1 / 47

Chapter 2

Chapter 2. Chemical Foundations for Cells. Chapter Outline. Review of elements and atomic structure Radioactive elements and health/medicine Chemical bonding Ionic Covalent: nonpolar and polar Hydrogen “bonding” Properties of water Acids, bases, and buffers Chemical change.

lilly
Download Presentation

Chapter 2

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 2 Chemical Foundations for Cells

  2. Chapter Outline • Review of elements and atomic structure • Radioactive elements and health/medicine • Chemical bonding • Ionic • Covalent: nonpolar and polar • Hydrogen “bonding” • Properties of water • Acids, bases, and buffers • Chemical change

  3. Elements (2.1, 2.3) • Living organisms are composed of matter • Matter is composed of elements • Element - substance that cannot be broken down into other substances by chemical means • Elements are made up of atoms. • Atoms join together to make compounds. Atoms Compounds

  4. Elements (2.1) • 92 naturally occurring elements • Life requires ~25 of these • ~96% of human body is made up of: • Carbon (C) • Hydrogen (H) • Oxygen (O) • Nitrogen (N)

  5. Compounds (2.1) • Atoms of one element can join with atoms of other elements to form compounds. • A given compound is always made of the same elements combined in the same ways. • NaCl – table salt • H2O - water • C6H12O6 - glucose

  6. Compounds of Life • Only living organisms have the ability to make the compounds of life: • Carbohydrates: C, H, O • Lipids: C, H, O • Proteins: C, H, O, N, S • Nucleic acids: C, H, O, N, P

  7. Atoms (2.3) • An atom is the smallest unit of an element • Atoms are composed of 3 subatomic particles: • Protons • Neutrons • Electrons

  8. Subatomic Particles

  9. Subatomic Particles and the Elements • Each element has a unique number of protons. • Atomic number - number of protons in an atom • Elements are arranged by atomic number on the periodic table. • Atoms are neutral, therefore # p = # e

  10. Isotopes • Number of neutrons is NOT on the periodic table for most elements…. • Isotopes - atoms of a given element that differ in the number of neutrons in the nucleus • Mass number – sum of the protons and neutrons in an atoms’ nucleus • The periodic table shows the average of the mass numbers for the isotopes of an element.

  11. Describing Isotopes Mass number 12C • Isotopes of carbon • 12C carbon-12 __ neutrons • 13C carbon-13 __ neutrons • 14C carbon-14 __ neutrons • All contain ____ protons and electrons. Carbon on the Periodic table

  12. Isotopes and Radioactivity (RA) • RA isotope has an unstable nucleus • Nucleus emits energy and particles in an effort to become more stable • May change the number of protons in the nucleus and become a different element.

  13. Radioactive Isotopes • Possible to target the energy and detect the radioactivity. • RA isotopes are used: • in research to track/follow molecules • in medicine to treat cancer and diagnose disease • Radiation therapy – treatment of localized cancer • PET - diagnosis

  14. Radioactive Isotopes • Overexposure to RA isotopes is HARMFUL. • Energy emitted damages cells. • radiation therapy takes advantage of this, goal is to damage and kill cancer cells • Exposure to RA can also cause mutations that lead to cancers • Eg – exposure to RA element radon is the 2nd leading cause of lung cancer

  15. Diagnosis - PET Scans • A radioactive tracer is put into the body. • Often RA glucose • The RA glucose goes to the parts of the body that use glucose for energy. • Cancers use glucose differently from normal tissue • As the radiotracer is broken down positrons are made. This energy appears as a 3-dimensional image on a computer monitor.

  16. Electron Arrangement (2.5) • When compounds form, the electrons of the bonding atoms interact in attempt to obtain a more stable state. • Some electron arrangements are more stable than others…….see board

  17. Chemical Bonding (2.6-2.7) • Chemical bonding – atoms gain, lose, or share electron(s) to obtain a stable number of electrons • Can be ionic bond or covalent bond

  18. Chemical Bonding - Ionic • Ionic Bond – strong attractive force between oppositely charged ions • Atoms form ions by losing or gaining enough electron(s) to obtain a stable # of electrons in their outer shell

  19. electron transfer SODIUM ATOM 11 p+ 11 e- CHLORINE ATOM 17 p+ 17 e- SODIUM ION 11 p+ 10 e- CHLORINE ION 17 p+ 18 e-

  20. Ionic Bonding

  21. Chemical Bonding - Covalent • Covalent Bond – bonded atoms share pair(s) of electrons and form molecules. • Occurs between nonmetals such as: C, O, H, N, P, S • Covalent bonding occurs in • H2 • O2 • H2O

