Section 1 of Chapter 2

1. Chapter 2 Chemistry Section 1 of Chapter 2

2. Why study chemistry? • A clear understanding of chemistry is essential for the study of physiology. • This is because organ functions depends on cellular functions, which occur as a result of chemical reactions. Watson & Crick first proposed the double helix structure of DNA

3. Definitions Matter = Anything that has mass and takes up space (Solids, liquids, gasses) Biochemistry = Chemistry of living things Element = Fundamental substance of matter (e.g. Carbon, Hydrogen, Oxygen) Compound = Two or more different elements chemically bonded together (e.g. H2O = water, C6H12O6 = glucose) Molecule = two or more atoms chemically joined together. Molecules may be compounds (H2O = water molecule), or Molecules may be of the same element (H2= hydrogen molecule)

4. Our body consists of 11 bulk elements and 7 trace elements. • Bulk elements make up 99.9% of our body: • Hydrogen (H)Oxygen (O)Carbon (C) • Nitrogen (N)Sulfur (S)Magnesium (Mg) • Sodium (Na)Potassium (K)Calcium (Ca) • Chlorine (Cl)Phosphorus (P) • Trace elements make up less than 0.1% of our body: • Cobalt (Co)Zinc (Zn)Manganese (Mn) • Iron (Fe)Iodine (I)Copper (Cu) • Fluorine (F) Learn each bulk element and trace element along with their atomic symbols shown in parentheses

5. All elements are arranged onto a Periodic table

6. Atoms Atoms are the smallest particles of an element that still have the properties of that element. • Atoms are composted of 3 subatomic particles: • Proton – carries a single positive charge • Neutron– carries no electrical charge • Electron – carries a single negative charge An atom contains a central nucleus composed of protons and neutrons. Electrons orbit the nucleus.

7. Subatomic Particles Electrical Charge: Proton: +1 charge. Electron: -1 charge. Neutron: 0 charge Atomic Mass: Proton: 1 dalton Neutron: 1 dalton Electron: 0 • Most atoms contain equal number of protons and electrons, so an atom contains no overall net charge and is neutral.

8. Subatomic Particles Atomic Number: The number of protons in one atom. Atomic number identifies an element. Example. The atomic number of oxygen is 8. Oxygen, and only oxygen has 8 protons. • Atomic Weight: The sum of protons and neutrons in one atom. Remember, the weight of electrons is negligible.

9. Examples of atomic numbers and atomic weight Hydrogen has: 1 proton and 1 electron Atomic number = 1 Atomic weight = 1 Carbon has: 6 protons, 6 neutrons, and 6 electrons Atomic number = 6 Atomic weight = 12

10. Table 2.1 Some Particles of Matter

11. Isotopes • Isotopes are atoms with the same atomic number, but different atomic weights. Isotopes occur because the number of neutrons of an element varies between atoms. Two isotopes of oxygen: Oxygen 16 (O16) Oxygen 17 (O17) protons: 8 protons: 8 electrons: 8 electrons: 8 neutrons: 9 neutrons: 8 8 Atomic Number: 8 17 Atomic Weight: 16 • *The atomic weight of an element is an average of the isotopes present.

12. Understand the notations on a periodic table. End of Section 1, Chapter 2

13. Section 2 of Chapter 2 Bonding of Atoms

14. Properties of electrons Electron Shells: Electrons encircle the nucleus in discrete orbits, called electron shells. Each shell can contain only a fixed number of electrons. 1st shell holds 2 electrons 2nd shell holds 8 electrons 3rd shell holds 8 electrons Octet rule: Except for the 1st shell, each electron shell holds up to 8 electrons * Lower shells are filled first.

15. Examples of filling electron shells Helium Atomic number = 2 Atomic weight = 4 (2 electrons fill the 1st electron shell) Carbon Atomic number = 6 Atomic weight = 12 (The first 2 electrons fill the inner shell, and the remaining 4 electrons are placed the 2nd electron shell).

16. Ions Ions are atoms that readily gain or loose electrons • Cation: an ion that looses electrons • Cations are positively charged ions • Anion: an ion that gains electrons • Anions are negatively charged ions

17. Example of a cation Sodium (Na) atomic number = 11 atomic weight = 23 Na+ = Sodium cation Only 1 lone electron sits in the outer shell. This electron is unpaired and is easily lost, forming the sodium cation.

18. Example of an anion Chlorine (Cl) atomic number = 17 atomic weight = 35 Cl- = Chloride anion 7 electrons fill the outer shell of chlorine, leaving room for 1 more electron. Chlorine readily accepts one electron, creating the chloride anion.

19. Ionic Bond Ionic bonds are formed when the oppositely charged particles attract. Figure 2.4 (a) An ionic bond forms when on atom gains and another atom looses electrons, and then (b) oppositely charged ions attract. Na+ + Cl- → NaCl

20. Ionic bonds do not form molecules Cations & anions attract in all directions, forming organized arrays, such as crystals. They do not form molecules. Figure 2.4 (c) salt crystal formation occurs because of the ionic bonds of sodium and potassium.

21. Covalent Bonds Covalent bonds are formed when atoms share electrons. Example: A hydrogen molecule (H2) is formed when two hydrogen atoms share their single electron. H + H H2

22. Covalent Bonds of water Water consist of oxygen covalently bonded to two hydrogen atoms. O H H Structural Formula: depicts the covalent bonds of a molecule as lines. Molecular Formula: is a shorthand notation for representing molecules. H2O

23. Types of covalent bonds A Single Bond occur when two atoms share one pair of electrons Oxygen joined to two hydrogen atoms by single bonds A Double Bond occurs when atoms are joined by two pairs of electrons Two oxygen atoms joined by a double bond. A Triple Bond occurs when atoms are joined by three pairs of electrons A Carbon atom joined to hydrogen by a single bond and to nitrogen by a triple bond.

