acids and Bases
4. Inorganic Molecules
5. Organic Molecules
Chapter 1: The Human Organism

LEARNING OUTCOMES
▪ Define chemistry and state its relevance to anatomy and physiology
▪ Define matter, mass, and weight
▪ Distinguish between an element and an atom
▪ Enumerate the four most abundant elements in the human body
▪ Define atomic number, mass number, isotope, atomic mass, and mole
▪ Name the subatomic particles of an atom, and indicate their mass, charge, and location
▪ Compare and contrast ionic and covalent bonds
▪ Explain how intermolecular forces are formed, and relate its importance
▪ Differentiate between a molecule and a compound
Section 2-1. Basic Chemistry

LEARNING OUTCOMES
▪ Explain hydrogen bond formation, and relate its importance
▪ Describe solubility and the process of dissociation
▪ Predict if a compound or molecule is an electrolyte or a nonelectrolyte

CHEMISTRY: AN INTRODUCTION
▪ Chemistry is the scientific discipline concerned with the atomic composition and
structure of substances and the reactions they undergo
▪ Why is chemistry important in the study of human anatomy and physiology?
▪ Chemicals make up the body’s structures, and the interactions of chemicals with one
another are responsible for the body’s functions
▪ Nerve impulse generation
▪ Digestion
▪ Muscle contraction
▪ Metabolism
▪ Many abnormal conditions and their treatments can also be explained in chemical
terms, even though their symptoms appear as malfunctions in organ systems
▪ Parkinson disease – lack of dopamine in nerve cells of the brain

MATTER, MASS, AND WEIGHT
▪ Matter is anything that occupies space and has mass
▪ All living and non-living things are composed of matter
▪ Mass is the amount of matter in an object
▪ International unit for mass is the kilogram (kg), which is the mass of a platinum-
iridium cylinder kept at the International Bureau of Weights and Measurements in
France
▪ The mass of all other objects is compared with this cylinder
▪ For example, a 2.2-lb lead weight and 1 liter (L) of water each have a mass of
approximately 1 kg
▪ Weight is the gravitational force acting on an object of a given mass

ELEMENTS AND ATOMS
▪ An element is the simplest type of matter having unique chemical properties
▪ It is a substance that cannot be separated into simpler substances by chemical means
▪ About 96% of the body’s weight results from the following elements:
▪ Carbon (C)
▪ Hydrogen (H)
▪ Oxygen (O)
▪ Nitrogen (N)
▪ However, many other elements also play important roles in the human body
▪ Calcium helps form bones
▪ Sodium ions are essential for neuronal activity
▪ Some of these elements are present in only trace amounts but are still essential in life

ELEMENTS AND ATOMS
▪ An atom is the smallest particle of an element that has the chemical characteristics of
that element
▪ An element is composed of atoms of only one kind
▪ The element carbon is composed of only carbon atoms
▪ The element oxygen is composed of only oxygen atoms

ATOMIC STRUCTURE
▪ The characteristics of matter result from the structure, organization, and behavior of
atoms
▪ Characteristics of subatomic particles of an atom:
▪ Neutrons have no electrical charge
▪ Protons have positive charges
▪ Electrons have negative charges
▪ Positive charge of a proton is equal in magnitude to the negative charge of an
electron
▪ Number of protons and number of electrons are equal, and the individual charges
cancel in each other; therefore, each atom is electrically neutral

ATOMIC STRUCTURE
▪ Characteristics of subatomic particles of an atom:
▪ Protons and neutrons form the nucleus at the center of an atom, which accounts for
99.97% of an atom’s mass, but only 1-ten-trillionth of its volume
▪ Electrons, which account for most of the atom’s volume, move around the nucleus in
a region known as the electron cloud
▪ Each element is uniquely defined by the number of protons in the atoms of that
element
▪ Only hydrogen atoms have one proton
▪ Only carbon atoms have 6 protons
▪ Only oxygen have 8 protons

ATOMIC STRUCTURE
▪ The number of protons in each atom is called the atomic number
▪ Because the number of protons and number of electrons are equal, the atomic number
is also the number of electrons
▪ Protons and neutrons have the same mass, and they are responsible for most of the
mass of atoms
▪ Electrons, on the other hand, have very little mass
▪ The mass number of an element is the number of protons plus the number of neutrons
in each atom
▪ For example, the mass number for carbon is 12 because it has 6 protons and 6 neutrons

Question 2-1
The atomic number of fluorine is 9, and the mass number is 19.
What is the number of protons, neutrons, and electrons in an atom
of fluorine?

ISOTOPES
▪ Isotopes are two or more forms of the same element that have the same number of
protons and electrons but a different number of neutrons
▪ Isotopes have the same atomic number but different mass numbers

ISOTOPES
▪ For example, there are three isotopes of hydrogen:
▪ Hydrogen
▪ Deuterium
▪ Tritium
▪ All three isotopes have 1 proton and 1 electron, but hydrogen has no neutrons in its
nucleus, deuterium has 1 neutron, and tritium has 2 neutrons
▪ Isotopes can be denoted using the symbol of the element preceded by the mass number
(number of protons and neutrons) of the isotope
▪ Therefore, hydrogen is 1H, deuterium is 2H, and tritium is 3H

ISOTOPES
▪ Individual atoms have very little mass
▪ A hydrogen atom has a mass of 1.67 × 10−24 g
▪ To avoid working with such small numbers, scientists use a system of relative atomic
mass
▪ In this system, a dalton (Da), or unified atomic mass unit (u), is 1/12 the mass of 12C, a
carbon atom with 6 protons and 6 neutrons
▪ Therefore, 12C has an atomic mass of exactly 12 Da
▪ However, a naturally occurring sample of carbon contains mostly 12C and a small
quantity of other carbon isotopes, such as 13C, which has 6 protons and 7 neutrons

ISOTOPES
▪ The atomic mass of an element is the average mass of its naturally occurring isotopes,
taking into account the relative abundance of each isotope
▪ For example, the atomic mass of the element carbon is 12.01 Da, which is slightly more
than 12 Da because of the additional mass of the small amount of other carbon isotopes
▪ Because the atomic mass is an average, a sample of carbon can be treated as if all the
carbon atoms had an atomic mass of 12.01 Da

MOLE AND MOLAR MASS
▪ Just as a grocer sells eggs in lots of a dozen, a chemist groups atoms in lots of 6.022 ×
1023, which is called Avogadro’s number, or 1 mole (abbreviated mol)
▪ A mole of a substance contains Avogadro’s number of entities, such as atoms, ions, or
molecules
▪ Molar mass is the mass of 1 mole of a substance expressed in grams
▪ Molar mass is a convenient way to determine the number of atoms in a sample of an
element
▪ Because 12 g of 12C is used as the standard, the atomic mass of an entity expressed in
unified atomic mass units is the same as the molar mass expressed in grams

MOLE AND MOLAR MASS
▪ Therefore, carbon atoms have an atomic mass of 12.01 Da, and 12.01 g of
carbon has Avogadro’s number(1 mol) of carbon atoms
▪ By the same token, 1.008 g of hydrogen (1 mol) has the same number of atoms
as 12.01 g of carbon (1 mol)

ASSESS YOUR PROGRESS
1. Define matter. How are the mass and the weight of an object different?
2. Differentiate between element and atom. What four elements are found in
the greatest abundance in the human body?
3. For each subatomic particle of an atom, state its charge and location. Which region of
an atom is most responsible for the mass of the atom? Its volume?
4. Which subatomic particle determines the atomic number? What determines the mass
number?
5. What is an isotope? How are isotopes denoted?
6. What is Avogadro’s number? How is it related to a mole and molar mass?

