The Chemistry of Life

1
The Chemistry of Life
2
The Nature of Matter

• Life depends on chemistry
• When you eat food or inhale oxygen, your body
  uses these materials in chemical reactions that
  keep you alive
• Just as buildings are made from bricks, steel,
  glass, and wood, living things are made from
  chemical compounds
• If the first task of an architect is to
  understand building materials, then the first job
  of a biologist is to understand the chemistry of
  life
• YOU ARE WHAT YOU EAT!!!!!!!!!

3
Atoms

• The study of chemistry begins with the basic unit
  of matter, the atom
• The Greek word atomos, which means unable to be
  cut, was first used to refer to matter by the
  Greek philosopher Democritus nearly 2500 years
  ago
• Democritus asked a simple question If you take
  an object like a stick of chalk and break it in
  half, are both halves still chalk?
• The answer, of course, is yes
• But what happens if you go on? Suppose you break
  it in half again and again and again
• Can you continue to divide without limit, or does
  there come a point at which you cannot divide the
  fragment of chalk without changing it into
  something else?
• Democritus thought that there had to be a limit
• He called the smallest fragment the atom, a name
  scientists still use today

4
Atoms

• Atoms are incredibly small
• Placed side by side, 100 million atoms would make
  a row only about 1 centimeter longabout the
  width of your little finger!
• Despite its extremely small size, an atom
  contains subatomic particles that are even smaller

5
Atoms

• The figure TO THE RIGHT shows the subatomic
  particles in a helium atom
• The subatomic particles that make up atoms are
  protons, neutrons, and electrons
• Protons and neutrons have about the same mass
• However, protons are positively charged particles
  () and neutrons carry no charge
• Their name is a reminder that they are neutral
  particles
• Strong forces bind protons and neutrons together
  to form the nucleus, which is at the center of
  the atom

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Atoms
7
Atoms

• Helium atoms contain protons, neutrons, and
  electrons
• The positively charged protons and uncharged
  neutrons are bound together in the dense nucleus,
  while the negatively charged electrons move in
  the space around the nucleus

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Atoms

• The electron is a negatively charged particle (-)
  with 1/1840 the mass of a proton
• Electrons are in constant motion in the space
  surrounding the nucleus
• They are attracted to the positively charged
  nucleus but remain outside the nucleus because of
  the energy of their motion
• Because atoms have equal numbers of electrons and
  protons, and because these subatomic particles
  have equal but opposite charges, atoms are
  neutral before they react!!!!!!

9
Elements

• A chemical element is a pure substance that
  consists entirely of one type of atom
• More than 100 elements are known, but only about
  two dozen are commonly found in living organisms
• Elements are represented by a one- or two-letter
  symbol
• Example
• C stands for carbon
• H for hydrogen
• Na for sodium
• The number of protons in an atom of an element is
  the element's atomic number
• Carbon's atomic number is 6, meaning that each
  atom of carbon has six protons and, consequently,
  six electrons

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• ATOM
• IS THE FUNDAMENTAL UNIT OF MATTER
• COMPOSED OF SUBATOMIC PARTICLES
• ATOMIC NUCLEUS
• PROTON
• NEUTRON
• ELECTRONSORBIT (OUTSIDE) THE NUCLEUS 

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• ATOMIC NUMBER THE NUMBER OF PROTONS IN THE
  NUCLEUS OF THE ATOM.
• BEFORE AN ATOM REACTS THE NUMBER OF PROTONS AND
  ELECTRONS ARE EQUAL 

14

• MASS NUMBER (ATOMIC MASS) IS THE SUM OF THE
  PROTONS AND NEUTRONS IN THE NUCLEUS OF THE ATOM.
• amu .000,000,000,000,000,000,000,001,67 g
• Electron mass is so small (negligible) that when
  calculating MASS NUMBER of an atom the mass of
  the electron is considered zero ( 0).

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• LOCATION MASS CHARGE
• ( in amu)
• _____________________________________
• PROTON nucleus 1
  1
• _____________________________________
• NEUTRON nucleus 1 0
• _____________________________________
• ELECTRON outside 1/2000 -1
• nucleus
• _____________________________________

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Isotopes

• Atoms of an element can have different numbers of
  neutrons
• Example
• Some atoms of carbon have six neutrons, some have
  seven, and a few have eight
• Atoms of the same element that differ in the
  number of neutrons they contain are known as
  isotopes
• The sum of the protons and neutrons in the
  nucleus of an atom is called its mass number
• Isotopes are identified by their mass numbers.
• The figure at right shows the subatomic
  composition of carbon-12, carbon-13, and
  carbon-14 atoms
• The weighted average of the masses of an
  element's isotopes is called its atomic mass
• Weighted means that the abundance of each
  isotope in nature is considered when the average
  is calculated
• Because they have the same number of electrons,
  all isotopes of an element have the same chemical
  properties

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Isotopes of Carbon

• Because they have the same number of electrons,
  these isotopes of carbon have the same chemical
  properties
• The difference among the isotopes is the number
  of neutrons in their nuclei

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Isotopes of Carbon
21
Radioactive Isotopes 

• Some isotopes are radioactive, meaning that their
  nuclei are unstable and break down at a constant
  rate over time
• The radiation these isotopes give off can be
  dangerous, but radioactive isotopes have a number
  of important scientific and practical uses
• Geologists can determine the ages of rocks and
  fossils by analyzing the isotopes found in them
• Radiation from certain isotopes can be used to
  treat cancer and to kill bacteria that cause food
  to spoil
• Radioactive isotopes can also be used as labels
  or tracers to follow the movements of
  substances within organisms

