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Marieb Chapter 2: Part A
It’s Alive!
(Chemistry Comes Alive)
Student Version
© 2013 Pearson Education, Inc.
Energy vs. Matter
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•
•
•
Major concepts in this chapter
What are they?
How do they work in the body?
Lets compare their properties:
Energy
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Matter
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•
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Matter
• Matter—anything that has mass and
occupies space
• 3 states of matter
–
–
–
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Energy
• Capacity to do work
• Types of energy
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–
—energy in action
—stored (inactive) energy
• Energy can be converted from
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Forms of Energy
•
energy
– Stored in bonds of chemical substances
•
energy
– Results from movement of charged particles
•
energy
– Directly involved in moving matter
•
energy
– Travels in waves (e.g., visible light, ultraviolet
light, and x-rays)
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Energy Conversions
• Energy may be converted from one form
to another
• Energy conversion is inefficient
– Some energy is “lost” as heat (partly unusable
energy)
• Gasoline is ___________ energy; we can only
remove 8 - 10% of the available energy to do work
• Lance Armstrong’s muscles are 23% efficient
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The Energy Name Game!
A
?
B
?
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Name the Energy!
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Video
• Potential vs Kinetic Energy … Ole!
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Challenge Yourself! What Type of Energy?
Potential or Kinetic?
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Composition of Matter: Elements
• Elements
– Matter is composed of elements
– Elements cannot be broken into simpler
substances by ordinary chemical methods
– Each has its own unique properties
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Composition of Matter
• Atoms
– We will use the term atom rather than element
• Atomic symbol
– One- or two-letter chemical shorthand for
each element
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The Periodic Table
Lanthanides
Actinides
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Major Elements of the Human Body
• Four elements make up ?% of body mass
Element
Carbon
Hydrogen
Oxygen
Nitrogen
Atomic symbol
C
H
O
N
Know these symbols!
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Lesser Elements of the Human Body
9 elements make up 3.9% of body mass
Element
Calcium
Phosphorus
Potassium
Sulfur
Sodium
Chlorine
Magnesium
Iodine
Iron
Atomic symbol
Ca
P
K
S
Na
Cl
Mg
I
Fe
Yep, you need to know these too!
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Trace Elements of the Human Body
Occur in very minute amounts
• 11 elements make up < 0.01% of body mass
– Many are part of, or activate, enzymes
• For example:
Element
Chromium
Copper
Fluorine
Manganese
Silicon
Zinc
Atomic symbol
Cr
Cu
F
Mn
Si
Zn
These are important but you won’t find them
mentioned as often. Don’t worry about these symbols.
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How Important Are These Elements?
This is a generic form of _____________
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Table 2.1.1
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Table 2.1.2
Atomic Structure
• Atoms are composed of subatomic
particles
–
–
–
• Protons and neutrons are in the nucleus
• Electrons orbit the nucleus in an electron
cloud
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Atomic Structure: The Nucleus
• Almost entire mass of the atom
– Neutrons
• Carry no charge (Neutral!)
– Protons
• Carry positive charge
• Mass = 1 amu
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Atomic Structure: Electrons
• Electrons in orbitals within electron cloud
– Circle outside the nucleus
– Carry negative charge
– Much smaller than protons or neutrons
– Number of protons and electrons are always
equal
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Figure 2.1 Two models of the structure of an atom.
Nucleus
Nucleus
Helium atom
Helium atom
2 protons (p+)
2 neutrons (n0)
2 electrons (e–)
2 protons (p+)
2 neutrons (n0)
2 electrons (e–)
Planetary model
Proton
Neutron
Electron
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Electron
cloud
Orbital model
Identifying Elements
• Different atoms contain different numbers
of subatomic particles
– Hydrogen has 1 proton, 0 neutrons, and 1
electron
– Lithium has 3 protons, 4 neutrons, and 3
electrons
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Figure 2.2 Atomic structure of the three smallest atoms.
Proton
Neutron
Electron
Hydrogen (H)
(1p+; 0n0; 1e–)
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Helium (He)