ELECTRONS AND CHEMICAL BONDING
▪ The chemical behavior of an atom is determined largely by its outermost electrons
▪ Chemical bonding occurs when outermost electrons are transferred or shared between
atoms
▪ Two major types of chemical bonding:
▪ Ionic bonding (ionic bond)
▪ Covalent bonding (covalent bond)

▪ Ionic bonding
▪ An atom is electrically neutral because it has an equal number of protons and
electrons
▪ If an atom loses or gains electrons, the numbers of protons and electrons are no
longer, and a charged particle called an ion is formed
▪ After an atom loses an electron, it has one more proton than it has electrons and is
positively charged; thus, it becomes a cation (positively charged ion)
▪ Sodium (Na) – sodium ion (Na+)
▪ Potassium (K) – potassium ion (K+)
▪ Calcium (Ca) – calcium ion (Ca2+)

▪ Ionic bonding
▪ After an atom gains an electron, it has one more electron than it has protons and is
negatively charged; thus, it becomes an anion (negatively charged ion)
▪ Chlorine (Cl) – chloride ion (Cl-)
▪ Fluorine (F) – fluoride ion (F-)
▪ Because oppositely charged ions are attracted to each other, cations tend to remain
close to anions
▪ An ionic bond occurs when electrons are transferred between atoms, creating
oppositely charged ions, thus forming an ionic compound
▪ Na+ and Cl- are held together by ionic bonding to form an array of ions called
sodium chloride (NaCl), or table salt
▪ Ca2+ and Cl- forms CaCl2

Question 2-2
If an iron (Fe) atom loses three electrons, what is the charge of the
resulting ion?

▪ Covalent bonding
▪ A covalent bond forms when atoms share one or more pairs of electrons
▪ The resulting combination of atoms is called a molecule
▪ An example is the covalent bond between two hydrogen atoms to form a hydrogen
molecule (H2)
▪ Sharing of one pair of electrons by two atoms results in a single covalent bond
▪ A single line between the symbols of the atoms involved (for example H-H)
represents a single covalent bond
▪ A double covalent bond results when two atoms share two pair of electrons
▪ When a carbon atom combines with two oxygen atoms to form carbon dioxide (CO2),
two double covalent bonds are formed

▪ Double covalent bonds are indicated by a double line between atoms (O=C=O)
▪ Electrons can be shared unequally in covalent bonds
▪ When there is unequal, asymmetrical sharing of electrons, the bond is called a
polar covalent bond because the unequal sharing of electrons results in one end
(pole) of the molecule having a partial electrical charge opposite to that of the other
end (e.g., water)
▪ In water, hydrogen atoms do not share equally the electrons with the oxygen
atom, and the electrons tend to spend more time around the oxygen atom than
around the hydrogen atoms
▪ Molecules with asymmetrical distribution of electrical charges are called polar
molecules

▪ When there is equal sharing of electrons between atoms of a molecule, a nonpolar
covalent bond is formed
▪ A good example is methane (CH4)
▪ Molecules with a symmetrical distribution of electrical charges are called nonpolar
molecules

MOLECULES AND COMPOUNDS
▪ A molecule is formed when two or more atoms chemically combine to form a structure
that behaves as an independent unit
▪ Sometimes, the atoms that combine are of the same time, such as the hydrogen gas
molecule (H2)
▪ More typically, a molecule consists of two or more different types of atoms, such as
two hydrogen atoms and one oxygen atom combining to form water
▪ Thus, a glass of water consists of a collection of individual water molecules positioned
next to one another

▪ A compound is a substance resulting from the chemical combination of two or more
different types of atoms
▪ Water is an example of a substance that is a compound and a molecule
▪ Not all molecules are compounds (e.g. H2 molecule)
▪ Some compounds are molecules and some are not
▪ Covalent compounds, in which different types of atoms are held together by
covalent bonds, are molecules because the sharing of electrons results in distinct
units
▪ Ionic compounds, in which ions are held together by the force of attraction
between opposite charges, are not molecules because they do not consist of
distinct units

▪ The molecular mass of a molecule or compound can be determined by adding up the
atomic masses of its atoms (or ions)
▪ The term molecular mass is used for convenience for ionic compounds, even though
they are not molecules
▪ For example, the atomic mass of sodium is 22.99 and that of chloride is 35.45
▪ The molecular mass of NaCl is therefore 58.44 (22.99 + 35.45)

INTERMOLECULAR FORCES
▪ Intermolecular forces are attractive forces between molecules
▪ Intramolecular forces hold atoms together in a molecule (chemical bonds)
▪ Intramolecular forces stabilize individual molecules and are much stronger than
intermolecular forces
▪ Van der Waals forces
▪ Dipole-dipole forces or interactions – between polar molecules
▪ Dipole-induced dipole forces or interactions – between a polar and nonpolar
molecule
▪ Dispersion forces – attractive forces that arise as a result of temporary dipoles
induced in atoms or molecules

INTERMOLECULAR FORCES
▪ Ion-dipole forces or interactions – between an ion and a polar molecule
▪ Hydrogen bond
▪ Not a true bond
▪ It is a special type of dipole-dipole interaction between the hydrogen atom in a polar
bond, such as N-H, O-H, or F-H, and an electronegative O, N, or F atom
▪ For example, the positively charged hydrogen of one water molecule is weakly
attracted to a negatively charged oxygen of another molecule
▪ Water molecules are held together by hydrogen bonds

SOLUBILITY AND DISSOCIATION
▪ Solubility is the ability of one substance to dissolve in another – for example, sugar
dissolving in water
▪ Charged substances, such as sodium chloride, and polar substances, such as glucose,
readily dissolve in water
▪ Nonpolar substances, such as oils, do not dissolve in water
▪ Substances dissolve in water when they become surrounded by water molecules
▪ If the positive and negative ends of the water molecules are more attracted
to the charged ends of other molecules than to each other, the hydrogen bonds
between the ends of the water molecules break, and water molecules surround the
other molecules, which become dissolved in water

▪ When molecules (covalent compounds) dissolve in water, they usually remain intact,
even though they are surrounded by water molecules
▪ Thus, in a glucose solution, glucose molecules are surrounded by water molecules
▪ When ionic compounds dissolve in water, their ions dissociate, or separate, from each
other because:
▪ Positively charged ions are attracted to the negative ends of the H2O
▪ Negatively charged ions are attracted to the positive ends of the H2O
▪ A good example is sodium chloride dissociating in water
▪ The dissociated ions (cations and anions) are sometimes called electrolytes because
they have the capacity to conduct an electrical current

