Covalent bonds – where electrons are shared

1. Covalent bonds – where electrons are shared • Typically the strongest bonds in biological systems. • Can be polar (where electrons are not equally shared) or non-polar (electrons are equally shared).

2. Hydrogen atoms (2 H) In each hydrogen atom, the single electron is held in its orbital by its attraction to the proton in the nucleus. + + 1 2 3 When two hydrogen atoms approach each other, the electron of each atom is also attracted to the proton in the other nucleus. + + The two electrons become shared in a covalent bond, forming an H2 molecule. + + Hydrogen molecule (H2) • Formation of a covalent bond Figure 2.10

3. A molecule • Consists of two or more atoms held together by covalent bonds • A single bond • Is the sharing of one pair of valence electrons • A double bond • Is the sharing of two pairs of valence electrons

4. Name (molecular formula) Electron- shell diagram Space- filling model Structural formula Hydrogen (H2). Two hydrogen atoms can form a single bond. H H Oxygen (O2). Two oxygen atoms share two pairs of electrons to form a double bond. O O Figure 2.11 A, B • Single and double covalent bonds (a) (b)

5. Name (molecular formula) Electron- shell diagram Space- filling model Structural formula (c) Water (H2O). Two hydrogen atoms and one oxygen atom are joined by covalent bonds to produce a molecule of water. H O H (d) Methane (CH4). Four hydrogen atoms can satisfy the valence of one carbon atom, forming methane. H H H C H Figure 2.11 C, D Covalent bonding in compounds

6. Electronegativity • Is the attraction of a particular kind of atom for the electrons in a covalent bond • The more electronegative an atom • The more strongly it pulls shared electrons toward itself

7. A nonpolar covalent bond • The atoms have similar electronegativities • Share the electron equally • Common in hydrocarbons

8. A polar covalent bond Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen. d– This results in a partial negative charge on the oxygen and a partial positive charge on the hydrogens. O H H d+ d+ H2O • The atoms have differing electronegativities • Share the electrons unequally Figure 2.12

9. Ionic Bonds • Electron transfer between two atoms creates ions • Ions • Are atoms with more or fewer electrons than usual • Are charged atoms • An anion • Is negatively charged ions • A cation • Is positively charged

10. The lone valence electron of a sodium atom is transferred to join the 7 valence electrons of a chlorine atom. Each resulting ion has a completed valence shell. An ionic bond can form between the oppositely charged ions. – + 1 2 Cl Na Na Cl Cl– Chloride ion (an anion) Na+ Sodium on (a cation) Na Sodium atom (an uncharged atom) Cl Chlorine atom (an uncharged atom) Sodium chloride (NaCl) An ionic bond An attraction between anions and cations These bonds are strong in crystal form, but weak in water Figure 2.13

11. Na+ Cl– Figure 2.14 Ionic compounds • Are often called salts, which may form crystals

12. H Water (H2O) O A hydrogen bond results from the attraction between the partial positive charge on the hydrogen atom of water and the partial negative charge on the nitrogen atom of ammonia. H  +  – Ammonia (NH3) N H H d+ + H Figure 2.15 Weak Chemical Bonds – form due to differences in polarity • Hydrogen bonds • Form when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom  –  +  +

13. Van der Waals Interactions • Van der Waals interactions • Occur when transiently positive and negative regions of molecules attract each other • Weak chemical bonds • Reinforce the shapes of large molecules • Help molecules adhere to each other

14. BSC 2010 - Exam I Lectures and Text Pages • I. Intro to Biology (2-29) • II. Chemistry of Life • Chemistry review (30-46) • Water (47-57) • Carbon (58-67) • Macromolecules (68-91) • III. Cells and Membranes • Cell structure (92-123) • Membranes (124-140) • IV. Introductory Biochemistry • Energy and Metabolism (141-159) • Cellular Respiration (160-180) • Photosynthesis (181-200)

15. Water – The Solvent of Life (Ch. 3) Cells are made of 70-95% water, the “SOLVENT OF LIFE”.All living things require water more than any other substance. • Solvent - • Solute - • Aqueous -

16. Figure 3.1 • Three-quarters of the Earth’s surface is submerged in water • The abundance of water is the main reason the Earth is habitable

17. – Hydrogenbonds + H – + H + –  – + Figure 3.2 The water molecule is a polar molecule • The polarity of water molecules • Allows them to form hydrogen bonds with each other (negative O ends are attracted to positive H ends) • Contributes to the various properties water exhibits

18. Emergent Properties of Water Contribute to Life • A. cohesion: (related properties: surface tension and adhesion) • B. Water tends to resist rupturing. (related to cohesion) • C. Water resists changes in temperature. • D. Water expands when it freezes. • E. Water is a versatile solvent.