(2p+; 2n0; 2e–)
Lithium (Li)
(3p+; 4n0; 3e–)
Combining Matter: Forming Molecules
• Most atoms chemically combined with
other atoms to form molecules
– Molecule
• Two or more atoms bonded together (e.g., H2 or
C6H12O6)
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Mixtures
• Most matter exists as mixtures
– Two or more components physically intermixed
• Three types of mixtures
– Solutions
– Colloids
– Suspensions
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Types of Mixtures: Solutions
• Homogeneous mixtures
• Most are true solutions in body
– Gases, liquids, or solids dissolved in water
– Usually transparent, e.g., atmospheric air or salt
solution
• Solvent
–
– Usually a liquid; usually water
• Solute(s)
–
• For example:
– If glucose is dissolved in blood, glucose is the
; blood is the
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How Do We Make Solutions?
• Can be expressed as
– Percent of solute in total solution
(solvent assumed to be water)
• Parts solute per 100 parts solvent
– Milligrams per deciliter (mg/dl)
– (weight per volume)
– Molarity, or moles per liter (M)
• 1 mole of an element or compound = Its atomic or
molecular weight (sum of atomic weights) in grams
• 1 mole of any substance contains 6.02  1023
molecules of that substance (Avogadro’s number)
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Colloids and Suspensions
• Colloids (also called emulsions)
– Heterogeneous mixtures, e.g., cytosol
– Large solute particles do not settle out
•
• Suspensions
– Heterogeneous mixtures
– Large, visible solutes settle out
•
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Figure 2.4 The three basic types of mixtures
Colloid
Suspension
Solute particles are larger than
in a solution and scatter light;
do not settle out.
Solute particles are very large,
settle out, and may scatter light.
Solution
Solute particles are very tiny,
do not settle out or scatter light.
Solute
particles
Solute
particles
Example
Mineral water
Solute
particles
Example
Example
Jello
Whole
Blood
Plasma
Settled red
blood cells
Unsettled Settled
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What About Those Electrons?
• Electrons have potential energy
– Each shell corresponds to a specific level of potential
energy
– Shells that are farther from the nucleus contain
electrons with more potential energy
– Up to seven electron shells (energy levels) can exist
• Atoms are most stable when their outermost
(valence) shell is full
• Atoms will interact with other atoms to fill their
outermost shells (via chemical bonds)
• Octet rule (rule of eights)
– Except for the first shell (full with two electrons) atoms
interact to have eight electrons in their valence shell
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Chemically Inert Elements
• Valence shell fully occupied or contains
eight electrons
• Stable and unreactive because of this
• The “loners”
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Figure 2.5a Chemically inert and reactive elements.
Chemically inert elements
Outermost energy level (valence shell) complete
2e
Helium (He)
(2p+; 2n0; 2e–)
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8e
2e
Neon (Ne)
(10p+; 10n0; 10e–)
Chemically Reactive Elements
• Outer shell not full
• Tend to gain, lose, or share electrons
(form bonds) with other atoms to achieve
stability
• The “party animals” - like to interact with
other atoms
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Figure 2.5b Chemically inert and reactive elements.
Chemically reactive elements
Outermost energy level (valence shell) incomplete
1e
Hydrogen (H)
(1p+; 0n0; 1e–)
6e
2e
Oxygen (O)
(8p+; 8n0; 8e–)
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4e
2e
Carbon (C)
(6p+; 6n0; 6e–)
1e
8e
2e
Sodium (Na)
(11p+; 12n0; 11e–)
Types of Chemical Bonds
• We will discuss three major types:
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–
–
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Ionic Bonds
• Ions are formed
– An atom gains or loses electrons and
becomes charged
• The transfer of valence shell electrons
from one atom to another forms ions
– One becomes an anion (negative charge)
• Atom that gained one or more electrons
– One becomes a cation (positive charge)
• Atom that lost one or more electrons
• Attraction of opposite charges results in an
ionic bond
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Figure 2.6a–b Formation of an ionic bond.
(17p+; 18n0; 18e–))
(11p+; 12n0; 10e–))
+
Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
Sodium gains stability by losing
one electron, and chlorine becomes
stable by gaining one electron.
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Sodium ion (Na+)
—
Chloride ion (Cl–)
Sodium chloride (NaCl)
After electron transfer,
the oppositely charged ions
formed attract each other.
Ionic Compounds
• Most ionic compounds are salts
– When dry, salts form crystals instead of
