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The Chemistry of Life Chapter 2 Why should we study chemistry in 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. • Wouldn’t you want an architect to understand building materials? Same idea applies to geneticists, ecologists, zoologists, botanists, biologists, and etc. • The study of chemistry begins with the basic unit of matter…the History • Greeks were first to try to explain chemical reactions • 400 BC: thought all matter composed of: – – – – Fire Earth Water Air • Democritus first used word “atomos”, meaning indivisible • Atoms are composed of 3 main particles: (subatomic particles) – Protons (+) – Neutrons – Electrons (-) Protons and Neutrons • Strong forces bind protons and neutrons together to form the nucleus, which is at the center of the atom. • Both particles have about the same mass. Electrons • Electrons are negatively charged with about 1/1840 the mass of a proton. • They are in constant motion in the space surrounding the nucleus. • Atoms have equal numbers of electrons and protons. • Because these subatomic particles have equal but opposite charges, atoms are neutral. Elements • Elements are the building blocks of all matter. • Elements cannot be decomposed into simpler matter. Group Number: 1 2 3 4 5 6 7 8 The Elements • 110 known elements • 88 occur naturally The 110 elements form a plethora of compounds, just as 26 letters of the alphabet make a seemingly endless number of words. Atomic Number Counts the number of protons in an atom Atomic Number on the Periodic Table Atomic Number Symbol 11 Na All atoms of an element have the same number of protons 11 protons Sodium 11 Na Atomic Mass • Mass of an atom. • Approximately equal to the number of protons and neutrons • Find number of neutrons by subtracting the number of protons from the mass. Review: • An element's atomic number tells how many protons are in its atoms. • An element's mass number tells how many protons and neutrons are in its atoms. Learning Check 1 State the number of protons for atoms of each of the following: A. Nitrogen 1) 5 protons 2) 7 protons 3) 14 protons B. Sulfur 1) 32 protons 2) 16 protons 3) 6 protons C. Barium 1) 137 protons 2) 81 protons 3) 56 protons The Periodic Table Isotopes • Isotopes are atoms that have the same atomic number but different mass number. • Most elements have two or more isotopes. • Same chemical properties because the electron number does not change. Isotope symbols Mass number A Z Atomic number X Example 11 B 5 •How many protons does this have? •How many neutrons does this have? •Is the “5” necessary ? More about isotopes: • Some isotopes have unstable nuclei which break down over time. • They are called radioactive isotopes • Some radiation is harmful. • Radiation can also be useful Radioactive Dating Cancer Treatment Tracers with X-rays Kill bacteria More About Atomic Structure • The center of the atom is called the nucleus. • Electrons live in something called shells. • Shells are areas that surround the center of an atom. • A shell is sometimes called an orbital or energy level. More About Electrons • Every shell can hold only so many electrons • The further from the nucleus, the more electrons a shell can hold Valence Electrons • The electrons on the outside edge of the atom • This is where the action is- where bonding takes place • Atoms have no more than 8 valence electrons The Octet Rule: • Atoms will combine to form compounds in order to reach eight electrons in their outer energy level. This is very stable! • Atoms with less than 4 electrons tend to lose electrons. •Atoms with more than 4 electrons tend to gain electrons. Compound • Two or more elements chemically combined in specific proportions • Examples: – Water – Salt – Sugar H 2O NaCl C6H12O6 Chemical Formulas are used to represent compounds Two types of compounds: Ionic Covalent Ionic Compounds • Form when electrons are transferred from one atom to another. Ions - Atoms with a net charge due to gaining or losing electrons – Gaining electrons gives an ion a negative charge – Losing electrons gives an ion a positive charge **If they have to choose, atoms would rather be stable (with a full “octet”) than neutral. How Does This Happen? Some atoms have a few too many electrons Some atoms only need a few electrons What do you do if you are a sodium (Na) atom with one extra electron? Go look for an atom that wants it! Ionic Bonding • Negative ions and positive ions are held together by ionic bond. • Ionic compounds form between metals and nonmetals What If No One Will Give Up An Electron? • Atoms with less than 8 valence electrons can move close to each other and share their electrons • The electrons spend their time around both atoms. • And they lived happily ever after! Covalent Bonds • Formed when a pair of electrons is shared between two atoms. • Sometimes the atoms share two pairs of electrons and form a double bond, or three pairs of electrons to form a triple bond. • Structures formed by covalent bonds are molecules. • Covalent compounds form between 2 nonmetals Van der Waals Forces • There are small attractive forces between all atoms • Help to hold molecules to each other – Ex: Gecko Let’s summarize what we know! Why do compounds form? • Atoms are trying to get 8 valence electrons How do compounds form? • By ionic (e- transfer) or covalent (e- sharing) bonding How can you tell if a compound is ionic or covalent? • By the types of elements in the compound (ionic = M + M covalent = M + NM) Learning Check 2: Indicate