▪ An electrocardiogram (ECG) is a recording of electric currents produced by the heart
▪ These currents can be detected by electrodes on the surface of the body because the
ions in the body fluids conduct electric currents
▪ Molecules that do not dissociate in solution and do not conduct electricity are called
nonelectrolytes
▪ Pure water is a nonelectrolyte
▪ Maintaining the proper balance of electrolytes is important for keeping the body
hydrated, controlling blood pH, and ensuring the proper function of muscles and nerves

7. Describe how an ionic bond is formed. What are a cation and an anion?
8. What occurs in the formation of a covalent bond? What is the difference between
polar and nonpolar covalent bonds?
9. Distinguish between a molecule and a compound. Give an example of each.
Are all molecules compounds? Are all compounds molecules?
10. What are intermolecular forces, and how do they create a hydrogen bond?
11. What is meant by the statement “table sugar is soluble in water?”
12. Describe what occurs during the dissociation of NaCl in water.
13. What occurs when glucose (C6H12O6 ) dissolves in water?

14. Explain the difference between electrolytes and nonelectrolytes. Classify each of the
following water solutions as an electrolyte or a nonelectrolyte:
▪ Potassium iodide (KCl)
▪ sucrose (C12H22O11)
▪ Magnesium bromide (MgBr2)
▪ Lactose (C12H22O11)

CHEMICAL REACTIONS
SECTION 2-2

LEARNING OUTCOMES
▪ Summarize the characteristics of synthesis, decomposition, reversible reactions, and
oxidation-reduction reactions
▪ Illustrate what occurs in dehydration and hydrolysis reactions
▪ Explain how reversible reactions produce chemical equilibrium
▪ Contrast potential and kinetic energy
▪ Distinguish between chemical reactions that release energy and those that take in
energy
▪ Describe the factors that can affect the rate of chemical reactions
Section 2-2. Chemical Reactions

DEFINITION
▪ In a chemical reaction, atoms, ions, molecules, or compounds interact either to form or
break chemical bonds
▪ The substances that enter into a chemical reaction are called reactants
▪ The substances that result from the chemical reaction are called products
▪ All chemical reactions in a living organism are called biochemical reactions

SYNTHESIS REACTIONS
▪ When two or more reactants combine to form a larger, more complex product, the
process is called a synthesis reaction
▪ An example is synthesis of adenosine triphosphate (ATP) in the human body
▪ All synthesis reactions that occur in the body are collectively referred to as anabolism
▪ Growth, maintenance and repair of the body could not take place without anabolic
reactions
A + B AB
A-P-P + Pi A-P-P-P
(ADP) (ATP)

DECOMPOSITION REACTIONS
▪ In a decomposition reaction, reactants are broken down into smaller, less complex
products
▪ Decomposition reaction is the reverse of a synthesis reaction and can be represented
this way:
▪ Example is the breakdown of ATP to ADP and phosphate
▪ Decomposition reactions that occur in the body are collectively called catabolism
▪ The sum of anabolic and catabolic reactions in the body is called metabolism
AB A + B
A-P-P-P A-P-P + Pi
(ATP) (ADP)

REVERSIBLE REACTIONS
▪ A reversible reaction is a chemical reaction that can proceed from reactants to products
and from products to reactants
▪ When the rate of product formation is equal to the rate of reactant formation, the
reaction is said to be at equilibrium
▪ At equilibrium, the amount of the reactants relative to the amount of products remain
constant
▪ An important reversible reaction in the human body occurs when carbon dioxide (CO2)
and water (H2O) form hydrogen ions (H+) and bicarbonate ions (HCO3
ˉ)
CO2 + H2O H+ + HCO3
ˉ

REVERSIBLE REACTIONS
▪ Maintaining a constant level of H+ in body fluids is necessary for the body to function
properly
▪ This level can be maintained, in part, by controlling blood CO2 levels
▪ For example, slowing the respiration rate causes blood CO2 levels to increase, which
causes an increase in H+ concentration [H+] in the blood
Question 2-3
If the respiration rate increases, CO2 is removed from the blood.
What effect does it have on blood H+ levels?

OXIDATION-REDUCTION REACTIONS
▪ Chemical reactions that result from the exchange of electrons between the reactants are
called oxidation-reduction reactions
▪ When sodium and chlorine react to form sodium chloride, the sodium atom loses an
electron and the chlorine atom gains an electron
▪ Loss of an electron by an atom is called oxidation
▪ Gain of an electron is called reduction
▪ Transfer of the electron can be complete, resulting in an ionic bond, or it can be partial,
resulting in a covalent bond
▪ Because one atom partially or completely loses an electron and another atom gains that
electron, these reactions are called oxidation-reduction reactions
▪ Synthesis and decomposition reactions can be oxidation-reduction reactions

Question 2-4
When hydrogen gas combines with oxygen gas to form water, is
the hydrogen reduced or oxidized? Explain.

15. Using the terms reactant and product, describe what occurs in a chemical reaction
16. Contrast synthesis and decomposition reactions, and explain how catabolism and
anabolism relate to these two types of reactions
17. Describe the role of water in dehydration and hydrolysis reactions
18. What is a reversible reaction? How does this type of reaction lead to chemical
equilibrium?
19. What are oxidation-reduction reactions?

ENERGY AND CHEMICAL REACTIONS
▪ Energy is defined as the capacity to do work – that is, to move matter
▪ Energy can be subdivided into potential energy and kinetic energy:
▪ Potential energy is the stored energy that can do work but is not doing so
▪ Kinetic energy is energy caused by the movement of an object and is the form of
energy that actually does work
▪ When potential energy is released, it becomes kinetic energy, thus doing work
▪ Chemical energy of a substance is a form of potential energy stored in chemical bonds
▪ As similarly charged particles move closer together, their potential energy increases,
and as they move farther apart, their potential energy decreases
▪ Chemical bonding is a form of potential energy because of the charges and positions
of the subatomic particles bound together

▪ Chemical reactions are important because of the products they form and the energy
changes that result as the relative positions of subatomic particles change
▪ If the products of a chemical reaction contain less potential energy than the reactants,
energy is released
▪ Food molecules contain more potential energy than waste products
▪ The difference in potential energy between food and waste products is used by the
human body to drive activities such as growth, repair, movement and heat production
▪ Example of a reaction that releases energy is breakdown of ATP to ADP and phosphate
▪ In ATP, the phosphate group is attached to ADP by a covalent bond, which has
potential energy
▪ After ATP breakdown, some energy is released as heat, and some is available for
cellular activities

▪ According to the law of conservation of energy, the total energy of the universe is
constant; therefore, energy is neither created nor destroyed
▪ However, one type of energy can be changed into another, such as potential energy
being converted to kinetic energy
▪ Since conversion between energy states is not 100% efficient, heat energy is released
▪ Mechanical energy is energy resulting from the position or movement of objects
▪ Many of the activities of the body, such as movement, breathing or blood circulation,
involve mechanical energy

▪ If the products of a chemical reaction contain more energy than the reactants, energy
must be added from another source
▪ The energy released during the breakdown of food molecules is the source of energy for
this kind of reaction in the body
▪ Energy from food is used to synthesize molecules such as ATP, fats, and proteins
ADP + Pi + Energy (from food breakdown) ATP
Question 2-5
Why does body temperature increase during exercise?