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Chemical Compounds

• In nature, most elements are found combined with
  other elements in compounds
• A chemical compound is a substance formed by the
  chemical combination of two or more elements in
  definite proportions
• Scientists show the composition of compounds by a
  kind of shorthand known as a chemical formula
• Water, which contains two atoms of hydrogen for
  each atom of oxygen, has the chemical formula H2O
• The formula for table salt, NaCl, indicates that
  the elements from which table salt formssodium
  and chlorinecombine in a 1 1 ratio

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Chemical Compounds

• The physical and chemical properties of a
  compound are usually very different from those of
  the elements from which it is formed
• Example
• Hydrogen and oxygen, which are gases at room
  temperature, can combine explosively and form
  liquid water
• Sodium is a silver-colored metal that is soft
  enough to cut with a knife
• It reacts explosively with cold water
• Chlorine is very reactive, too
• It is a poisonous, greenish gas that was used to
  kill many soldiers in World War I.
• Sodium and chlorine combine to form sodium
  chloride (NaCl), or table salt
• Sodium chloride is a white solid that dissolves
  easily in water. As you know, sodium chloride is
  not poisonous
• In fact, it is essential for the survival of most
  living things

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Chemical Bonds

• The atoms in compounds are held together by
  chemical bonds
• Much of chemistry is devoted to understanding how
  and when chemical bonds form
• Bond formation involves the electrons that
  surround each atomic nucleus
• The electrons that are available to form bonds
  are called valence electrons
• The main types of chemical bonds are ionic bonds
  and covalent bonds

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Ionic Bonds 

• An ionic bond is formed when one or more
  electrons are transferred from one atom to
  another
• Recall that atoms are electrically neutral
  because they have equal numbers of protons and
  electrons
• Electrically stable BUT chemically unstable
• An atom that loses electrons has a positive
  charge
• An atom that gains electrons has a negative
  charge
• These positively and negatively charged atoms are
  known as ions

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• ATOMIC NUMBER THE NUMBER OF PROTONS IN THE
  NUCLEUS OF THE ATOM
• BEFORE AN ATOM REACTS THE NUMBER OF PROTONS AND
  ELECTRONS ARE EQUAL 

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• ELECTRONS
• Do not move about an atom in definite orbits
• Only the probability of finding an electron at a
  particular place in an atom can be determined
• Each electron seems to be locked into a certain
  area in the electron cloud

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• Modern Atomic Theory
• Electrons are arranged in energy levels
• An energy level represents the most likely
  location in the electron cloud in which an
  electron can be found

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• MODERN ATOMIC THEORY
• Electrons with the lowest energy are found in the
  energy level closest to the nucleus
• Electrons with the highest energy are found in
  the energy levels farther from the nucleus
• Each energy level can hold only a maximum number
  of electrons

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• MODERN ATOMIC THEORY
• First energy levelmaximum of 2 electrons
• Second energy level maximum of 8 electrons
• Third energy levelmaximum of 18 electrons (three
  sublevels of 2/8/8)

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• Chemical activitydepends on the arrangement of
  electrons in the outermost energy level

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• FORCES WITHIN THE ATOM
• Electromagneticnegative charged electrons are
  attracted to the positive charged protons 
• Strong Forceprevents the positively charged
  protons from repelling each other
• keeps protons together 
• Weak Forceresponsible for radioactive decay
• neutron changes into a proton and an electron 
• Gravity force of attraction that depends on the
  mass of two objects and the distance between them
  

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• ATOMS AND BONDING 
• Before an atom reacts it is electrically neutral
  (same number of protons and electrons)
  electrically stable

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• ATOMS AND BONDING
• However the atom might not be chemically stable
• Chemical stability depends on the valence
  electrons (outermost energy level)
• 1/2/3 electrons will lose electrons
• 5/6/7 electrons will gain electrons
• after losing/gaining electrons the atom will be
  chemically stable but now will become
  electrically unstable

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• ATOMS AND BONDING
• lose 1 electron 1
• lose 2 electrons 2
• lose 3 electrons 3
• gain 3 electrons -3
• gain 2 electrons -2
• gain 1 electron -1
• When the outermost energy level (valence
  electrons) is filled, the atom is chemically
  stable

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• IONIC BONDSinvolves the transfer of electrons
• One atom gains electrons and the other atom loses
  electrons resulting in filled outer energy levels
• Ions are formed (charged atom or group of
  atoms polyatomic)

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Ionic Bonds

• The figure above shows how ionic bonds form
  between sodium and chlorine in table salt
• A sodium atom easily loses its one valence
  electron and becomes a sodium ion (Na)
• A chlorine atom easily gains an electron and
  becomes a chloride ion (Cl-)

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Ionic Bonds
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• IONIC BONDS
• The process of removing electrons and forming
  positive ions (more protons than electrons) is
  called ionization (energy is absorbed)
• Energy is needed for ionizationionization
  energy
• Low for atoms with few valence electrons (metals)
• High for atoms with many valence electrons
  (nonmetals)

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• IONIC BONDS
• The process of gaining electrons and forming
  negative ions (more electrons than protons) is
  called electron affinity (energy is released)
• Low for atoms with few valence electrons (metals)
• High for atoms with many valence electrons
  (nonmetals)

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• IONIC BONDS
• It is much easier to gain 1 or 2 electrons than
  to lose 6 or 7 electrons!!!!
• Positive () ions attract negative (-) ions
  resulting in ionic bonds 

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Ionic Bonds

• The chemical bond in which electrons are
  transferred from one atom to another is called an
  ionic bond
• The compound sodium chloride (NaCl) forms when
  sodium loses its valence electron to chlorine

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Ionic Bonds

• In a salt crystal, there are trillions of sodium
  and chloride ions
• These oppositely charged ions have a strong
  attraction
• The attraction between oppositely charged ions is
  an ionic bond.

