Biochemistry

1. Biochemistry

2. Background information relevant to your study of biochemistry (you learned it in your previous science classes)

3. Properties of Matter • Matter – any substance that has mass and volume. • Mass - the quantity of matter in an object. • Volume - the amount of space that the matter takes up. • The more properties we can identify for a substance, the better we can understand its nature!

4. Properties of Matter • Physical properties of matter can be observed and measured without changing the identity of the matter. • Physical change - can affect size, shape, or color of a substance but does NOT affect the composition of the matter.

5. Physical Properties of Matter Luster Hardness Melting Point Boiling Point Phase Density • Mass • Color • Volume • Odor • Texture • Taste

6. Chemical Properties of Matter • Chemical properties of matter describe a substance’s ability to change into a NEW substance as a result of a chemical change. • Chemical change - bonds are broken and new bonds form between atoms. • Substances display different physical and chemical properties after the change. • A chemical change is irreversible!

7. Signs a chemical change occurred include… • Production of light • Production of heat • Color change • Gas production (bubbles) • Odor • Sound • A substance was created that wasn’t there before! • Examples: burning coal, ripening banana, baking a cake. • Examples: food is metabolized in body, photosynthesis.

8. Physical Change: Reversible You can “un-freeze water” No new substance is formed Water and ice are both H2O molecules Chemical Change: Not reversible You can’t “un-burn” wood A new substance is formed Burning wood results in CO2, ash, etc. Physical vs. Chemical Changes

9. Phases/States of Matter • Phase of matter - physical property of matter that describes one of a number of different states of the same substance.

10. Phases/States of Matter • Solid - definite shape, definite volume • Liquid - no definite shape, definite volume • Gas - no definite shape, no definite volume • Plasma - no definite shape, no definite volume • highly ionized gas that occurs at high temps

11. Chapter 2 - Biochemistry

12. Atoms • Atom - basic unit of matter. • Greek word “atomos” – unable to cut. • Atoms are the smallest component of a cell.

13. Atoms • Atoms compose all living and non living things. • Atoms contain subatomic particles: protons (+), neutrons (neutral), and electrons (-). • Protons and neutrons are found in the center of the atom in the atomic nucleus. • Electrons float around the nucleus in energy levels and are attracted to the nucleus by the protons (+’s attract –’s).

14. Atoms • Atoms are electrically neutral because their proton number and electron number balance out their charges.

15. Atomic Structure Protons Neutrons Electrons

16. Atomic Structure

17. The Nature of Atoms • Protons determine the identity of an atom! • Atomic number – number of protons in the nucleus of an atom. Each atom has a different proton number (identity). • Electrons determine how an atom behaves! • Electrons float around the nucleus in energy levels; most of an atom is empty space.

18. The Nature of Atoms • Mass number is the total number of protons and neutrons in the nucleus. • Most of the mass of an atom is in the nucleus!

19. Atomic Number Symbol Name Atomic Mass (Mass #) Atomic number equals the number of protons in nucleus. Atomic mass or mass number equals the number of protons + neutrons.

20. Atoms have Energy • Electrons in an atom have energy. • Energy is needed to keep electrons in the clouds so that they are not pulled into the nucleus.

21. Atoms have Energy • Each energy level can hold a certain number of electrons. • First level: 2 electrons • Second level: 8 electrons • Third level: 8 electrons • Fourth level: 10 electrons

22. Elements • Elements - substances that are composed of only one type of atom. • Cannot be chemically broken down to any other substances. • Are represented by chemical symbols on periodic table. • More than 100 elements are known, about 25 are found in living organisms. • 6 most abundant include: O, C, H, N, P, S

23. Isotopes • Isotopes - atoms of the same element that differ in the number of neutrons. • Still have the same number of protons - (proton number identifies the substance). • Isotopes of an element have the same chemical properties. They differ by the number of neutrons (a physical property).

24. Radioactive Isotopes • Radioactive isotopes - are unstable and from time to time breakdown releasing radiation from their nucleus. • Used to study organisms, diagnose disease (as tracers), treat disease (kill cancer cells), sterilize food, measure the ages of rocks. • Radiation is dangerous! It can kill or damage living things (i.e. Chernobyl’s radioactive fallout).

25. Chemical Compounds • A chemical compound is a group of atoms held together by chemical bonds. • Compounds are represented by chemical formulas.- show the proportion of atoms in a compound • Examples of chemical formulas: • NaCl – table salt • H2O – water • NH3 – ammonia • C6H12O6 - glucose

26. Interactions of Matter • Atoms want to achieve stability – full outermost energy level (valence shell). • In order to achieve stability, atoms will either gain, lose, or share electrons with other atoms in a process called chemical bonding. • Atoms will bond with other atoms if the bonding will give both atoms complete outermost energy levels. • Valence electrons- electrons in outermost energy level

27. Chemical Reactions • Chemical reaction – a process that changes one set of chemicals into another set of chemicals; involves the breaking and reforming of chemical bonds. • Reactants - chemicals that undergo a change (left side of equation). • Products - chemicals that are the result of a change (right side of equation). A + B ---------> C + D

28. Energy in Chemical Reactions • Energy is stored within chemical bonds. • When bonds are broken, energy is released. • All living organisms must have a source of energy to carry out chemical reactions! • Two types or reactions deal with the energy stored in chemical bonds: • Endergonic reactions • Exergonic reactions

29. Endergonic Reactions • Endergonic reactions – reactions that absorb energy. • Need a source of energy to trigger the reaction (don’t occur spontaneously). • Reactions tend to feel cold.

