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The Chemistry of Life Chapter 2 Section 2.1 The Nature of Matter Atoms - the basic unit that makes up all matter • 2 main parts: • Nucleus – the center of the atom • Electron cloud – area around the nucleus • Contains smaller particles called subatomic particles. • Protons – found in nucleus • Positively charged • Mass = 1 atomic mass unit (1 u) • Neutrons – found in nucleus • No charge • Mass = 1 atomic mass unit (1 u) • Electrons – found in electron cloud • Negatively charged • No significant mass (1/1840 mass of proton) Section 2.1 The Nature of Matter Element • Pure substance consisting of one type of atom • Represented by 1-3 letter symbols • First letter capitalized • Other letters lower case • Examples: H, Na, Uub • All atoms of the same element have the same number of protons • Neutral atoms have the same number of protons and electrons • Isotopes • Atoms of the same element that have different numbers of neutrons and different masses • All have the same chemical & physical properties Section 2.1 The Nature of Matter Atomic Number • Equals the number of protons in an element Mass Number • Equals the sum of the protons and neutrons in an element Atomic Mass • Weighted average of the masses of an element’s isotopes • The abundance of each isotope in nature is considered when the average is calculated Section 2.1 The Nature of Matter Radioactive Isotopes • Nuclei are unstable and break down at a constant rate over time • Radiation released may be dangerous • Uses for radioisotopes: • Determination of the ages of rocks and fossils • Detection and treatment of cancer • Killing of bacteria that cause food to spoil • Used as “tracers” to follow the movement of substances within organisms Section 2.1 The Nature of Matter Chemical Compounds • Substance formed by the chemical combination of two or more elements in definite proportions. • Chemical formula • Shows the composition of a compound by a kind of short hand • Examples: H2O, NaCl • The physical and chemical properties of a compound are usually very different from those of the elements from which it is formed. Section 2.1 The Nature of Matter Section 2.1 The Nature of Matter Chemical Bonds • Hold atoms together • 2 main types: • Ionic bonds • Formed when electron(s) transferred from one atom to another • Transferred electrons are outer or valence electrons • Smallest part: ion • Charged atom • Negative ion: atom gains electron(s) • Positive ion: atom loses electron(s) • Very strong attraction • Example: Sodium chloride (table salt) Section 2.1 The Nature of Matter Chemical Bonds • Hold atoms together • 2 main types: • Covalent bonds • Formed when one or more pairs of valence electrons is shared between two atoms. • 1 pair: single bond • 2 pairs: double bond • 3 pairs: triple bond • Smallest part: molecule • Examples: • Water (H2O) • Oxygen gas (O2) Section 2.1 The Nature of Matter Chemical Bonds • van der Waals forces • Weak intermolecular attraction between the oppositely charged regions of nearby molecules • Not as strong as ionic or covalent bonds • Helps hold molecules together • Especially when large • Named after the scientist who discovered them Section 2.2 Properties of Water (pages 40 - 44) Water Properties: • Neutral molecule (10 protons and 10 electrons) • Polar molecule • Oxygen attracts electrons more than hydrogen • Oxygen has more protons to attract electrons • Shared electrons shared unequally • Has a partial positive area (around hydrogens) • Has a partial negative area (around oxygen) • Electrons stay closer to oxygen • Hydrogen bonds • Attraction between hydrogen atom on one water molecule and oxygen atom on another water molecule • Not as strong as ionic or covalent bonds • Multiple hydrogen bonds can form Section 2.2 Properties of Water (pages 40 - 44) Water Properties due to Hydrogen bonding: • Cohesion • Attraction between molecules of the same substance • Causes water molecules to be drawn together • Causes “beading” of water drops • Causes surface tension • Allows some insects and spiders to won on the surface of a pond. Section 2.2 Properties of Water (pages 40 - 44) Water Properties due to Hydrogen bonding: • Adhesion • Attraction between molecules of different substances • Water more attracted to glass of graduated cylinder than to other water molecules • Slight dip in water (meniscus) forms due to attraction of water to glass • Water rises in a narrow glass tube due to attraction of water to glass • Called capillary action Section 2.2 Properties of Water (pages 40 - 44) Water Properties due to