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• IONIC BONDS
• The placement of ions in an ionic compound
  results in a regular, repeating arrangement
  called a crystal lattice
• Gives great stability/high melting points
• Chemical formula shows the ratio of ions not the
  actual number present
• Each ionic compound has a characteristic crystal
  lattice arrangement

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• COVALENT BONDS sharing of electrons
• Results in filled outer energy levels of both
  sharing atoms
• The positively charged nucleus of each atom
  simultaneously attracts the negatively charged
  electrons that are being shared

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Covalent Bonds 

• Sometimes electrons are shared by atoms instead
  of being transferred
• What does it mean to share electrons?
• It means that the moving electrons actually
  travel in the orbitals of both atoms
• A covalent bond forms when electrons are shared
  between atoms
• When the atoms share two electrons, the bond is
  called a single covalent bond
• Sometimes the atoms share four electrons and form
  a double bond
• In a few cases, atoms can share six electrons and
  form a triple bond

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Covalent Bonds

• The structure that results when atoms are joined
  together by covalent bonds is called a molecule
• The molecule is the smallest unit of most
  compounds
• The diagram, to the right, of a water molecule
  shows that each hydrogen atom forms a single
  covalent bond with the oxygen atom

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Covalent Bonds

                                                                   
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• COVALENT BONDS
• Combination of atoms formed by a covalent bond
  are called molecules
• Molecule is the smallest particle of a
  covalently bonded substance that has all the
  properties of that substance
• Chemical formula for a molecule shows the exact
  number of atoms of each element involved in the
  bond
• Tend to have low melting points

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• COVALENT BONDS
• Represented by electron dot diagrams
• Chemical symbol represents the nucleus and all
  inner energy levels
• Dots surrounding the symbol represent the valence
  (outermost) electrons

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Covalent Bonds

• Sharing is NOT always equal
• O2 and H2 sharing is equal
• H2O sharing is NOT equal
• Results in the molecule having a slight
  electrical charge
• POLAR

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Van der Waals Forces 

• Because of their structures, atoms of different
  elements do not all have the same ability to
  attract electrons
• Some atoms have a stronger attraction for
  electrons than do other atoms
• Therefore, when the atoms in a covalent bond
  share electrons, the sharing is not always equal
• Even when the sharing is equal, the rapid
  movement of electrons can create regions on a
  molecule that have a tiny positive or negative
  charge

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Van der Waals Forces

• When molecules are close together, a slight
  attraction can develop between the oppositely
  charged regions of nearby molecules
• Chemists call such intermolecular forces of
  attraction van der Waals forces, after the
  scientist who discovered them
• Although van der Waals forces are not as strong
  as ionic bonds or covalent bonds, they can hold
  molecules together, especially when the molecules
  are large

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Van der Waals Forces

• People who keep geckos as pets have already seen
  van der Waals forces in action
• These remarkable little lizards can climb up
  vertical surfaces, even smooth glass walls, and
  then hang on by a single toe despite the pull of
  gravity
• How do they do it?
• No, they do not have some sort of glue on their
  feet and they don't have suction cups
• A gecko foot is covered by as many as half a
  million tiny hairlike projections
• Each projection is further divided into hundreds
  of tiny, flat-surfaced fibers.
• This design allows the gecko's foot to come in
  contact with an extremely large area of the wall
  at the molecular level
• Van der Waals forces form between molecules on
  the surface of the gecko's foot and molecules on
  the surface of the wall permits the animal to
  climb vertical structures
• The combined strength of all the van der Waals
  forces allows the gecko to balance the pull of
  gravity
• When the gecko needs to move its foot, it peels
  the foot off at an angle and reattaches it at
  another location on the wall

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Properties of Water  

• Water is also the single most abundant compound
  in most living things
• Water is one of the few compounds that is a
  liquid at the temperatures found over much of
  Earth's surface
• Unlike most substances, water expands as it
  freezes
• Thus, ice is less dense than liquid water, which
  explains why ice floats on the surface of lakes
  and rivers
• If the ice sank to the bottom, the situation
  would be disastrous for fish and plant life in
  regions with cold winters, to say nothing of the
  sport of ice skating!

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Water Molecule

• Like all molecules, a water molecule (H2O) is
  neutral (BUT conducts electricity!!!!!)
• The positive charges on its 10 protons balance
  out the negative charges on its 10 electrons
• However, there is more to the story

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• WATER
• 70 of earths surface is covered by water
• 65 of your body mass is water
• Thousands of substances dissolve (soluble) in
  water (UNIVERSAL SOLVENT)
• Certain substances will not dissolve in water
  (insoluble)

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Polarity

• With 8 protons in its nucleus, an oxygen atom has
  a much stronger attraction for electrons than
  does the hydrogen atom with a single proton in
  its nucleus
• Thus, at any moment, there is a greater
  probability of finding the shared electrons near
  the oxygen atom than near the hydrogen atom
• Because the water molecule has a bent shape, as
  shown to the right, the oxygen atom is on one end
  of the molecule and the hydrogen atoms are on the
  other
• As a result, the oxygen end of the molecule has a
  slight negative charge and the hydrogen end of
  the molecule has a slight positive charge

