The Chemistry of Life

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 ( ) Where are they located?

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: • H20 • 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

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