The chemical basis of life

1. The Chemical Basis of Life
Allen Titus G. Paz, RN

2. Chemistry
• Chemicals make up the body’s structures, and the
interactions of chemicals with one another are responsible for
the body’s functions.
• The processes of nerve impulse generation, digestion, muscle
contraction, and metabolism can all be described in chemical
terms. Many abnormal conditions and their treatments can
also be explained in chemical terms.

3. Definitions of Terms
• Matter - all living and nonliving things
• Mass - amount of matter in an object
• Weight - gravitational force acting on an object of a given
mass.
• Element - simplest type of matter having unique chemical
properties

5. Definitions of Terms
• 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. For example, the
element carbon is composed of only carbon atoms, and the
element oxygen is composed of only oxygen atoms.
• Neutron has no electrical charge
• Proton has one positive charge
• Electron has one negative charge.
• Nucleus - Protons and neutrons are found in the center of the
atom.

6. Definitions of Terms
• Atomic Number – The number of protons in each atom
• Mass Number - determined by adding the number of protons
Neutrons in each atom. For example, because carbon has 6
protons and 6 neutrons, its mass number is 12(6 + 6 = 12 ).
• Isotopes - two or more forms of the same element that have
the same number of protons and electrons, but have a
different number of neutrons.

7. Elements
and Atom

8. Electrons and Chemical Bonding
• Chemical Bonds - this level is the interaction between atoms
to form molecules by either sharing or transferring their
outermost electrons
• Electron shells – the energy levels are often drawn as
concentric rings
• Valence Shell - outermost shell, The number of electrons in
the valence shell determines an atom’s chemical behavior.

9. Electrons and Chemical Bonding
• Octet - If an atom’s valence shell is not at its maximum, it
will form bonds with other atoms to achieve a full valence
shell
• Electronegativity is the ability of the atom’s nucleus to pull
electrons toward it. Atoms with a strong electronegativity are
those with 6 or 7 electrons in the valence shell—these are
very close to achieving an octet. Atoms with a weak
electronegativity are those with only 1 or 2 electrons in the
valence shell—these are farther away from achieving an
octet.

10. Relationship Between Electronegativity and
Chemical Bonding
• Electronegativity - a measure of how much an atom attracts
electrons from another atom to form a chemical bond.
Covalent bonds are formed by the sharing of electrons
between atoms that have the same electronegativity (nonpolar
covalent bond; e.g., H2) or a relatively small difference in
electronegativities (polar covalent bonding; e.g., H2O). Ionic
bonds are formed by the transfer of electrons between two
atoms that have very different electronegativities (e.g., NaCl).

11. Ionic bonding
• Recall that an atom is electrically neutral because it has an
equal number of protons and electrons
• Ion - numbers of protons and electrons are no longer equal,
and a charged particle
• Sodium atom (Na) can lose an electron to become a
positively charged sodium ion (Na+)

12. Ionic bonding
• After an atom gains an electron, it has one more electron than
it has protons and is negatively charged. A chlorine atom (Cl)
can accept an electron to become a negatively charged
chloride ion (Cl−). After this transfer of electrons, both
chlorine and sodium have full valence shells.

13. Ionic
Bonding

16. Covalent
Bond

17. Chemical Reaction
• Synthesis Reactions - When two or more reactants combine to
form a larger, more complex product, the process: A + B → AB
• A-P-P + Pi → A-P-P-P
• Decomposition reaction - reactants are broken down into
smaller, less complex products. A decomposition reaction is the
reverse of a synthesis reaction and can be represented in this way:
AB → A + B
• A-P-P-P → A-P-P + P

18. Chemical Reaction
• Reversible reaction - can run in the opposite direction, so
that the products are converted back to the original reactants.
• CO2 + H2O ⇄ H+ + HCO3

19. Definitions of Terms
• Energy is defined as the capacity to do work.
• Potential energy is stored energy.
• Kinetic energy is energy caused by the movement of an
object and is the form of energy that actually does work.
• Mechanical energy is energy resulting from the position or
movement of objects
• Chemical energy is a form of potential energy stored within
the chemical bonds of a substance.

20. Energy and
Chemical
Reactions
(a) The input of energy is
required for the synthesis of
ATP.
(b) Energy is released as a
result of the breakdown of
ATP.

21. Definition of Terms
• Catalyst - increases the rate of a chemical reaction, without
itself being permanently changed or depleted.
• Enzyme - is a protein catalyst that increases the rate at which
a chemical reaction proceeds, without the enzyme being
permanently changed.