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Water Molecule
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• WATER STRUCTURE
• Two hydrogen atoms bond covalently with one
  oxygen atom (electrons are shared)
• Sharing is unequal
• Oxygenslight negative charge
• Hydrogenslight positive charge

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Water Molecule

• The unequal sharing of electrons causes a water
  molecule to be polar
• The hydrogen end of the molecule is slightly
  positive, and the oxygen end is slightly negative
  
• A molecule in which the charges are unevenly
  distributed is called a polar molecule because
  the molecule is like a magnet with poles
• A water molecule is polar because there is an
  uneven distribution of electrons between the
  oxygen and hydrogen atoms
• The negative pole is near the oxygen atom and the
  positive pole is between the hydrogen atoms

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Hydrogen Bonds 

• Because of their partial positive and negative
  charges, polar molecules such as water can
  attract each other, as shown to the right
• The charges on a polar molecule are written in
  parentheses, (-) or (), to show that they are
  weaker than the charges on ions such as Na and
  Cl-
• The attraction between the hydrogen atom on one
  water molecule and the oxygen atom on another
  water molecule is an example of a hydrogen bond
• Hydrogen bonds are not as strong as covalent or
  ionic bonds, but water's ability to form multiple
  hydrogen bonds is responsible for many of its
  special properties

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Hydrogen Bonds
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Cohension

• A single water molecule may be involved in as
  many as four hydrogen bonds at the same time
• The ability of water to form multiple hydrogen
  bonds is responsible for many of water's
  properties
• Cohesion is an attraction between molecules of
  the same substance
• Because of hydrogen bonding, water is extremely
  cohesive
• Water's cohesion causes molecules on the surface
  of water to be drawn inward, which is why drops
  of water form beads on a smooth surface
• Cohesion also explains why some insects and
  spiders can walk on a pond's surface

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ADHESION

• Adhesion is an attraction between molecules of
  different substances
• Have you ever been told to read the volume in a
  graduated cylinder at eye level?
• The surface of the water in the graduated
  cylinder dips slightly in the center because the
  adhesion between water molecules and glass
  molecules is stronger than the cohesion between
  water molecules
• Adhesion between water and glass also causes
  water to rise in a narrow tube against the force
  of gravity
• This effect is called capillary action
• Capillary action is one of the forces that draw
  water out of the roots of a plant and up into its
  stems and leaves
• Cohesion holds the column of water together as it
  rises

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• METALLIC BONDSouter electrons of the atoms form
  a common electron cloud
• The electrons become the property of all the
  atoms
• The positive nuclei of atoms of metals are
  surrounded by free-moving electrons that are all
  attracted by the nuclei at the same time
• Electrons are free to flow
• Excellent conductors of both heat and electricity
• High melting points

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• OXIDATION NUMBER
• Describes the combining capacity of an atom
• Indicates the number of electrons an atom gains,
  loses, or shares when it forms chemical bonds
• Positive number indicates a lose of electrons
• Negative number indicates a gain of electrons
• Number indicates how many electrons
• Used to predict how atoms will combine and what
  the formula for the resulting compound will be
• The sum of the oxidation numbers of the atoms in
  a compound must be ZERO

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Solutions and Suspensions

• Water is not always pureit is often found as
  part of a mixture
• A mixture is a material composed of two or more
  elements or compounds that are physically mixed
  together but not chemically combined
• Salt and pepper stirred together constitute a
  mixture
• So do sugar and sand
• Earth's atmosphere is a mixture of gases
• Living things are in part composed of mixtures
  involving water
• Two types of mixtures that can be made with water
  are solutions and suspensions

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Solutions

• If a crystal of table salt is placed in a glass
  of warm water, sodium and chloride ions on the
  surface of the crystal are attracted to the polar
  water molecules
• Ions break away from the crystal and are
  surrounded by water molecules, as illustrated to
  the right
• The ions gradually become dispersed in the water,
  forming a type of mixture called a solution

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Solutions

• All the components of a solution are evenly
  distributed throughout the solution
• In a saltwater solution, table salt is the
  solutethe substance that is dissolved
• Water is the solventthe substance in which the
  solute dissolves
• Water's polarity gives it the ability to dissolve
  both ionic compounds and other polar molecules,
  such as sugar
• Without exaggeration, water is the greatest
  solvent on Earth.

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Solutions

• NaCl Dissolving in Water
• When an ionic compound such as sodium chloride is
  placed in water, water molecules surround and
  separate the positive and negative ions

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Solutions
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• WATER STRUCTURE
• Molecule has oppositely charged ends
• Charged ends give the property of Polarity
• A force of attraction is set up between the
  solute and solvent
• Separates the molecules of the solute, causing
  the solute to dissolve
• NONPOLAR substances will not dissolve in water
  but will dissolve in nonpolar solvents
• LIKE DISSOLVES LIKE 

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Suspensions

• Some materials do not dissolve when placed in
  water but separate into pieces so small that they
  do not settle out
• The movement of water molecules keeps the small
  particles suspended
• Such mixtures of water and nondissolved material
  are known as suspensions
• Some of the most important biological fluids are
  both solutions and suspensions
• The blood that circulates through your body is
  mostly water, which contains many dissolved
  compounds
• However, blood also contains cells and other
  undissolved particles that remain in suspension
  as the blood moves through the body

