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Chapter 2 – Chemistry

Chapter 2 – Chemistry. Why Study Chemistry in Biology?. Chemical changes in matter are the foundation for all life processes Living things are composed of the same kinds of matter that make up nonliving things. Chemical breakdown of human body is:. 65% Oxygen

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Chapter 2 – Chemistry

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  1. Chapter 2 – Chemistry

  2. Why Study Chemistry in Biology? • Chemical changes in matter are the foundation for all life processes • Living things are composed of the same kinds of matter that make up nonliving things

  3. Chemical breakdown of human body is: • 65% Oxygen • 18% Carbon Makes up 96% of living things • 10% Hydrogen • 3% Nitrogen • 1.5% Calcium • .35% Potassium • .25% Sulfur • .15% Sodium • .15% Chlorine • .05% Magnesium • .0004% Iron • .00004% Iodine • Traces of fluorine, silicon, manganese, zinc, copper, aluminum, and arsenic

  4. How are these elements all put together to make up the human body? 2.1 “Composition of Matter” • Matter – • anything that occupies space and has … • Mass – • quantity of matter an object has. (Weight is not the same – downward force of gravity factored in) • Elements – • pure substances that cannot be broken down chemically into simpler kinds of matter. • 118 elements as of 2006. • N, C, H, O, P, S are important elements in Biology

  5. The Periodic Table of Elements Table that lists all known elements and their important information The elements are organized in a specific way • Table

  6. ALL ELEMENTS ARE ELECTRICALLY NEUTRAL TO START!!! • Atomic number = • the number of protons in atom (= electrons also) • Element symbol • (Carbon) • Mass number= • # of protons + # of neutrons • Atomic Mass = relative average mass of element – a decimal 6 C 12 12.01

  7. Element symbols to be familiar with… • Carbon __________ • Hydrogen __________ • Oxygen _________ • Nitrogen __________ • Calcium __________ • Phosphorus __________ • Potassium __________ • Sulfur __________ • Sodium __________ • Chlorine __________ • Magnesium _________ • Iron __________ • Iodine __________ • Fluorine __________ • Silicon __________ • Zinc __________ • Copper __________ • Aluminum __________ • Arsenic __________ 20. Manganese __________

  8. Atoms • simplest particles of an element that retain all the properties of that element. Properties of atoms determine the properties of matter they compose

  9. Parts of an Atom • Protons: Positive electrical charge Mass: 1AMU Location: in nucleus • Neutrons: No electrical charge Mass: 1AMU Location: in nucleus 3. Electrons: Negative electrical charge Mass: 1/2000 (so it is counted as 0 AMU) Location: surrounding nucleus

  10. Isotopes Element that have the same number of protons but different number of neutrons They vary by their atomic mass and mass number For example: C-12, C-13, C-14 The decimal you see on the PT = average of the relative amounts in nature of the various isotopes

  11. Models of the atom Bohr model – electrons appear to “orbit” the nucleus – aka Planetary Model Electron cloud model - protons and neutrons concentrated in the nucleus and electrons occupying various energy levels around the nucleus – not sure where the electrons are at any time

  12. How are the electrons arranged in the energy levels? • First energy level • will get a maximum of 2 electrons • Second energy level = 8 electrons • Third energy level = 8 electrons Electrons “sit” in these levels in ONLY this order!!!

  13. Filling Energy Levels • Filling rules – 2,8,8 electrons • Orbitals - probability Yellow = nucleus Blue = level 1 = 2 Red = level 2 = 8 Green = level 3 = 8

  14. 1 2

  15. Diagram the following atoms • 1H1 • 12C6atomic number • 14N7 • 16O8 mass number • 23Na11 • 35Cl17 • 39K19 • 40Ca20 NOTE – THESE ARE ISOTOPE DESIGNATIONS

  16. Why Do Atoms Combine? • Atoms will combine chemically to produce compounds • Compounds form due to arrangement of electrons in outermost energy level= VALENCE ELECTRONS • Atoms are most stable when outer energy level is filled • Chemical bonds are broken, atoms are rearranged, and new chemical bonds are formed. • Chemical Bonds = attractive forces holding atoms together ALL OF THESE CHANGES INVOLVE AN EXCHANGE OF ENERGY

  17. When atoms combine you get… Compounds are: Pure substances made up of atoms of 2 or more different elements • i.e. Water, Glucose, Salt • can be ionic or covalent Molecules are: Pure substances made up of atoms of 2 or more similar elements, i.e. O2 - Can only be covalent

  18. Atoms in molecules and compounds are arranged in fixed proportions • The chemical formula: 2H2O coefficient CH4 (Methane) • 1 C 4 H • C6H12O6 (Glucose) • 6 C 12 H 6 O • (NH4)2SO4 (Ammonium Sulfate) • 2 N 8H 1S 4 O subscript

  19. Types Of Chemical Bonds, Overview • https://www.youtube.com/watch?v=_M9khs87xQ8

  20. Physical and Chemical properties of compounds differ from the elements that make them up. • i.e. NaCl, H2O http://www.bing.com/videos/search?q=ionic+and+covalent+bonds&qs=n&form=QBVR&pq=ionic+and+covalent+bonds&sc=8-24&sp=-1&sk=#view=detail&mid=FC661AB5D4927AD1FDD7FC661AB5D4927AD1FDD7

  21. Valence electrons = those in the outermost energy level

  22. 1. Ionic Bonds Bonding produced when atoms transfer electrons. Not as strong as covalent bonds. Produce charged atoms (ions). • i.e. NaCl