RATE OF CHEMICAL REACTIONS
▪ The rate at which a chemical reaction proceeds is influenced by several factors, including
how easily the substances (reactants) react with one another, their concentration, the
temperature, and the presence of a catalyst
▪ Reactants
▪ Reactants differ from one another in their ability to undergo chemical reactions
▪ For example, iron corrodes much more rapidly than does stainless steel
▪ Concentration
▪ Within limits, the greater the concentration of the reactants, the greater the rate at
which a chemical reaction will occur because, as the concentration increases, the
reacting molecules are more likely to come in contact with other molecules
▪ For example, increasing oxygen concentration during exercises increases the rate of
ATP production in muscle cells which are needed for movement

RATE OF CHEMICAL REACTIONS
▪ Temperature
▪ Because molecular motion changes as environmental temperature changes, the rate
of chemical reactions is partially dependent on temperature
▪ Increasing temperature increase movement of molecules making them more likely to
come in contact with other molecules
▪ Catalyst
▪ At normal body temperatures, most chemical reactions would take place too slowly to
sustain life if not for substances called catalysts
▪ A catalyst increases the rate of a chemical reaction, without itself being permanently
changed or depleted
▪ An enzyme is a protein molecule that acts as a catalyst
▪ Many of the biochemical reactions require enzymes

20. Define energy. How are potential and kinetic energies different from each other?
21. Summarize the characteristics of mechanical, chemical, and heat energies
22. Use ATP and ADP to illustrate the release or input of energy in chemical reactions
23. Define activation energy, catalyst, and enzymes; then explain how they affect the rate
of chemical reactions
24. What effect does increasing temperature or increasing concentration of reactants
have on the rate of a chemical reaction?

LEARNING OUTCOMES
▪ Describe the pH scale and its relationship to acidic and basic solutions
▪ Explain the importance of buffers in organisms
Section 2-3. Acids and Bases

DEFINITION
▪ The body has many molecules and compounds called acids and bases
▪ An acid is a proton donor
▪ Because a hydrogen atom without its electron is a proton, any substance that releases
hydrogen ions (H+) in water is an acid
▪ For example, hydrochloric acid (HCl) in the stomach forms H+ and Clˉ (chloride ions)
▪ A base is a proton acceptor
▪ For example, sodium hydroxide (NaOH) forms sodium ions (Na+) and hydroxide or
hydroxyl ions (OHˉ)
▪ It is a base because the OHˉ is a proton acceptor that binds with H+ to form water
HCl H+ + Clˉ

Sodium hydroxide (NaOH) is a base because the hydroxide
ion (OHˉ) is a proton acceptor that binds with H+ to form
water

THE pH SCALE
▪ The pH (power of hydrogen) scale indicates the hydrogen ion concentration [H+] of a
solution
▪ The scale ranges from 0 to 14
▪ A neutral solution has an equal number of H+ and OH- and thus a pH of 7.0.
▪ An acidic solution has a greater concentration of H+ than of OH- and thus a pH less 7.0.
▪ A basic or alkaline solution has fewer H+ than OH- and thus a pH greater than 7.0.
▪ [H+] and pH are inversely related, meaning the lower the [H+], the higher the pH; the
higher the [H+], the lower the pH

THE pH SCALE
▪ As the pH value becomes smaller, the solution
becomes more acidic
▪ As the pH value becomes larger, the solution
becomes more basic or alkaline
▪ A change of one unit on the pH scale represents a
10-fold change in the [H+]
▪ A solution with a pH of 6.0 has 10 times more H+
than a solution with a pH of 7.0
▪ Small changes in pH represent large changes in
[H+]

THE pH SCALE
▪ The normal pH range for human blood is 7.35 to 7.45.
▪ If blood pH drops below 7.35, a condition called acidosis results
▪ Nervous system is depressed
▪ Individual becomes disoriented and possibly comatose
▪ If blood pH rises above 7.45, alkalosis results
▪ Nervous system becomes over excitable
▪ Individual can be extremely nervous or have convulsions
▪ Both acidosis and alkalosis can result in death

SALTS
▪ A salt is a compound consisting of a positive ion (cation) other than H+ and a negative
ion (anion) other than OH-
▪ Salts are formed by the reaction of an acid and a base
▪ For example, hydrochloric acid (HCl) combines with sodium hydroxide (NaOH) to form
the salt, sodium chloride (NaCl)
HCl + NaOH NaCl + H2O
(acid) (base) (salt) (water)

BUFFERS
▪ The chemical behavior of many molecules changes as the pH of the solution in which
they are dissolved changes
▪ The survival of an organism depends on its ability to maintain homeostasis by keeping
body fluid pH within a narrow range
▪ One way normal body fluid pH is maintained is through the use of buffers
▪ A buffer is a chemical solution that resists changes in pH when either an acid or a base is
added to a solution containing the buffer
▪ Buffer solution is a solution of (1) a weak acid or a weak base, and (2) its conjugate
base or conjugate acid in salt form

BUFFERS
▪ When an acid is added to a buffered solution, the buffer binds to the H+, preventing
these ions from causing a decrease in the pH of the solution
▪ When a base is added to a buffered solution, the buffer binds to the OH-, preventing
these ions from causing an increase in the pH of the solution
Question 2-6
If a base is added to a solution, will the pH of the solution increase
or decrease?
If the solution is buffered, what response from the buffer prevents
the change in pH?

BUFFERS
▪ Important buffers in living systems are composed of bicarbonate, phosphates, amino
acids, and proteins
▪ Buffers prevent large changes in pH values by acting as conjugate acid-base pairs
▪ A conjugate base is what remains of an acid after the H+ (proton) is lost
▪ A conjugate acid is formed when a H+ is transferred to the conjugate base
▪ Two substances related in this way are a conjugate acid-base pair
▪ A major buffer in our body fluids is the bicarbonate system
▪ A bicarbonate ion (HCO3ˉ) is formed by the dissociation of carbonic acid (H2CO3):

BUFFERS
▪ The greater the buffer concentration, the more effectively it can resist a change in pH
▪ However, buffers cannot entirely prevent some change in the pH of a solution
▪ For example, when an acid is added to a buffered solution, the pH decreases, but not to
the extent it would have without the buffer

Question 2-7
Dihydrogen phosphate ion (H2PO4ˉ) and monohydrogen phosphate
ion (HPO4
2−) form the phosphate buffer system. Identify the
conjugate acid and the conjugate base in the phosphate buffer
system:
Explain how they function as a buffer when either H+ or OHˉ are
added to the solution.

25. Define acid and base, and describe the pH scale
26. What is the difference between a strong acid or base and a weak acid or base?
27. The blood pH of a patient is 7.30. What condition does this patient have, and what are
the symptoms?
28. How are salts related to acids and bases?
29. What is a buffer, and why are buffers important in the body?
30. What is a conjugate acid-base pair?

