Chapter 2 - Chemistry of Life

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Chapter 2 - Chemistry of Life

  1. 1. • • • • • • • • • • • Biosphere Ecosystem Community Population Organism Organ system Organ Tissue Cell Molecule Atom
  2. 2. • Everything in the universe is made of matter. • Matter is anything that occupies space and has mass and is made of atoms! • Mass is the measurement of the amount of matter in an object. Matter in our galaxy
  3. 3. • Elements are pure substances that cannot be broken down chemically into simpler kinds of matter. (some things don’t count, for example: protons, electrons, etc.) • Many elements will be very familiar to you. For example… • Others will not be familiar. 1.Helium • For example… Einsteinium, Americium, Nobelium 2.Oxygen 3.Gold 4.Platinum 5.Aluminum
  4. 4. • Over 100 elements have been identified • 30 are important to living things • 90% of the mass of living things are composed of 4 elements 1. Oxygen (O) 2. Carbon (C) 3. Hydrogen (H) 4. Nitrogen (N)
  5. 5. • Elements have chemical symbols • Composed of one or two letters • Usually taken from letters in the common name. • Sometimes taken from Latin or Greek name • Aurum (Au) for gold • Natrium (Na) for sodium • Elements are also identified by their atomic number. • Elements are arranged in the Periodic Table.
  6. 6. • The atom is the simplest particle of an element that retains all of the properties of that element. • In other words, an atom is one single “piece” of an element. • For example, the smallest amount of carbon = 1 atom of carbon • Atoms are too small to be observed by conventional means • Scientists show in model form – Models do not show exactly what an atom looks like. – Used to predict how they will act
  7. 7. • The atom is broken down into 2 major components 1. Nucleus - contains protons and neutrons 2. Electron cloud - contains electrons
  8. 8. • The nucleus has 99.999% of the mass of an atom but little volume • Contains protons and neutrons • Protons are positively (+) charged particles • Neutrons are particles with no charge (0) • All atoms of the SAME element have the SAME number of protons. • Atomic number = number of protons in the atom • Mass Number = total number of protons and neutrons in the atom Atomic Number 6 C Carbon 12 Chemical Symbol Chemical Name Atomic Mass or Mass Number
  9. 9. • Electrons have a negative charge (-) • Very low mass, but high energy • Electrons are NOT in the nucleus, but in the electron cloud • The net electrical charge of an atom is zero (not positive or negative) because the atom has an equal number of electrons (-) and protons (+). The equal but opposite charges cancel each other out. • Number of Electrons = Number of Protons = Atomic Number • Atoms can gain or lose electrons in chemical reactions and become ions (more on this later)
  10. 10. • Isotopes are atoms of the same element that have a different number of neutrons. • The Average atomic mass of an element takes into account the relative amounts of each isotope in the element.
  11. 11. • Electron cloud • We do not really know exactly where the electrons are at any time in the atom. We only know where they might be. • Fortunately, for most chemistry it doesn't really matter where the electron actually is, we only care about how much energy it has. • It’s convenient to think that electrons move around the nucleus in orbits, like the planets in the solar system. • Orbitals may be misleading about where an electron is, but they tell us how much energy it has. We call this the Energy Level of the electron.
  12. 12. • Bigger Orbit = Bigger Energy Level = Higher Energy • Energy levels can hold different numbers of electrons Level 1: 2 electrons Level 2: 8 electrons • The number of electrons in the outer energy level determines the characteristics of the element. • Energy Levels are usually not filled (except noble gases) • Goal: have outer energy level full
  13. 13. • Very few elements exist by themselves naturally, instead, they are usually combined with other elements. • Compounds are made up of atoms of two or more elements in fixed proportions – Chemical formula shows the kinds and proportions of atoms of each element that forms a particular compound (Ex: H20)
  14. 14. • Compounds are usually very different from the elements they form from – Sodium – a reactive, soft , silvery metal – Chlorine – a reactive, poisonous green gas – Sodium chloride (or table salt) – stable, colorless crystals +
  15. 15. • Most atoms are not stable in their natural state, so they tend to react with other atoms in different ways to form compounds and become more stable. – Remember, the goal to stability is having a full outer energy level. This is accomplished when atoms bond together to form compounds. • Chemical bonds are the attractive forces that hold atoms together • Electrons from the outermost energy level of the atoms are SHARED or TRANSFERRED whenever a bond is made.
