1. The Chemistry of Life

2. Learning Objectives Describe the basic structure of an atom. Understand how electrons determine atom interaction. Define the term isotope and understand how they are used in science. Describe ionic, covalent and hydrogen bonds.

3. Learning Objectives Understand the biologically important characteristics of water. Understand the four major classes of organic molecules, their functions, and their building blocks.

4. Atoms All matter is composed of atoms; the smallest particles that can be divided and still maintain its chemical properties. Atom: Protons (+) Electrons (-) Neutrons ( )

5. 3.1 Atoms An atom can be characterized by the number of protons it has or by its overall mass atomic number the number of protons in the nucleus atoms with the same atomic number exhibit the same chemical properties and are considered to belong to the same element mass number the number of protons plus neutrons in the nucleus electrons have negligible mass

6. 3.1 Atoms Electrons determine the chemical behavior of atoms these subatomic components are the parts of the atom that come close enough to each other in nature to interact

7. Ions…What Are They? In an electronically neutral atom, there are an equal number of protons and electrons. Ions form when atoms do not have an equal number. Therefore, all ions are electrically charged.

8. Isotopes Isotopes – atoms that have the same number of protons but different numbers of neutrons most elements in nature exist as mixtures of different isotopes

9. What’s an Isotope? The number of protons of an atom is its atomic number. Atomic mass of an atom includes the number of protons and neutrons. The number of protons never varies, but the number of neutrons may change…giving rise to an isotope.

10. Ions and Isotopes Short-lived isotopes decay rapidly and do not harm the body Can be used as tracers to study how the body functions Figure 3.7 Using a radioactive tracer to identify cancer

11. 3.2 Ions and Isotopes Some isotopes are unstable and break up into particles with lower atomic numbers this process is known as radioactive decay Radioactive isotopes have multiple uses Medicine: detection and treatment of disorders dating fossils

12. 3.2 Ions and Isotopes Dating fossils the rate of decay of a radioactive element is constant by measuring the fraction of radioactive elements that have decayed, scientists can date fossils the older the fossil, the greater the fraction of its radioactive atoms that have decayed

13. Figure 3.8 Radioactive isotope dating

14. How Electrons Determine Actions Electrons carry energy (potential energy) Energy levels surrounding the nucleus reflects the amount of energy possessed by the electron existing there. Less energy is present in electrons close to the nucleus.

15. Oxidation and Reduction Oxidation occurs when there is an interaction between an electron and two atoms..the loss of an electron is oxidation. Reduction occurs when one atom gains an electron from another atom.

16. Electron Orbitals The path that an electron follows around a nucleus is called an orbital. Each orbital holds 2 electrons. The first energy level has one orbital..2 electrons. The second and third energy level has 4 orbitals..8 electrons each. When orbitals are not filled with electrons, the atoms are ready to react to fill those orbitals.

17. Molecules A molecule is made up of two or more atoms held together by a chemical bond. Chemical bonds firm between atoms through the interaction of electrons. What types of chemical bonds do you know about?

18. Ionic Bonds Ionic bonds form when ions are electrically attracted to each other by opposite charges. NaCl is an example of ionic bonding. Na gives up an electron to Na; Na + Cl - Ionic bonds are strong and not directional, supports the formation of crystal.

19. Covalent Bonds Covalent bonds form when electrons are shared between atoms. Most organic molecules are formed from covalent bonds. Two key properties make covalent bonds ideal for use in biological molecules: They are strong They are very directional

20. Hydrogen Bonds Hydrogen bonds result from weak attractions between hydrogen atoms and the larger atoms of polar molecules. Hydrogen bonds are weak and highly directional. They play an important role in maintaining the conformation of large, biologically important molecules.

21. Hydrogen Bonds and Water All organisms are made up of a large amount of water. Water is biologically important because it is a polar molecule and forms hydrogen bonds between its own molecules. Unique properties given to water by hydrogen include: Cohesion Heat storage Ice formation High polarity High heat of vaporization

22. Heat Storage Water has the capacity for heat storage because of its hydrogen bonds. Water changes temperature slowly…good for living organisms. Assists in regulating homeostasis.

23. Ice Formation When water freezes, the hydrogen bonds space water molecules apart. Makes ice less dense than water

24. High Heat of Vaporization A large amount of energy is needed to break hydrogen bonds in water and turn the water into vapor. This mechanism of energy use is why evaporative cooling removes heat from the body.

25. Cohesion The attraction of water molecules to other water molecules. The surface tension of water is a result of cohesion. When other polar molecules are attracted to water molecules it is called adhesion.

26. High Polarity Hydrophilic (water-loving) molecules attach to water molecules making them water soluble. (Polar molecules) Hydrophobic (water-fearing) molecules are not readily soluble in water (non-polar). Why is this important?

27. Water Chemistry Water ionizes spontaneously, forming hydrogen ions (H + ) and hydroxyl ions (OH - ). This process is called ionization: H 2 0 Hydrogen ion concentration in solution can be described using the pH scale.

28. pH Scale Acids are solutions with an increased concentration of hydrogen ions. Bases are substances that combines the hydrogen ion concentration in solution. The pH inside the cells of most living organisms is close to 7. A buffer is a substance that acts as a reservoir for hydrogen ions. It resists change to pH. Blood has the acid-base pair of carbonic acid and bicarbonate as a buffer (pH of blood is 7.4)

29. Figure 3.15 The pH scale

30. Examples of pH in nature Hydrangea macrophylla blossoms are either pink or blue, depending on a pH- dependent mobilization and uptake of soil aluminium into the plants.

31. Examples of pH in living systems The pH of different cellular compartments, body fluids, and organs is usually tightly regulated in a process called acid-base homeostasis. The pH of blood is usually slightly basic with a value of pH 7.4. This value is often referred to as physiological pH in biology and medicine.

32. Examples of pH in living systems CompartmentpH Gastric acid0.7 Lysosomes4.5 Granules of chromaffin cells5.5 Urine6.0 Neutral H 2 O at 37 °C6.81 Cytosol7.2 Cerebrospinal fluid (CSF)7.3 Blood7.34 – 7.45 Mitochondrial matrix7.5 Pancreas secretions8.1

33. 3.5 Water Ionizes The pH in most living cells and their environments is fairly close to 7 proteins involved in metabolism are sensitive to any pH changes Acids and bases are routinely encountered by living organisms from metabolic activities (i.e., chemical reactions) from dietary intake and processing Organisms use buffers to minimize pH disturbances

34. 3.5 Water Ionizes Buffer – a chemical substance that takes up or releases hydrogen ions buffers don’t remove the acid or the base affecting pH but minimize their effect on it most buffers are pairs of substances, one an acid and one a base