individual molecules
– Example is NaCl (sodium chloride)
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Figure 2.6c Formation of an ionic bond.
Cl–
Na+
Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
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Do Salts Exist In the Body?
• Rarely, because most dissolve in water!
• Most ionic bonds are easily broken by
water; ions are released
• Their ions do exist in the body =
• Exceptions:
–Calcium phosphate in bone
–Kidney stones (ouch!)
–Gall stones (also ouch!)
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Ionic Bonds are
Weak most salt crystals
dissolve in water
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Figure 2.12
Common Ions You Need To Know
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OHPO43HCO3NH4+
H+
Na+
K+
Ca2+
ClICOO-
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Hydroxide
Phosphate
Bicarbonate
Ammonium
Hydrogen (proton)
Sodium
Potassium
Calcium
Chloride
Iodide
Carboxyl
Covalent Bonds
• Formed by
of two or more
valence shell electrons
• Allows each atom to fill its valence shell at
least part of the time
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Figure 2.7a Formation of covalent bonds.
Reacting atoms
Resulting molecules
+
or
Structural formula
shows single bonds.
Carbon atom
Hydrogen atoms
Formation of four single covalent bonds:
Carbon shares four electron pairs with
four hydrogen atoms.
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Molecule of methane gas (CH4)
Figure 2.7b Formation of covalent bonds.
Reacting atoms
Resulting molecules
+
Oxygen atom
Oxygen atom
Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
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or
Structural formula
shows double bond.
Molecule of oxygen gas (O2)
Figure 2.7c Formation of covalent bonds.
Reacting atoms
Resulting molecules
or
+
Nitrogen atom
Nitrogen atom
Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
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Structural formula
shows triple bond.
Molecule of nitrogen gas (N2)
Two Different Types of Covalent Bonds
• Non-polar Covalent bonds
• Electrons equally shared
• Polar Covalent Bonds
• Electrons UNequally shared
• Come in single, double, or triple forms
• More lines means more electrons shared
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Single, Double, and Triple Covalent Bonds
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Chemical and Structural Formulas
C2H5OH
Structural Formula
Chemical Formula
(chemist’s shorthand)
Sometimes you only see the “skeleton” of a molecule!
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What Do I Mean?
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More “Shorthand” - The Missing Atoms
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More “Shorthand” - The Missing Atoms
This is Cholesterol
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Here Are The Missing Atoms!
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Nonpolar Covalent Bonds
• Electrons shared equally
• Which bonds are nonpolar?
• C-C
• C-H
• If we have a lot of these bonds in a
biological molecule, the molecule is
non-polar (Most lipids= fats!)
• It would be oily, greasy, and hate water
– (hydrophobic =
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)
Polar Covalent Bonds
• Unequal sharing of electrons produces polar bonds
• Examples of polar covalent bonds:
•
•
•
•
•
C-O
N-H
O-O
N-O
O-H
• A molecule with lots of polar covalent bonds loves to
interact with the polar covalent bonds in water
– Hydrophilic =
• Most biological molecules are polar.
• Carbohydrates
• Proteins
• Nucleic acids
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• Some lipids
Water Is A Polar Molecule
–
+
–
+
+
+
V-shaped water (H2O) molecules have two
poles of charge—a slightly more negative
oxygen end (–) and a slightly more positive
hydrogen end (+).
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Figure 2.9 Ionic, polar covalent, and nonpolar covalent
bonds compared.
Ionic bond
Polar covalent
bond
Nonpolar
covalent bond
Complete
transfer of
electrons
Unequal sharing
of electrons
Equal sharing of
electrons
Separate ions
(charged
particies)
form
Slight negative
charge (–) at
one end of
molecule, slight
positive charge (+)
at other end
Charge balanced
among atoms
–
+
Sodium chloride
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+
Water
Carbon dioxide
Let’s Watch A Video or Two or Three…!
• http://bit.ly/VkdqRJ
• http://bit.ly/SouL10
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Hydrogen Bonds
• Attractive force between a + charged
hydrogen atom of one molecule and an
- charged atom of another molecule
– Common between molecules with polar bonds
such as water
– Also found within a molecule, holding it in a
three-dimensional shape
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Figure 2.10a Hydrogen bonding between polar water
molecules.
+