  16. 16. • Covalent bonds are formed from atoms sharing electrons. – Share 2 electrons (1 pair) = Single Covalent Bond – Share 4 electrons (2 pairs) = Double Covalent Bond – Share 6 electrons (3 pairs) = Triple Covalent Bond • Usually occurs between two non-metal elements. Nonmetals Metals The animation above shows what happens in the formation of a covalent bond. The individual atoms are atoms of chlorine with only their outer level of electrons shown. Note that each chlorine atom has only seven outer electrons, but really wants eight.
  17. 17. • Ionic bonds involve the transfer of electrons between atoms. • When an atom gains or loses an electron it is called and ion. • Usually occurs between a metal and a nonmetal. Nonmetals Metals
  18. 18. • Transferring electrons causes charges ( + or -) to develop • Opposites attract and compounds are formed!
  19. 19. • Energy is defined as the ability to do work and to cause change. • There are many different forms of energy: 1. Chemical Energy- stored in the bonds of atoms and molecules. The food we eat, natural gas, etc. 2. Thermal Energy- heat. 3. Electrical Energy- movement of electrical charges, lightening, light. 4. Mechanical Energy- movement of objects and substances from one place to another.
  20. 20. • Energy is never created or destroyed, but it can change form. • A car engine burns gasoline, converting the chemical energy in gasoline into mechanical energy and thermal energy. • Scientists try to understand this FLOW of energy in living systems and the universe.
  21. 21. • Living things depend on energy • Chemical- food • Thermal-your body regulates your body temperature • Electrical- nerves use electrical impulses to communicate information to your brain • Mechanical- moving your muscles
  22. 22. • Atoms in all states of matter are constantly moving. The RATE of this motion determines the STATE of matter. 1. Solid:  Least energetic, definite volume and shape  Particles closely packed together, vibrate, and have high attraction to each other
  23. 23. 2. Liquid:  Definite volume, variable shape (takes shape of container)  Has ability to flow, particles have freedom to move
  24. 24. 3. Gas:  High Energy  Takes volume and shape of container  Ability to flow, low attraction to each other
  25. 25. 4. Plasma:  Most energetic  Forms when a gas' temperature is raised to over 10,000°  Its molecules collide so violently that the electrons are knocked off the atoms.  Plasma is different from a gas because it is VERY hot and is influenced by electric and magnetic fields.  Exists above the earth’s atmosphere The influence of magnetic fields on the behavior of plasmas is shown here by loop of plasma on the sun
  26. 26. Gelatin is made of protein. Proteins are solids at most temperatures. When you mix the Jell-O powder into the hot water, the protein dissolves in the water. As the solution cools down, the protein comes out of solution and turns back into a solid. But it doesn’t just form a solid that settles onto the bottom, the protein molecules stick to each other forming a giant mixed-up jungle gym of protein molecules all sticking together. So, Jell-O is basically a solid suspended in a liquid. The scientific word for this is a ’colloid’. If you heat it up enough, the solid will become dissolved again, and it will become a liquid all through. But, if you cool it down enough, the liquid water will freeze, and it will become a solid all through.
