Chemical Foundations for Cells Chapter 6

Chemical Benefits and Costs Understanding of chemistry provides fertilizers, medicines, etc. Chemical pollutants damage ecosystems

Elements Fundamental forms of matter Can’t be broken apart by normal means 92 occur naturally on Earth

Most Common Elements in Living Organisms Oxygen = 65% Hydrogen = 9.5% Carbon = 18.5% Nitrogen = 3.3%

What Are Atoms? Smallest particles that retain properties of an element Made up of subatomic particles: –Protons (+) –Electrons (-) –Neutrons (no charge)

HYDROGEN HELIUM electron proton neutron Hydrogen and Helium Atoms

Atomic Number Number of protons All atoms of an element have the same atomic number Atomic number of hydrogen = 1 Atomic number of carbon = 6

Mass Number Number of protons + Number of neutrons Isotopes vary in mass number

Isotopes Atoms of an element with different numbers of neutrons (different mass numbers) Carbon 12 has 6 protons, 6 neutrons Carbon 14 has 6 protons, 8 neutrons

What Determines Whether Atoms Will Interact? The number and arrangement of their electrons Atoms seek to be more stable – complete orbitals

Electrons Carry a negative charge Repel one another Are attracted to protons in the nucleus Move in orbitals - volumes of space that surround the nucleus Z X When all p orbitals are full y

Electron Orbitals Orbitals can hold up to two electrons Atoms differ in the number of occupied orbitals Orbitals closest to nucleus are lower energy and are filled first

Shell Model First shell –Lowest energy –Holds 1 orbital with up to 2 electrons Second shell –4 orbitals hold up to 8 electrons CALCIUM 20p+, 20e -

Electron Vacancies Unfilled shells make atoms likely to react Hydrogen, carbon, oxygen, and nitrogen all have vacancies in their outer shells CARBON 6p+, 6e - NITROGEN 7p+, 7e - HYDROGEN 1p+, 1e -

Chemical Bonds, Molecules, & Compounds Bond is union between electron structures of atoms Atoms bond to form molecules Molecules may contain atoms of only one element - O 2 Molecules of compounds contain more than one element - H 2 O

Chemical Bookkeeping Use symbols for elements when writing formulas Formula for glucose is C 6 H 12 O 6 – 6 carbon atoms – 12 hydrogen atoms – 6 oxygen atoms

Chemical Bookkeeping Chemical equation shows reaction Reactants ---> Products Equation for photosynthesis: 6CO 2 + 6H 2 O ---> + C 6 H 12 O H 2 O

Important Bonds in Biological Molecules Ionic Bonds Covalent Bonds Hydrogen Bonds

Covalent Bonding Atoms share a pair or pairs of electrons to fill outermost shell Single covalent bond Double covalent bond Triple covalent bond

Nonpolar Covalent Bonds Atoms share electrons equally Nuclei of atoms have same number of protons Example: Hydrogen gas (H-H)

Polar Covalent Bonds Number of protons in nuclei of participating atoms is NOT equal Electrons spend more time near nucleus with most protons Water - Electrons more attracted to O nucleus than to H nuclei

Ion Formation Atom has equal number of electrons and protons - no net charge Atom loses electron(s), becomes positively charged ion Atom gains electron(s), becomes negatively charged ion

Ionic Bonding One atom loses electrons, becomes positively charged ion Another atom gains these electrons, becomes negatively charged ion Charge difference attracts the two ions to each other

Formation of NaCl Sodium atom (Na) –Outer shell has one electron Chlorine atom (Cl) –Outer shell has seven electrons Na transfers electron to Cl forming Na + and Cl - Ions remain together as NaCl

7mm SODIUM ATOM 11 p + 11 e - SODIUM ION 11 p + 10 e - electron transfer CHLORINE ATOM 17 p + 17 e - CHLORINE ION 17 p + 18 e - Fig. 2.10a, p. 26 Formation of NaCl

Hydrogen Bonding Molecule held together by polar covalent bonds has no NET charge However, atoms of the molecule carry different charges Atom in one polar covalent molecule can be attracted to oppositely charged atom in another such molecule

one large molecule another large molecule a large molecule twisted back on itself Fig. 2.12, p. 27 Examples of Hydrogen Bonds

Hydrogen Ions: H + Unbound protons Have important biological effects Form when water ionizes

