Chemistry

1. Chapter 2 Chemistry Anatomy & Physiology ivyanatomy.com

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.

3. Definitions Chemistry = Science that exams the composition and interactions of matter Biochemistry = Chemistry of living things • Matter = Anything that has mass and takes up space • (Solids, liquids, gasses) Elements are listed on a Periodic Table • Element = Fundamental substance of matter • (e.g. Carbon, Nitrogen, Oxygen) 6 7 8 C N O 12.01 14.01 15.99

4. Definitions • Atom = Smallest functional particle of an element Molecule = Particle of two or more atoms chemically joined together. Compound = Two or more different elements chemically bonded together (e.g. H2O = water, C6H12O6 = glucose) Moleculeof an element = two or more identical atoms chemically bonded together. (e.g. H2 = hydrogen molecule, O2 = oxygen molecule)

5. Elements of the Body • 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)

6. Periodic Table

7. Atomic Structure electron • atom - the smallest particles of an element that still have the properties of that element. neutron proton + + + + • subatomic particles: • Proton – single positive charge • Neutron – carries no electrical charge • Electron – single negative charge - - - nucleus - An atom contains a central nucleus composed of protons and neutrons. Electrons orbit the nucleus.

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

9. Atomic Number and Atomic Weight • 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 (mass): • The sum of protons and neutrons in one atom. • Remember, the weight of electrons is negligible. 8 O 15.99

10. atomic number and atomic weight Hydrogen (H) + + - + + + - Carbon (C) + - - - - + -

11. Isotopes • Isotopes - atoms of an element with the same atomic number, but different atomic weights. • number of neutrons of an element varies between atoms. ≈ 99% ≈ 0.1% ≈ 1% Isotopes of carbon 6 6 6 C C C 12 14 13 Atomic weight (mass) of carbon 6 • *The atomic weight of an element is an average of the isotopes present. C 12.01

12. Periodic Table

13. Properties of Electrons 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 - - - - - - - nucleus - - - - Octet rule: rule of thumb that except for the 1st shell, each atom reacts to have 8 electrons in its outer (valence) shell. - - - * Lower shells are filled first. - -

14. 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).

15. Ions Ion - atom that readily gain or loose electrons + • Cation: an ion that is deficient in electrons • positively charged ions • Anion: an ion that has additional electrons • negatively charged ions -

16. Example of a cation Sodium readily loses an electron. - An unpaired electron in its own shell is highly unstable. 11 Na 23 Na+ Sodium cation

17. Example of an anion Chlorine readily accepts an additional electron - Chlorine only needs one more electron to fill its outer shell and become stable. 17 Cl 35 Cl- Chloride anion The anion of chlorine is called Chloride.

18. Chemical Bonds Ionic Bond - Oppositely charged ions attract Na+ + Cl- → NaCl (cation + anion)

19. Ionic bond = complete exchange of electrons _ + 11 17 Na Cl 23 35 Na+ Cl- Sodium cation Opposite charges attract Chloride anion

20. Ionic bonds form crystals Cations & anions attract in all directions, forming organized arrays, such as crystals. - - + + + + - - - - + + + + - Salt crystal formation occurs because of the ionic bonds of sodium and chloride.

21. Chemical Bonds 2. Covalent Bonds - atoms share electrons • non-polar – atoms equally share electrons • carry no charge • non-polar molecules do not dissolve in water (hydrophobic)

22. non-polar 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. + H H H-H Structural formula of the hydrogen molecule Hydrocarbons are an important group of non-polar molecules CH4 molecular formula Structural formula

23. Chemical Bonds 2. Covalent Bonds - atoms share electrons • polar – unequal sharing of electrons • Electrons spend most of their time on one side of the molecule • Polar molecules dissolve in water (hydrophilic)

24. polar covalent bonds • Water is a polar molecule because the larger oxygen nucleus 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. δ- partial negative charge O H H Structural formula • H2O molecular formula δ+ δ+ partial positive charge

25. Types of covalent bonds Single Bond - atoms share one pair of electrons O • (H2O) H H Double Bond – atoms share two pairs of electrons • (O2) O O Triple Bond – atoms share three pairs of electrons • (N2) • N • N

26. Chemical Bonds 3. Hydrogen bond δ- O slightly positive (hydrogen) end of a polar molecule weakly attracts to the slightly negative end of another molecule H H δ+ δ+ Hydrogen bonds

27. Chemical Bonds Hydrogen bonds • weak bonds at room temperature, but are strong enough to form ice • Stabilize large proteins, DNA, and RNA

28. Types of chemical bonds 1. Ionic Bond – opposite charges attract 2. Covalent Bond – atoms share their electrons non-polar – equal sharing polar – unequal sharing 3. Hydrogen Bond – weak attraction between a positive hydrogen and a negative portion of another molecule

