The Chemistry of Life The Nature of Matter Life depends on chemistry When you eat food or inhale oxygen, your body uses these materials in chemical reactions.

The Chemistry of Life

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!!!!!!!!!

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

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

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

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

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!!!!!!

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

ATOM –IS THE FUNDAMENTAL UNIT OF MATTER –COMPOSED OF SUBATOMIC PARTICLES ATOMIC NUCLEUS –PROTON –NEUTRON –ELECTRONS:ORBIT (OUTSIDE) THE NUCLEUS

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

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).

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

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

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

Isotopes of Carbon

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

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 H 2 O –The formula for table salt, NaCl, indicates that the elements from which table salt forms—sodium and chlorine—combine in a 1 : 1 ratio

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

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

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

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

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

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

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

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

Chemical activity—depends on the arrangement of electrons in the outermost energy level

FORCES WITHIN THE ATOM –Electromagnetic—negative charged electrons are attracted to the positive charged protons –Strong Force—prevents the positively charged protons from repelling each other keeps protons together –Weak Force—responsible 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

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

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

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

IONIC BONDS—involves 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)

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−)

Ionic Bonds

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 ionization— ionization energy: Low for atoms with few valence electrons (metals) High for atoms with many valence electrons (nonmetals)

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)

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

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

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.

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

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

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

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

Covalent Bonds

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

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

Covalent Bonds Sharing is NOT always equal O 2 and H 2 sharing is equal H 2 O sharing is NOT equal –Results in the molecule having a slight electrical charge –POLAR

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

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

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

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!

Water Molecule Like all molecules, a water molecule (H 2 O) 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

WATER: –70% of earth’s 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)

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

Water Molecule

WATER STRUCTURE: –Two hydrogen atoms bond covalently with one oxygen atom (electrons are shared) Sharing is unequal –Oxygen—slight negative charge –Hydrogen—slight positive charge

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

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

Hydrogen Bonds

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

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

METALLIC BONDS—outer 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

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

Solutions and Suspensions Water is not always pure—it 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

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

Solutions All the components of a solution are evenly distributed throughout the solution In a salt–water solution, table salt is the solute—the substance that is dissolved Water is the solvent—the 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.

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

Solutions

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

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

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

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

pH SCALE –Measure of the hydronium ion ( H 3 O + ) 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( H 2 O )

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

The pH scale

pH –10 -7 MOLES OF H 3 O + IONS IN 1 LITER OF H 2 O –pH 7 –0.0000001 moles H 3 O + / liter H 2 O

pH: MOLES H 3 O + / LITER H 2 O 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 -14 moles/l)

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.

ACIDS: –Physical property—sour 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 paper—red –phenolphthalein—clear (colorless)

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

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 (H 2 O) molecule and a hydrogen ion (H + ) results in the formation of a hydronium ion (H 3 O + )

ACIDS: –PROTON DONOR

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

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

ACIDS H 2 SO 4 HCl HNO 3

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.

BASES: –Physical property—bitter taste (never taste in lab) and slippery to the touch –Can be poisonous and corrosive

BASES: –Affect color of indicators ( compounds that show a definite color change when mixed with an acid or a base) Litmus paper—blue Phenolphthalein—pink

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

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

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

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

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

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

ACID-BASE BALANCE Because of the abundance of hydrogen bonds in the body’s 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

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 H 2 CO 3, a weak acid, is very incomplete, and some molecules of H 2 CO 3 remain undissociated in solution

COMPARISON OF DISSOCIATION OF STRONG AND WEAK ACIDS

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 + NaHCO 3 → H 2 CO 3 + 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 + H 2 CO 3 → NaHCO 3 + H 2 O Strong base + weak acid → weak base + water pH rises very little –Bicarbonates buffer system: sodium, potassium, and magnesium

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 (H 2 PO 4 - ) –Sodium salts of monohydrogen phosphate (HPO 4 2- ) –NaH 2 PO 4 acts as a weak acid – HCl + Na 2 HPO 4 → NaH 2 PO 4 + NaCl Strond acid + weak base → weak acid + salt –H + released by strong acids is tied up in weak acids NaOH + NaH 2 PO 4 → Na 2 HPO 4 + H 2 O Strong base weak acid → weak base + water –Strong bases are converted to weak bases