INORGANIC MOLECULES
SECTION 2-4

LEARNING OUTCOMES
▪ Distinguish between inorganic and organic molecules
▪ Describe how the properties of O2, CO2, and water contribute to their physiological
functions
Section 2-4. Inorganic Molecules

DEFINITION
▪ Inorganic chemistry deals with substances that do not contain carbon
▪ Organic chemistry is the study of carbon-containing substances
▪ A few exceptions are carbon dioxide (CO2), carbon monoxide (CO), and bicarbonate
ion (HCO3ˉ) which are classified as inorganic molecules, even though they contain
carbon
▪ Inorganic substances play many vital roles in human anatomy and physiology
▪ Water (H2O)
▪ Oxygen (O2) – for breathing
▪ Calcium phosphate (Ca3(PO4)2) – makes up our bones
▪ Metals that are required for protein functions, such as Fe2+ in hemoglobin and Zn2+ in
alcohol dehydrogenase

WATER
▪ Water has remarkable properties due to its polar nature
▪ A molecule of water is formed when an atom of oxygen forms polar covalent bonds with
two atoms of hydrogen
▪ This gives a partial positive charge to the hydrogen atoms and a partial negative charge
to the oxygen atom
▪ Because of water’s polarity, hydrogen bonds form between the positively charged
hydrogen atoms of one water molecule and the negatively charged oxygen atoms of
another water molecule
▪ These hydrogen bonds organize the water molecules into a lattice, which holds the
water molecules together and are responsible for many unique properties of water

WATER
▪ Cohesion
▪ The attraction of water to another water molecule is called cohesion
▪ An example of cohesion is the surface tension exhibited when water bulges over the
top of a full glass without spilling over
▪ Adhesion
▪ The same attractive force of hydrogen bonds with water will also attract other
molecules
▪ This process is called adhesion
▪ The combination of cohesion and adhesion helps hold cells together and move fluids
through the body

WATER
▪ Water accounts for approximately 50% of the weight of a young adult female and 60%
of a young adult male
▪ Females have a lower percentage of water than males because they typically have more
body fat, which is relatively free of water
▪ Plasma, the liquid portion of blood, is 92% water
▪ Water has physical and chemical properties well suited for its many functions in living
organisms

FUNCTIONS OF WATER
Stabilizing Body Temperature
▪ Water can absorb large amounts of heat and remain at a fairly stable temperature
▪ Therefore, it tends to resist large temperature fluctuations
▪ Because of this property, blood, which is mostly water, can transfer heat from deep in
the body to the surface, where the heat is released
▪ In addition, when water evaporates, it changes from a liquid to a gas
▪ Because heat is required for that process, the evaporation of water from the surface of
the body rids the body of excess heat

FUNCTIONS OF WATER
Protection
▪ Water is an effective lubricant that provides protection against damage resulting
from friction
▪ For example, tears protect the surface of the eye from rubbing of the eyelids
▪ Water also forms a fluid cushion around organs, helping protect them from trauma
▪ The cerebrospinal fluid (CSF) that surrounds the brain is an example

FUNCTIONS OF WATER
Facilitating Chemical Reactions
▪ Many of the chemical reactions necessary for life do not take place unless the reacting
molecules are dissolved in water
▪ For example, sodium chloride (NaCl) must dissociate in water into Na+ and Clˉ, which
can then react with other ions
▪ Water also directly participates in many chemical reactions
▪ Dehydration reaction is a synthesis reaction that produces water
▪ Hydrolysis reaction is a decomposition reaction that requires water

FUNCTIONS OF WATER
Mixing Medium
▪ Mixture is a combination of two or more substances physically blended together,
but not chemically combined
▪ A solution is any mixture of liquids, gases, or solids in which the substances are
uniformly distributed with no clear boundary between them
▪ Salt solution consists of salt dissolved in water
▪ Air is a solution containing a variety of gases
▪ Wax is a solid solution composed of several fatty substances
▪ Solute dissolves in the solvent
▪ In a salt solution, water is the solvent and the dissolved salt is the solute
▪ Sweat is a salt solution in which NaCl and other solutes are dissolved in water

FUNCTIONS OF WATER
Mixing Medium
▪ A suspension is a mixture containing materials that separate from each other
unless they are continually, physically blended together
▪ Blood is a suspension – i.e., red blood cells are suspended in a liquid called
plasma
▪ As long as the red blood cells and plasma are mixed together as they pass
through blood vessels, the red blood cells remain suspended in the plasma
▪ However, if the blood is allowed to sit in a container, the red blood cells and
plasma separate from each other

FUNCTIONS OF WATER
Mixing Medium
▪ Colloid is a mixture in which a dispersed (solute-like) substance is distributed
throughout a dispersing (solvent-like) substance
▪ Dispersed particles are larger than a simple molecule but small enough that they
remain dispersed and do not settle out
▪ Proteins, which are large molecules, are common dispersed particles
▪ Proteins and water form colloids
▪ For instance, the plasma portion of blood and the liquid interior of cells are colloids
containing many important proteins

FUNCTIONS OF WATER
Mixing Medium
▪ In living organisms, the complex fluids inside and outside cells consist of solutions,
suspensions, and colloids
▪ Blood is an example of all of these mixtures
▪ A solution containing dissolved nutrients, such as sugar
▪ A suspension holding red blood cells
▪ A colloid containing proteins
▪ Water’s ability to mix with other substances enables it to act as a medium for transport,
moving substances from one part of the body to another
▪ Body fluids, such as plasma, transport nutrients, gases, waste products, and a variety
of molecules involved in regulating body functions

SOLUTION CONCENTRATIONS
▪ Percent of solute by weight per volume of solution (w/v)
▪ A 10% solution of sodium chloride can be made by dissolving 10 g of sodium chloride
into enough water to make 100 mL of solution
▪ Osmolality
▪ Often used by physiologists
▪ Express the number of particles in a solution
▪ A particle can be an atom, an ion, or a molecule
▪ An osmole (Osm) is Avogadro’s number of particles of a substance in 1 kilogram (kg)
of water
▪ Osmolality of a solution reflects the number, not the type, of particles in a solution

SOLUTION CONCENTRATIONS
▪ Osmolality
▪ For example, a 1 Osm glucose solution and a 1 Osm NaCl solution both contain
Avogadro’s number of particles per kg of water
▪ The glucose solution has 1.0 Osm of glucose molecules
▪ On the other hands, the NaCl solution has 0.5 Osm of Na+ and 0.5 Osm of Clˉ
because NaCl dissociates into Na+ and Clˉ in water
▪ Because the concentration of particles in body fluids is so low, physiologists use the
measurement milliosmole (mOsm), 1/1000 of an osmole
▪ Most body fluids have a concentration of about 300 mOsm and contain many
different ions and molecules
▪ The concentration of body fluids is important because it influences the movement of
water into or out of cells