30. Exergonic Reactions • Exergonic reactions – reactions that release energy. • Energy is released as heat, light, or gas. • Can occur spontaneously. • Often feel warm.

31. Ionic bonds • Ionic bonds – chemical bonds that transfer electrons from one atom to another forming charged particles called ions.

32. Example: NaCl is a compound formed by ionic bonds. • Na has 1 electron in its outermost energy level. • When Na looses an electron, it becomes positively charged (Na+, or a sodium ion). • Cl needs 1 electron to fill its outermost energy level. • When Cl gains an electron from Na, it produces a negatively charged ion, Cl-. • The two oppositely charged ions are attracted to one another and form NaCl through transferring electrons in ionic bonding.

33. Covalent bonds • Covalent bonds – chemical bond formed by the sharing of electrons so that each atom fills its outermost energy level. • Most bonds in living organisms are covalent. • Examples: H2O, CO2, NH3, C6H12O6. • Molecule – smallest particle of a covalently bonded compound. Dogs Teaching Bonding: http://www.youtube.com/watch?v=_M9khs87xQ8&sns=em

34. Intermolecular Forces • Intermolecular forces - also called molecular attraction. • Are forces of attraction between stable molecules. • Example: hydrogen bonds (see section 3-3).

35. Two types of IM forces: • Cohesion - intermolecular force of attraction between LIKE molecules. • Adhesion - intermolecular force of attraction between DIFFERENT molecules.

36. Intermolecular Forces • Why do they occur? • Due to differences in charge densities or uneven distribution of electrons!

37. Acids, Bases, and pH • Acid – a substance that releases hydrogen ions (H+ ) when dissolved in water. • Example: HCl ---> H+ + Cl- • Base– a substance that releases hydroxide ions (OH-) when dissolved in water. • Example: NaOH ---> Na+ + OH-

38. Acids, Bases, and pH • Water is a neutral solution - water separates forming an equal number of hydrogen and hydroxide ions. • Neutralization reaction - Hydrogen ions and hydroxide ions react to form water. • Occurs when H+ ions from strong acids are mixed in perfect ratios with OH- ions from strong bases. H+ + OH- -----> H2O

40. pH Scale • pH – measures the amount of hydrogen in a solution, each measurement of pH represents ten times. • pH Scale - ranges from 0 to 14. • Less than 7 is for acids (more H+ than OH-). • Greater than 7 is for bases (more OH- than H+). • 7 is neutral (equal amounts of H+ and OH- in solution). • Most cells have a pH of 6.5-7.5. • Controlling pH is an example of homeostasis.

41. pH Scale • Acids - substances that forms hydrogen ions when dissolved in water. • The more hydrogen ions (less hydroxide) the more acidic. • Bases - substances that forms hydroxide ions when dissolved in water. • The more hydroxide ions (less hydrogen) the more basic or alkaline.

42. pH Scale • What happens when acid is added to a solution? • As more acid is added the pH will go down, but the H+ concentration goes up. • What happens when base is added to a solution? • As more base is added the pH will go up, but the H+ concentration goes down.

43. Buffers • Buffers – weak acids or bases that can react with strong acids or bases to prevent sharp, sudden changes in pH. • Are important for maintaining homeostasis in living organisms. • Ex. Carbonic acid and sodium bicarbonate buffer your blood’s pH.

44. Properties of Water • All cells contain water. • About two thirds of the molecules in our body are water. • Water provides a medium in which other molecules can interact. • Water exists as all three states/phases of matter. • Water expands when it freezes!!!!

45. Water is Polar • Water is a polarmolecule - molecule has slight charge (+ or -) on each end due to uneven distribution of electrons. • Oxygen pulls hydrogen’s electrons closer to it therefore the oxygen atom is slightly negative and the hydrogen becomes slightly positive. • This is the most important property of water! • Allows a strong attraction between water molecules or between water and other polar molecules!

46. Polar vs. Non-Polar Molecules • Polar - unequal distribution of charge means a great amount of attraction between molecules. • Non-Polar - equal distribution of charge means a weak attraction between molecules.

47. Do Polar and Non-Polar Solutions Mix? • Polar solutions mix with other polar solutions! • Example: Milk and water. • Non-polar solutions mix with other non-polar solutions! • Example: Oil and grease. • Polar solutions will NEVER mix with non-polar solutions! • Example: Italian salad dressing.

48. Water clings to itself & other molecules -Cohesion– Intermolecular force of attraction between like molecules. • Water molecules cling to other WATER molecules (hydrogen bonding) – Beading of water on a smooth surface. • Adhesion – Intermolecular force of attraction between different molecules. • Water molecules cling to other molecules – Meniscus in a graduated cylinder.

49. Water is good at forming mixtures • Due to slight charge of water molecules. • Mixture - substance composed of two or more elements or compounds that are mixed together but not chemically combined (are not linked by chemical bonds). • Examples: salt and pepper stirred together; atmosphere. • Two types of mixtures: Solutions & Suspensions

50. Water’s role in suspensions Suspension– a mixture where the solute does not fully dissolve. • Solute will settle out. • Example blood (plasma and blood cells).