  23. Na has lost an electron (thus the + charge) • Cl has gained an electron (thus the – charge) Na+ Cl- ----NaCl

  24. 2. Covalent Bonds Covalent Bonds • form when 2 atoms share 1 or more electrons H2O, CH4, C6H12O6 • Strong bonds • All organic (have C & H) substances H2O

  25. Covalent Bonds Covalent bonds are strong bonds due to the sharing of electrons. • In order to break a covalent bond, • High heat • Electrical current • High Pressure • Enzymes (catalysts produced by living things)

  26. Molecule • simplest part of a substance that retains all of the properties of the substance • only covalent compounds form molecules water

  27. Using the model kits, make the following*: • Colors of spheres on board • Use wooden sticks for single bonds; springs for double and triple bonds • H2O • O2 • CH4 • HCl • O3 • CO2 • H2 8. C2H6 9. C2H4 10. C2H2 11. C3H8 12. NH3 * Draw Structural Diagrams of each

  28. 2.2 “Energy” Energy – Ability to do work or cause change All living things need a constant flow of energy – types??? 1st Law of Energy: Energy is neither created nor destroyed but transferred Free energy – energy in a system that is available to do work. Body contains glucose to provide free energy and stored energy (glycogen and fat). To tap into these, you must break these down (digestion) into their simplest form (glucose).

  29. States of Matter Solids, liquids and gases Particle movement (greater in gases) – higher energy Shape and volume (fixed in solids) Concentration of particles (tighter in a solid)

  30. Chemical Reactions Yields CO2 + H2O H2CO3 (Carbonic acid) Reactants Product The above reaction occurs in one direction and is non-reversible. The reaction below occurs in both directions and is reversible. CO2 + H2O H2CO3 (Carbonic acid)

  31. “THE WINTER BALL”

  32. Energy Transfer Body continuously goes through series of chemical reactions = METABOLISM • Exergonic – net release of free energy • Temperature increases to indicate a release of energy

  33. Endergonic – net absorption of free energy • Temperature decreases to indicate an absorption of energy

  34. Activation Energy Amount of energy necessary to begin a reaction Catalysts – Chemicals that lower the amount of needed Activation Energy Enzymes – Organic catalysts (found only in living things) – help without being changed Lactase breaks down the milk sugar Lactose

  35. Redox Reactions INVOLVE Transfer of electrons Oxidation reactions – Reactants lose one or more electrons forming + ions • Reduction reactions – Reactants gain one or more electrons. Form – ions • Na + Cl Na+Cl- • Oxidized Reduced • These reactions always occur together

  36. 2.3 “Water and Solutions” • Most mass of living things is water (Universal solvent) • About 65% of the total mass of our cells is water • Chemical reactions occur in water • Must understand chemistry of water

  37. DEMO • Add water to a beaker, about half full • Into another beaker, add the same amount of rubbing alcohol • Into a third beaker, do the same with cooking oil • Sprinkle a small amount of salt into each and swirl • Let stand for a moment • Name the solute and solvent in each beaker. • Solute Solvent Salt Water, Alcohol, cooking oil • In which beaker did the salt dissolve (go into solution)? • The water and somewhat in alcohol • Which solution is an aqueous solution (one in which the solvent is water)? • Only the water

  38. Solutions A mixture of one or more substances uniformly distributed in another substance – physically combined • Solute – substance being dissolved (sugar) • Solvent – substance that does the dissolving (water)

  39. Solutions, cont. • Physically but not chemically combined • Solutions can vary in concentration of solute; 5% sugar solution has 5% sugar and 95% water • A solution is said to be saturated when the solvent can no longer dissolve all the solute • Aqueous solution – solvent is water

  40. Concentration

  41. What Makes Water Such a Good Solvent? • The chemical nature of water is called POLARITY

  42. The Hydrogen Bondhttp://www.youtube.com/watch?v=aH2IbYs_XjY Occurs between H and O and between H and N

  43. Properties of Water(due to its polar nature) • Cohesion – water sticking to itself – a barrel of monkeys- surface tension • http://www.youtube.com/watch?v=ynk4vJa-VaQ • Adhesion – water sticking to another polar substance – glass slide demo - capillarity • Thermal regulation – high heat capacity, evaporative cooling • Density of ice

  44. Acids and Bases • Acid – sour, corrosive - lemons • Alkaline – bitter, smooth - bleach • Chemical significance???

  45. Ionization or Dissociation • the production of ions when atoms or molecules break apart NaCl Na+ and Cl- ionic dissociate H20 H+ + OH –covalent ionize • H+ = Hydrogen ion • OH- = Hydroxide ion • H3O+ = Hydronium ion

  46. Dissociation of Salt in Waterhttp://www.youtube.com/watch?v=CLHP4r0E7hg

  47. Production of the Hydronium ion • Due to the high kinetic energy of the molecules of water, there are numerous collisions. Some of these collisions are strong enough to dislodge protons (H+) from a water molecule or from an ionized acid molecule such as H+Cl-. • Other water molecules will pick up these stray protons • H2O H+ + OH- • H20 + H+ H3O+ (Hydronium ion)

  48. Acids • Acid Solution • # of H30+ ions outnumbers the OH- ions in a solution • HCl = H+ + Cl- • H+ + H2O H3O+ Hydronium ion • Acids are sour and corrosive • Acid rain - pH of normal rain ~ 5.0 – 5.6 on pH scale • SO3 + H2O H2SO4

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