1. Chapter 2: The Chemistry of Life BASIC CHEMISTRY:  Matter & Energy Matter & Energy  Chemical Reactions Chemical Reactions  Properties of Water Properties of Water  Acids & Bases Acids & Bases

2.  Matter  Refers to anything that takes up space and has mass  Can exist in a solid, liquid, or gaseous state  Is composed of elements The Composition of Matter

3. States of Matter  3 States:  Solid - Definite shape and volume Ex: Bones  Liquid - Definite volume only Ex: Plasma  Gas - No definite shape or volume Ex: O 2, CO 2  PHYSICAL CHANGES - Do NOT alter the nature of the substance, only the state. Ex: Ice melting  CHEMICAL CHANGES - Do alter the composition of the substance. Ex: Digestion of food

4.  Element – substance that cannot be broken down into another substance by ordinary chemical means  Only 92 naturally occurring elements  6 elements make up about 98% of the body weight of most living organisms – CHNOPS Elements

5.  Atomic Theory – states that elements consist of atoms  Atom – the simplest particle of an element that displays all the properties of the element  Atomic symbol – name of the atom or element Atomic Structure

6.  PROTONS – positive charge (+)  mass = 1 amu  located in nucleus of an atom  NEUTRONS – no charge (neutral)  mass = 1 amu  located in nucleus of an atom  ELECTRONS – negative charge (-)  mass = 0 amu (negligible mass)  occupy region around nucleus in orbital shells Subatomic Particles

9. Reading The Periodic Table 11 Na Sodium 22.99  Atomic No. is # of Protons (also equals the # of electrons in a neutral atom with no charge)  Mass number is equal to sum of protons and neutrons – electrons have about zero mass Element Name & Abbreviation  All atoms of an element have this same number of protons

10.  Isotopes – atoms of the same element that differ in the number of neutrons  Isotopes have the same number of protons but a different number of neutrons (different mass numbers)  Unstable and may decay, emitting radiation  Can be used as tracer – PET scan  Can cause damage to cells leading to cancer  Can be used to sterilize medical equipment Isotopes

11. Isotopes in Biology  All isotopes of an element have essentially the same chemical characteristics  Some are unstable and break down or decay  These are called radioisotopes  Isotopes of a given element are usually metabolized by an organism in a similar way  This makes radioisotopes (ex: 3 H, 14 C, 32 P) extremely useful research tools  Used to study age of fossils, DNA synthesis, sugar transport in plants, medical diagnoses, etc.

12. Pet Scan

13. High Levels of Radiation

14.  Electrons are constantly moving  Chemical properties of atoms are largely determined by the arrangement of their electrons  Each shell contains a certain number of electrons  It is useful to construct models of atoms with energy levels or electron shells Arrangements of electrons

15.  For atoms up through number 20  2 electrons fill first shell  8 electrons fill each additional shell  Octet rule for valence shell  Valence shell – outermost shell  If an atom has more than 2 shells, the outer shell is most stable with 8 electrons  Atoms can give up, accept, or share electrons to have 8 Arrangements of electrons

16. Bohr Models of Atoms

17. Atoms have Energy 18 e- 8 e- 2 e- Nucleus If the outermost shell is full, then the atom is stable or INERT (Ex: He) Electrons closest to the nucleus have the LEAST energy Electrons in the outermost shells have the MOST energy and tend to react with other atoms’ electrons VALENCE SHELL – outermost shell

18.  Elements combine depending on the number and arrangement of electrons in their orbitals  An atom is chemically stable when its highest level orbital is filled (satisfied) with the maximum number of electrons  At this point the atom will not react with other atoms II) How Atoms Combine

19.  Molecule – simplest part of a substance that retains all of the properties of that substance and can exist in a free state (ex: H 2 O or O 2 ) Molecules

20.  Compound – substance made up of atoms of two or more elements in fixed proportions  Chemical formula – shows the kinds and number of atoms of each element in a compound (ex: H 2 O)  Physical and chemical properties of a compound are different than the elements it is made of (ex: H 2 O) Compounds

21.  Chemical bonds – attractions holding atoms close together  Form molecules  Result of sharing or completely transferring valence electrons Chemical Bonds

22.  Covalent Bonds – form when two atoms share one or more pairs of electrons (ex: H 2 O)  Strongest type of bond Covalent Bonds

24.  Ion – an atom or molecule with an electrical charge (protons ≠ electrons)  Cations - (+) charge Ex: Na +  Anions - (-) charge Ex: Cl -  Ionic bond – formed when oppositely charged ions attract each other  Ex: NaCl Ionic Bonds

25.  Energy – the ability to do work  Takes several forms  Radiant energy (light)  Thermal energy (heat)  Chemical energy (in bonds)  Electrical energy  Mechanical energy (movement)  Most are significant to biology in some way III) Energy and Matter

26.  Chemical reaction – one or more substances change to produce one or more different substances  Energy is absorbed when chemical bonds are formed  Energy is released when chemical bonds are broken Chemical Reactions

27.  Exhibited in a chemical equation Ex: CO 2 + H 2 0  H 2 CO 3  Reactants are shown on the left  Products are shown on the right Chemical Reactions

28.  For most chemical reactions to begin, energy must be added to the reactants  Activation energy – amount of energy needed to get the reaction started  Catalyst – a chemical substance which reduces the amount of activation energy needed Getting Reactions Started