23. Acid and Base
• Acid - proton donor. (HCl → H+ + Cl−)
• base - proton acceptor

24. pH Scale
• a means of referring to the H+ concentration in a solution
• Neutral Solution pH 7
• Basic fewer H than OH = ph above 7
***Lower the pH number, the higher the H concentration.

27. Normal Values (Human blood)
• Acidosis – Below 7.35
• Alkalosis – Above 7.45

28. Salts and Buffers
• Salt is a compound consisting of a positive ion other than H+
and a negative ion other than OH−.
HCl + NaOH → NaCl + H2O
(Acid) (Base) (Salt) (Water)
• Buffer - a chemical that resists changes in pH when either an
acid or base is added to a solution containing the buffer

29. 1. When an acid is added to an
unbuffered water solution, the result is
increased H+ and decreased pH.
2. Buffers reduce large changes in pH by
binding and releasing H+.
3. 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.

30. Inorganic Molecules
• Inorganic chemistry - generally deals with substances that
do not contain carbon, although a more rigorous definition is
that it deals with substances that lack carbon-hydrogen bonds.
• Organic chemistry - is the study of carbon-containing
substances, with a few exceptions.

31. Oxygen and Carbon Dioxide
• Oxygen (O2) - a small, nonpolar, inorganic molecule
consisting of two oxygen atoms bound together by a double
covalent bond. About 21% of the gas in the atmosphere is O2,
and it is essential for most living organisms.
• 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. Carbon dioxide is produced
when food molecules, such as glucose, are metabolized
within the cells of the body.

32. Water
• Water stabilizes body temperature.
• Water protects the body.
• Water is required for many chemical reactions.

33. Organic Molecules
• Carbon’s ability to form covalent bonds with other atoms
makes possible the formation of the large, diverse,
complicated molecules necessary for life.

34. Purpose of Organic Molecules
(1)as energy molecules for synthesis of ATP,
(2)as structural components of the cell, and
(3)as regulatory molecules. Carbon atoms bound together by
covalent bonds constitute the “backbone” of many large
molecules

35. Four Major Groups of Organic
Molecule
• Carbohydrate
• Lipid
• Protein
• Nucleic Acid

37. Carbohydrates
(1)they are parts of other organic molecules,
(2)they are broken down to provide energy, and
(3) when undigested, they provide bulk (fiber) in feces. In most
carbohydrates, for each carbon atom there are two hydrogen
atoms and one oxygen atom. Note that this two-to-one ratio is
the same as in water (H2O).

38. Kinds of Carbohydrates
• Monosaccharides simple sugars.
• Monosaccharides are considered the building blocks of
carbohydrates because larger, more complex carbohydrates
are formed by covalent bonds that link many
monosaccharides together.
• Disaccharides are formed when two monosaccharides are
joined by a covalent bond. For example, glucose and fructose
combine to form the disaccharide sucrose (table sugar)

40. Polysaccharides
• long chains of monosaccharides
• Glycogen - main storage form of glucose in humans. Glucose is
quickly broken down by cells to make ATP. Glycogen serves as a
ready supply of more glucose for ATP production.
• Starch stores energy for plants in the same way as glucose stores
energy for animals.

41. Lipids
• Insoluble in water
• composed mainly of carbon, hydrogen, and oxygen
(1)provide protection and insulation
(2)help regulate many physiological processes
(3)form membranes
(4)are major energy storage molecules, which can be broken
down and used as an energy source

42. Major Classes of Lipids
(1)Fats, which are mostly triglycerides
(2)Phospholipids
(3)Eicosanoids
(4)Steroids
• Fats are important energy-storage molecules; they also pad
and insulate the body.
• Triglycerides - most common type of fat molecules and have
three fatty acids bound to a glycerol molecule.

43. Major Classes of Lipids
• Phospholipids - composed of a polar region containing
phosphate and a nonpolar region consisting of two fatty acid
chains
• Eicosanoids - are a group of important chemicals derived
from fatty acids. 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.
• Steroids - derived from cholesterol include bile salts (lipid
digestion), estrogen, progesterone, and testosterone
(reproductive hormones). In addition, cholesterol is an
important component of cell membranes.

45. Proteins
(1)Regulation of body processes
(2)Acting as a transportation system
(3)Providing protection
(4)Helping muscles contract
(5) Providing structure and energy

46. Nucleic Acids: DNA and RNA
• large molecules composed of carbon, hydrogen, oxygen,
nitrogen, and phosphorus
• Deoxyribonucleic acid - genetic material of cells. DNA
contains genes, which determine the structure of proteins.
Ribonucleic acid - which is structurally related to DNA and
important in protein synthesis.