Hydrogen bonding: • High Heat Capacity • It takes a large amount of heat energy to change the temperature of water. • Allows large bodies of water to absorb a great deal of heat with little temperature change • Oceans and lakes • Protects water organisms from drastic temperature changes • Allows heat from chemical reactions inside cells to be absorbed in the water of the cell to keep cell temperatures regulated. Section 2.2 Properties of Water (pages 40 - 44) Mixture: • Material composed of 2 or more elements or compounds that are physically mixed together but not chemically combined • Solutions • All parts of the mixture are evenly distributed throughout the solution • Solute: part of a solution being dissolved • Example: sugar in sugar water • Solvent: part of a solution doing the dissolving • Example: water in sugar water • Water’s polarity gives it the ability to dissolve both ionic compounds and other polar compounds. Section 2.2 Properties of Water (pages 40 - 44) Mixtures: • Suspensions • Mixture of water and undissolved material • Movement of water keeps undissolved materials suspended so they don’t settle out • When movement of water stops, materials will settle to the bottom • Blood in your body contains solutions and suspensions: • The liquid part of your blood is mostly water and contains many dissolved compounds • Blood also contains cells that do not dissolve in water and stay suspended as long as the water continues to move. Section 2.2 Properties of Water (pages 40 - 44) Acids, Bases, and pH: • pH scale • Ranges from 0 to 14 • Indicates concentration of hydrogen ions (H+) in solution • pH less than 7 is acidic • More H+ than OH• pH greater than 7 is basic • Less H+ than OH• pH equal to 7 is neutral • Equal amounts H+ and OH- Section 2.2 Properties of Water (pages 40 - 44) HONORS pH: • Means “function of hydrogen ion concentration” • Hydrogen ions (H+) can also be written as hydronium ions (H3O+) • As the pH increases, the concentration of hydrogen or hydronium ions decreases • Each step on the pH scale represents a negative power of 10. Section 2.2 Properties of Water (pages 40 - 44) Acids: • Any compound that forms extra hydrogen ions (H+) ions in solution • Sour taste • pH less than 7 • Strong acids: • pH = 0 – 3 • Can be corrosive to skin • Stomach acid is HCl (hydrochloric acid) Section 2.2 Properties of Water (pages 40 - 44) Bases: • Any compound that forms extra hydroxide (OH-) ions in solution • Bitter taste • Used as cleansers because they dissolve grease • pH greater than 7 • Strong bases: • pH = 11 – 14 • Can be corrosive to skin • Lye is NaOH • Chemical name is sodium hydroxide Section 2.2 Properties of Water (pages 40 - 44) Buffers: • Weak acid or base that can react with strong acids or bases to prevent sharp, sudden changes in pH • pH of most body fluids usually between 6.5 and 7.5 • Blood normal pH = 7.4 • Buffers found in our bodies keep us from large changes in pH to maintain homeostasis • Antacids are sometimes used to buffer stomach acids Section 2.3 Carbon Compounds (pages 45 - 49) Carbon Chemistry • Organic Chemistry • Study of compounds that contain bonds between carbon atoms • Inorganic Chemistry • Study of all other compounds • Carbon can combine with many elements to form the molecules of life • Including: hydrogen, oxygen, phosphorus, sulfur, and nitrogen • Remember CHONPS • Living organisms are made up of molecules that consist of carbon and these other elements Section 2.3 Carbon Compounds (pages 45 - 49) Carbon Atoms • Have 4 valence electrons • Carbon atoms can bond to each other and other elements • Can form 4 bonds • Carbon compounds can form single, double, or triple covalent bonds • Many different shapes of carbon compounds can be seen • Chains • Branched structures • Ring structures Section 2.3 Carbon Compounds (pages 45 - 49) Macromolecules • Large organic compounds • Means “giant” molecules • Most formed by the process of polymerization • Large compounds being built by joining smaller ones together • Monomers are the smaller compounds • Polymers are the large compounds being built • The monomers in a polymer may all be identical or may be different • There are 4 groups of macromolecules based on their chemical composition: • Carbohydrates, lipids, nucleic acids, proteins Section 2.3 Carbon Compounds (pages 45 - 49) Carbohydrates • Macromolecules • Made of carbon, hydrogen, and oxygen • Hydrogen and oxygen in 2:1 ratio • Means carbon + water • Functions: • Main source of energy for organisms • Structural purposes • Like the strings in celery • Polymers • Made of monomers called