+

Hydrogen bond
(indicated by
dotted line)
+
+

+
+

+

+

+
+
The slightly positive ends (+) of the water molecules
become aligned with the slightly negative ends ()
of other water molecules.

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Figure 2.10b Hydrogen bonding between polar water
molecules causes water to be “strong”.
A water strider can walk on a pond because of the high
surface tension of water, a result of the combined
strength of its hydrogen bonds.
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DNA Has Lots Of Hydrogen Bonds
Sugar:
Phosphate Deoxyribose
Base:
Adenine (A)
Thymine (T)
Thymine nucleotide
Adenine nucleotide
Hydrogen
bond
Sugarphosphate
backbone
Deoxyribose
sugar
Phosphate
Adenine (A)
Thymine (T)
Cytosine (C)
Guanine (G)
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Sugar
Phosphate
Water vs. Ice
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Hydrogen Bond Videos
• http://www.youtube.com/watch?v=
LK7ERiCy5b8
• http://www.youtube.com/watch?v=
PyC5r2mB4d4
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Chemical Groups In Biological Molecules
• Hydroxyl
• Amino
• Phenyl
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Chemical Groups In Biological Molecules
• Methyl
• Carboxyl
• Phosphate
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Chemical Groups In Biological Molecules
• Hydrocarbon chain
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Polar Or Non-polar Molecule?
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Polar Or Non-polar Molecule?
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Polar Or Non-polar Molecule?
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Chemical Reactions
• Occur when chemical bonds are formed,
rearranged, or broken
• Represented as chemical equations using
molecular formulas
• Chemical equations contain
– Reactant(s) = what you start with
– Chemical composition of the product(s)
= what you end up with
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Examples of Chemical Equations
Reactant(s)
H+H
4H + C
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

Product(s)
H2 (Hydrogen gas)
CH4 (Methane)
Types of Chemical Reactions
• Synthesis (combination) reactions
• Decomposition reactions
• Exchange (swapping) reactions
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Synthesis Reactions
• A + B  A-B
– Atoms or molecules combine to form larger,
more complex molecule
– Always involve bond formation
– Anabolic reaction
– SMALLER TO BIGGER
• Example:
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Figure 2.11a Patterns of chemical reactions.
Synthesis reactions
Smaller particles are bonded
together to form larger,
more complex molecules.
Example
Amino acids are joined together to
form a protein molecule.
Amino acid
molecules
Protein
molecule
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Peptide bond
Decomposition Reactions
• A-BA+B
– Molecule is broken down into smaller
molecules or its constituent atoms
• Reverse of synthesis reactions
– Involve breaking of bonds
– Catabolic reaction
– BIGGER TO SMALLER
– Example:
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Figure 2.11b Patterns of chemical reactions.
Decomposition reactions
Bonds are broken in larger
molecules, resulting in smaller,
less complex molecules.
Example
Glycogen is broken down to release
glucose units.
Glycogen
Glucose
molecules
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Exchange Reactions
• A-B + C  A-C + B
– Also called displacement reactions
– Involve both synthesis and decomposition
– Bonds are both made and broken
– SWAPPING ATOMS
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Figure 2.11c Patterns of chemical reactions.
Exchange reactions
Bonds are both made and broken
(also called displacement reactions).
Example
ATP transfers its terminal phosphate
group to glucose to form glucosephosphate.
+
Adenosine triphosphate (ATP)
Glucose
+
Adenosine diphosphate
(ADP)
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Glucosephosphate
Oxidation-Reduction (Redox) Reactions
• Let’s talk about these later when we
discuss metabolism in detail…
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Reversibility of Chemical Reactions
• Almost all biochemical reactions are
reversible
A + B  AB
AB  A + B
• Chemical equilibrium occurs when a
reaction is reversible
• Many biological reactions appear
irreversible
– Due to energy requirements
– Due to removal of products
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Rate of Chemical Reactions
• Affected by
–  Temperature   Rate
–  Concentration of reactant   Rate
–  Particle size   Rate
– Catalysts:  Rate without being chemically
changed or part of product
• Enzymes are biological catalysts
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The Law of Mass Action
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The Law of Mass Action
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The Law of Mass Action
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