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Acids, Bases, and pH

• A water molecule can react to form ions
• This reaction can be summarized by a chemical
  equation in which double arrows are used to show
  that the reaction can occur in either direction
• How often does this happen?
• In pure water, about 1 water molecule in 550
  million reacts and forms ions
• Because the number of positive hydrogen ions
  produced is equal to the number of negative
  hydroxide ions produced, water is neutral

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The pH scale 

• Chemists devised a measurement system called the
  pH scale to indicate the concentration of H ions
  in solution
• As the figure at right shows, the pH scale ranges
  from 0 to 14
• At a pH of 7, the concentration of H ions and
  OH- ions is equal
• Pure water has a pH of 7
• Solutions with a pH below 7 are called acidic
  because they have more H ions than OH- ions
• The lower the pH, the greater the acidity
• Solutions with a pH above 7 are called basic
  because they have more OH- ions than H ions.
• The higher the pH, the more basic the solution.
• Each step on the pH scale represents a factor of
  10
• Example
• liter of a solution with a pH of 4 has 10 times
  as many H ions as a liter of a solution with a
  pH of 5

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• pH SCALE 
• Measure of the hydronium ion ( H3O)
  concentration (hydrogen ion H ion
  concentration)
• Hydronium ion is formed by the attraction between
  a hydrogen ion (H ) from an acid and a water
  molecule( H2O )

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• pH
• Indicates HOW ACIDIC a solution is
• Series of numbers 0 to 14
• middle is 7 (neutral point)
• below 7 (acid)
• above 7 (base)

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The pH scale
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• pH
• 10-7 MOLES OF H3O IONS IN 1 LITER OF H2O
• pH 7
• 0.0000001 moles H3O / liter H2O  

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• pH MOLES H3O / LITER H2O
  pH 1 ( 0.1 moles/liter ) ( 10 -1 moles/l)
  pH 2 ( 0.01 moles/liter ) (
  10 -2 moles/l) pH 3 (
  0.001 moles/liter ) ( 10 -3 moles/l)
  pH 4 ( 0.0001 moles/liter ) ( 10 -4
  moles/l) pH 5 ( 0.00001
  moles/liter ) ( 10 -5 moles/l) pH
  6 ( 0.000001 moles/liter ) ( 10 -6 moles/l)
  pH 7 ( 0.0000001 moles/liter) ( 10
  -7 moles/l) pH 8 ( 0.00000001
  moles/liter ) ( 10 -8 moles/l) pH 9 (
  0.000000001 moles/liter ) ( 10 -9 moles/l)
  pH 10 ( 0.0000000001 moles/liter ) ( 10 -10
  moles/l) pH 11 ( 0.00000000001 moles/liter )
  ( 10 -11 moles/l) pH 12 ( 0.000000000001
  moles/liter ) ( 10 -12 moles/l) pH 13 (
  0.0000000000001 moles/liter )( 10 -13 moles/l) pH
  14 ( 0.00000000000001 moles/liter )( 10
  -14moles/l)

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Acids

• Where do all those extra H ions in a low-pH
  solution come from?
• They come from acids
• An acid is any compound that forms H ions in
  solution
• Acidic solutions contain higher concentrations of
  H ions than pure water and have pH values below
  7
• Strong acids tend to have pH values that range
  from 1 to 3
• The hydrochloric acid produced by the stomach to
  help digest food is a strong acid.

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• ACIDS
• Physical propertysour taste (never taste in lab)
• Affects color of indicators
• compounds that show a definite color change
  when mixed with an acid or a base)
• litmus paperred
• phenolphthaleinclear (colorless)

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• ACIDS
• React with active metals to produce hydrogen gas
  and a metal compound (corrodes the metal and
  produces a residue)
• Lab
• Car battery (danger)

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• ACIDS
• ALL contain HYDROGEN
• When dissolved in water, acids ionize to produce
  positive () hydrogen ions (H)
• Hydrogen ion is a PROTON
• Acids are defined as proton producers
• The attraction between a water (H2O) molecule and
  a hydrogen ion (H ) results in the formation of
  a hydronium ion (H3O)

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• ACIDS
• PROTON DONOR

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• STRONG ACIDS
• Ionize to a high degree in water and
  produce hydrogen ions
• Strong electrolytes

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• WEAK ACIDS
• Do not ionize to a high degree in water
• Produce few hydrogen ions
• Poor electrolytes
• Good BUFFERS

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ACIDS

• H2SO4
• HCl
• HNO3

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Bases

• A base is a compound that produces hydroxide ions
  (OH- ions) in solution
• Basic, or alkaline, solutions contain lower
  concentrations of H ions than pure water and
  have pH values above 7
• Strong bases, such as lye, tend to have pH values
  ranging from 11 to 14.

