Human Anatomy & Physiology
Chapter 02
The Chemistry of Life
PowerPoint® Lectures created by Suzanne Pundt, University of Texas at Tyler
Copyright © 2019, 2016 Pearson Education, Inc. All Rights Reserved
Second Edition
Module 2.1 Atoms and Elements
Matter
•	Matter – anything that has mass and occupies space; exists in three states: solid, liquid, or gas 
•	Chemistry – study of matter and its interactions
Atoms and Atomic Structure (1 of 3)
•	Atom – smallest unit of matter that retains original properties
•	Made up of even smaller structures called subatomic particles
Figure 2.1 Structure of a representative atom.
Atoms and Atomic Structure (2 of 3)
•	Subatomic particles:
•	Protons (p+) – in central core 
of atom (atomic nucleus); positively 
charged
•	Neutrons (n0) – in atomic 
nucleus; slightly larger than protons; 
no charge
•	Electrons (e-) – outside atomic 
nucleus; negatively charged
•	Atoms are electrically neutral – no charge; number of protons and electrons are equal; cancel each other’s charge; neutron number does not have to equal proton number
Figure 2.1 Structure of a representative atom.

Atoms and Atomic Structure (3 of 3)
•	Electron shells – regions surrounding atomic nucleus where electrons exist; each holds certain number of electrons:
•	1st shell (closest to nucleus) – 
2 electrons
•	2nd shell – 8 electrons
•	3rd shell – 18 electrons; “satisfied” with 8
•	Some atoms may have more than 3 shells
Elements in the Periodic Table and the Human Body (1 of 3)
•	Atomic number – number of protons in atomic nucleus; defines every element:
•	Element – substance that cannot be broken down into simpler substance by chemical means
•	Each element is made of atoms with same number of protons
Elements in the Periodic Table and the Human Body (2 of 3)
•	Periodic table of elements – lists elements by increasing atomic numbers:
•	Groups – organized by certain properties
•	Chemical symbol – represents each element 
The Periodic Table
Figure 2.2 Elements in the human body and their positions in the periodic table.
Elements in the Periodic Table and the Human Body (3 of 3)
•	Human body is made up of four major elements: 
•	Hydrogen
•	Oxygen
•	Carbon
•	Nitrogen
•	Also 7 mineral elements and 13 trace elements
Isotopes and Radioactivity
•	Mass number – equal to sum of all protons and neutrons in atomic nucleus
•	Isotope – atom with same atomic number (number of protons) but different mass number (number of neutrons)

•	Radioisotopes – unstable isotopes; high energy or radiation released by radioactive decay; allows isotope to assume more stable form
Nuclear Medicine
Common applications of radioisotopes:
•	Cancer radiation therapy – radiation damages structure of cancer cells; interferes with functions
•	Radiotracers – injected into patient; radiation they produce is 
detected by a camera; image analyzed by 
computer; shows size, shape, and activity of organs and cells
•	Treatment of thyroid disorders – high doses of iodine-131 treat overactive or cancerous thyroid tissue; radioisotope accumulates and damages cells
Module 2.2 Matter Combined: Mixtures and Chemical Bonds
Matter Combined
•	Matter can be combined physically to form mixtures – atoms of two or more elements physically intermixed without changing chemical nature of atoms
•	Three basic types of mixtures: suspensions, colloids, and solutions (Figure 2.3) 
Mixtures (1 of 3)
•	Suspension – mixture containing two or more components 
•	Large, unevenly distributed particles
•	Will settle out when left undisturbed
Figure 2.3a The three types of mixtures.
Mixtures (2 of 3) 
•	Colloids – two or more components 
•	Small, evenly distributed particles
•	Will not settle out
Figure 2.3b The three types of mixtures.
Mixtures (3 of 3) 
•	Solutions – two or more components; extremely small, evenly distributed particles; will not settle out:
•	Solute – substance dissolved
•	Solvent – substance that dissolves solute
Figure 2.3c The three types of mixtures.
Chemical Bonds (1 of 4) 
•	Matter can be combined chemically when atoms are combined by chemical bonds
•	Chemical bond is not physical structure; energy relationship or attractive force between atoms
•	Molecule – formed by chemical bonding 
between two or more atoms of same element
•	Compound – formed when two or more atoms from different elements combine by chemical bonding
Figure 2.5  Formation of a covalent bond.
Chemical Bonds (2 of 4) 
Table 2.1 Electron Sharing in Covalent Bonds.
Chemical Bonds (3 of 4)
•	Chemical bonds are formed when valence electrons (outermost shell) of atoms interact
•	Valence electrons determine how an atom interacts with other atoms; whether it will form bonds with specific atom
•	Octet rule – atom is most stable when 8 electrons in valence shell (CO2)
Chemical Bonds (4 of 4)