29.  The amazing biological usefulness of water is due to its chemical structure  Water is a polar molecule  Polar – having unequal distribution of charges IV) Water & Solutions

30.  In water’s covalent bond, the oxygen has more protons so has a greater electronegativity  Electronegativity – ability to attract electrons Polarity

31. 31 Nonpolar Covalent Bonds  Nonpolar covalent bonds – sharing of electrons is equal  One atom “wants” (with a specific intensity) to donate electron(s)  The other atom “wants” (with the same intensity) to receive electron(s)  The bond electrons will spend about equal time with both atoms

32. 32 Polar Covalent Bonds  Polar covalent bonds – sharing of electrons is unequal  One atom “wants” to donate or receive electron(s) with a specific intensity  The other atom “wants” to donate or receive electron(s) with a different intensity  The bond electrons will spend more time with one atom than the other

33. VSVS

34.  In Water - sharing of electrons by oxygen and hydrogen is not equal  The oxygen atom with more protons attracts the electrons closer therefore assumes a partial negative charge  The oxygen atom, which gets the most time with the electrons, will be slightly negative

35. Water’s Importance to Life  Life began in water  Single most important molecule on Earth  All organisms are 70-90% water  Water has unique properties that make it a life-supporting substance  Properties stem from the structure of the molecule

36. Structure of water  Hydrogen bonds – slightly positive hydrogen of one water molecule attracted to the slightly negative oxygen in another water molecule

37. Properties of water that support life:  Solvency  Cohesion and adhesion  High surface tension  High heat capacity  High heat of vaporization  Varying density

38. Water is a solvent  Due to polarity, water dissolves many substances making it the Universal Solvent  Hydrophilic – molecules attracted to water  Hydrophobic – molecules not attracted to (fear) water

39. Water is Cohesive and Adhesive  Cohesion – ability of water molecules to cling to each other due to hydrogen bonding  Adhesion – ability of water molecules to cling to other polar surfaces  Allows water to be excellent transport system both in and outside of living organisms  Contributes to water transport in plants

40. Water has a high surface tension  Mainly due to hydrogen bonding, water molecules at the surface cling more tightly to each other than to the air above

41. Water has a high heat capacity  Temperature of water rises and falls slowly – has a high specific heat  The many hydrogen bonds linking water molecules allow water to absorb heat without greatly changing its temperature

42. High heat of vaporization  Takes a great deal of energy to break hydrogen bonds for evaporation  Heat is dispelled as water evaporates

43. Water is less dense than ice  Unlike other substances, water expands as it freezes due to hydrogen bonds  Ice floats rather than sinks and acts as an insulator  It makes life possible in water

44. A Pond in Winter ice layer Protists provide food for fish. River otters visit ice-covered ponds. Aquatic insects survive in air pockets. Freshwater fish take oxygen from water. Common frogs and pond turtles hibernate.

45. IV) ACIDS & BASES  Contrast acids and bases, and discuss their properties  Convert the hydrogen ion concentration (moles per liter) of a solution to a pH value  Describe how buffers help minimize changes in pH

46. Ionization of Water  When water dissociates, it releases an equal number of ions.  Hydrogen ions (H + )  Hydroxide ions (OH - ) 

47. Acids  Acid – a substance that releases H + when dissolved in water.  Acid  H + + anion  Acidic solutions have a high H + concentration.  Ex)

48. Acids

49. Bases  Base – a substance that releases OH - when dissolved in water.  Base  OH - + cation  Basic solutions have a low H + concentration.  Ex)

50. Bases

51. pH and the pH scale  pH – a mathematical way of indicating the number of H + ions in a solution.  The pH scale is used to express acidity or basicity (alkalinity).

52. The pH Scale  The negative logarithm (base 10) of the hydrogen ion concentration [H + ] (expressed in moles per liter):  pH = −log 10 [H + ]  The negative logarithm corresponds to a positive pH value  Pure water has a hydrogen ion concentration of 0.0000001 (10 —7 mol/L)  Logarithm = − 7; pH is 7

53. Calculating pH Values and Hydroxide Ion Concentrations

54. pH of Solutions  Neutral Solution (pH 7)  Equal concentrations of hydrogen ions and hydroxide ions (concentration of each is 10 −7 mol/L)  Acidic Solution (pH <7)  Hydrogen ion concentration is higher than hydroxide ion concentration  Basic Solution (pH >7)  Hydrogen ion concentration is lower than hydroxide ion concentration

55. **pH of most plant and animal cells around 7.2 to 7.4

56. Buffers and pH  Buffer – substance that resists changes in pH when an acid or base is added  A buffering system includes a weak acid or a weak base  When H + is added to pure water at pH 7, pH goes down and water becomes acidic  A buffer may or counter by adding OH -  When OH - is added to pure water at pH 7, pH goes up and water becomes alkaline  A buffer may counter by adding H +

57. 57 Buffers in Biology  Health of organisms requires maintaining pH of body fluids within narrow limits  Buffers help maintain this pH level  Human blood normally 7.4 (slightly alkaline)  Many foods and metabolic processes add or subtract H + or OH - ions  Reducing blood pH to 7.0 results in acidosis  Increasing blood pH to 7.8 results in alkalosis  Both life threatening situations  Bicarbonate ion ( - HCO 3 ) in blood buffers pH to 7.4