Chemical Principles

1. Chemical Principles Chapter 2

2. The Structure of Atoms

3. Atoms and Compounds • Matter is made up of atoms • Atoms cannot be broken down any further by normal chemical reactions • 92 natural elements recognized by chemists • Gold (Au), Copper (Cu), Carbon (C), Oxygen (O) • Molecules are two or more atoms combined in a fixed ratio • Table salt (NaCl), Water (H2O) • Properties of compounds are different than its individual atoms

4. Subatomic Particles • Atomic Nucleus • Protons (positive charge; mass of 1 Dalton) • Neutrons (no charge; mass of 1 Dalton) • Orbitals are where an atom’s electrons are found 90% of the time • Electrons (negative charge; no mass) Carbon

5. Atomic Mass and Atomic Number • Atomic Number denotes the number of protons • Neutral (uncharged) atoms have equal number of protons and electrons • Mass Number denotes the number of protons and neutrons in the nucleus • Atomic Mass ≈ Mass Number • Also referred to as atomic weight

6. The Elements of Life • Elements are atoms with the same chemical behavior • Found on periodic table • C, H, O, and N makes up 96% of all living matter • Essential Elements are required by all organisms • Trace Elements are only required in small amounts • Often used as cofactors for enzymes

7. Isotopes • Possess more neutrons than normally present • Alters the atomic mass • Many isotopes are stable • Radioactive Isotopes are inherently unstable • Gives off particles and energy • Often used to tag proteins, to track the movement of fluids in the body, to identify tissues (PET scan)

8. Energy Levels • Energy is the capacity to cause change • Potential Energy is a type of energy due to position or structure • Electrons have potential energy due to distance from atomic nucleus • Electron Shells

9. Electron Configuration • Valence Electrons are the outermost electrons in an atom’s orbital • Valence Number is 2 for the first orbital and 8 for the second and third orbital • Valence Shell is the outermost orbital of a particular atom • An atom’s chemical behavior is determined by the interactions of its valence electrons

10. Check Your Understanding • Describe the structure of an atom and its relation to the physical properties of elements. • What are isotopes and how are they sometimes used in medicine? • What are electron orbitals? What is the valence number of an atom? What is the relationship between these?

11. How atoms form molecules: chemical bonds

12. Chemical Bonds • Attractions that keep atoms close together • Due to interactions between electrons of neighboring atoms • Forms compounds when the molecules are made of two or more different elements • Strongest chemical bonds in aqueous solutions are covalent bonds • Typical of biological compounds • Ionic Bonds are generally strong when NOT present in aqueous solutions

13. Electronegativity • Strength of attraction for a given atom to the electrons in a chemical bond • Determines if electrons are shared, how they are shared, or if they are lost completely

14. Covalent Bonds • Sharing of a pair of valence electrons between 2 atoms • Single Bonds • Double Bonds • Molecules are created by a covalent bond between 2 or more atoms

15. Types of Covalent Bonds Nonpolar Covalent Polar Covalent Equal Sharing Unequal Sharing

16. Ionic Bonds • One atom is so electronegative, it takes an electron from its partner • Result is one negatively-charged and one positively-charged atom (Ion) • Cationsare positively charged • Anions are negatively charged • Resulting ions are held together by electrostatic interactions • Often results in salts such as NaCl, KCl, MgCl2

17. Hydrogen Bonds • Occur between polar molecules • Hydrogen of one molecule is attracted to electronegative atom of another

18. Check Your Understanding • Differentiate between covalent and ionic bonds. • Differentiate between a polar covalent and a non-polar covalent bond. • What is a hydrogen bond and how are they different from covalent or ionic bonds?

19. Chemical Reactions

20. Chemical Reactions • Making and breaking of chemical bonds • Reactants are the starting materials • Products are the end materials • Many biological reactions are Reversible • Most reactions occur until an Equilibrium is reached Products Reactants

21. Energy in Chemical Reactions • Exergonic reactions occur with a net release of energy • Products have less free energy than the reactants • Endergonic reactions require energy be added to the system before the reaction can occur • Products have more free energy than the reactants

22. Synthesis Reactions • Result in the formation of larger molecules by joining two or more smaller ones together • Also called anabolic reactions A + B AB

23. Decomposition Reactions • Break down larger molecules into two or more smaller molecules • Also referred to as catabolic reactions AB A + B

24. Exchange Reactions • Occur when reactants interact to form new molecules NaOH + HClNaCl + H2O

25. Check Your Understanding • Differentiate between exergonic and endergonic reactions. • Differentiate between synthesis, decomposition, and exchange reactions.