•	Duet rule (atoms with 5 or fewer electrons) – atom is most stable when valence electron shell holds 
2 electrons
Figure 2.5 Formation of a covalent bond.
Ions and Ionic Bonds
•	Ionic bond – electrons are transferred from metal atom to nonmetal atom; results in formation of ions: cations and anions (Figure 2.4) 
•	Cation – positively charged ion; metal loses one or more electrons
•	Anion – negatively charged ion; nonmetal gains one or more electrons
•	Attraction between opposite charges bonds ions to one another forming salt
Ionic Bonds
Figure 2.4 Formation of an ionic bond.
Covalent Bonds (1 of 2) 
•	Covalent bonds – strongest bond; two or more nonmetals share electrons (Figures 2.5, 2.6; Table 2.1) 
•	Two atoms can share one (single bond), two (double bond), or three (triple bond) electron pairs:
Table 2.1 Electron Sharing in Covalent Bonds.
Covalent Bonds (2 of 2)
Protons attract electrons; known as electronegativity:
•	Element’s electronegativity increases from bottom left to upper right of periodic table; fluorine (F) is most electronegative element
•	Electronegative elements strongly attract electrons; pull away from less electronegative elements
Nonpolar Covalent Bonds (1 of 2)
Nonpolar covalent bonds – two nonmetals in molecule with similar or identical ectronegativities pull with equal force; share electrons equally (Figure 2.6a)
Nonpolar molecules occur in 3 situations:
•	Atoms sharing electrons are same element 
•	Arrangement of atoms makes one atom unable to pull more strongly than another atom (CO2)
•	Bond is between carbon and hydrogen
Nonpolar Covalent Bonds (2 of 2)
Figure 2.6a Nonpolar vs. polar covalent bonds.
Polar Covalent Bonds (1 of 2)
•	Polar covalent bonds – nonmetals with different electronegativities share electrons unequally; form polar molecules (Figure 2.6b)
•	Atom with higher electronegativity becomes partially negative (δ−); pulls shared electrons close to itself
•	Atom with lower electronegativity becomes partially positive (δ+); shared electrons are pulled toward other atom 
•	Dipoles – polar molecules with partially positive and partially negative ends
Polar Covalent Bonds (2 of 2)
Figure 2.6b Nonpolar vs. polar covalent bonds.
Hydrogen Bonds (1 of 2)
Hydrogen bonds – weak attractions between partially positive hydrogen atoms and partially negative nonmetal atoms in polar covalent molecules.
•	Responsible for key property of water – surface tension
Figure 2.7a Hydrogen bonding and surface tension between water molecules.
Hydrogen Bonds (2 of 2)
•	Polar water molecules are more strongly attracted to one another than to nonpolar air molecules at surface
Figure 2.7b Hydrogen bonding and surface tension between water molecules.
Concept Boost: Determining the Type of Bonds in a Molecule (1 of 2)
Basic “rules”:
•	A compound is ionic when: the compound contains both a metal and a nonmetal. 
•	If the molecule or compound contains two or more nonmetals; bond is covalent (hydrogen behaves like nonmetal):
•	Molecule contains two identical nonmetals; nonpolar covalent 
(e.g., O2)
•	Molecule contains only (or primarily) carbon and hydrogen; nonpolar covalent (e.g., CH4)
Concept Boost: Determining the Type of Bonds in a Molecule (2 of 2)
Basic “rules” (continued):
•	A compound is polar when: The molecule contains two nonmetals of significantly different electronegativities, then polar covalent (hydrogen and carbon have low electronegativities; oxygen, nitrogen, and phosphorus have high electronegativities); compounds with proportionally high number of C-O, C-N, P-O, H-N, C-Cl, and H-O bonds are polar (e.g., C5H10O5, CH3Cl, H3PO4)

Concept Boost: Determining the Type of Bonds in a Molecule
•	Remember that these rules don’t apply to every situation, but in general, they are a good place to start.

Module 2.3 Chemical Reactions
Chemical Notation (1 of 2) 
•	Chemical reaction – when chemical bond is formed, broken, or rearranged, or electrons are transferred between two or more atoms (or molecules)
•	Chemical notation – series of symbols and abbreviations; demonstrates what occurs in a reaction; chemical equation (basic form of chemical notation) has two parts: 
•	Reactants (left side of equation) – starting ingredients; will undergo reaction 
•	Products (right side of equation) – results of chemical reaction
Chemical Notation
•	Reversible reactions can proceed in either direction as denoted by two arrows that run in opposite directions (as below)
•	Irreversible reactions proceed from left to right as denoted by a single arrow
CO2 + H2O      H2CO3	
	Reactants (carbon dioxide + water)	Product (carbonic acid)

