1. 1 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. PowerPoint to accompany CONCEPTS IN BIOLOGY TWELFTH EDITION Enger Ross Bailey CHAPTER 2

2. 2 2.1 Matter, energy and life Matter is anything that has mass and occupies space. Energy is the ability to do work. – There are two types of energy: – Potential energy Stored energy, available to do work – Kinetic energy Energy of motion – Potential energy can be converted to kinetic energy to do work.

3. 3 Law of conservation of energy Energy is never created or destroyed. – The first law of thermodynamics Energy can be converted from one form to another, but the total energy remains constant. – An object at the top of a hill has potential energy based on its location. – When the object rolls down the hill, the potential energy is converted to kinetic energy.

4. 4 Forms of energy There are five forms of energy: 1. Mechanical energy Energy of movement 2. Nuclear energy Energy from reactions involving atomic nuclei 3. Electrical energy Flow of charged particles 4. Radiant energy Energy in heat, light, x-rays and microwaves 5. Chemical energy Energy in chemical bonds

5. 5 2.2 What is the nature of matter? Atoms – The smallest units of matter that can exist separately. Elements – Chemical substances composed of the same kind of atoms. – Listed on the periodic table. – Each element is represented by a symbol of one or two letters. – The principal elements that comprise living things are: C, H, O, P, K, I, N, S, Ca, Fe, and Mg.

6. 6 The periodic table of the elements

7. 7 Atomic structure Atoms are composed – The atomic nucleus Protons - positively charged – Atomic number-the number of protons – All atoms of the same element have the same # of protons. Neutrons – no charge – Electrons Orbit the nucleus in energy levels Are constantly in motion

8. 8 Atomic structure

9. 9 Elements Atoms of the same element have equal number of electrons and protons. – Thus, they have a neutral charge. Isotopes – Atoms of the same element that have different numbers of neutrons – Atomic weight-the average of all of the isotopes in a mixture Mass number – The sum of protons and neutrons in the nucleus.

10. 10 Isotopes of hydrogen

11. 11 Electrons Electrons occupy specific energy levels around the nucleus. – Electrons closest to the nucleus have the lowest energy. Energy levels hold specific numbers of electrons. – The first energy level can have up to 2 electrons. – All other energy levels can have up to 8 electrons. Atoms seek to have a full outer energy level. – Atoms that have full outer energy levels are inert. – Other atoms seek to fill their outer energy levels through chemical bonds.

12. 12 Electrons

13. 13 2.3 The formation of molecules Molecules consist of two or more atoms joined by a chemical bond. A compound is a chemical substance made of two or more elements combined in chemical bonds. – The formula of a compound describes the nature and proportions of the elements that comprise the compound. H 2 0

14. 14 2.4 Molecules and kinetic energy Molecules are constantly in motion. Temperature is a measure of the average speed of the molecules in a substance. – The greater the speed, the higher the temperature. – Measured in Fahrenheit or Celsius Heat is a measure of the total kinetic energy of molecules. – Measured in calories (amount of heat that will raise 1g of water 1 degree Celsius). Heat and Temperature are related. – Add heat energy to a substance and the molecules will speed up, and the temperature will rise.

15. 15 2.5 Kinetic energy, physical changes and phases of matter Three phases of matter – Solid – Liquid – Gas The phase in which a substance exists depends on it’s kinetic energy and the strength of its attractive forces. – Solids-strong attractive forces, low kinetic energy, little to no molecular movement – Liquid-enough kinetic energy to overcome the attractive forces; more molecular movement. – Gas-high kinetic energy, little to no attractive forces; maximum movement

16. 16 2.6 Chemical changes-Forming new kinds of matter Chemical reactions – Creating different chemical substances by forming and breaking chemical bonds. – Remember: Atoms form chemical bonds to fill their outermost electron energy levels, achieving stability. There are several types of chemical bonds. – We will discuss: Ionic bonds Covalent bonds

17. Ionic bonds Atoms can gain or lose electrons to achieve a full outermost energy level. – Atoms with charge are called ions. – When an atom gives away an electron, it ends up with more protons than electrons and gains a positive charge; cation – When an atom accepts an electron, it ends up with more electrons than protons and gains a negative charge; anion – This process is called ionization. An ionic bond – The attraction between oppositely charged ions Example: NaCl – Sodium (Na) has one electron in its outer energy level. – Chloride has seven electrons in its outer energy level. – Sodium donates an electron to chloride, each achieving stability. – The positively charged sodium is attracted to the negatively charged chloride.

18. 18 Ion formation

19. 19 Covalent bonds Atoms can achieve full outermost energy levels by sharing electrons instead of exchanging them. A covalent bond is formed by the sharing of electrons. – The atoms sharing electrons sit close enough together so that their outer energy levels overlap. – Single covalent bond-one pair of electrons is shared. H 2 – Double covalent bond-two pairs of electrons are shared. ethylene – Triple covalent bond-three pairs of electrons are shared. N 2

20. 20 Covalent bonds

21. 21 Fig. 2.10 Ethylene and the Ripening Process

22. 22 2.7 Water: The essence of life Water has special properties that make it an essential molecule for life. –H2O–H2O – Electrons are shared unequally by hydrogen and oxygen. This is a polar covalent bond. Oxygen has more protons than hydrogen. – The electrons spend more time around oxygen than around hydrogen. – The oxygen end of water is more negative. – The hydrogen end of water is more positive.

