Chemistry Comes Alive

Transcription

1 Chapter 2 Part A Chemistry Comes Alive Annie Leibovitz/Contact Press Images PowerPoint Lecture Slides prepared by Karen Dunbar Kareiva Ivy Tech Community College

2 Why This Matters: Understanding chemistry and biochemistry helps to determine the most effective solutions to use to treat dehydration and fluid loss

3 Video: Why This Matters

4 Chemistry and Physiological Reactions Body is made up of many chemicals Chemistry underlies all physiological reactions: Movement, digestion, pumping of heart, nervous system Chemistry can be broken down into: Basic chemistry Biochemistry

5 Part 1 Basic Chemistry 2.1 Matter and Energy Matter Matter is anything that has mass and occupies space Matter can be seen, smelled, and/or felt Weight is mass plus the effects of gravity

6 Matter States of matter Matter can exist in three possible states: Solid: definite shape and volume Liquid: changeable shape; definite volume Gas: changeable shape and volume

7 Energy Energy is the capacity to do work or put matter into motion Energy does not have mass, nor does it take up space The greater the work done, the more energy it uses up

8 Animation Energy Concepts

9 Energy (cont.) Kinetic versus potential energy Energy exists in two possible forms: Kinetic energy in action Potential stored (inactive) energy Energy can be transformed from potential to kinetic energy Stored energy can be released, resulting in action

10 Energy (cont.) Forms of energy Chemical energy Stored in bonds of chemical substances Electrical energy Results from movement of charged particles Mechanical energy Directly involved in moving matter Radiant or electromagnetic energy Travels in waves (example: heat, visible light, ultraviolet light, and X rays)

11 Energy (cont.) Energy form conversions Energy may be converted from one form to another Example: turning on a lamp converts electrical energy to light energy Energy conversion is inefficient Some energy is lost as heat, which can be partly unusable energy

12 2.2 Atoms and Elements All matter is composed of elements Elements are substances that cannot be broken down into simpler substances by ordinary chemical methods Four elements make up 96% of body: Carbon, oxygen, hydrogen, and nitrogen 9 elements make up 3.9% of body 11 elements make up <0.01% Periodic table lists all known elements

13 2.2 Atoms and Elements All elements are made up of atoms, which are: Unique building blocks for each element Smallest particles of an element with properties of that element What give each element its particular physical & chemical properties

14 2.2 Atoms and Elements Atomic symbol One- or two-letter chemical shorthand for each element Example: O for oxygen, C for carbon Some symbols come from Latin names: Na (natrium) is sodium; K (kalium) is potassium

15 Table Common Elements Composing the Human Body

16 Table Common Elements Composing the Human Body (continued)

17 Table Common Elements Composing the Human Body (continued)

18 Structure of Atoms Atoms are composed of three subatomic particles: Protons Carry a positive charge (+) Weigh an arbitrary 1 atomic mass unit (1 amu) Neutrons Have no electrical charge (0) Also weigh 1 amu Electrons Carry a negative charge ( ) Are so tiny they have virtually no weight (0 amu)

19 Structure of Atoms (cont.) Number of positive protons is balanced by number of negative electrons, so atoms are electrically neutral Protons and neutrons are found in a centrally located nucleus; electrons orbit around the nucleus Chemists devise models of how subatomic particles are put together Planetary model Orbital model

20 Structure of Atoms (cont.) Planetary model: simplified and outdated because it incorrectly depicts electrons in orbits, fixed circular paths Still useful for illustrations Orbital model: current model used that depicts orbitals, probable regions where an electron is most likely to be located (rather than fixed orbits) Shading in regions of greatest electron density results in an electron cloud around nucleus Useful for predicting chemical behavior of atoms

21 Figure 2.1 Two models of the structure of an atom. Nucleus Nucleus Helium atom 2 protons (p + ) 2 neutrons (n 0 ) 2 electrons (e ) Planetary model Helium atom 2 protons (p + ) 2 neutrons (n 0 ) 2 electrons (e ) Orbital model Proton Neutron Electron Electron cloud