Energy and Chemical Reactions (1 of 3)
•	Energy is defined as capacity to do work, can put matter into motion or fuel chemical reactions; two general forms of energy: 
•	Potential energy is stored; can be released to do work at some later time
•	Kinetic energy is potential energy that has been released or set in motion to perform work; 
all atoms have kinetic energy 
as they are in constant motion; 
the faster they move the
greater that energy
Energy and Chemical Reactions (2 of 3)
Three forms of energy in human body: chemical, electrical, and mechanical, each may be potential or kinetic depending on location or process
•	Chemical energy – in bonds between atoms; drives nearly all chemical processes
•	Electrical energy – generated by movement of charged particles or ions
•	Mechanical energy – energy directly transferred from one object to another
Energy and Chemical Reactions (3 of 3)
Energy, inherent in all chemical bonds, must be invested any time chemical reaction occurs:
•	Endergonic reactions – require input of energy from another source; products contain more energy than reactants (energy invested so reaction could proceed)
•	Exergonic reactions – release excess energy; products have less energy than reactants
Homeostasis and Types of Chemical Reactions (1 of 5)
Three fundamental processes maintain homeostasis (breaking down molecules, converting energy in food to usable form, and building new molecules); carried out by three types of chemical reactions:
•	 Catabolic reactions (decomposition reactions)
•	 Exchange reactions
•	 Anabolic reactions (synthesis reactions)
Homeostasis and Types of Chemical Reactions (2 of 5)
•	Catabolic reactions (decomposition reactions) – when a large substance is broken down into smaller substances
•	General chemical notation for reaction is
AB  A + B
•	Usually exergonic because chemical bonds are broken
Homeostasis and Types of Chemical Reactions (3 of 5)
•	Exchange reactions occur when one or more atoms from reactants are exchanged for another
•	General chemical notation for reaction is  
AB + CD  AD + BC
HCL + NaOH  H2O + NaCL
Homeostasis and Types of Chemical Reactions (4 of 5)
•	Oxidation-reduction reactions (redox reactions) – special kind of exchange reaction; occur when electrons are exchanged instead of atoms
•	Reactant that loses electrons is oxidized 
•	Reactant that gains electrons is reduced
•	Redox reactions are usually exergonic reactions capable of releasing large amounts of energy
Homeostasis and Types of Chemical Reactions (5 of 5)
•	Anabolic reactions (synthesis reactions) occur when small simple subunits are united by chemical bonds to make large more complex substances
•	General chemical notation for reaction is
A + B  AB
•	These reactions in the human body are endergonic; fueled by chemical energy
Reaction Rates and Enzymes (1 of 9)
•	For a reaction to occur atoms must collide with enough force to overcome the repulsion of their electrons
•	This energy required for all chemical reactions is called the activation 
energy (Ea)

Reaction Rates and Enzymes (2 of 9)
Analogy – activation energy must be supplied so that reactants reach transition states 
(i.e., get to top of energy “hill”) to react and form products (i.e., roll down hill)
Figure 2.8  Activation energy.
Reaction Rates and Enzymes (3 of 9)
•	Following factors increase reaction rate by reducing activation energy or increasing likelihood of strong collisions between reactants: 
•	Concentration
•	Temperature
•	Reactant properties
•	Presence or absence of catalyst
Reaction Rates and Enzymes (4 of 9)
•	When reactant concentration increases, more reactant particles are present; increases chance of successful collisions between reactants
•	Raising temperature of reactants increases kinetic energy of atoms; leads to more forceful and effective collisions between reactants
Reaction Rates and Enzymes (5 of 9)
•	Both particle size and phase (solid, liquid, or gas) influence reaction rates:
•	Smaller particles move faster; more energy than larger particles
•	Reactant particles in gaseous phase have higher kinetic energy than either solid or liquid phase
Reaction Rates and Enzymes (6 of 9)
•	Catalyst – increases reaction rate by lowering activation energy; not consumed or altered in reaction
•	Enzymes – biological catalysts; most are proteins with following properties:
•	Speed up reactions by lowering activation energy (Figure 2.9)
•	Highly specific for individual substrates (substance that can bind to enzyme’s active site)
•	Do not alter reactants or products
•	Not permanently altered in reactions catalyzed
Reaction Rates and Enzymes (7 of 9)
Figure 2.9 The effect of enzymes on activation energy.
Reaction Rates and Enzymes (8 of 9)
•	 Induced-fit mechanism – enzyme’s interaction with substrate(s)
•	Binding of substrate causes small shape change that brings substrate to transition state
Reaction Rates and Enzymes (9 of 9)
•	 Induced-fit mechanism (continued):
•	Reduces energy of activation; allows transition state to  proceed to final products
Figure 2.10  Enzyme–substrate interaction.
Enzyme Deficiencies
Examples of common enzyme deficiencies:
•	Tay-Sachs Disease – deficiency of hexosaminidase; gangliosides accumulate around neurons of brain; death usually by age 3
•	Severe Combined Immunodeficiency Syndrome (SCIDs) – may be due to adenosine deaminase deficiency; nearly complete absence of immune system; patients must live in sterile “bubble”
•	Phenylketonuria – deficiency of phenylalanine hydroxylase; converts phenylalanine into tyrosine; symptoms include seizures and intellectual disability; can be prevented by dietary modification