Chemical Buffer System The Protein Buffer System Proteins in plasma and in cells are the body’s 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 –R—COOH → R—COO - + H + Bind excess H + when the pH declines –R—COO - + H + → R—COOH –Consists of an amide group that can act as a base and accept H +: The exposed –NH 2 group can bind with hydrogen ions, becoming –NH 3 - –R—NH 2 + H + → R—NH 3 + »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

Respiratory Regulation of H + Carbon dioxide from cellular metabolism enters erythrocytes and is converted to bicarbonate ions for transport in the plasma: – carbonic – anhydrase –CO 2 + H 2 O ↔ H 2 CO 3 ↔ H + + HCO 3 - – 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 CO 2 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 CO 2 is removed from the blood, pushing the reaction to the left and reducing the H + concentration

Carbon Compounds Until the early 1800s, many chemists thought that compounds created by organisms—organic compounds—were 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

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

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

The Chemistry of Carbon

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 bond—2 electrons Double bond—4 electrons Triple bond—6 electrons

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

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)

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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)

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

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

Lipids

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

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

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

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 can be joined by covalent bonds to form a polynucleotide, or nucleic acid

Nucleotide

Nucleic Acids The monomers that make up a nucleic acid are nucleotides Each nucleotide has a 5-carbon sugar, a phosphate group, and a nitrogenous base Nucleic acids store and transmit hereditary, or genetic, information There are two kinds of nucleic acids: –Ribonucleic acid (RNA): contains the sugar ribose –Deoxyribonucleic acid (DNA): contains the sugar deoxyribose

NUCLEIC ACIDS Complex organic molecules that store important information in the cell Composed of thousands of monomers called nucleotides –Three components: phosphate group, five carbon sugar, and a ring-shaped nitrogen base Two types: –DNA (Deoxyribonucleic Acid) Stores information that is essential for almost all cell activities Replicated in cell division –RNA (Ribonucleic Acid) Stores and transfers information that is essential for the manufacturing of proteins

Proteins Macromolecules that contain nitrogen as well as carbon, hydrogen, and oxygen Proteins are polymers of molecules called amino acids Amino acids are compounds with an amino group (−NH2) on one end and a carboxyl group (−COOH) on the other end

PROTEINS Organic compounds composed mainly of hydrogen, oxygen, carbon, and nitrogen Formed from the linkage of monomers called amino acids in a chemical reaction of condensation (dehydration synthesis) Structural and functional compounds of importance in cells (plants and animals)

PROTEINS Amino acids –20 different types –Monomers that form proteins –Share the same basic structure –Each amino acid contains a central carbon atom to which four other atoms or groups of atoms bond covalently A single hydrogen atom bonds at one site A carboxyl group (COOH) bonds at a second site A amine group (NH 2 ) bonds at the third site A “R” group bonds at the fourth site –The difference between amino acids results from different R groups

Proteins The figure below shows one reason why proteins are among the most diverse macromolecules More than 20 different amino acids are found in nature All amino acids are identical in the regions where they may be joined together by covalent bonds This uniformity allows any amino acid to be joined to any other amino acid — by bonding an amino group to a carboxyl group

Proteins

The portion of each amino acid that is different is a side chain called an R-group –Some R-groups are acidic and some are basic –Some are polar and some are nonpolar –Some contain carbon rings The instructions for arranging amino acids into many different proteins are stored in DNA Each protein has a specific role Some proteins control the rate of reactions and regulate cell processes Some are used to form bones and muscles Others transport substances into or out of cells or help to fight disease

PROTEIN Dipeptides: –Two amino acids chemically bonding together in a condensation reaction (dehydration synthesis) The amino group of one amino acid releases a hydrogen ion (H + ) and the carboxyl group of the second amino acid releases a hydroxide ion (OH - ) –Producing HOH (H 2 O) The nitrogen atom from the amine group and the carbon atom from the carboxyl group bond covalently Covalent bond between the amine group of one amino acid and the carboxyl group of another amino acids forms a peptide bond