22 Identifying Elements Different elements contain different numbers of subatomic particles Hydrogen has 1 proton, 0 neutrons, and 1 electron Helium has 2 protons, 2 neutrons, and 2 electrons Lithium has 3 protons, 4 neutrons, and 3 electrons Identifying facts about an element include its atomic number, mass number, isotopes, and atomic weight

23 Figure 2.2 Atomic structure of the three smallest atoms. Proton Neutron Electron Hydrogen (H) (1p + ; 0n 0 ; 1e ) Helium (He) (2p + ; 2n 0 ; 2e ) Lithium (Li) (3p + ; 4n 0 ; 3e )

24 Identifying Elements (cont.) Atomic number Number of protons in nucleus Written as subscript to left of atomic symbol Example: 3 Li Mass number Total number of protons and neutrons in nucleus Total mass of atom Written as superscript to left of atomic symbol Example: 7 Li

25 Identifying Elements (cont.) Isotopes Structural variations of same element Atoms contain same number of protons but differ in the number of neutrons they contain Atomic numbers are same, but mass numbers different Atomic weight Average of mass numbers of all isotope forms of an atom

26 Figure 2.3 Isotopes of hydrogen. Proton Neutron Electron Hydrogen ( 1 H) (1p + ; 0n 0 ; 1e ) Deuterium ( 2 H) (1p + ; 1n 0 ; 1e ) Tritium ( 3 H) (1p + ; 2n 0 ; 1e )

27 Radioisotopes Radioisotopes are isotopes that decompose to more stable forms Atom loses various subatomic particles Sometimes loss results in an isotope becoming a different element As isotope decays, subatomic particles that are being given off release a little energy This energy is referred to as radioactivity Can be detected and measured with scanners

28 Radioisotopes (cont.) Radioisotopes are a valuable tool for biological research and medicine Share same chemistry as their stable isotopes so will be taken up by body Can then be used for diagnosis of disease All radioactivity can damage living tissue Some types can be used to destroy localized cancers Some types cause cancer Radon from uranium decay causes lung cancer

29 2.3 Combining Matter Molecules and Compounds Most atoms chemically combine with other atoms to form molecules and compounds Molecule: general term for 2 or more atoms bonded together Compound: specific molecule that has 2 or more different kinds of atoms bonded together Example: C 6 H 12 O 6 Molecules with only one type of atom (H 2 or O 2 ) are just called molecules

30 Mixtures Most matter exists as mixtures: two or more components that are physically intermixed Three basic types of mixtures Solutions Colloids Suspensions

31 Figure 2.4 The three basic types of mixtures. Solution Colloid Suspension Solute particles are very tiny, do not settle out or scatter light. Solute particles are larger than in a solution and scatter light; do not settle out. Solute particles are very large, settle out, and may scatter light. Solute particles Solute particles Solute particles Example Example Example Mineral water Jell-O Blood Unsettled Settled Plasma Settled red blood cells

32 Mixtures (cont.) Solutions Are homogeneous mixtures, meaning particles are evenly distributed throughout Solvent: substance present in greatest amount Usually a liquid, such as water Solute(s): substance dissolved in solvent Present in smaller amounts Example: blood sugar glucose is solute, and blood (plasma) is solvent

33 Mixtures (cont.) Solutions (cont.) True solutions are usually transparent Example: air (gas solution), salt solution, sugar solution Most solutions in body are true solutions of gases, liquids, or solids dissolved in water

34 Mixtures (cont.) Concentration of true solutions Three common ways to express concentrations: 1. Percent of solute in total solution How many parts of solute are in 100 total parts of solution Solvent is usually water Example: 10 parts salt to 90 parts water is a 10% salt solution 2. Milligrams per deciliter (mg/dl) Deciliter equals 1/100th of a liter Example: normal fasting blood glucose levels are around 80 mg/dl