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• BASES
• Physical propertybitter taste (never
  taste in lab) and slippery to the touch
• Can be poisonous and corrosive

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• BASES
• Affect color of indicators ( compounds
  that show a definite color change when mixed with
  an acid or a base)
• Litmus paperblue
• Phenolphthaleinpink

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• BASES
• Emulsify, or dissolve fats and oils
• React with the fat or oil to form a soap

99

• BASES
• ALL contain the HYDROXIDE ION (OH-)
• Since the hydroxide ion ( OH-) can combine with a
  hydrogen ion ( H) and form water, a base is
  often called a PROTON ACCEPTOR

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• STRONG BASES
• Ionize to a high degree in water and produce
  large number of ions
• Good electrolytes

101

• WEAK BASES
• Do not ionize to a high degree in water
• Produce few ions
• Poor electrolytes
• Good BUFFERS

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BASES

• NaOH
• LiOH
• Ca(OH)2
• Ba(OH)2
• Al(OH)3

103
Buffers

• The pH of the fluids within most cells in the
  human body must generally be kept between 6.5 and
  7.5
• If the pH is lower or higher, it will affect the
  chemical reactions that take place within the
  cells
• Thus, controlling pH is important for maintaining
  homeostasis
• One of the ways that the body controls pH is
  through dissolved compounds called buffers
• Buffers are weak acids or bases that can react
  with strong acids or bases to prevent sharp,
  sudden changes in pH

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ACID-BASE BALANCE

• Because of the abundance of hydrogen bonds in the
  bodys functional proteins (enzymes, hemoglobin,
  cytochromes, and others) they are strongly
  influenced by hydrogen ion concentration
• It follows then that nearly all biochemical
  reactions are influenced by the pH of their fluid
  environment, and the acid-base balance of body
  fluids is closely regulated
• Optimal pH varies from one body fluid to another
  
• When arterial blood pH rises above 7.45, the body
  is in alkalosis (alkalemia) when arterial pH
  falls below 7.35, the body is in acidosis
  (acidemia)
• Between 7.0 and 7.35 is called physiological
  acidosis even though the value is slightly basic
• Most hydrogen ions originate as metabolic by
  products, although they can also enter the body
  via ingested foods

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ACID-BASE BALANCE

• Dissociation of strong and weak acids
• (a) when added to water, the strong acid HCl
  dissociates completely into its ions (H and Cl-)
• (b) dissociation of H2CO3, a weak acid, is very
  incomplete, and some molecules of H2CO3 remain
  undissociated in solution

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COMPARISON OF DISSOCIATION OF STRONG AND WEAK
ACIDS
107
Chemical Buffer System Bicarbonate Buffer Systems

• A chemical buffer is a system of one or two
  molecules that acts to resist changes in pH by
  binding H when the pH drops, or releasing H
  when the pH rises
• The bicarbonate buffer system is the main buffer
  of the extracellular fluid, and consists of
  carbonic acid and its salt, sodium bicarbonate
• When a strong acid is added to the solution,
  carbonic acid is mostly unchanged, but
  bicarbonate ions of the salt bind excess H,
  forming more carbonic acid
• HCl NaHCO3 ? H2CO3
  NaCl
• Strong acid weak base ? weak acid salt
• pH lowered slightly
• When a strong base is added to solution, the
  sodium bicarbonate remains relatively unaffected,
  but carbonic acid dissociates further, donating
  more H to bind the excess hydroxide
• NaOH H2CO3 ? NaHCO3
  H2O
• Strong base weak acid ? weak base water
• pH rises very little
• Bicarbonates buffer system sodium, potassium,
  and magnesium

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Chemical Buffer System Phosphate Buffer System

• The phosphate buffer system operates in the urine
  and intracellular fluid similar to the
  bicarbonate buffer system
• The components of the phosphate system are the
• Sodium salts of dihydrogen phosphate (H2PO4-)
• Sodium salts of monohydrogen phosphate (HPO42-)
• NaH2PO4 acts as a weak acid
• HCl Na2HPO4 ? NaH2PO4
  NaCl
• Strond acid weak base ? weak acid
  salt
• H released by strong acids is tied up in weak
  acids
• NaOH NaH2PO4 ? Na2HPO4
  H2O
• Strong base weak acid ? weak base
  water
• Strong bases are converted to weak bases

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Chemical Buffer SystemThe Protein Buffer System

• Proteins in plasma and in cells are the bodys
  most plentiful and powerful source of buffers,
  and constitute the protein buffer system
• At least ¾ of all the buffering power of body
  fluids resides in cells, and most of this
  reflects the buffering activity of intracellular
  proteins
• Proteins are polymers of amino acids
• Consists of organic acids containing carboxyl
  groups that dissociate to
• Release H when the pH begins to rise
• RCOOH ? RCOO- H
• Bind excess H when the pH declines
• RCOO- H ? RCOOH
• Consists of an amide group that can act as a base
  and accept H
• The exposed NH2 group can bind with hydrogen
  ions, becoming NH3-
• RNH2 H ? RNH3
• Because this removes free hydrogen ions from the
  solution, it prevents the solution from becoming
  too acidic
• Consequently, a single protein molecule can
  function reversibly as either an acid or a base
  depending on the pH of its environment
• Molecules with this ability are called amphoteric
  molecules
• Example hemoglobin

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Respiratory Regulation of H

• Carbon dioxide from cellular metabolism enters
  erythrocytes and is converted to bicarbonate ions
  for transport in the plasma
• carbonic
• anhydrase
  
• CO2 H2O ? H2CO3 ? H
  HCO3-
• carbonic acid
  bicarbonate ion
• When hypercapnia (increased amount of carbon
  dioxide in the blood) occurs, blood pH drops,
  activating medullary respiratory centers,
  resulting in increased rate and depth of
  breathing and increased unloading of CO2 in the
  lungs
• The reaction is pushed to the right
• A rising plasma H concentration resulting from
  any metabolic process excites the respiratory
  center indirectly (peripheral chemoreceptors) to
  stimulate deeper, more rapid respiration
• As ventilation increases, more CO2 is removed
  from the blood, pushing the reaction to the left
  and reducing the H concentration