Proteins Proteins can have up to four levels of organization –First: is the sequence of amino acids in a protein chain –Second: the amino acids within a chain can be twisted or folded –Third: the chain itself is folded –Fourth: If a protein has more than one chain Each chain has a specific arrangement in space as shown by the red and blue structures in the figure at right Van der Waals forces and hydrogen bonds help maintain a protein's shape In the next section, you will learn why a protein's shape is so important

Proteins

Chemical Reactions and Enzymes Chemical Reactions: –Process that changes one set of chemicals into another set of chemicals –Some chemical reactions occur slowly, such as the combination of iron and oxygen to form an iron oxide called rust –Other reactions occur quickly –When hydrogen gas is ignited in the presence of oxygen, the reaction is rapid and explosive –The elements or compounds that enter into a chemical reaction are known as reactants –The elements or compounds produced by a chemical reaction are known as products. –Chemical reactions always involve the breaking of bonds in reactants and the formation of new bonds in products

Chemical Reactions One example of an important chemical reaction that occurs in your body involves carbon dioxide Your cells constantly produce carbon dioxide as a normal part of their activity This carbon dioxide is carried to your lungs through the bloodstream, and then is eliminated as you exhale However, carbon dioxide is not very soluble in water. The bloodstream could not possibly dissolve enough carbon dioxide to carry it away from your tissues were it not for a chemical reaction As it enters the blood, carbon dioxide reacts with water to produce a highly soluble compound called carbonic acid, H 2 CO 3

Chemical Reactions CO 2 + H 2 O → H 2 CO 3 The reaction shown above enables the bloodstream to carry carbon dioxide to the lungs In the lungs, the reaction is reversed: H 2 CO 3 → CO 2 + H 2 O –This reverse reaction produces carbon dioxide gas, which is released as you exhale

Energy in Reactions Energy is released or absorbed whenever chemical bonds form or are broken Because chemical reactions involve breaking and forming bonds, they involve changes in energy

Energy Changes Some chemical reactions release energy (exothermic), and other reactions absorb energy (endothermic) –Energy changes are one of the most important factors in determining whether a chemical reaction will occur Chemical reactions that release energy often occur spontaneously Chemical reactions that absorb energy will not occur without a source of energy An example of an energy-releasing reaction is hydrogen gas burning, or reacting, with oxygen to produce water vapor : –2H 2 + O 2 → 2H 2 O The energy is released in the form of heat, and sometimes—when hydrogen gas explodes—light and sound

Energy Changes 2H 2 O → 2H 2 + O 2 The reverse reaction, in which water is changed into hydrogen and oxygen gas, absorbs so much energy that it generally doesn't occur by itself In fact, the only practical way to reverse the reaction is to pass an electrical current through water to decompose water into hydrogen gas and oxygen gas –Thus, in one direction the reaction produces energy, and in the other direction the reaction requires energy

Energy Changes What significance do these energy changes have for living things? –In order to stay alive, organisms need to carry out reactions that require energy Thus, every organism must have a source of energy to carry out chemical reactions Plants get their energy by trapping and storing the energy from sunlight in energy-rich compounds (Photosynthesis) Animals get their energy when they consume plants or other animals (Respiration) Humans release the energy needed to grow tall, to breathe, to think, and even to dream through the chemical reactions that occur when humans metabolize, or break down, digested food

Activation Energy Even chemical reactions that release energy do not always occur spontaneously –That's a good thing because if they did, the pages in textbooks might burst into flames The cellulose in paper burns in the presence of oxygen and releases heat and light –However, the cellulose will burn only if you light it with a match, which supplies enough energy to get the reaction started Chemists call the energy that is needed to get a reaction started the activation energy As the figure to the right shows, activation energy is a factor in whether the overall chemical reaction releases energy or absorbs energy

Activation Energy

Chemical reactions that release energy often occur spontaneously Chemical reactions that absorb energy will occur only with a source of energy The peak of each graph represents the energy needed for the reaction to go forward The difference between this required energy and the energy of the reactants is the activation energy

Enzymes Some chemical reactions that make life possible are too slow or have activation energies that are too high to make them practical for living tissue These chemical reactions are made possible by a process that would make any chemist proud—cells make catalysts –A catalyst is a substance that speeds up the rate of a chemical reaction –Catalysts work by lowering a reaction's activation energy