35 Mixtures (cont.) 3. Molarity (M) is number of moles of solute per liter of solvent (water) 1 mole of a compound is equal to its molecular weight (sum of atomic weights) in grams Example: glucose (C 6 H 12 O 6 ) has a molecular wt of amu, so grams of glucose added to enough H 2 O to make 1 liter is a 1 M solution of glucose 1 mole of any substance always contains molecules of that substance This number is called Avogadro s number Molarities in the body are so small (can be M), they are expressed in millimoles (mm) so 1000 mm = 1 M

36 Mixtures (cont.) Colloids Also known as emulsions; are heterogeneous mixtures, meaning that particles are not evenly distributed throughout mixture Can see large solute particles in solution, but these do not settle out Gives solution a cloudy or milky look Some undergo sol-gel (solution to gel) transformations Example: Jell-O goes from liquid to gel Cytosol of cell is also a sol-gel type solution

37 Mixtures (cont.) Suspensions Heterogeneous mixtures that contain large, visible solutes that do settle out Example: mixture of water and sand Blood is considered a suspension because if left in a tube, the blood cells will settle out

38 Difference Between Mixtures and Compounds Three main differences: Unlike compounds, mixtures do not involve chemical bonding between components Mixtures can be separated by physical means, such as straining or filtering; compounds can be separated only by breaking their chemical bonds Mixtures can be heterogeneous or homogeneous; compounds are only homogeneous

39 2.4 Chemical Bonds Chemical bonds are energy relationships between electrons of reacting atoms Chemical bonds are not actual physical structures Electrons are the subatomic particles that are involved in all chemical reactions They determine whether a chemical reaction will take place and if so, what type of chemical bond is formed

40 Role of Electrons in Chemical Bonding Electrons can occupy areas around nucleus called electron shells Each shell contains electrons that have a certain amount of kinetic and potential energy, so shells are also referred to as energy levels Depending on its size, an atom can have up to 7 electron shells Shells can hold only a specific number of electrons; the shell closest to nucleus is filled first Shell 1 can hold only 2 electrons Shell 2 holds a maximum of 8 electrons Shell 3 holds a maximum of 18 electrons

41 Role of Electrons in Chemical Bonding (cont.) Outermost electron shell is called valence shell Electrons in valence shell have the most potential energy because they are farthest from nucleus These are electrons that are involved in chemical reactions

42 Role of Electrons in Chemical Bonding (cont.) Octet rule (rule of eights) Atoms desire 8 electrons in their valence shell Exceptions: smaller atoms (examples: H and He) want only 2 electrons in shell 1 Desire to have 8 electrons is driving force behind chemical reactions Noble gases already have full 8 valence electrons (or 2 for He) so are not chemically reactive Most atoms do not have full valence shells Atoms will gain, lose, or share electrons (form bonds) with other atoms to achieve stability of 8 electrons in valence shell

43 Figure 2.5a Chemically inert and reactive elements. Chemically inert elements Outermost energy level (valence shell) complete He 2e Ne 2e 8e Helium (He) (2p + ; 2n 0 ; 2e ) Neon (Ne) (10p + ; 10n 0 ; 10e )

44 Figure 2.5b Chemically inert and reactive elements. Chemically reactive elements Outermost energy level (valence shell) incomplete H 1e C 4e 2e Hydrogen (H) (1p + ; 0n 0 ; 1e ) Carbon (C) (6p + ; 6n 0 ; 6e ) O 2e 6e Na 8e 1e 2e Oxygen (O) (8p + ; 8n 0 ; 8e ) Sodium (Na) (11p + ; 12n 0 ; 11e )

45 Types of Chemical Bonds Three major types of chemical bonds Ionic bonds Covalent bonds Hydrogen bonds

46 Types of Chemical Bonds (cont.) Ionic bonds Ions are atoms that have gained or lost electrons and become charged Number of protons does not equal number of electrons

47 Types of Chemical Bonds (cont.) Ionic bonds involve the transfer of valence shell electrons from one atom to another, resulting in ions One becomes an anion (negative charge) Atom that gained one or more electrons One becomes a cation (positive charge) Atom that lost one or more electrons Attraction of opposite charges results in an ionic bond

48 Figure 2.6ab Formation of an ionic bond. + Na Cl Na Cl Sodium atom (Na) (11p + ; 12n 0 ; 11e ) Chlorine atom (Cl) (17p + ; 18n 0 ; 17e ) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron. Sodium ion (Na + ) Chloride ion (Cl ) Sodium chloride (NaCl) After electron transfer, the oppositely charged ions formed attract each other.