26. Important Biological molecules

27. Inorganic Compounds

28. Water is a Polar Molecule • Oxygen atoms are more electronegative than hydrogen atoms • Electrons spend more time around oxygen than hydrogen • Creates a polar covalent bond between the atoms

29. Water Forms Hydrogen Bonds • All polar molecules can form hydrogen bonds with other polar molecules • Hydrogen bonds are constantly in flux • Form, break, and reform constantly in aqueous solutions • Gives water its unique properties

30. Water is An Excellent Solvent • Solutions are homogenous mixture of two substances • Solute is substance dissolved • Solvent is substance acting as the dissolving agent • Aqueous Solutions Solution = Solute + Solvent • Water is an excellent solvent for polar or ionic compounds

31. Dissociation of Water • Hydrogen Ions(H+) within hydrogen bonds may shift from one molecule to another resulting in a: • Hydroxide Ion (OH-) • Hydronium Ion (H3O+) • H+ are simply protons • Reversible reactions • Measured using pH scale

32. Acids • Increase H+concentrationof a solution • Protondonor HCl H+ + Cl- Hydrochloric Acid

33. Bases • Reduce H+ concentration of a solution • Proton acceptor NH3+ H+ NH4+ NaOH Na+ + OH- Ammonia Ammonium Sodium Hydroxide

34. Salts • Substances that dissolve into cations and anions, neither of which is H+ or OH-

35. pH Scale • [H+] and [OH-] always equal 10-14 at 25°C • pH is a logarithmic scale of hydrogen ion concentration [H+] of 10-7 = pH of 7 • pH declines as [H+] increases • Each pH unit represents a 10-fold change in [H+]

36. Buffers • Buffers minimize changes in [H+] and [OH-] of a solution • Consist of a weak acid and a base • Accept H+ from a solution when in excess and donate H+ into a solution when depleted H2O + CO2 H2CO3 HCO3- + H+ Water Carbon Dioxide Carbonic Acid Bicabonate

37. Check Your Understanding • List several properties of water that are important to living systems. • Why is the polarity of a water molecule important? • Define acid, base, salt, and pH. • What is a buffer and how do they work?

38. Organic Compounds

39. The Molecules of Life • Carbohydrates, Lipids, Proteins, and Nucleic Acids • Large, complex organic molecules • Macromolecules • Often created from the covalent bonding of many, smaller monomers to create a complex polymer • Carbohydrates, proteins, and nucleic acids

40. Macromolecules are polymers, built from monomers

41. Synthesis of Polymers • Reactions are facilitated by enzymes • Biological catalysts • Very often are proteins • Increase reaction rates • Monomers are joined by removing a molecule of H2O from between them • Dehydration Reaction • Forms a covalent bond

42. Breakdown of Polymers • Reactions are facilitated by enzymes • Polymers are broken apart by adding a molecule of H2O to the covalent bond • Hydrolysis Reaction • Releases individual monomers

43. Diversity of Polymers • Inherent differences among a single species is often the result of subtle variations in polymers • Complex carbohydrates, proteins, and nucleic acids may be several hundred to several thousand monomers in length • Proteins alone have 20 amino acid building blocks to create a given protein • Actual possible combinations may be almost limitless given a polymer’s size

44. Carbohydrates serve as fuel and building material

45. Monosaccharides • Single sugars with a molecular ratio of (CH2O) • Glucose is a crucial monosaccharide for energy • A hexose • Ribose anddeoxyribose are components of nucleotides • Each is a pentose

46. Alternate Forms of Monosaccharides Most monosaccharides form rings when placed in an aqueous solution

47. Dissacharides • Double sugars • Two monosaccharides linked by a glycosidic bond through the removal of H2O • Sucrose is common table sugar • Glucose + Fructose • Lactose is the sugar found in milk • Glucose + Galactose

48. Storage Polysaccharides • Starch: a polysaccharide used by plants to store sugars • Polymer of glucose • Amylose • Amylopectin • Glycogen: a polysaccharide used by animals to store sugars • Polymer of glucose

49. Structural Polysaccharides • Cellulose: polysaccharide used by plants for structure • Cell Walls • Polymer of glucose

50. Starch vs. Cellulose • Starch is held together via 1-4 α-linkages • Amylase: enzyme capable of hydrolyzing starch • Present in most animals • Cellulose is held together via 1-4 β-linkages • Cellulase: enzyme used to hydrolyze cellulose • Absent in most animals