29. 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

30. Activation energy • Activation Energy – Energy required to initiate a reaction • Catalysts - increases the rate of a reaction without being consumed. Activation energy without catalyst Activation energy with a catalyst energy of reactants A Catalyst lowers the activation energy required to initiate a reaction energy of products (unchanged)

31. Acids, Bases, and Salts Electrolytes – dissociate in water to release ions. Hydration shell Hydration shell + Na+ Cl- NaCl Electrolytes: Na+, Cl-, K+,

32. Acids, Bases, and Salts Electrolytes: Na+, Cl-, K+, Ca2+, Mg2+ Phosphate (PO42-) Sulfate (SO42-) Bicarbonate (H2CO3-)

33. Acids, Bases, and Salts Acid - electrolyte that releases H+ in solution + HCl H+ Cl- • Example: Hydrochloric acid Hydrogen ion Chloride anion H+ = (hydrogen ion/proton)

34. Acids, Bases, and Salts • Many bases release hydroxide ions (OH-) Base (alkaline) - electrolyte that removes H+ from solution + • Example: + NaOH OH- Na+ Sodium Hydroxide The hydroxide reacts with H+ to form water + H+ OH- H2O

35. Acids, Bases, and Salts Salt – electrolyte formed by the reaction between an acid and a base (alkaline) Acid + Base → Salt + water HCl + NaOH→ NaCl + H2O

36. Acid and Base Concentrations • pH - measures the concentration of hydrogen ions [H+] in a solution • pH measures the –log [H+] • logarithm means the pH tracks the number of decimal places • -log = as pH decreases, [H+] increases = solution is more acidic acidic property increasing alkaline property increasing 0 7 14 pH neutral

37. Hydrogen ion concentration vs. pH • Small changes in pH reflect large changes in [H+] • 1 pH = 10 fold change in [H+] • 2 pH = 100 fold change in [H+] • 3 pH = 1000 fold change in [H+] • 4 pH = 10,000 fold change in [H+] Increasingly acidic Increasingly alkaline

38. pH of Blood • Average pH of blood = 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

39. Chemicals of the Cell 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

40. Inorganic Chemicals • 1. Water (H2O) - 2/3 of weight in a person • Transports gasses, nutrients, wastes, hormones, ect. • 2. Oxygen (O2) - Used in cellular respiration • 3. Carbon Dioxide (CO2) - Waste of metabolic reactions • 4. Inorganic Salts: Na+, Cl-, K+, Ca2+, Mg2+, HCO3-, PO42-

41. Organic Chemicals • 1. Carbohydrates • 2. Lipids • 3. Proteins • 4. Nucleic Acids

42. Organic Molecules C 6 Carbon 12 + • Organic molecules contain carbon + - + + + • Carbon forms 4 covalent bonds - + C • Carbon-to-Carbon bonds can form long hydrocarbon chains and hydrocarbon rings H - - H H H H H H H H H H H H H H H H H H H H O C C C C C C C C C C Fatty acid – example of an organic molecule Glucose (C6H12O6) - -

43. Organic Synthesis Organic Synthesis Small molecules (monomers) join together to form larger molecules (polymers) Several Monomers Polymer

44. Organic Chemicals Polymer Monomer Monosaccharides • 1. Carbohydrates • 2. Lipids • 3. Proteins • 4. Nucleic Acids Polysaccharide 1 Glycerol + 3 Fatty Acids (for fats only) Triglyceride (for fats only) Polypeptide Amino Acids Polynucleotide Nucleotides

45. Carbohydrates Monosaccharides (simple sugars) • 2:1 ratio of Hydrogen to Oxygen (eg. C6H12O6) • Polar molecules – water soluble Glucose Fructose Disaccharides (double sugars) Maltose Lactose Sucrose

46. Carbohydrates Polysaccharides (complex carbohydrates) • Large molecules composed of several monosaccharides and disaccharides joined together • Starch – easily digested • Cellulose- Plant polysaccharide, indigestible by humans • i.e. Dietary Fiber • Glycogen – storage form of energy, synthesized by liver Each ring represents a monosaccharide

47. Lipids Lipids include Fats Phospholipids Steroids • Lipids are mostly non-polar • Water insoluble (hydrophobic)

48. Lipids Fat (triglyceride) Fats are composed of: 1 glycerol + 3 fatty acid molecules

49. Lipids Fat (triglyceride) • Fatty Acids may be saturated or unsaturated • Saturated fatty acid • All carbon-to-carbon single bonds • Unsaturated fatty acid • Contains one or more carbon-to-carbon double bonds

50. Lipids Fat (triglyceride) 3 fatty acids glycerol Saturated Fat Unsaturated Fat