MOLECULAR OXYGEN (O2)
▪ Molecular oxygen is an inorganic molecule consisting of two oxygen atoms bound
together by a double covalent bond
▪ About 21% of the gas in the atmosphere is O2
▪ Essential for most living organisms
▪ Humans require O2 in the final step of a series of reactions that extract energy from food
molecules

CARBON DIOXIDE (CO2)
▪ Consists of one carbon atom bound to two oxygen atoms
▪ Each oxygen atom is bound to the carbon atom by a double covalent bond
▪ CO2 is produced when organic molecules, such as glucose, are metabolized within the
cells of the body
▪ Much of the energy stored in the covalent bonds of glucose is transferred to other
organic molecules when glucose is broken down and CO2 is released
▪ Once CO2 is produced, it is eliminated from the cell as a metabolic by-product,
transferred to the lungs by the blood, and exhaled during respiration
▪ If CO2 were allowed to accumulate within cells, it will become toxic

31. What is the difference between inorganic and organic chemistry?
32. What two properties of water are the result of hydrogen bonding, and how are these
two properties different?
33. List and briefly describe the four functions that water performs in living organisms
34. Using the terms solute and solvent, summarize the properties of solutions,
suspensions, and colloids
35. How is the osmolality of a solution determined? What is a milliosmole?
36. What are the functions of oxygen and carbon dioxide in living systems?

LEARNING OUTCOMES
▪ Describe the structural organization and major functions of carbohydrates, lipids,
proteins, and nucleic acids
▪ Explain how enzymes work
▪ Describe the roles of nucleotides in the structures and functions of DNA, RNA, and ATP
Section 2-5. Organic Molecules

OVERVIEW
▪ Carbon’s ability to form covalent bonds with other atoms makes possible the formation
of the large, diverse, complicated molecules necessary for life
▪ Carbon atoms bound together by covalent bonds constitute the “framework” of many
large molecules
▪ Two mechanisms that allow the formation of a wide variety of molecules are:
▪ Variation in the length of the carbon chains; and
▪ Combination of the atoms involved
▪ For example, some protein molecules have thousands of carbon atoms bound by
covalent bonds to one another or to other atoms, such as nitrogen, sulfur, hydrogen,
and oxygen

OVERVIEW
▪ Four major groups of organic molecules essential to living organisms:
▪ Carbohydrates
▪ Lipids
▪ Proteins
▪ Nucleic acids
▪ In addition, a high energy form of a nucleic acid building block, called ATP (adenosine
triphosphate), is an important organic molecule in cellular processes
▪ Each of these groups and ATP have specific structural and functional characteristics

CARBOHYDRATES
▪ Composed primarily of carbon, hydrogen, and oxygen atoms, and range in size from
small to very large
▪ In most carbohydrates, there are approximately two hydrogen atoms and one oxygen
atom for each carbon atom
▪ Note that this two-to-one ratio is the same as in water (H2O)
▪ The molecules are called carbohydrates because carbon toms are combined with the
same atoms that form water (hydrated)
▪ Large number of oxygen atoms in carbohydrates makes them relatively polar molecules
▪ Consequently, they are soluble in polar solvents, such as water

CARBOHYDRATES
▪ Carbohydrates are important parts of other organic molecules, and they can be broken
down to provide the energy necessary for life
▪ Undigested carbohydrates also provide bulk in feces, which helps maintain the normal
function and health of the digestive tract

CARBOHYDRATES
Monosaccharides
▪ Simple sugars
▪ Building blocks of large carbohydrates
▪ Commonly contain 3 carbons (trioses), 4 carbons (tetroses), 5 carbons (pentoses), and 6
carbons (hexoses)
▪ Monosaccharides most important to humans include both 5- (pentose) and 6-carbon
(hexose) sugars
▪ Common hexoses, such as glucose, fructose, and galactose, are isomers, which are
molecules that have the same number and types of atoms but differ in their three-
dimensional arrangement

CARBOHYDRATES
Monosaccharides
▪ Glucose, or blood sugar, is the major carbohydrate in the blood and a major nutrient for
most cells of the body
▪ Blood glucose levels are tightly regulated by insulin and other hormones
▪ In people with diabetes, the body is unable to regulate glucose levels properly
▪ Diabetics need to monitor their blood glucose carefully to minimize the deleterious
effects of this disease
▪ Fructose and galactose are also important dietary nutrients
▪ Important pentoses include ribose and deoxyribose, which are components of
ribonucleic acid (RNA) and deoxyribonucleic acid (DNA), respectively

CARBOHYDRATES
Disaccharides
▪ Composed of two simple sugars bound together through a dehydration reaction
▪ For example, glucose and fructose combine to form a disaccharide called sucrose (table
sugar) plus a molecule of water
▪ Several disaccharides are important to humans, including sucrose, lactose, and maltose
▪ Lactose, or milk sugar, is glucose combined with galactose
▪ Maltose, or malt sugar, is two glucose molecules joined together

Figure 2.15 Carbohydrates
(a) Sucrose, a disaccharide, forms by a
dehydration reaction involving glucose
and fructose (monosaccharides). (b)
Glycogen is a polysaccharide formed by
combining many glucose molecules. (c)
The transmission electron micrograph
shows glycogen granules in a liver cell.

CARBOHYDRATES
Polysaccharides
▪ Consist of many monosaccharides bound together to form long chains that are either
straight or branched
▪ Glycogen, or animal starch, is a polysaccharide composed of many glucose molecules
▪ Because glucose can be metabolized rapidly and the resulting energy can be used by
cells, glycogen is an important energy-storage molecule
▪ A substantial amount of the glucose that is metabolized to produce energy for muscle
contraction during exercise is stored in the form of glycogen in the cells of the liver and
skeletal muscles
▪ Starch and cellulose are two important polysaccharides found in plants, and both are
composed of long chains of glucose

CARBOHYDRATES
Polysaccharides
▪ Starch
▪ Storage form of carbohydrates in plants
▪ Plants use starch as an energy-storage molecule
▪ Glycogen
▪ Storage form of carbohydrates in man
▪ Animals use glycogen as an energy-storage molecule
▪ Cellulose is an important structural component of plant cell walls

CARBOHYDRATES
Polysaccharides
▪ When humans ingest plants, the starch can be broken down and used as an energy
source
▪ However, humans do not have the digestive enzymes necessary to break down cellulose
▪ Thus, cellulose is eliminated in the feces, where it provides bulk

37. Why is carbon such a versatile element?
38. What is the building block of carbohydrates? What are isomers?
39. List the 5- and 6-carbon sugars that are important to humans
40. What are disaccharides and polysaccharides, and what type of reaction is used to
make them?
41. Which carbohydrates are used for energy? What is the function of starch and cellulose
in plants? What is the function of glycogen and cellulose in animals?