The pH Scale Measures H + concentration of fluid Change of 1 on scale means 10X change in H + concentration Highest H + Lowest H Acidic Neutral Basic

Acids & Bases Acids –Donate H + when dissolved in water –Acidic solutions have pH < 7 Bases –Accept H + when dissolved in water –Acidic solutions have pH > 7

Properties of Water Polarity Temperature-Stabilizing Cohesive Solvent

Water Is a Polar Covalent Molecule Molecule has no net charge Oxygen end has a slight negative charge Hydrogen end has a slight positive charge O HH

O H H O H H + _ _ + + Liquid Water

Water Cohesion Hydrogen bonding holds molecules in liquid water together Creates surface tension Allows water to move as continuous column upward through stems of plants

Temperature-Stabilizing Effects Liquid water can absorb much heat before its temperature rises Why? Much of the added energy disrupts hydrogen bonding rather than increasing the movement of molecules

Why Ice Floats In ice, hydrogen bonds lock molecules in a lattice Water molecules in lattice are spaced farther apart then those in liquid water Ice is less dense than water

Water Is a Good Solvent Ions and polar molecules dissolve easily in water When solute dissolves, water molecules cluster around its ions or molecules and keep them separated

Spheres of Hydration

Diffusion Brownian motion – molecules are in constant motion Diffusion – movement from area of high concentration to area of low concentration –Affected by Concentration Temperature or agitation Pressure

Dynamic Equilibrium Molecules are still in motion No net gain or loss of molecules Living systems seek to achieve

Organic Compounds Hydrogen and other elements covalently bonded to carbon Carbohydrates Lipids Proteins Nucleic Acids

Carbon’s Bonding Behavior Outer shell of carbon has 4 electrons; can hold 8 Each carbon atom can form covalent bonds with up to four atoms

Bonding Arrangements Carbon atoms can form chains or rings Other atoms project from the carbon backbone

Condensation Reactions Form polymers from subunits Enzymes remove -OH from one molecule, H from another, form bond between two molecules Discarded atoms can join to form water

Fig. 3.4a, p. 37 enzyme action at functional groups CONDENSATION

Hydrolysis A type of cleavage reaction Breaks polymers into smaller units Enzymes split molecules into two or more parts An -OH group and an H atom derived from water are attached at exposed sites

enzyme action at functional groups HYDROLYSIS Fig. 3.4b, p. 37

Carbohydrates – energy source Monosaccharides (simple sugars) Disaccharides (two simple sugars) Polysaccharides (complex carbohydrates)

Monosaccharides Simplest carbohydrates Most are sweet tasting, water soluble Most have 5- or 6-carbon backbone Glucose (6 C)Fructose (6 C) Ribose (5 C)Deoxyribose (5 C)

Two Monosaccharides glucosefructose

Disaccharides Two monosaccharides covalently bonded Formed by condensation reaction + H 2 O glucosefructose sucrose

Polysaccharides Straight or branched chains of many saccharides Most common are composed entirely of glucose –Cellulose tough, indigestible structural material in plants –Starch easily digested storage form in plants –Glycogen sugar storage form in animals –Chitin structural material for hard parts of invertebrates cell walls of many fungi

Most include fatty acids –Fats –Phospholipids –Waxes Tend to be insoluble in water Energy source, insulation & protection Lipids

Fatty Acids Carboxyl group (-COOH) at one end Carbon backbone (up to 36 C atoms) –Saturated - Single bonds between carbons –Unsaturated - One or more double bonds

Three Fatty Acids What difference does the double bond make? stearic acidoleic acidlinolenic acid

Fats Fatty acid(s) attached to glycerol Triglycerides are most common

Proteins Carbon, hydrogen, oxygen, nitrogen & sulfur Amino acid building blocks AA linked by peptide bonds Enzymes Build tissue

Enzymes Protein Act as catalyst –Helps reaction happen faster or at lower temperatures Substrate specific shapes –Lock & key system –Recycled; not used up

Denaturation Disruption of three- dimensional shape Breakage of weak bonds Causes of denaturation: –pH –Temperature Destroying protein shape disrupts function

Nucleic Acids Carbon, hydrogen, oxygen, nitrogen & phosphorus Nucleotides – building blocks DNA, RNA Genetic information