49 Types of Chemical Bonds (cont.) Most ionic compounds are salts When dry, salts form crystals instead of individual molecules Example is NaCl (sodium chloride)

50 Figure 2.6c Formation of an ionic bond. Cl Na + Large numbers of Na + and Cl ions associate to form salt (NaCl) crystals.

51 Types of Chemical Bonds (cont.) Covalent bonds Covalent bonds are formed by sharing of two or more valence shell electrons between two atoms Sharing of 2 electrons results in a single bond Sharing of 4 electrons is a double bond Sharing of 6 electrons is a triple bond Allows each atom to fill its valence shell at least part of the time Two types of covalent bonds: Polar and nonpolar covalent bonds

52 Figure 2.7a Formation of covalent bonds. Reacting atoms Resulting molecules H H H C H C H or H H H Structural formula shows single bonds. Hydrogen atoms Carbon atom Molecule of methane gas (CH 4 ) Formation of four single covalent bonds: Carbon shares four electron pairs with four hydrogen atoms.

53 Figure 2.7b Formation of covalent bonds. Reacting atoms Resulting molecules O O O O or Structural formula shows double bond. Oxygen atom Oxygen atom Molecule of oxygen gas (O 2 ) Formation of a double covalent bond: Two oxygen atoms share two electron pairs.

54 Figure 2.7c Formation of covalent bonds. Reacting atoms Resulting molecules N N N N or Structural formula shows triple bond. Nitrogen atom Nitrogen atom Molecule of nitrogen gas (N 2 ) Formation of a triple covalent bond: Two nitrogen atoms share three electron pairs.

55 Types of Chemical Bonds (cont.) Covalent bonds (cont.) Nonpolar covalent bonds Equal sharing of electrons between atoms Results in electrically balanced, nonpolar molecules such as CO 2

56 Figure 2.8a Carbon dioxide and water molecules have different shapes, as illustrated by molecular models. Carbon dioxide (CO 2 ) molecules are linear and symmetrical. They are nonpolar.

57 Types of Chemical Bonds (cont.) Polar covalent bonds Unequal sharing of electrons between 2 atoms Results in electrically polar molecules Atoms have different electron-attracting abilities, leading to unequal sharing Atoms with greater electron-attracting ability are electronegative, and those with less are electropositive

58 Types of Chemical Bonds (cont.) Polar covalent bonds (cont.) H 2 O is a polar molecule Oxygen is more electronegative, so it exerts a greater pull on shared electrons, giving it a partial negative charge and giving H a partial positive charge Having two different charges is referred to as dipole

59 Figure 2.8b Carbon dioxide and water molecules have different shapes, as illustrated by molecular models. d d + d + V-shaped water (H 2 O) molecules have two poles of charge a slightly more negative oxygen end (d ) and a slightly more positive hydrogen end (d + ).

60 Figure 2.9 Ionic, polar covalent, and nonpolar covalent bonds compared along a continuum. Ionic bond Polar covalent bond Nonpolar covalent bond Complete transfer of electrons Unequal sharing of electrons Equal sharing of electrons Separate ions (charged particles) form Slight negative charge (d ) at one end of molecule, slight positive charge (d + ) at other end Charge balanced among atoms Na + Cl Sodium chloride Water Carbon dioxide

61 Types of Chemical Bonds (cont.) Hydrogen bonds Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule Not true bond, more of a weak magnetic attraction Common between dipoles such as water What makes water liquid Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape

62 Animation Hydrogen Bonds

63 Figure 2.10a Hydrogen bonding between polar water molecules. H d + O H d d + Hydrogen bond (indicated by dotted line) d O H H d + d O H H d d + The slightly positive ends (d + ) of the water molecules become aligned with the slightly negative ends (d ) of other water molecules. d + d O H H d +

64 Figure 2.10b Hydrogen bonding between polar water molecules. A water strider can walk on a pond because of the high surface tension of water, a result of the combined strength of its hydrogen bonds.