24. Nonpolar covalent bonds Nonpolar covalent bonds occur when the atoms share the electrons equally, so the molecule has no overall charge. Two hydrogen atoms share their electrons equally. Thus, the hydrogen molecule has no overall charge and is nonpolar.

25. Polar covalent bonds Polar bonds have an unequal distribution of electrons. One portion of the atom has a higher affinity for electrons than the rest of the molecule (electronegative). Slightly negative end Slightly positive end Water is a polar molecule because the oxygen atom (with 8 protons) tends to pull the electrons away from hydrogen. The oxygen end has a slight negative charge, while the hydrogen end has a slight positive charge.

26. Hydrogen bonds Occur when the slightly positive (hydrogen) end of a polar molecule weakly attracts to the slightly negative end of another molecule. • Hydrogen Bonds: • Form weak bonds at room temperature, but are strong enough to form ice • Stabilize large proteins, DNA, and RNA End of Section 2, Chapter 2

27. section 3 of chapter 2 chemical reactions

28. Chemical Reactions Reactants (starting chemicals) are on the left → Products are on the right • Synthesis Reaction – joins molecules together • A + B → AB • Decomposition Reaction – breaks chemical bonds • AB → A + B • Exchange Reaction – reactants are swapped • AB + CD → AC + BD • Reversible Reaction – products can also yield reactants • A + B ↔ AB

29. activation energy: energy required to start a reaction A catalyst reduces the amount of energy needed to initiate a reaction. Catalysts increase the rate of reactions, but are not consumed by the reaction- reusable

30. Acids, Bases, and Salts Electrolytes – are substances that dissociate in water to release ions. Example: NaCl → Na+ + Cl-

31. Acids - electrolytes that dissociate to release protons (H+) in water Example: HCl→ H++ Cl- Bases- electrolytes that absorb H+ from water, or electrolytes that dissociate to release hydroxide ions (OH-) in water Examples: NaOH→ Na+ + OH- Salt – electrolyte formed by the reaction between an acid and base Acid + Base → Salt + water HCl + NaOH→ NaCl + H2O Example:

32. acid and base concentrations • pH • pH measures the concentration of hydrogen ions [H+] in a solution. • As pH decreases, [H+] increases = solution is more acidic acidic property increasing alkaline property increasing 0 7 14 pH neutral

33. Small changes in pH reflect large changes in [H+] • change of 1 pH = 10 fold change in [H+] • change of 2 pH = 100 fold change in [H+] • change of 3 pH = 1000 fold change in [H+]

34. Blood • Average blood pH = 7.35 - 7.45 • Acidosis = blood pH less than 7.3 • Symptoms include fatigue, disorientation, and difficulty breathing. • Alkalosis = blood pH greater than 7.5 • Symptoms include agitation and dizziness Blood contains several buffers Buffer = resists changes to pH

35. Chemical components of cells Organic Vs. Inorganic Molecules Organic molecules Compounds with carbon May form macromolecules Includes proteins, carbohydrates, lipids, nucleic acids Inorganic molecules Compounds that lack Carbon (exception is CO2) Usually dissociate in water

36. Inorganic Substances • Water (H2O) • 2/3 of weight in a person • Transports gasses, nutrients, wastes, hormones, ect. • Oxygen (O2) • Used in cellular respiration • Carbon Dioxide (CO2) • Waste of metabolic reactions • Inorganic Salts • Na+, Cl-, K+, Ca2+, HCO3-, PO42- End of Section 3, Chapter 2

37. Section 4, Chapter 2 Organic Molecules

38. Organic Molecules Molecules that contain carbon Organic Synthesis Small molecules (monomers) join together to form larger molecules (polymers) portion of a polymer Monomer

39. Covalent Bonds formed by Carbon Atomic Number of Carbon = 6 2 electrons in 1st shell 4 electrons in 2nd shell C 6 12.01 Note there are 4 empty spaces in the 2nd shell available for covalent bonds. Empty space for covalent bonding

40. Examples of covalent bonds formed by carbon Carbon can form 4 covalent bonds Carbon can also form double or even triple bonds Carbon to Carbon bonds can form long chains hydrocarbon

41. Polymers and Monomers Large organic molecules, called polymers consist of repeating subunits, called monomers. Example: Starch is a polysaccharide composed of many glucose molecules (monosaccharides) joined together.

42. major organic macromolecules of the cell

43. Carbohydrates • Simple carbohydrates = sugars • Monosaccharides • Disaccharides • Complex Carbohydrates • Also called Polysaccharides • Composed of several simple carbohydrates

44. monosaccharides • Twice as many Hydrogen as Oxygen atoms • Example: Glucose (C6H12O6)

45. disaccharides • 2 monosaccharides bonded together Examples of disaccharides

46. polysaccharide • Built of simple carbohydrates

47. examples of polysaccharides Starch – easily digested Cellulose- Plant polysaccharide, indigestible by humans Glycogen – storage form of energy, synthesized by liver Glycogen

48. LIPIDS Includes: Fats, Phospholipids, and Steroids

49. Fats (Triglycerides) nonpolar molecules They are soluble in oils, but insoluble in water Building blocks of fats 1 glycerol + 3 fatty acid molecules

50. Fatty Acids 1. Carboxyl end 2. Hydrocarbon chain Carboxyl group