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Carbon Compounds

• Until the early 1800s, many chemists thought that
  compounds created by organismsorganic
  compoundswere distinctly different from
  compounds in nonliving things
• In 1828, a German chemist was able to synthesize
  the organic compound urea from a mineral called
  ammonium cyanate
• Chemists soon realized that the principles
  governing the chemistry of nonliving things could
  be applied to living things
• Scientists still use the term organic chemistry,
  but now it describes something a little different
• Today, organic chemistry is the study of all
  compounds that contain bonds between carbon atoms

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The Chemistry of Carbon

• Carbon atoms have four valence electrons
• Each electron can join with an electron from
  another atom to form a strong covalent bond
• Carbon can bond with many elements, including
  hydrogen, oxygen, phosphorus, sulfur, and
  nitrogen

113
The Chemistry of Carbon

• Carbon atoms can bond to other carbon atoms,
  which gives carbon the ability to form chains
  that are almost unlimited in length
• These carbon-carbon bonds can be single, double,
  or triple covalent bonds
• Chains of carbon atoms can even close upon
  themselves to form rings, as shown below
• Carbon has the ability to form millions of
  different large and complex structures
• No other element even comes close to matching
  carbon's versatility

114
The Chemistry of Carbon
115

• CARBON
• 90 of all known compounds contain carbon
• Forms an important family of compounds called
  ORGANIC COMPOUNDS
• Forms covalent bonds (single, double, triple)
  with other carbon atoms
• Straight chains, branched chains, single rings,
  or joined rings
• Single bond2 electrons
• Double bond4 electrons
• Triple bond6 electrons

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CARBON

• A carbon atom has 6 protons
• Therefore, 6 electrons
• 2 electrons in the 1st energy level
• 4 electrons in the 2nd energy level
• Can form 4 covalent bonds

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Carbon Compounds

• Bonding diagram shows how the electrons in the
  outer energy level form covalent bonds
• Molecular formula tells the number of each kind
  of atom in a molecule of the compound
• Structural formula shows the bonds connecting
  the atoms and the arrangement of the atoms within
  each molecule
• Space-filling model shows how the atoms in the
  molecule are arranged in space (notice that the
  methane molecule is not flat but is shaped like a
  pyramid)

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• CARBON
• STRUCTURAL FORMULAS
• Shows the kind, number, and arrangement
  of atoms in a molecule
• A dash ( ) is used to represent the pair of
  shared electrons forming the covalent bond

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• CARBON
• ISOMERS
• Compounds with the same molecular formula but
  different structures
• Can have different physical and chemical
  properties
• As the number of carbon atoms increases, the
  number of isomers increases 

124

• HYDROCARBONS contain only hydrogen and carbon
• Most abundant source is Petroleum
• SATURATED all bonds between carbon atoms are
  single covalent bonds 
• UNSATURATED one or more of the bonds between
  carbon atoms is a double covalent or triple
  covalent bond

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Macromolecules

• Many of the molecules in living cells are so
  large that they are known as macromolecules,
  which means giant molecules
• Macromolecules are made from thousands or even
  hundreds of thousands of smaller molecules

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Macromolecules

• Macromolecules are formed by a process known as
  polymerization (pah-lih-mur-ih-ZAY-shun), in
  which large compounds are built by joining
  smaller ones together
• The smaller units, or monomers, join together to
  form polymers
• The monomers in a polymer may be identical, like
  the links on a metal watch band or the monomers
  may be different, like the beads in a
  multicolored necklace
• The figure below illustrates the formation of a
  polymer from more than one type of monomer

127
Macromolecules
128
Macromolecules

• Four groups of organic compounds found in living
  things are
• Carbohydrates
• Lipids
• Nucleic acids
• Proteins
• Sometimes these organic compounds are referred to
  as biomolecules

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Carbohydrates

• Compounds made up of carbon, hydrogen, and oxygen
  atoms, usually in a ratio of 1 2 1
• Living things use carbohydrates as their main
  source of energy
• Plants and some animals also use carbohydrates
  for structural purposes
• The breakdown of sugars, such as glucose,
  supplies immediate energy for all cell activities
• Living things store extra sugar as complex
  carbohydrates known as starches
• The monomers in starch polymers are sugar
  molecules

130
CARBOHYDRATES

• Composed of carbon, hydrogen, and oxygen
• Generalized formula C x H 2 O
• Types
• Monosaccharides
• Disaccharides
• Polysaccharides

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CARBOHYDRATES

• Monosaccharides
• Simple sugars
• Ratio C H 2 O
• Most common (glucose, fructose, galactose) are
  isomers
• All have the same chemical molecular formula ( C
  6 H 12 O 6 ) but different structural formulas
• Glucose (dextrose)
• Produced by plants (photosynthesis)
• Main source of energy in plants and animals
• Metabolized in cellular respiration releasing
  energy
• Fructose
• Found in fruits
• Sweetest of the monosaccharides
• Galactose
• Found in milk
• Usually in combination with glucose and fructose
  making disaccharides

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CARBOHYDRATES

• Disaccharides
• Double sugar
• Combination of two monosaccharides
• Formed by the chemical linking of two
  monosaccharides in a condensation reaction
  (dehydration synthesis)
• Examples
• Sucrose
• Found in sugarcane and sugar beets
• Composed of fructose and glucose
• Lactose
• Found in milk
• Composed of glucose and galactose
• Maltose
• Malt sugar