19. Cohesion • Water molecules exhibit cohesion • Cohesion • Is the bonding of a high percentage of the molecules to neighboring molecules • Water molecules stick together due to hydrogen bonding • Causes surface tension and adhesion.

20. Water conducting cells 100 µm Figure 3.3 Cohesion Helps pull water up through the microscopic vessels of plants. Water molecules stick to each other and to the walls of the xylem.

21. Figure 3.4 Surface tension Is a measure of how hard it is to break the surface of a liquid.

22. Moderation of Temperature • Water moderates air temperature • This is very important for the maintenance of homeostasis by living organisms. • Also - ~75% of the earth is covered with water, this helps stabilize climate. • Water absorbs heat from air that is warmer and releases the stored heat to air that is cooler

23. Water’s High Specific Heat • The specific heat of a substance • Is the amount of heat that must be absorbed or lost for 1 gram of that substance to change its temperature by 1ºC

24. Water’s High Specific Heat • Water has a high specific heat, which allows it to minimize temperature fluctuations to within limits that permit life. • Heat is absorbed when hydrogen bonds break. • Heat is released when hydrogen bonds form.

25. Evaporative Cooling • Heat of vaporization • Is the quantity of heat a liquid must absorb for 1 gram of it to be converted from a liquid to a gas • Evaporative cooling • Is due to water’s high heat of vaporization • Allows water to cool a surface

26. Hydrogen bond Liquid water Hydrogen bonds constantly break and re-form Ice Hydrogen bonds are stable Figure 3.5 Ice Floats • The hydrogen bonds in ice • Are more “ordered” than in liquid water, making ice less dense

27. Insulation of Bodies of Water by Floating Ice • Solid water, or ice • Is less dense than liquid water • Floats in liquid water • Allows life to exist in frozen lakes and ponds.

28. The Solvent of Life • Water is a versatile solvent due to its polarity • It can form aqueous solutions

29. Negative oxygen regions of polar water molecules are attracted to sodium cations (Na+). – Na+ + + – + – – Positive hydrogen regions of water molecules cling to chloride anions (Cl–). Na+ – + + Cl – Cl– + – – + – + – – Figure 3.6 Forming solutions with ionic solutes. • The different regions of the polar water molecule can interact with ionic compounds and dissolve them.

30. This oxygen is attracted to a slight positive charge on the lysozyme molecule. – + This hydrogen is attracted to a slight negative charge on the lysozyme molecule. (b) Lysozyme molecule (purple) in an aqueous environment such as tears or saliva (a) Lysozyme molecule in a nonaqueous environment (c) Ionic and polar regions on the protein’s Surface attract water molecules. Figure 3.7 Forming solutions with polar solutes. • Water can also interact with polar molecules such as proteins

31. Hydrophilic and Hydrophobic Substances • Some substances are attracted to water and others are not. • A hydrophilic substance • Has an affinity for water. Ions and polar molecules. • A hydrophobic substance is not attracted to water. • Nonpolar molecules.

32. + – H H H + H H H H H Hydroxide ion (OH–) Hydronium ion (H3O+) Figure on p. 53 of water dissociating Life is sensitive to pH (Acids and Bases) • Water can dissociate • Into hydronium ions and hydroxide ions • Changes in the concentration of these ions • Can have a great affect on living organisms

33. Acids and Bases • An acid • Is any substance that increases the hydrogen ion concentration of a solution (donates protons) • A base • Is any substance that reduces the hydrogen ion concentration of a solution (accepts protons)

34. The pH Scale • The pH of a solution • Is determined by the relative concentration of hydrogen ions • Is low in an acid • Is high in a base Most biological solutions range from pH of 6-8, but there are exceptions (stomach acids pH 1-2)

35. pH Scale 0 1 Battery acid 2 Digestive (stomach) juice, lemon juice Vinegar, beer, wine, cola 3 Increasingly Acidic [H+] > [OH–] 4 Tomato juice 5 Black coffee Rainwater 6 Urine Neutral [H+] = [OH–] 7 Pure water Human blood 8 Seawater 9 10 Increasingly Basic [H+] < [OH–] Milk of magnesia 11 Household ammonia 12 Household bleach 13 Oven cleaner 14 Figure 3.8 • The pH scale and pH values of various aqueous solutions

36. Buffers • The internal pH of most living cells • Must remain close to pH 7

37. Buffers • Are substances that minimize changes in the concentrations of hydrogen and hydroxide ions in a solution • Consist of a weak acid-base pair that reversibly combines with hydrogen ions