whether a bond between the following would be 1) Ionic 2) covalent ___ A. sodium and oxygen ___ B. nitrogen and oxygen ___ C. phosphorus and chlorine ___ D. calcium and sulfur ___ E. chlorine and bromine 2-2 Water is a Polar Molecule • Polar: Molecule in which electrons are shared unevenly between atoms, causing each end of the molecule to have a slight charge Negative end Positive end • This causes water to be attracted to other polar or charged particles – Water is attracted to ions – Water is attracted to itself, forming hydrogen bonds Hydrogen Bonds In Water Are Responsible For: • Adhesion – Attraction between molecules of different substances – Graduated cylinder • Cohesion – Attraction between molecules of the same substance – Drops of water on a penny • Ex: Surface Tension • Jesus Lizard Polarity of Water • In a water molecule two hydrogen atoms form single polar covalent bonds with an oxygen atom. Gives water more structure than other liquids – Because oxygen is more electronegative, the region around oxygen has a partial negative charge. – The region near the two hydrogen atoms has a partial positive charge. • A water molecule is a polar molecule with opposite ends of the molecule with opposite charges. • Water has a variety of unusual properties because of attractions between these polar molecules. – The slightly negative regions of one molecule are attracted to the slightly positive regions of nearby molecules, forming a hydrogen bond. – Each water molecule can form hydrogen bonds with up to four neighbors. Fig. 3.1 Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings HYDROGEN BONDS • Hold water molecules together • Each water molecule can form a maximum of 4 hydrogen bonds • The hydrogen bonds joining water molecules are weak, about 1/20th as strong as covalent bonds. • They form, break, and reform with great frequency • Extraordinary Properties that are a result of hydrogen bonds. – Cohesive behavior – Resists changes in temperature – High heat of vaporization – Expands when it freezes – Versatile solvent Organisms Depend on Cohesion Hydrogen bonds hold the substance together, a phenomenon called cohesion • Cohesion is responsible for the transport of the water column in plants • Cohesion among water molecules plays a key role in the transport of water against gravity in plants • Adhesion, clinging of one substance to another, contributes too, as water adheres to the wall of the vessels. • Surface tension, a measure of the force necessary to stretch or break the surface of a liquid, is related to cohesion. – Water has a greater surface tension than most other liquids because hydrogen bonds among surface water molecules resist stretching or breaking the surface. – Water behaves as if covered by an invisible film. – Some animals can stand, walk, or run on water without breaking the Fig. 3.3 surface. Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings Moderates Temperatures on Earth Water stabilizes air temperatures by absorbing heat from warmer air and releasing heat to cooler air. Water can absorb or release relatively large amounts of heat with only a slight change in its own temperature. Celsius Scale at Sea Level 100oC Water boils 37oC Human body temperature 23oC Room temperature 0oC Water freezes • What is kinetic energy? • Heat? • Temperature? • Calorie? • What is the difference in cal and Cal? • What is specific heat? Specific Heat is the amount of heat that must be absorbed or lost for one gram of a substance to change its temperature by 1oC. Three-fourths of the earth is covered by water. The water serves as a large heat sink responsible for: 1. Prevention of temperature fluctuations that are outside the range suitable for life. 2. Coastal areas having a mild climate 3. A stable marine environment Evaporative Cooling • The cooling of a surface occurs when the liquid evaporates • This is responsible for: – Moderating earth’s climate – Stabilizes temperature in aquatic ecosystems Density of Water • Most dense at 4oC • Contracts until 4oC • Expands from 4oC to 0oC The density of water: 1. Prevents water from freezing from the bottom up. 2. Ice forms on the surface first—the freezing of the water releases heat to the water below creating insulation. – When water reaches 0oC, water becomes locked into a crystalline lattice with each molecule bonded to to the maximum of four partners. – As ice starts to melt, some of the hydrogen bonds break and some water molecules can slip closer together than they can while in the ice state. – Ice is about 10% less dense than water at 4oC. Fig. 3.5 Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings Solvent for Life • Solution – Solute – solvent • Aqueous solution • Hydrophilic – Ionic compounds dissolve in water – Polar molecules (generally) are water soluble • Hydrophobic – Nonpolar compounds Acids and Bases • An acid is a substance that increases the hydrogen ion concentration in a solution. • Any substance that reduces the hydrogen ion concentration in a solution is a base. – Some bases reduce H+ directly by accepting hydrogen ions. • Strong acids and bases complete dissociate in water. • Weak acids and bases dissociate only partially and reversibly. pH Scale • The pH scale in any aqueous solution : – [ H+ ] [OH-] = 10-14 • Measures the degree of acidity (0 – 14) • Most biologic fluids are in the pH range from 6 – 8 • Each pH unit represents a tenfold difference (scale is logarithmic) – A small change in pH actually indicates a substantial change in H+ and OHconcentrations. Problem How