65 2.5 Chemical Reactions Chemical Equations Chemical reactions occur when chemical bonds are formed, rearranged, or broken These reactions can be written in symbolic forms called chemical equations Chemical equations contain: Reactants: substances entering into reaction together Product(s): resulting chemical end products Amounts of reactants and products are shown in balanced equations

66 Chemical Equations (cont.) Compounds are represented as molecular formulas Example: H 2 O or C 6 H 12 O 6 Subscript indicates atoms joined by bonds Prefix denotes number of unjoined atoms or molecules

67 Chemical Equations (cont.) Compounds are represented as molecular formulas Example: H 2 O or C 6 H 12 O 6 or H 2 or CH 4 In chemical equations, subscripts indicate how many atoms are joined by bonds, whereas prefix means number of unjoined atoms (example: 4H) Reactants Product H + H H 2 (Hydrogen gas) 4H + 1C CH 4 (Methane)

68 Types of Chemical Reactions Three main types of chemical reactions: 1. Synthesis (combination) reactions involve atoms or molecules combining to form larger, more complex molecule Used in anabolic (building) processes A + B AB

69 Figure 2.11a Types of chemical reactions. Synthesis reactions Smaller particles are bonded together to form larger, more complex molecules. Example Amino acids are joined together to form a protein molecule. Amino acid molecules Protein molecule

70 Types of Chemical Reactions (cont.) 2. Decomposition reactions involve breakdown of a molecule into smaller molecules or its constituent atoms (reverse of synthesis reactions) Involve catabolic (bond-breaking) reactions AB A + B

71 Figure 2.11b Types of chemical reactions. Decomposition reactions Bonds are broken in larger molecules, resulting in smaller, less complex molecules. Example Glycogen is broken down to release glucose molecules. Glycogen Glucose molecules

72 Types of Chemical Reactions (cont.) 3. Exchange reactions, also called displacement reactions, involve both synthesis and decomposition Bonds are both made and broken AB + C AC + B and AB + CD AD + CB

73 Figure 2.11c Types of chemical reactions. Exchange reactions Bonds are both made and broken (also called displacement reactions). P P P Adenosine triphosphate (ATP) Glucose P P P Adenosine diphosphate (ADP) Example ATP transfers its terminal phosphate group to glucose to form glucosephosphate. Glucosephosphate

74 Types of Chemical Reactions (cont.) In living systems, these reactions are also referred to as reduction-oxidation or redox reactions Atoms are reduced when they gain electrons and oxidized when they lose electrons Example: C 6 H 12 O 6 + 6O 2 6CO 2 + 6H 2 O + ATP In this example, glucose is oxidized, and oxygen molecule is reduced

75 Energy Flow in Chemical Reactions All chemical reactions are either exergonic or endergonic Exergonic reactions result in a net release of energy (give off energy) Products have less potential energy than reactants Catabolic and oxidative reactions Endergonic reactions result in a net absorption of energy (use up energy) Products have more potential energy than reactants Anabolic reactions

76 Reversibility of Chemical Reactions All chemical reactions are theoretically reversible A + B AB Chemical equilibrium occurs if neither a forward nor a reverse reaction is dominant Many biological reactions are not very reversible Energy requirements to go backward are too high, or products have been removed

77 Rate of Chemical Reactions The speed of chemical reactions can be affected by: Temperature: increased temperatures usually increase rate of reaction Concentration of reactants: increased concentrations usually increase rate Particle size: smaller particles usually increase rate

78 Rate of Chemical Reactions Catalysts Catalysts increase the rate of reaction without being chemically changed or becoming part of the product Enzymes are biological catalysts