LIPIDS
▪ Second major group of organic molecules common to living systems
▪ Composed principally of carbon, hydrogen, and oxygen
▪ Some lipids contain small amounts of other elements, such as phosphorus (P) and
nitrogen (N)
▪ Lipids have a lower ratio of oxygen to carbon than do carbohydrates, which makes
them less polar
▪ Consequently, lipids can be dissolved in nonpolar organic solvents, such as alcohol or
acetone, but they are relatively insoluble in water

LIPIDS
▪ Lipids have many important functions in the body:
▪ Provide protection and insulation
▪ Help regulate many physiological processes
▪ Form plasma membranes
▪ Major energy-storage molecules and can be broken down and used as a source of
energy
▪ Several kinds of molecules, such as fats, phospholipids, eicosanoids, steroids, and fat-
soluble vitamins, are classified as lipids

LIPIDS
Fats
▪ Major type of lipid
▪ Like carbohydrates, the fats humans ingest are broken down by hydrolysis reactions in
cells to release energy for use by those cells
▪ Conversely, if fat intake exceeds need, excess chemical energy from any source can be
stored in the body as fat for later use
▪ Fats also provide protection by surrounding and padding organs, and under-the-skin fats
act as an insulator to prevent heat loss

LIPIDS
Triglycerides
▪ Constitute 95% of the fats in the human body
▪ Triglycerides consist of two different types of building blocks:
▪ One glycerol – a 3-carbon molecule with a hydroxyl group attached to each carbon
atom
▪ Three fatty acids – each fatty acid consists of a straight chain of carbon atoms with a
carboxyl group attached at one end
▪ A carboxyl group (-COOH) consists of both an oxygen atom and a hydroxyl group
attached to a carbon atom

LIPIDS
Triglycerides
▪ The carboxyl group is responsible for the acidic nature of the molecule because it
releases hydrogen ions into solution
▪ Glycerides can be described according to the number and kinds of fatty acids that
combine with glycerol through dehydration reactions
▪ Monoglycerides have one fatty acid
▪ Diglycerides have two fatty acids
▪ Triglycerides have three fatty acids bound to glycerol

LIPIDS
Fatty Acids
▪ Consists of a straight chain of carbon atoms with a carboxyl group attached at one end
▪ Differ from one another according to the length and the degree of saturation of their
carbon chains
▪ Most naturally occurring fatty acids contain an even number of carbon atoms, with 14-
to 18-carbon chains the most common
▪ Saturation refers to the number of hydrogen atoms in the carbon chain
▪ Saturated fatty acid
▪ Contains only single covalent bonds between the carbon atoms
▪ Sources include beef, pork, whole milk, cheese, butter, eggs, coconut oil, and palm oil

LIPIDS
Fatty Acids
▪ Unsaturated fatty acid
▪ One or more double covalent bonds between carbon atoms
▪ The double covalent bond introduces a kink (or bend) into the carbon chain, which
tends to keep them liquid at room temperature
▪ Monounsaturated fats, such as olive and peanut oils, have one double covalent bond
between carbon atoms
▪ Polyunsaturated fats, such as safflower, sunflower, corn, and fish oils, have two or
more double covalent bonds between carbon atoms
▪ Unsaturated fats are the best type of fats in the diet because, unlike saturated fats,
they do not contribute to the development of cardiovascular disease

Naturally occurring
unsaturated fatty acids are
nearly all cis-, the molecules
being “bent” 120 degrees at
the double bond
Increase in the number of cis-
double bonds leads to various
spatial configurations of the
molecule, e.g. arachidonic acid
may have “kinks” or U-shape

Melting points increase with
chain length and decrease
according to unsaturation
Membrane lipids are more
unsaturated than storage lipids
Tissue lipids subjected to
cooling (hibernation) or in
extremities are more
unsaturated

LIPIDS
Trans Fats
▪ Unsaturated fats that have been chemically altered by the addition of H atoms
▪ The process makes the fats more saturated and hence more solid and stable (longer
shelf-life)
▪ However, the double covalent bonds that do not become saturated are changed from
the usual cis configuration (H on the same side of the double bond) to a trans
configuration (H on different sides)
▪ This change in structure makes the consumption of trans fats an even greater factor
than saturated fats in the risk for cardiovascular disease

cis- or trans- geometric
isomerism occurs in
unsaturated fatty acids
cis- if the acyl chains are on the
same side of the double bond
trans- if the acyl chains are on
the opposite side of the double
bond

LIPIDS
Phospholipids
▪ Similar to triglycerides, except that one of the fatty acids bound to the glycerol is
replaced by a molecule containing phosphate and, usually, nitrogen
▪ Phospholipid is polar at the end of the molecule to which the phosphate is bound and
nonpolar at the other end
▪ Polar end of the molecule is attracted to water and is said to be hydrophilic (water-
loving)
▪ Nonpolar end is repelled by water and is said to be hydrophobic (water-fearing)
▪ Phospholipids are important structural components of the membranes of cells

H2C
HC
H2C
O
O
O
C
O
CH2(CH2)12CH2CH3
C
O
CH2(CH2)14CH2CH3
C
O H
(CH2)7C=C-(CH2)7CH3
H
H2C
HC
H2C
O
O
O
C
O
CH2(CH2)12CH2CH3
C
O H
(CH2)7C=C-(CH2)7CH3
H
P
O
O
O-
X
Triacylglycerol Phospholipid

LIPIDS
Eicosanoids
▪ Group of important chemicals derived from fatty acids
▪ Include prostaglandins, thromboxanes, and leukotrienes
▪ Eicosanoids are made in most cells and are important regulatory molecules
▪ Among their numerous effects is their role in the response of tissues to injuries
▪ Prostaglandins have been implicated in regulating the secretion of certain hormones,
blood clotting, some reproductive functions, and many other processes
▪ Many of the therapeutic effects of aspirin and other anti-inflammatory drugs result
from their ability to inhibit prostaglandin synthesis

LIPIDS
Steroids
▪ All steroid molecules are composed of carbon atoms bound together into four ring-like
structures (cyclopentanoperhydrophenanthrene ring)
▪ Important steroid molecules include cholesterol, bile salts, estrogen, progesterone, and
testosterone
▪ Cholesterol is an especially important steroid because other steroid molecules are
synthesized from it
▪ Bile salts, which increase fat absorption in the intestines, are derived from cholesterol
▪ Reproductive hormones – estrogen, progesterone, and testosterone

LIPIDS
Steroids
▪ Cholesterol is an important component of plasma membranes
▪ Although high levels of cholesterol in the blood increase the risk for cardiovascular
disease, a certain amount of cholesterol is vital for normal function

LIPIDS
Fat-Soluble Vitamins
▪ Their structures are not closely related to one another, but they are nonpolar molecules
▪ Essential for many normal body functions

42. State six roles of lipids in the body, and give an example of each
43. What is the most common fat in the body, and what are its basic building blocks?
44. What is the difference between a saturated fat and an unsaturated fat? What is a
trans fat?
45. Describe the structure of a phospholipid. Which end of the molecule is hydrophilic?
Explain why
46. What are three examples of eicosanoids and their general functions?
47. Why is cholesterol an important steroid?

