Matter

1. Matter • The “stuff” that makes up the universe • anything that takes up space • States of matter • Solid • has definite shape and volume • Liquid • has definite volume, changeable shape • Gas • has changeable shape and volume

2. Composition of Matter • Elements • unique substances that cannot be broken down by ordinary chemical means • carbon, oxygen, helium, uranium, gold, iron… • Only 24 elements have a role in our body • 98.5% of body weight consists of • O, C, H, N, Ca, P • Atoms • building blocks for each element

3. The Chemical Elements • Periodic table • atomicsymbols of elements arranged by atomic number • Atomic number of each element • number of protons in its nucleus • Only 24 elements have a role in our body • 98.5% of body weight consists of • O, C, H, N, Ca, P

5. Atomic Structure (Bohr or Planetary Model) • Nucleus, the center of atom which contains: • protons • positive charge, mass (weight) of 1 atomic mass unit (amu) • determines atomic number • neutrons • neutral (no) charge, mass of 1 amu • Electron shells that surround the nucleus and contain: • electrons • negative charge, mass of 0 amu • valence electrons are in the outermost shell • furthest from the nucleus • interact with other atoms • All atoms have: • an equal number of protons and electrons • atoms are neutral (have no net charge) • an atomicmass = total mass of protons + neutrons

6. Planetary Model of an Atom

7. Electron Shells • The electron shell closest to the nucleus can hold up to 2 electrons • additional electrons are located in shells outside of the first shell • All other electron shells outside of the first shell can hold up to 8 electrons • octetrule • an electron shell is full when there are: • 2 electrons in the first shell • 8 electrons in 2nd, 3rd, 4th,… shell • If the valence shell of an atom is not completely FULL, then that atom is UNSTABLE • valance electrons of unstable atoms interact with valance electrons of other unstable atoms to create chemical bonds • allows both atoms to become stable

8. Atomic Structure • Nucleus, the center of atom which contains: • protons • positive charge, mass (weight) of 1 atomic mass unit (amu), determines atomic number • neutrons • neutral (no) charge, mass of 1 amu • atomic mass = total # of protons + neutrons • Electron shells • electrons: negative charge • # of electrons = # of protons, atoms have neutral charge • electrons further from nucleus have higher energy • valence electrons are in the outermost shell • interact with other atoms • determine chemical behavior

9. Planetary Model of an Atom

10. Electron Shells • The electron shell closest to the nucleus can hold up to 2 electrons • additional electrons are located in shells out side of the first shell • All other electron shells outside of the first shell can hold up to 8 electrons • Valance electrons of an atom interact with valance electrons of other atoms in order to obey the octet rule • octet rule: atoms react to obtain a full valence shell • 2 electrons in the first shell or 8 electrons in 2nd, 3rd, 4th,… shell • an atom with a full valence shell is STABLE

11. Planetary Models of Atoms p+ represents protons, no represents neutrons

12. Chemical Bonds • The reaction between 2 atoms results in the formation of a chemical bond • Bonds are formed between 2 atoms using the electrons in the valence shell of each atom • An atom is stable when the valence shell is completely full (satisfying the “octet rule”) • Octet rule – except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their valence shell (thus becoming stable)

13. Chemically Inert (Non-reactive) Elements • Inert elements have their valence shell fully occupied by electrons

14. Chemically Reactive Elements • Reactive elements do not have their outermost energy level fully occupied by electrons

15. Types of Chemical Bonds • Ionic • Covalent • Hydrogen

16. Molecules and Compounds • Molecules • two or more atoms of same element covalently bonded • Compounds • two or more atoms of different elements covalently bonded • Structural formula • shows arrangement of atoms • Molecular formula • itemizes each element present and its quantity

17. Molecular and Structural Formula

18. Formation of an Ionic Bond • Ionic bonds form between atoms after the transfer of one or more electrons from one atom to another atom in a process called ionization • results in the creation of 2 ions • Ions • atoms that have an unequal numbers of protons and electrons • also known as electrolytes

19. Anions and Cations • Cation • atom lost one or more electrons • more protons than electrons = net positive charge • Anion • atom gained one or more electrons • more electrons than protons = net negative charge

20. Anions and Cations • The anion and the cation are held together by the ATTRACTION between a positively charged substance and a negatively charged substance • 2 substances that have the same charge REPELL one another • Atoms bound by ionic bonds form crystals (salts) • NaCl (sodium chloride) • Ionic bonds cannot exist in water • the bond breaks (dissociates) • the components of the molecule exist in their ionic form (anion and cation) in the body • NaCl + H2O→ Na+ and Cl- +H2O

21. Covalent Bonds • Formed by sharingvalence electrons between 2 atoms • Types of covalent bonds include: • polar and nonpolar covalent bonds • Are stable (Do NOT dissociate) in water • Nearly all biologically important molecules are made from atoms bonded together with covalent bonds

22. electrons shared unequally Polar Covalent Bonds • Polar (electrically charged) bonds occur between an electronegative atom (O or N) and an atom that is neither O nor N • the nucleus of an electronegative atom has a stronger “pull” on the shared electrons, pulling the shared electron(s) closer to it • this causes the electronegative atom to become partiallynegative while the other atom in the covalent bond becomes partiallypositive • similar to a battery

23. Nonpolar Covalent Bonds • Nonpolar bonds occur between two atoms neither of which are O or N • the shared electrons are pulled equally between the nuclei of the 2 atoms • both atoms remain neutral • the nonpolar covalent bond is neutral electrons shared equally