simple sugars or monosaccharides • Called polysaccharides Section 2.3 Carbon Compounds (pages 45 - 49) Section 2.3 Carbon Compounds (pages 45 - 49) Carbohydrates • Monosaccharides • Simple sugars • Examples: glucose, fructose, galactose • Disaccharides • Double sugars • Made by joining two monosaccharides • Example: sucrose • Made of glucose + fructose • Polysaccharides • Complex carbohydrates • Polymers • 3 main types: • Starches, glycogen, and cellulose Section 2.3 Carbon Compounds (pages 45 - 49) Polysaccharides (Complex sugars) • Starches • Polysaccharides used by plants to store excess sugar • Cellulose • Tough, flexible fibers that give plants much of their strength and rigidity • Glycogen • Used by animals to store excess sugars • We store in our liver • When blood sugar is low, glycogen is released into the blood • Used by muscles for contractions and movement Section 2.3 Carbon Compounds (pages 45 - 49) Lipids • Macromolecules • No monomers • Not a polymer • Made mostly of carbon and hydrogen • Few oxygen • Common categories: fats, oils, and waxes • Generally not soluble in water • Functions: • Store energy • Important parts of biological membranes • Waterproof coverings • Chemical messengers (steroids produced by the body) Section 2.3 Carbon Compounds (pages 45 - 49) Lipids • Many formed when a glycerol molecule combines with compounds called fatty acids • Circulating fats are called triglycerides • Saturated Fats • Contain the maximum possible number of hydrogen atoms • No double bonds between the carbon atoms on the fatty acids • Less healthy Section 2.3 Carbon Compounds (pages 45 - 49) Lipids • Unsaturated Fats • Have at least one double bond between the carbon atoms on the fatty acids • Polyunsaturated fats have more than one double bond between the carbon atoms on the fatty acids • More healthy • Trans Fats • Unsaturated fat that has had air whipped into it until it becomes saturated • Example: shortening • Found in some cookies, crackers, and doughnuts • Very unhealthy! Section 2.3 Carbon Compounds (pages 45 - 49) Lipids • Fats • Usually solid at room temperature • From animals • Less healthy • Some plants • Coconut oil can be solid at room temp • More healthy • Usually saturated fats Section 2.3 Carbon Compounds (pages 45 - 49) Lipids • Oils: • Usually liquid at room temperature • Plants • Usually more healthy • Usually unsaturated • Can be monounsaturated or polyunsaturated Section 2.3 Carbon Compounds (pages 45 - 49) Lipids • Waxes: usually solid at room temperature • Sources: Animals and plants • Honeycomb, wax on outside of cucumbers Section 2.3 Analyzing Data (page 48) Comparing Fatty Acids 1. Which of the 4 fatty acids is saturated? Which are unsaturated? • Stearic acid is saturated • The other 3 fatty acids are unsaturated. 2. How does melting point change as the number of carbon-carbon double bonds increases? • Melting point decreases as the number of double bonds increases. Fatty Acid Number of Carbons Number of Double Bonds Melting Point (°C) Stearic Acid 18 0 69.6 Oleic Acid 18 1 14 Linoleic Acid 18 2 -5 Linolenic Acid 18 3 -11 3. If room temperature is 25°C, which fatty acid is a solid at room temperature? Which is liquid at room temperature? • Stearic acid is solid at room temperature • The other 3 fatty acids are liquid at room temperature. Section 2.3 Carbon Compounds (pages 45 - 49) Nucleic Acids • Macromolecules • Made of hydrogen, oxygen, nitrogen, carbon, and phosphorus • Polymers: Nucleic acids or polynucleotides • Monomers: Nucleotides • 3 main parts: • 5-carbon sugar • Phosphate group • Nitrogen base (also nitrogenous) • Function: • Store and transmit hereditary or genetic information • 2 main types: DNA (deoxyribonucleic acid) and RNA (ribonucleic acid) Section 2.3 Carbon Compounds (pages 45 - 49) Section 2.3 Carbon Compounds (pages 45 - 49) Proteins • Macromolecules • Made of nitrogen, carbon, hydrogen, and oxygen • Polymers: Proteins or polypeptides • Monomers: Amino acids • Compounds with an amino group (NH2) on one end and a carboxyl group (-COOH) on the other end • Linked together by peptide bonds to form a protein or polypeptide (polymer) Section 2.3 Carbon Compounds (pages 45 - 49) Proteins • Functions: • Control the rate of chemical reactions • Regulate cell processes • Form important cellular structures • Transport substances into or out of cells • Help to fight disease Section 2.3 Carbon Compounds (pages 45 - 49) Proteins • More than 20 different amino acids found in nature • Since all amino acids have an amino group and a carboxyl group, they can join to any other amino acids • Makes