  22. Two Classes of Covalent Bonds • Nonpolar Covalent Bond – bonded atoms share electrons equally • Occurs between like atoms or between atoms with a similar ability to attract shared electrons • Polar Covalent Bond – unequal sharing of electrons by the bonded atoms • Occurs between atoms with very different ability to attract shared electrons

  23. Two hydrogen atoms, each with one proton, share two electrons in a single nonpolar covalent bond. molecular hydrogen (H2) H—H Fig. 2-8b(1), p.25

  24. water (H2O) H—O—H Two oxygen atoms share four electrons in a nonpolar double covalent bond. molecular oxygen (O2) O=O

  25. Types of Covalent Bonds • Nonpolar covalent – bonded atoms share the electrons equally • Examples of nonpolar bonds: • H2

  26. Atoms with different electronegativity values form polar covalent bonds. • Electronegativity (EN) – measure of an atom’s ability to attract shared electrons in a covalent bond • Oxygen and nitrogen have fairly large EN values – often d - • Carbon and hydrogen have low EN values – often d +

  27. Polar Covalent – unequal pull on shared electrons by the bonded atoms • Results in partial charges on the bonded atoms d - O H H d + d +

  28. Common Polar Covalent Bonds O-H N-H C-O C=O Label the polarity in each bond.

  29. Forces between Molecules • Molecules are weakly attracted to each other by intermolecular (IM) forces, • The most important IM force in biology is the hydrogen “bond” (2.8) • Attractive force between d + H and d – O, N or F

  30. Hydrogen “bond” is a weak attractive force between a d + hydrogen and a d-O, N, or F in a second polar bond Water is a polar molecule capable of hydrogen bonding.

  31. Properties of Water • Water is cohesive and has high surface tension. • Cohesion – ability of molecules to stick together • Surface tension - ability to resist rupturing when under tension

  32. Properties of Water • Water resists changes in temperature. • When heat is applied to an aqueous solution much of the heat (energy) is used to break hydrogen bonds, not to increase the movement of the molecules.

  33. Properties of Water • Solid water (ice) is less dense than liquid water • Ice floats • Therefore, ice forms on the top of lakes and insulates the liquid water below.

  34. Water is a good solvent for ionic compounds and small polar molecules. Water H bonds to polar molecules like ethanol

  35. Water as solvent • Water pulls ions apart and hydrates them

  36. Related Terms • Hydrophilic • Water loving • Capable of hydrogen bonding to water (polar) • Hydrophobic • Water “fearing” • Cannot hydrogen bond to water (nonpolar)

  37. Acids, Base, and Buffers (2.14) • Many ions are dissolved in the fluids in/outside of cells – called electrolytes • Na+, Ca+2, K+ • H+ • Level of each ion is critical • Our focus is on H+ (hydrogen ions)

  38. Acids, Base, and Buffers • Acid: Substance that produces H+ when dissolved in water………. • Examples: • Hydrochloric acid – stomach acid • Lactic acid – made when cells run out of oxygen • Amino acids – building blocks of proteins

  39. Acids, Base, and Buffers • Base: substance that accepts H+1 (hydrogen ions) in water • Examples: • Sodium hydroxide - NaOH • Most nitrogen containing compounds • Ammonia – NH3 • Urea – in urine • Amino acids – building blocks for proteins

  40. Acids, Base, and Buffers • Classify substances as acid, base or neutral by their pH • Acids: pH < 7 • Base: pH > 7 • Neutral: pH = 7 • Pure water has a pH of 7 • See page 28

  41. Acids, Base, and Buffers • How the pH scale works • The lower the pH the more acidic • The higher the pH the more basic (alkaline) • A difference of 1 pH unit is a 10-fold difference in acidity or alkalinity

  42. Why is pH important? • Most cells require a pH near 7. • Above or below this pH for too long and they die. • Proteins function only at specific pHs. • In lab you will determine the optimal pH for a protein that is needed to breakdown hydrogen peroxide in cells

  43. Acids, Base, and Buffers • Buffers: solution that resists changes in pH even when acid or base is added • Buffers can both produce H+ and neutralize H+ • Buffers are key to maintaining pH homeostasis • Most body solutions are buffered

  44. Why is pH important? • Blood has a pH of 7.3 – 7.4 • If the pH is above or below this range for more than a couple of days death occurs. • The blood buffer system helps keep blood pH in a range that supports life.

More Related