24. Ions • Ions – atoms that carry a charge (unequal numbers of protons and electrons) • Ionization - transfer of electrons from one atom to another (provides stability of the atom)

25. Anions and Cations • Anion - atom gained electron, net negative charge • Cation - atom lost an electron, net positive charge

26. Formation of an Ionic Bond • Ionic bonds form between atoms by the transfer of one or more electrons • One atom donates one or more electrons, becoming a cation (+), the other atom accepts the donated electron(s), becoming an anion (-) • The anion and the cation are attracted to one another and are held together by an ionic bond • Example: NaCl (sodium chloride) • In water (such as in your body) the ionic bond dissociates (breaks) and the components of the molecule exist in their ionic form (anion and cation) • NaCl → Na+ and Cl-

27. Ionic Bonds • Ionic bonds are weak and dissociate (break) in water • These compounds tend to form crystals (salts)

28. Formation of an Ionic Bond

29. Formation of an Ionic Bond

30. Formation of an Ionic Bond

31. Sodium Chloride Crystal

32. Sodium Chloride Crystals

33. Covalent Bonds • Formed by sharing valence electrons between 2 atoms • Different types of covalent bonds • single covalent bond • double covalent bond • nonpolar covalent bond • polar covalent bond • DO NOT dissociate in water • because all molecules that are formed by ionic bonds dissociate in water, only molecules that are formed by covalent bonds exist (and function) in the body

34. Single Covalent Bond • One pair of electrons are shared

35. Single Covalent Bond

36. Double Covalent Bond • Two pairs of electrons are shared

37. Sharing of electrons • The sharing of electrons in a covalent bond can be either “equal” or “unequal” • When the shared electrons between 2 atoms is such that the electrons are located exactly between the nuclei of 2 atoms, the sharing is said to be equal • In other words, the shared electrons are not closer to the nucleus of either atom • However, when the shared electrons between 2 atoms is such that they are closer to one of the 2 atoms, the sharing is said to be unequal • In other words, the shared electrons are closer to the nucleus of one atom and further away from the nucleus of the other atom

38. Nonpolar and Polar Covalent Bonds • Electrons shared equally between atoms produce nonpolar bonds because the negative charge of the electron is spaced evenly between the 2 atoms • Unequal sharing of electrons produces polar bonds • Polar bonds occur between an electronegative atom (O or N) and an atom that is not O or N • Electronegative atoms have a stronger “pull” on the shared electrons, pulling the electrons closer to it • This causes the electronegative atom to become partiallynegative while the other atom in the covalent bond becomes partiallypositively

39. Nonpolar and Polar Covalent Bonds electrons shared equally electrons shared unequally

40. Functional Groups • small groups of covalently bonded atoms arranged in a very specific manner • parts of large compounds (carbohydrates, proteins, fats and nucleotides) • determine the chemical properties of large compounds (polar vs nonpolar, acid vs base) • react with functional groups on other compounds

41. Acid/Base Concentration (pH) • pH is the measurement on a scale ranging from 0 to 14 of H+ concentration in a solution • H+ is the ionized form of a hydrogen atom • the only electron has been removed, leaving a single proton • H+ = hydrogen ion = proton • pH = -log[H+] • [H+] = molar concentration of H+ in a solution • the greater the [H+] the lower the pH, the lower the [H+] the higher the pH • Acidic solutions have higher [H+] • a lower pH • Alkaline (basic) solutions have lower [H+] • a higher pH

42. pH Scale • Acidic: pH 0 – 6.99 • Basic: pH 7.01 – 14 • Neutral: pH 7.00

43. Acids and Bases • Acids are molecules that are capable of increasing the number of H+ in a solution • called protondonors • decrease the pH of a solution • Bases are molecules that are capable of decreasing the number of H+ in a solution are • called proton acceptors • increase the pH of a solution

44. Buffers • Substances that are capable of resisting large changes in the pH of a solution • allow pH to remain relatively constant • buffers in the body allow the body pH to remain at 7.4 (slightly basic)

45. Water as a Solvent • water molecules overpower the ionic bond above between Na+Cl- by forming hydration spheres • note orientation of water molecules: negative pole faces Na+, positive pole faces Cl-

46. Liquid Mixtures Substances that are physically blended but not chemically combined Water is the most abundant compound in biological mixtures • Solutions • Colloids • Suspensions

47. Solutions • Consists of: • Solute (less abundant substance in mixture) which is dissolved in the solvent • Solvent (more abundant substance in mixture) • alwayswater in the body • Transparent • e.g. copper sulfate solution

48. Concentration of Solutions Measurement of the amount of solute(s) in a solution • Weight per Volume • weight of solute in a given volume of solution • 8.5 grams of dextrose in 1 liter of water = 8.5 g/L • Percentages • weight or volume of solute in solution • 2 grams of dextrose + 8 grams of H2O is a solution containing a total of 10 grams • 2 of the 10 grams is dextrose, therefore this solution is a 20% dextrose solution

49. Concentration of Solutions • Molarity • number of moles of solute/liter in solution • based on molecular weight (MW) • the addition of the atomic masses of a molecule • H2O = H (1 amu) x 2 + O (16 amu) = 18 amu MW • for a known MW, weigh out that many grams, this gives you its gram molecular weight or 1 mole • 1 mole always contains the same number of molecules (6.02 × 1023 = Avogadro’s number) • MW of glucose is 180, so one mole of glucose is 180g, a one molar solution of glucose contains 180g/L

50. Molar # of molecules equal weight of solute unequal Percentage # of molecules unequal weight of solute equal Percentage vs. Molar Concentrations