23. 23 Hydrogen bonds The positive hydrogen end of one polar molecule is attracted to the negative end of another polar molecule. – This attraction is a hydrogen bond. Hydrogen bonds hold molecules together. – Since they do not hold atoms together, they are not considered true chemical bonds. Hydrogen bonds are very important in biology. – They stabilize the structure of DNA and proteins. – Water molecules can “stick” together with hydrogen bonds.

24. 24 Hydrogen bonds

25. 25 Outlooks 2.1 Water and Life

26. Water and life The following properties of water make it essential for life: – High surface tension Water molecules stick to each other via hydrogen bonds. Capillary action moves water through streams, soil, animals and plants. – High heat of vaporization It requires a lot of heat to break the hydrogen bonds holding water together. Large bodies of water absorb a lot of heat. – Temperate climates – Evaporative cooling

27. 27 Water and life – Unusual density properties Ice is less dense than water, so ice floats. Allows aquatic life to survive in cold climates. – The universal solvent Water can form hydrogen bonds with any polar or ionic compound. Therefore, many things can be dissolved in water.

28. 28 Mixtures and solutions A mixture – matter that contains two or more substances that are not in set proportions. A solution is a homogeneous mixture of ions or molecules of two or more substances. – Components are distributed equally throughout. – The process of making a solution is called dissolving. – The solvent is the substance present in the largest amount. Frequently a liquid – The solutes are the substances present in smaller amounts. Aqueous solutions are solids, liquids or gases dissolved in water.

29. 29 Mixtures vs. Pure substances

30. 30 2.8 Chemical reactions A chemical change: – When the bonds of compounds are made or broken, new materials with new properties are produced. – Happens via chemical reactions. In a chemical reaction the elements remain the same, but the compounds they form and their properties are different.

31. 31 Chemical reactions and energy Chemical reactions produce new compounds with less or more potential energy. – Energy is released when compounds are made with less potential energy. – Energy is used to make compounds with more potential energy.

32. 32 Chemical equations A chemical equation is a method of describing what happens in a chemical reaction. – For example, photosynthesis is described by the following equation: Energy + 6CO 2 + 6H 2 O → C 6 H 12 O 6 + 6H 2 O Reactants-substances that are changed, usually on the left side of the equation. Products-new chemical substances formed, usually on the right side of the equation.

33. 33 Five important chemical reactions in biology 1. Oxidation–reduction 2. Dehydration synthesis 3. Hydrolysis 4. Phosphorylation 5. Acid–base reactions

34. Oxidation-reduction reactions – reactions in which electrons (and their energy) are transferred from one atom to another. – Oxidation An atom loses an electron. – Reduction An atom gains an electron. For oxidation to occur, reduction must also occur. Example: – Respiration Sugar is oxidized to form carbon dioxide and oxygen is reduced to form water. Energy is released in the process. C 6 H 12 O 6 + 6O 2 → 6H 2 O + 6CO 2 + Energy Sugar + oxygen → water+ carbon dioxide + energy

35. 35 Dehydration synthesis reaction When two small molecules are joined to form a larger molecule, – a molecule of water is released. Example: – Joining amino acids to form proteins. NH 2 CH 2 CO-OH + H-NH CH 2 CO-OH  NH 2 CH 2 CO-NH CH 2 CO-OH + H-OH amino acid 1 + amino acid 2 = protein + water

36. 36 Hydrolysis reactions When a larger molecule is broken down into smaller parts, – a water molecule is split. – Opposite of a dehydration synthesis. Example: – Digesting proteins into amino acids. NH 2 CH 2 CO-NH CH 2 CO-OH + H-OH  NH 2 CH 2 CO-OH + H-NH CH 2 CO-OH Protein + water = amino acid 1 + amino acid 2

37. Phosphorylation reactions When phosphate groups are added to other molecules, – phosphate groups are clusters of oxygen and phosphate atoms. Bonds between phosphate groups and other molecules contain high potential energy. – When these bonds are broken, the energy that is released can be used by the cell to do work. – Phosphorylation reactions are commonly used to transfer potential energy. Q-P + Z  Q + Z-P

38. 38 Fig. 2.13 Phosphorylation and Muscle Contractions

39. 39 Acid-base reactions Occurs when ions from an acid interact with ions from a base. This type of reaction allows harmful acids and bases to neutralize one another. H + Cl - + Na + OH - → Na + Cl - + H + OH - Hydrocloric + Sodium Sodium + Water acid hydroxide chloride

40. Acids, bases and salts An acid – ionic compounds that release hydrogen ions (H + ) into a solution. – Phosphoric acid, hydrochloric acid A base – Compounds that release hydroxide ions (OH-) into a solution. – Sodium hydroxide, ammonia Because bases are negatively charged, they will react with a positively charged hydrogen in solution. The strength of an acid or base is determined by how completely it will dissociate in water. – Strong acids release almost all of their hydrogen ions into water. – Strong bases release almost all of their hydroxide ions into water.

41. 41 Salts Neither acids or bases Usually formed when acids and bases react – The dissociated hydrogen ions and hydroxide ions join to form water. – The remaining ions form ionic bonds creating a salt. – This is an example of neutralization. H + Cl - + Na + OH - → Na + Cl - + H + OH - Hydrocloric + Sodium Sodium + Water acid hydroxide chloride

42. 42 Some common acids, bases and salts

43. 43 pH A measure of hydrogen ion concentration. Solutions with high hydrogen ion concentrations – have low pH. – are acidic. Solutions with low hydrogen ion concentrations – have a high pH. – are basic. There is a 10-fold difference in hydrogen ion concentration between solutions that differ by one pH unit. – A solution with pH 4 has ten times as many hydrogen ions as a solution with pH 5.

44. 44 The pH scale