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CARBOHYDRATES

• Polysaccharides
• Complex molecule composed of three or more
  monosaccharides
• Formed by the chemical linking of three or more
  monosaccharides in condensation reactions
  (dehydration synthesis)
• Examples
• Starch
• Storage form of glucose in plants
• Two basic forms
• Long unbranched chains that coil like a telephone
  cord
• Highly branched like glycogen
• Glycogen
• Storage form of glucose in animals
• Called animal starch
• Composed of hundreds of glucose molecules in a
  highly branched chain
• Cellulose
• Gives strength and rigidity to the plant cell
• Thousands of glucose monomers are linked in long,
  straight chains

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CARBOHYDRATES

• The large macromolecules formed from
  monosaccharides are known as polysaccharides
• Many animals store excess sugar in a
  polysaccharide called glycogen, or animal starch
• When the level of glucose in your blood runs low,
  glycogen is released from your liver
• Must be broken down to glucose before the body
  can utilize the energy stored
• Hydrolysis splitting by the addition of water
• The glycogen stored in your muscles supplies the
  energy for muscle contraction and, thus, for
  movement
• Must be broken down to glucose before the body
  can utilize the energy stored
• Hydrolysis splitting by the addition of water

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CARBOHYDRATES

• Plants use a slightly different polysaccharide,
  called plant starch, to store excess sugar
• Must be broken down to glucose before the body
  can utilize the energy stored
• Hydrolysis splitting by the addition of water
• Plants also make another important polysaccharide
  called cellulose
• Tough, flexible cellulose fibers give plants much
  of their strength and rigidity
• Cellulose is the major component of both wood and
  paper, so you are actually looking at cellulose
  when you are reading a textbook

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Lipids

• Lipids are a large and varied group of biological
  molecules that are generally not soluble in water
• Lipids are made mostly from carbon and hydrogen
  atoms
• The common categories of lipids are fats, oils,
  and waxes
• Lipids can be used to store energy
• Some lipids are important parts of biological
  membranes and waterproof coverings
• Steroids are lipids as well
• Many steroids serve as chemical messengers

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LIPIDS

• Fatty compound made up of a large number of
  carbon and hydrogen atoms but a smaller number of
  oxygen atoms
• Fats, oil, waxes, triglycerides, steroids
• Not soluble in water (insoluble)
• Major component of cell (plasma) membrane forming
  a barrier between the internal and external
  aqueous environments
• Store energy efficiently
• Large number of carbon-hydrogen bonds that store
  more energy than carbon-oxygen bonds

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LIPIDS

• Fatty acids
• Monomers that make of lipids
• Long-straight hydrocarbon chain with a carboxyl
  (COOH) group attached at one end
• Carboxyl end is polar which attracts water which
  is polar (hydrophilic)
• Hydrogen end is nonpolar which tends to repel
  water (hydrophobic)
• Cell membrane the hydrophilic ends are oriented
  to the aqueous side and the hydrophobic ends are
  oriented to the center

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LIPIDS

• Fatty acids
• Saturated
• All single bonds between carbon atoms
• No double bonds
• Maximum possible number of hydrogen atoms bonded
  to each carbon atom
• Molecule is saturated with hydrogen
• Unsaturated
• One or more double bonds between carbon atoms
• Fewer hydrogen atoms (not saturated)

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Lipids

• Many lipids are formed when a glycerol molecule
  combines with compounds called fatty acids, as
  shown below
• If each carbon atom in a lipid's fatty acid
  chains is joined to another carbon atom by a
  single bond, the lipid is said to be saturated
• The term saturated is used because the fatty
  acids contain the maximum possible number of
  hydrogen atoms
• The lipid represented has double bonds
  unsaturated

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Lipids

• If there is at least one carbon-carbon double
  bond in a fatty acid, the fatty acid is said to
  be unsaturated
• Lipids whose fatty acids contain more than one
  double bond are said to be polyunsaturated
• If the terms saturated and polyunsaturated seem
  familiar, you have probably seen them on food
  package labels
• Lipids such as olive oil, which contains
  unsaturated fatty acids, tend to be liquid at
  room temperature
• Cooking oils, such as corn oil, sesame oil,
  canola oil, and peanut oil, contain
  polyunsaturated lipids

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Lipids
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LIPIDS

• Triglycerides
• Lipid in which the macromolecule is composed of
  three molecules of fatty acids joined by chemical
  condensation reactions (dehydration synthesis) to
  one molecule of glycerol
• Two main types
• Oils liquid triglyceride at room temperature
• Found mainly in plants (seeds)
• Source of stored energy
• Fats solid triglycerides at room temperature
• Found mainly in animals
• Source of stored energy

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LIPIDS

• Wax
• Consists of long fatty acid chain joined to a
  long alcohol chain
• Highly waterproof
• In plants, forms a protective covering on the
  outer surfaces
• In animals, forms protective layer
• E.g. earwax barrier that keeps microorganisms
  from entering the middle ear

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LIPIDS

• Steroid
• Composed of four carbon rings
• No fatty acids
• Considered a lipid because they do not dissolve
  in water (insoluble in water)
• Some hormones, nerve tissue, toad venoms, and
  plant poisons

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Nucleic Acids

• Macromolecules containing hydrogen, oxygen,
  nitrogen, carbon, and phosphorus
• Nucleic acids are polymers assembled from
  individual monomers known as nucleotides
• Nucleotides consist of three parts a 5-carbon
  sugar, a phosphate group, and a nitrogenous base,
  as shown in the figure below
• Individual nucleotides c