  27. 27. • Energy needs to be added or removed from system to change the state of matter Add Energy Remove Energy
  28. 28. • http://www.harcourtschool.com/activity/stat es_of_matter/
  29. 29. • Chemical Reaction – one or more substances change to produce one or more different substances – Energy is absorbed or released when chemical bonds are broken and new ones are formed
  30. 30. • Living things perform thousands of chemical reactions during life processes • We write the reactions in equations CH4 + 2 O2 CO2 + 2 H2O yields • Broken down into two parts 1. Reactants – substances or molecules that participate in a chemical reaction; found on left side of an equation 2. Products – substances that form in a chemical reaction; found on right side of an equation • Atoms on each side of an equation have to equal in number and type
  31. 31. • Your body is fueled by carbohydrates (sugar and starch), proteins and fats! • Body breaks them down into carbon dioxide and water in a process called cellular respiration. This process releases energy for your body to use. Cellular Respiration: 6O2 + C6H12O6 --> 6H2O + 6CO2 + energy • Plants absorb energy from the sun and combine it with carbon dioxide and water to produce sugars and oxygen (photosynthesis) Photosynthesis: 6CO2 + 6H2O + energy --> 6O2 + C6H12O6 • Metabolism - all of the chemical reactions that occur in an organism
  32. 32. • Activation Energy – the amount of energy needed to start a chemical reaction • Catalysts – chemical substances that reduce the amount of activation energy that is needed for a reaction to begin
  33. 33. • Basically, by reducing the activation energy, catalysts speed up reactions and cause reactions to occur. Without them, reactions are very slow or never occur at all!!
  34. 34. • Catalysts in living things are called enzymes • Enzymes bind temporarily to one or more of the reactants of the reaction. This lowers the amount of activation energy needed and speeds up the reaction. • Enzymes are VERY specific (a different enzyme for every reaction!) – Ex. Sucrase catalyzes the breakdown of sucrose into glucose and fructose • Enzymes – speed up reactions without being permanently changed or destroyed
  35. 35. • Living things are mostly water • Most reactions in living things occur in water solutions • Water has several unique properties that make it one of the most important compounds found in living things • A solution is a mixture in which one or more substances are dissolved in another • Solvents and solutes make up solutions 1. Solute: part of a solution that is dissolved (Splenda) 2. Solvent: part of the solution that material is dissolved in (coffee)
  36. 36. • Liquids, Solids, and Gases can all be used in solutions.  Solutions can be solids dissolved in liquids Ex. Salt water  Solutions can be gases dissolved in liquids Ex. carbonated beverages  Solutions can be made from two solids Ex. Brass is a solution containing copper and zinc .
  37. 37. • Solutions can have varying amount of solute dissolved in varying amounts of solvent  Concentration - A measurement of the amount of solute dissolved in a fixed amount of solvent  2% salt solution = 2g of salt in 100mL of water  12% salt solution = 12g of salt in 100mL of water  The more solute dissolved in solution, the higher the concentration  Saturated solution - A solution where no more solute can be dissolved in the solvent
  38. 38. • Aqueous solutions are solutions that have water as the solvent…(aq) • Aqueous solutions are universally important to living things. • Fish depend on oxygen dissolved in water to survive. • Most nutrients plants need are in aqueous solutions in moist soil. • Body cells exist in an aqueous solution and are filled with aqueous solution.
  39. 39. • Water is made of 2 hydrogen atoms and 1 oxygen atom (H20) • They are covalently bonded so they share electrons. • The oxygen is “greedy” and pulls the electrons closer to it.  The oxygen is therefore a little bit negative, and the hydrogens are a little bit positive.  Because of this uneven distribution of charge, water is called a polar molecule.
  40. 40. • The polar nature of water allows it to dissolve polar substances, such as sugars, ionic compounds, and some proteins. • Water does not dissolve nonpolar substances, such as fats like oil.
  41. 41. • The polar nature of water causes water molecules to be attracted to one another. • Opposites attract: the oxygen of one molecule is attracted to a hydrogen of another. This attraction results in a hydrogen bond.
  42. 42. • Hydrogen bonds in water exert an attractive force strong enough that water “sticks” to itself and other substances!  Cohesion - An attractive force that holds molecules of a single substance together. - Ex: water molecules stick to each other  Cohesion causes “surface tension” or a thin “skin” on the surface of water.  Adhesion - An attractive force between two particles of different substances. - Ex: water molecules stick to glass molecules  Adhesion causes capillarity, which results in the rise of the surface of a liquid when in contact with a solid.
  43. 43. • Because of its hydrogen bonds, water has a high heat capacity, which means that water can absorb or release large amounts of energy in the form of heat with only a slight change in its temperature. • During a hot day, water can absorb heat (hydrogen bonds break) and cool the air. At night, the water cools (hydrogen bonds reform) and releases heat into the air. • Earth’s oceans stabilize global temperatures enough for life to exist. • Water’s high heat capacity allows your cells to keep an even temperature despite changes in the environment.