much greater is the [ H+ ] in a solution with pH 2 than in a solution with pH 6? Answer: pH of 2 = [ H+ ] of 1.0 x 10-2 = 1/100 M pH of 6 = [ H+ ] of 1.0 x 10-6 = 1/1,000,000 M 10,000 times greater Buffers • A substance that eliminates large sudden changes in pH. • Buffers help organisms maintain the pH of body fluids within the narrow range necessary for life. – Are combinations of H+ acceptors and donors forms in a solution of weak acids or bases – Work by accepting H+ from solutions when they are in excess and by donating H+ when they have been Acid Precipitation • Rain, snow or fog with more strongly acidic than pH of 5.6 • West Virginia has recorded 1.5 • East Tennessee reported 4.2 in 2000 • Occurs when sulfur oxides and nitrogen oxides react with water in the atmosphere – Lowers pH of soil which affects mineral solubility – decline of forests – Lower pH of lakes and ponds – In the Western Adirondack Mountains, there are lakes with a pH <5 that have no fish. Types of Chemical Substances • Compounds and Elements are called pure substances. • Most matter is neither of these. Mixtures • Mixtures are combinations of substances held together by physical forces, not chemical bonds. • Each substance keeps its own properties Mixtures may be either: Solutions Colloids Suspensions Solutions • Have small particles • Are transparent (not the same as colorless) • Do not separate • Water solutions are very common in biological systems – Examples: salt water, kool-aid, air, brass, vinegar Colloids • Have medium size particles • Do not separate – Examples: fog, whipped cream, milk, cheese, mayonnaise Suspensions • Have very large particles • Settle out (separates into layers) – Examples: blood platelets, muddy water, calamine lotion, oil & water, Italian salad dressing pH Scale • Measures concentration of hydrogen ions in a solution • Ranges from 0 to 14 • 7 is neutral • 0-7 have more hydrogen ions (H+) and are acidic • 7-14 have more hydroxide ions (OH-) and are basic Acids, Bases, and pH • Water molecules form ions – H2O H+ – Water hydrogen ion + + OHhydroxide ion • Very few ions are formed in pure water, but there are equal numbers of hydrogen and hydroxide ions • Water is neutral! pH of common substances pH and Homeostasis • Maintaining a pH between 6.5 and 7.5 is important in cells • Dissolved compounds called buffers control pH – Proteins – Phosphates – Hydrogen carbonate Chemical Reactions • When one set of chemicals changes into another set of chemicals, a chemical reaction occurs • Bonds are either broken or formed (or both!) Chemical Equations • Represent a reaction • Give the types and amounts of substances that react and form Reactants 2H2 + O2 “yields” “yields” Products 2H2O Evidence of a Chemical Reaction • Formation of a precipitate (a solid substance separated from a liquid) • Gas is evolved (seen by bubbles forming in a liquid) • Change in heat or light energy Organic Compounds Organic Compounds • Make up most of living organisms • Contain bonds between two or more carbon atoms • C can easily bond with up to 4 other elements 4 valence electrons = 4 covalent bonds Organic Compounds • Carbon atom is versatile, can be “backbone” of long chains or rings • Organic molecules can be extremely large and complex; these are called macromolecules Organic Compounds • Four main types of organic macromolecules: Carbohydrates Lipids Proteins Nucleic Acids Carbohydrates • Made of C, H, & O • Main energy source for living things • Breakdown of sugars supplies immediate energy for cell activities • Extra sugar is stored as complex carbs called starches Carbohydrates • Single sugar molecules are called monosaccharides • Examples: • glucose – in many plant and animal tissues, most common monosaccharide • fructose – in many fruits • galactose – component of milk Carbohydrates • Large molecules of many monosaccharide are polysaccharides • Examples: • glycogen – animals use to store excess sugar • plant starch – plants use to store excess sugar • cellulose – fibers that give plants their rigidity & strength Lipids • Store more energy than CHOs because the chains are longer • Ex: Fats, oils, waxes • Won’t dissolve in water Lipids • Important parts of biological membranes and waterproof coverings • Steroids are lipids that act as chemical messengers Lipids • Many lipids are made from a glycerol combined with fatty acids – If all carbons have single bonds, lipid is saturated – Ex: butter, lard, animal fat (usually solid at room temperature) – If any carbons have double or triple bonds, lipid is unsaturated – Ex: vegetable oil, fish oil, peanut oil room temperature) (usually liquid at Proteins • Contain C, H, O, plus nitrogen • Formed from amino acids joined together • More than 20 amino acids can be joined in any order or number to make countless proteins (think of how many words can be made from 26 letters!) Proteins • Chains are folded and twisted giving each protein a unique shape • Van der Waals forces and hydrogen bonds help maintain protein’s shape • Shape of protein is important to its function! Proteins • Provide structure – Ex: Collagen- makes up your skin, muscles & bones • Aid chemical activities in your body – Ex: Enzymes- work to speed up rxns in your body • Transport substances into or out of cells • Help fight diseases Nucleic Acids • Contain C, H, O, N plus phosphorus • Formed by bonding of individual units called nucleotides nucleotide Nucleic Acid Nucleic Acids • Store and transmit hereditary information –Ex: DNA (deoxyribonucleic acid) RNA (ribonucleic acid)