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Chapter 2 The Chemistry of Life. CP Biology. Chapter 2 : The Chemistry of Life. 2.1: The Nature of Matter 2.2: Properties of Water 2.3: Carbon Compounds 2.4: Chemical Reactions and Enzymes. 2.1 The Nature of Matter. What is an Atom???.
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Chapter 2The Chemistry of Life CP Biology
Chapter 2:The Chemistry of Life • 2.1: The Nature of Matter • 2.2: Properties of Water • 2.3: Carbon Compounds • 2.4: Chemical Reactions and Enzymes
What is an Atom??? the smallest unit of an element, having all the characteristics of that element
Structure of the Atom ________________ Center of the atom Contains protons & neutrons Positive charge ________________ Positive charge Located inside of the nucleus ________________ Neutral charge Located inside of nucleus ________________ Negative charge Located outside of the nucleus in energy levels
Size of Particles(smallest largest) Quarks & Leptons Quarks Positive charge Have antiquarks (opposite of quarks – negative charge) Make up protons and neutrons, but not electrons Leptons Most well known lepton is the electron Subatomic particles Protons, neutrons, and electrons Atoms Made up of subatomic particles Elements Made up of atoms Compounds 2 or more elements
Elements and Compounds • _____________ • Cannot be broken down into simpler substances • Organized on the Periodic Table • Fe, H, O, Hg, Cu • _____________ • 2 or more elements combined • Properties differ from the properties of their component elements • H2O, CO, FeCl3
Atomic Number Identifies an element Tells the number of protons Tells the number of electrons in a neutral atom How the P.T. is arranged Example: Neon 10 p, 10 e-, 10 n
Atomic Mass Number Mass of an element Tells the number of protons and neutrons in an element May change Isotopes: atoms of the same element with different masses (different # neutrons) Decimal Number Example: Argon 39.9amu 18, p, 18 e-, 22 n
Isotopes Have the same atomic number, different mass Same # protons Different # neutrons
Ions • ___________: charged atoms • ____________: • positive ion • Gives away electrons • Example: Ca+2 • ___________: • Negative ion • Gains electrons • Example: N-3
Types of Chemical Bonds • Ionic Bonds • Covalent Bonds • Van der Waals Forces • Metallic Bonds • Between 2 metals • __________: mixture of 2 or more metals
Ionic transfers e- between a metal and nonmetal (+ and -) force binds opposite charged ions together conduct electricity when in solution at room temp, most are crystalline in structure high melting and boiling point hard, rigid, brittle solids Covalent shares e- usually between 2 nonmetals (- and -) do not conduct electricity in solution Polar or Nonpolar Polar: when e- are NOT shared equally IONIC vs. COVALENT BONDS
Van der Waals Forces • When molecules are close together, a slight attraction develops between the oppositely charged atoms • Weaker than ionic and covalent bonds
H2O # Lone Pairs = 2 # Bonding Pairs = 2 Bond angle = 104.5° Shape = Bent Modification of tetrahedral Example: Water
How does the structure of water contribute to its unique properties? • _________________: does not share its electrons equally • ___________________: attraction between a hydrogen atom on one water molecule to a hydrogen atom on another water molecule
Water’s Special Properties • ________________: • Is an attraction between water molecules of the same substance • _______________: • Is an attraction between molecules of different substances • ________________: • Takes a large amount of heat energy to cause water molecules to move quickly to heat up
Mixtures • Can be solids, liquids, or gases • Types of mixtures • Homogeneous mixture • Heterogeneous mixture • Colloid • Suspension
Homogeneous Mixtures(SOLUTIONS) • uniform throughout • a.k.a. _____________ • ___________: part being dissolved • ___________: part doing the dissolving ***WATER IS THE UNIVERSAL SOLVENT • Other examples: • iced tea • lemonade
Heterogeneous Mixtures • not uniform throughout • individual substances remain unique • Examples: • chef salad • Sand and water
Colloids & Suspensions • ___________: Mixtures that contain tiny, fine particles that reflect light • _________________: when particles in a mixture reflect light • Example: milk • ______________: when mixtures separate • Example: Italian salad dressing
Acids • pH range 0-7 • Produces H+ ions • Taste sour • Turn litmus paper pink/red • Corrosive • React with bases to form salts • examples: • Most foods (tomatoes, lemons, etc) • Battery acid • Stomach acid
Bases • produces [OH- ]] ions • tastes bitter • slippery • pH: beyond 7.0 -14 but not including 7.0 • turns litmus paper blue • combines with acids to form salts • electrolytes • can be corrosive • examples: cleaning supplies, eggs, seawater
pH Scale The pH scale is a measure of the [H+] ion concentration in a substance pH colors – ROYGBIV (red – very acidic green-neutral violet –very basic)
pH scale continued. . . • pH 1-3/4 • very strong acids • pH 5-6 • weak acids • pH 7 • neutral solutions • pH 7-10/11 • Weak alkaline solutions • pH 11-14 • Strong alkaline solutions
Buffers • Buffers • Weak acids or bases that can react with strong acids or bases to prevent sharp, sudden changes in __________ • Example: • Blood has a pH of 7.4 • Changes in blood pH are usually prevented by blood buffers such as bicarbonate and phosphate ions
Organic Chemistry • Organic Chemistry: • The study of compounds that contain bonds __________________ atoms
Carbon • Has _____ valence electrons • Electrons in the outer most energy level involved in bonding • One carbon atom can bond to another to make chains and branches
4 Groups _________________ _________________ _________________ _________________ Macromolecules are made from thousands of smaller molecules ______________ join together to form ______________ Macromolecules
Carbohydrates • Made up of carbon, hydrogen and oxygen atoms usually in the ratio 1:2:1 • Living things use carbs as their main source of energy • Structure: rings or chains • Types: • __________________ (glucose) 5-6 carbons • _____________ (sucrose) 10 or more carbons
Lipids • Not soluble in water • Used to store energy • Make up some biological membranes • Examples: fats, steroids • Types: • ____________: all single bonds • ____________: double and/or triple bonds
Nucleic Acids • Contain carbon, oxygen, nitrogen, hydrogen, and phosphorus • Are polymers made from nucleotides • __________________: • A five carbon sugar • Phosphate group (PO4-3) • and a nitrogen base • Store and transmit heredity • Examples: ________________
Proteins • Contain carbon, nitrogen, hydrogen, and oxygen • Made up of long chains of amino acids • _______________: • Compounds with an amino group (-NH2) • More than 20 in nature • Some control the rates of reactions and regulate cell processes • Some form cell structures
What is a chemical reaction? • ____________________________________________________________________
Evidence of a Chemical Reaction • Color change • Effervescence • Evolution of gas
Evidence of a Chemical Reaction continued. . . • Evolution of light • Odor • New substance forms • Formation of a Precipitate
Evidence of a Chemical Reaction continued. . . • Temperature change Exothermic/Spontaneous Reaction • Heat is given off • Heat/energy are products • Example: hot pack • Endothermic/Nonspontaneous Reaction • Heat is absorbed • Heat/energy are reactants • Example: cold pack
Parts of a Chemical Reaction ~ 2H2 + O2 2H2O Reactants: elements or compounds to the left of the arrow that combine together in a chemical reaction Products: elements or compounds to the right of the arrow that are produced in a chemical reaction Coefficient: large number before a chemical symbol Subscript: small lowered number after a chemical symbol : yields or produces Catalyst: symbol written above the arrow; speeds up rxn ∆: heat ~: electricity Pt: platinum