PROTEINS
▪ All proteins contain carbon, hydrogen, oxygen, and nitrogen bound together by covalent
bonds
▪ Most proteins contain some sulfur
▪ In addition, some proteins contain small amounts of phosphorus, iron, and iodine
▪ Molecular mass of proteins can be very large
▪ For the purpose of comparison, the molecular mass of water is approximately 18,
sodium chloride 58, and glucose 180, but the molecular mass of proteins ranges from
approximately 1000 to several million
▪ Proteins regulate body processes, act as a transportation system, provide protection,
help muscles contract, and provide structure and energy

PROTEINS
Protein Structure
▪ Basic building blocks for proteins are the 20 amino acid molecules
▪ Each amino acid has an amine group (-NH2), a carboxyl group (-COOH), a hydrogen
atom, and a side chain designated by the symbol R attached to the same carbon atom
(α carbon)
▪ The side chain can be a variety of chemical structures, and the differences in the side
chains make the amino acids different from one another

PROTEINS
Protein Structure
▪ Covalent bonds formed between amino acid
molecules during protein synthesis are called
peptide bonds
▪ Dipeptide is two amino acids bound
together by a peptide bond
▪ Tripeptide is three amino acids bound
together by peptide bonds
▪ Polypeptide is many amino acids bound
together by peptide bonds
▪ Proteins are polypeptides composed of
hundreds of amino acids (≥50)

PROTEINS
Protein Structure
▪ Primary structure
▪ Determined by the sequence of the amino acids bound by peptide bonds
▪ The potential number of different protein molecules is enormous
▪ 20 different amino acids exist
▪ Each amino acid can be located at any position along a polypeptide chain
▪ Characteristics of the amino acids in a protein ultimately determine the 3D shape of
the protein, and the shape of the protein determines its function
▪ A change in one or a few amino acids in the primary structure can alter protein
function, usually making the protein less functional or even non-functional

PROTEINS
Protein Structure
▪ Secondary structure
▪ Results from the folding or bending of the polypeptide chain caused by the hydrogen
bonds between amino acids
▪ Two common shapes that result are β pleated (folded) sheets and α helices (sing.
helix, coil)
▪ If the hydrogen bonds that maintain the shape of the protein are broken, the protein
becomes non-functional
▪ This change in shape is called denaturation, and it can be caused by abnormally high
temperatures or changes in the pH of body fluids
▪ An everyday example of denaturation is the change in the proteins of egg whites
when they are cooked

PROTEINS
Protein Structure
▪ Tertiary structure
▪ Results from large-scale folding of the protein driven by interactions within the
protein and with the immediate environment
▪ These interactions allow the pleated sheets and helices of the secondary structure to
be arranged and organized relative to each other
▪ Tertiary structure determines the shape of a domain, which is a folded sequence of
100–200 amino acids within a protein
▪ The functions of proteins occur at one or more domains
▪ Therefore, changes in the primary or secondary structure that affect the shape of the
domain can change protein function

PROTEINS
Protein Structure
▪ Quaternary structure
▪ If two or more proteins associate to form a functional unit, the individual proteins are
called subunits
▪ The quaternary structure results from the spatial relationships between the
individual subunits

PROTEINS
Enzymes
▪ A protein catalyst that increases the rate at which a
chemical reaction proceeds without the enzyme
being permanently changed
▪ 3D shape of enzymes is critical for their normal
function because it determines the structure of the
enzyme’s active site
▪ At the active site, reactants are brought into close
proximity and the reaction occurs
▪ After the reactants combine, products are released
from the active site, and the enzyme is capable of
catalyzing additional reactions

PROTEINS
Enzymes
▪ Activation energy required for a chemical reaction to
occur is lowered by enzymes
▪ Why? Because they orient the reactants toward each
other in such a way that a chemical reaction is more
likely to occur
▪ Slight changes in the structure of an enzyme can
destroy the active site’s ability to function
▪ Enzymes are very sensitive to changes in
temperature or pH, which can break the hydrogen
bonds within them

PROTEINS
Enzymes
▪ As a result, the relationship between amino acids
changes, thereby producing a change in shape that
prevents the enzyme from functioning normally
▪ Enzymes control the rate at which most chemical
reactions proceed in living systems
▪ Consequently, they control essentially all cellular
activities
▪ At the same time, the activity of enzymes
themselves is regulated by several mechanisms
within the cells

48. What are the building blocks of proteins? What type of bond chemically connects
these building blocks? What is the importance of the R group?
49. What determines the primary, secondary, tertiary, and quaternary structures of a
protein?
50. What is denaturation? Name two factors that can cause it
51. Compare the lock-and-key and the induced fit models of enzyme activity. What
determines the active site of an enzyme? State the difference between a cofactor and
a coenzyme

NUCLEIC ACIDS
▪ Nucleic acids are large molecules composed of carbon, hydrogen, oxygen, nitrogen, and
phosphorus
▪ Deoxyribonucleic acid (DNA)
▪ Genetic material of cells
▪ Copies of DNA are transferred from one generation of cells to the next generation
▪ Contains the information that determines the structure of proteins
▪ Ribonucleic acid (RNA)
▪ Structurally related to DNA
▪ Three types of RNA also play important roles in protein synthesis – messenger RNA
(mRNA), transfer RNA (tRNA), and ribosomal RNA (rRNA)

NUCLEIC ACIDS
Nucleotides
▪ Building blocks of nucleic acids
▪ Each is composed of a 5-carbon
monosaccharide to which a nitrogenous
base and phosphate group are attached
▪ Deoxyribose for DNA
▪ Ribose for RNA
▪ Nitrogenous bases consist of carbon and
nitrogen atoms organized into rings
▪ They are called bases because N atoms
tend to take up H+ from solution

NUCLEIC ACIDS
Nucleotides
▪ Single-ringed bases are called
pyrimidines
▪ Cytosine
▪ Thymine – only found in DNA
▪ Uracil – only found in RNA
▪ Double-ringed bases are called purines
▪ Adenine
▪ Guanine

ADENOSINE TRIPHOSPHATE
▪ Adenosine triphosphate (ATP) is an especially important organic molecule in all living
organisms
▪ Consists of adenosine (the sugar ribose with the nitrogenous base adenine) and three
phosphate groups

ADENOSINE TRIPHOSPHATE
▪ The potential energy stored in the covalent bond between the second and third
phosphate groups of ATP is important to living organisms because it provides the energy
used in nearly all of the chemical reactions within cells
▪ Removal of the third phosphate generates adenosine diphosphate (ADP), which has only
two phosphate groups
▪ ATP is often called the energy currency of cells because it is capable of both storing and
providing energy
▪ The concentration of ATP is maintained within a narrow range of values, and essentially
all energy-requiring chemical reactions stop when the ATP levels become inadequate

52. Name two types of nucleic acids, and state their functions
53. What are the basic building blocks of nucleic acids? What kinds of sugars and bases
are found in DNA? In RNA?
54. DNA is like a twisted ladder. What forms the sides of the ladder? The rungs?
55. Name the complementary base pairs in DNA and RNA
56. What is meant by the statement “DNA strands are antiparallel”?
57. Describe the structure of ATP. Where does the energy to synthesize ATP come from?
What is the energy stored in ATP used for?

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