PROTEIN Enzymes –Organic chemical catalyst Speeds up chemical reactions without being affected (consumed) by the reactions themselves –Proteins that act as catalysts in intermediary metabolism –Essential for the functioning of any cell –Speed up a chemical reaction by using an alternate way to change reactant to product Reaction catalyzed by the enzyme has a lower activation energy than the reaction without the enzyme

ENZYMES Proteins that act as biological catalysts Enzymes speed up chemical reactions that take place in cells Like other catalysts, enzymes act by lowering the activation energies, as illustrated by the graph to the right

ENZYMES

Lowering the activation energy has a dramatic effect on how quickly the reaction is completed How big an effect does it have? –Consider the reaction in which carbon dioxide combines with water to produce carbonic acid: CO 2 + H 2 O → H 2 CO 3

Effect of Enzymes Enzymes speed up chemical reactions that take place in cells Notice how the addition of an enzyme lowers the activation energy in this reaction –This action speeds up the reaction: –CO 2 + H 2 O → H 2 CO 3 Left to itself, this reaction is so slow that carbon dioxide might build up in the body faster than the bloodstream could remove it Your bloodstream contains an enzyme called carbonic anhydrase that speeds up the reaction by a factor of 10 million

Effect of Enzymes With carbonic anhydrase on the job, the reaction takes place immediately and carbon dioxide is removed from the blood quickly Enzymes are very specific, generally catalyzing only one chemical reaction –For this reason, part of an enzyme's name is usually derived from the reaction it catalyzes –Carbonic anhydrase gets its name because it catalyzes the reaction that removes water from carbonic acid

Enzyme Action How do enzymes do their jobs? For a chemical reaction to take place, the reactants must collide with enough energy so that existing bonds will be broken and new bonds will be formed –If the reactants do not have enough energy, they will be unchanged after the collision

PROTEIN Enzyme –Shape of a particular enzyme allows it to hook up with a specific molecule (substrate) (reactant) Linkage of the enzyme and the substrate weakens some chemical bonds in the substrate lowering the activation energy producing the products which the enzymes releases Enzyme is unaltered and able to combine with another substrate continuing the chemical reactions

The Enzyme-Substrate Complex Enzymes provide a site where reactants can be brought together to react Such a site reduces the energy needed for reaction The reactants of enzyme-catalyzed reactions are known as substrates

The Enzyme-Substrate Complex The figure at right provides an example of an enzyme-catalyzed reaction The enzyme is hexokinase The substrates are glucose and ATP During the reaction, a phosphate group is transferred from ATP to the glucose molecule Recall that each protein has a specific, complex shape The substrates bind to a site on the enzyme called the active site The active site and the substrates have complementary shapes The fit is so precise that the active site and substrates are often compared to a lock and key

The Enzyme-Substrate Complex

The figure to the right shows a substrate fitting into an active site on an enzyme The enzyme and substrate are bound together by intermolecular forces and form an enzyme-substrate complex They remain bound together until the reaction is done Once the reaction is over, the products of the reaction are released and the enzyme is free to start the process again

The Enzyme-Substrate Complex

Regulation of Enzyme Activity Because they are catalysts for reactions, enzymes can be affected by any variable that influences a chemical reaction Enzymes, including those that help digest food, work best at certain pH values Many enzymes are affected by changes in temperature Not surprisingly, those enzymes produced by human cells generally work best at temperatures close to 37°C, the normal temperature of the human body

Regulation of Enzyme Activity Cells can regulate the activities of enzymes in a variety of ways Most cells contain proteins that help to turn key enzymes “on” or “off” at critical stages in the life of the cell Enzymes play essential roles in regulating chemical pathways, making materials that cells need, releasing energy, and transferring information

The Chemistry of Life The Nature of Matter Life depends on chemistry When you eat food or inhale oxygen, your body uses these materials in chemical reactions.

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The Chemistry of Life The Nature of Matter Life depends on chemistry When you eat food or inhale oxygen, your body uses these materials in chemical reactions.

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Presentation on theme: "The Chemistry of Life The Nature of Matter Life depends on chemistry When you eat food or inhale oxygen, your body uses these materials in chemical reactions."— Presentation transcript:

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