  44. 44. • Solid water is less dense than liquid water. – This is opposite of all other substances • Hydrogen bonding causes ice crystals to have large amounts of open space.
  45. 45. • When bodies of water freeze, they freeze from the top down and not the bottom up. • Ice insulates the water below from the cold air, which allows fish and other aquatic animals to survive under the icy surface.
  46. 46. • The alkalinity or acidity of a solution can determine the survival or death of organisms! • What do we mean when we say acidic and alkaline (basic)?
  47. 47. • Water molecules bump into each other and can actually break each other apart! This results in a hydroxide and hydrogen ion. • An ion is any atom(s) that have a positive or negative charge – Hydroxide ion is the OH– Hydrogen ion is the H+ • The hydrogen ion (H+) can become attracted to the oxygen in another water molecule resulting in a hydronium ion. – Hydronium ion is the H3O+ • This process is called the ionization of water
  48. 48. • The relative concentrations of hydronium (H 3O+) and hydroxide (OH-) ions in a solution determines if it is an acid or a base. • Pure water contains an equal number of both, so it is a NEUTRAL solution.
  49. 49. • Some compounds, when they dissolve in water will separate and form H+ ions. These compounds are called acids. • H+ will react with H20 to form H30+ – When HCl gas dissolves in water, it breaks up into H+ and Cl– H+ reacts with H20 to form H3O+. There is now more H3O+ ions than OH- ions. • There are always more H3O+ than OH- in acidic solutions. Some common examples: Hydrochloric Acid = HCl Sulfuric Acid = H2SO4 Nitric Acid = HNO3 Acetic Acid = HCH3OO Phosphoric Acid = H3PO4
  50. 50. • Called acidic • Acids tend to taste sour • Concentrated acids are very corrosive – HCl in your stomach helps to breakdown and digest proteins • Examples: Orange juice, vinegar, soda • Drinking acidic drinks over a long period of time can erode the tooth enamel!
  51. 51. • Some compounds, when they dissolve in water and separate will form OH- ions. These compounds are called Bases. • When the solid NaOH dissolves in water, it breaks up into Na+ and OH• There are now more OH- than H3O+ ions. • There are always more OH- than H3O+ in basic solutions. Some common examples: Barium Hydroxide = Ba(OH)2 Sodium Hydroxide = NaOH Potassium Hydroxide = KOH Calcium Hydroxide = Ca(OH)2
  52. 52. • Called alkaline • Slippery sensation when touched – Bases react with oils in skin to form soap • Bases have a bitter taste • Examples: Soap, antacids (Magnesium hydroxide and Aluminum hydroxide), ammonia
  53. 53. • Scientists have developed a scale for comparing hydronium and hydroxide ions in solution. • pH scale – Measures from 0 to 14 – Below 7 is acidic – Above 7 is basic – 7 is neutral – Logarithmic scale (tenfold change per number) – pH of 4 is 100 times more acidic (H3O+) than a pH of 6
  54. 54. • How much more acidic is vinegar compared to urine? • Difference of 3 pH levels = 10 x 10 x 10 = 1000 times more acidic!
  55. 55. Acid + Base = A Neutral Solution (water and a salt) Acid Base HCl + NaOH HBr + KOH Water H2O + H2O + Salt NaCl KBr
  56. 56. • Living systems are very sensitive to pH because enzymes can only function in very specific pH ranges. Uh oh, I think we’re gonna be in pH 4 soon! But how will we function? That is not acidic enough, Help!
  57. 57. • Living things use buffers to prevent pH from changing too much. • Buffers are chemical substances that neutralize small amounts of acids or bases in a solution. • Complex buffering systems keep you body’s pH values at the right level. • pH varies by body system  The pH of the human stomach is usually between 2 and 3.  Ideally, the pH of the blood should be maintained at 7.4. If the pH drops below 6.8 or rises above 7.8, death may occur.

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