proteins very diverse • Levels of organization in proteins: • Primary structure: sequence of amino acids • Secondary structure: folding or coiling of polypeptide chain • Tertiary structure: complete 3-D arrangement of polypeptide chain • Quaternary structure: proteins with more than one chain are arranged in a specific way Section 2.4 Chemical Reactions and Enzymes (pages 50-53) Chemical Reactions • Process that changes, or transforms, one set of chemicals into another. • Mass and energy are absorbed during chemical reactions • Including chemical reactions in living organisms • Some reactions happen slowly • Example: When iron and oxygen react to form rust • Some reactions happen quickly • Example: When baking soda and vinegar are mixed Section 2.4 Chemical Reactions and Enzymes (pages 50-53) Chemical Reactions • Reactants • The elements or compounds that enter into a chemical reaction • Products • The elements or compounds produced by a chemical reaction • Coefficients • Whole numbers placed in front of elements or compounds in a chemical equation • Used to balance chemical equations • Chemical reactions involve changes in the chemical bonds that join atoms in compounds. Section 2.4 Chemical Reactions and Enzymes (pages 50-53) Energy in Reactions • Energy is released or absorbed whenever chemical bonds are formed or broken • Chemical reactions involved changes in energy • Chemical reactions that release energy often occur spontaneously • Chemical reactions that absorb energy will not occur without a source of energy Section 2.4 Chemical Reactions and Enzymes (pages 50-53) Energy in Reactions • With reversible chemical reactions, if the reaction releases energy in one direction 2H2 + O2 2H2O • The reaction will absorb energy in the other direction, and will not occur spontaneously 2H2O 2H2 + O2 Energy Sources • In order to stay alive, organisms need to carry out reactions that require energy • Plants get their energy by trapping and storing energy from the Sun • Animals get their energy when they consume other animals or plants Section 2.4 Chemical Reactions and Enzymes (pages 50-53) Activation Energy • Energy needed to get a chemical reaction started • Needed regardless of whether a reaction releases or absorbs energy • Energy releasing reactions need a lower activation energy than energy absorbing reactions Section 2.4 Chemical Reactions and Enzymes (pages 50-53) • A catalyst is a substance that: • Speeds up the rate of a chemical reaction • Does not get used up in a chemical reaction • Lowers activation energy • Enzymes • Proteins that act as biological catalysts • Speed up reactions that take place in cells • Lowers activation energy • Are very specific • Usually only catalyze one chemical reaction • Part of the enzyme’s name comes from the reaction it catalyzes. • Usually ends in –ase • Example: Carbonic anhydrase gets it name because it catalyzes the reverse reaction that removes water from carbonic acid Section 2.4 Chemical Reactions and Enzymes (pages 50-53) • Without carbonic anhydrase, the reaction is very slow • Carbon dioxide can build up in the blood to dangerous levels • When carbonic anhydrase is present, the reaction takes place immediately • Carbon dioxide gets removed quickly from the blood Section 2.4 Chemical Reactions and Enzymes (pages 50-53) Enzyme-Substrate Complex • For chemical reactions to occur, reactants must collide with enough energy so that existing bonds are broken and new bonds will be formed. • If reactants do not have enough energy, they will not be changed after the collision • Enzymes provide a site where reactants can come together to react • The site reduces the activation energy • Called the active site • Reactants of enzyme-catalyzed reactions are called substrates Section 2.4 Chemical Reactions and Enzymes (pages 50-53) Enzyme-Substrate Complex • The substrate and active site on the enzyme have complementary shapes • The fit between the substrate and active site is so precise, it can be compared to a lock and a key • The active site on the enzyme is the lock • The substrate (reactant) is the key • The “key” will only fit into one “lock” Section 2.4 Chemical Reactions and Enzymes (pages 50-53) Regulation of Enzyme Activity • Enzymes play essential roles in: • Controlling chemical pathways • Making materials cells need • Releasing energy • Transferring information • Factors that affect enzyme activity: • Change in temperature • Enzymes work at specific temperatures • Change in pH • Enzymes work at specific pH • Regulatory molecules • Enzymes cannot work when certain substances are present