Life, Chemistry, and Water

Chapter 2

Why It Matters

Bioremediation of selenium

Fig. 2-1, p. 22

2.1 The Organization of Matter: Elements and Atoms

Living organisms are composed of about 25 key elements

Elements are composed of atoms, which combine to form molecules

Matter

Matter occupies space, has mass, is composed of elements

Elements cannot be broken down into simpler substances• 92 naturally-occurring elements

• 15+ artificially-synthesized elements

25 Key Elements in Living Organisms

96% of the weight of living organisms = carbon, hydrogen, oxygen, and nitrogen

4% = calcium, phosphorus, potassium, sulfur, sodium, chlorine, and magnesium

0.01% = nine trace elements vital to biological functions

Proportions of Elements

Fig. 2-2, p. 23

Atoms and Molecules

Atoms • Smallest units of elements

Molecules • Formed from atoms

• Combined in fixed numbers and ratios

Compounds • Molecules with different component atoms

2.2 Atomic Structure

The atomic nucleus contains protons and neutrons

The nuclei of some atoms are unstable and tend to break down to form simpler atoms

The electrons of an atom occupy orbitals around the nucleus

2.2 (cont.)

Orbitals occur in discrete layers around an atomic nucleus

The number of electrons in the outermost energy level of an atom determines its chemical activity

Atomic Structure

Atomic nucleus contains protons and neutrons Electrons travel around nucleus in orbitals

Fig. 2-3, p. 24

Animation: Electron arrangements in atoms

Atomic Nucleus

Protons• Positively charged

Atomic number • Number of protons in an element

Neutrons• Uncharged

Atomic Mass

Atomic mass = mass of protons + neutrons (electrons have insignificant mass)• Proton = 1 dalton = 1.66 X 10-24 grams

• Neutron = 1 dalton

Isotopes

Atoms of an element with differing numbers of neutrons

Differ in physical but not chemical properties

Fig. 2-4, p. 24

Fig. 2-4a, p. 24

3H (tritium)1 proton

2 neutronsatomic number = 1mass number = 3

Isotopes of hydrogen

1H1 proton

atomic number = 1mass number = 1

2H (deuterium)1 proton

1 neutronatomic number = 1mass number = 2

Fig. 2-4b, p. 24

Isotopes of carbon

12C6 protons6 neutrons

atomic number = 6mass number = 12

13C 6 protons7 neutrons

atomic number = 6mass number = 13

14C6 protons8 neutrons

atomic number = 6mass number = 14

Animation: Isotopes of hydrogen

Radioisotopes

Some isotope nuclei are unstable and break down (decay)• Release particles of matter and energy

(radioactivity)

Radioisotopes decay at a steady rate• Used to estimate the age of organic material,

rocks, fossils

• Used as tracers to label molecules in chemical reactions

Electrons

Electron are negatively charged Number of electrons = number of protons Electron mass = 1/1800 dalton

Electron Orbitals

Electrons are found in regions of space called energy levels (shells)

Within each energy level, electrons are grouped into electron orbitals

Electron Orbitals

1s = lowest energy level• 1 = closest to the nucleus

• s = spherical shape

• Holds up to 2 electrons

Electron Orbitals

Second energy level• One 2s orbital (spherical shape, up to 2

electrons)

• Three 2p orbitals (dumbbell shape, up to 2 electrons in each)

Electrons Orbitals

Neon atom = 10 electron

Animation: The shell model of electron distribution

Electron Orbitals

Third energy level• Up to 18 electrons in 9 orbitals

Fourth energy level• Up to 32 electrons in 16 orbitals

Outermost orbital typically has 1 to 8 electrons in 4 orbitals (valence electrons)

Electron Orbitals

Fig. 2-6, p. 27

Fig. 2-6, p. 27

Hydrogen (H)

First energy level

Neon (Ne)Lithium (Li) Beryllium (Be)

Boron (B) Carbon (C) Nitrogen (N) Oxygen (O) Fluorine (F)

Second energy level

Third energy levelArgon (Ar)Sodium (Na)

Magnesium (Mg)

Aluminum (Al) Silicon (Si)

Phosphorus (P) Sulfur (S) Chlorine (Cl)

Helium (He)

Energy level 1Energy level 2Energy level 3

Elements not found

Amount in living organisms

Common elementsTrace elements

Number of electrons in energy levelsAtomic numberNumber of electrons (e–)Number of protons (p+)

Animation: Predicting the number of bonds of elements

Valence Electrons

If outermost energy level filled, atoms are stable and unreactive

Atoms tend to lose, gain, or share electrons to fill the outermost energy level

Leads to chemical bonds and forces that hold atoms together in a molecule

2.3 Chemical Bonds

Ionic bonds are multidirectional and vary in strength

Covalent bonds are formed by electrons in shared orbitals

Unequal electron sharing results in polarity

Polar molecules tend to associate with each other and exclude nonpolar molecules

2.3 (cont.)

Hydrogen bonds also involve unequal electron sharing

Van der Waals forces are weak attractions over very short distances

Bonds form and break in chemical reactions

Ions (1)

Charged atoms• Cation: Positively charged ion

• Anion: Negatively charged ion

Ionic Bond• Forms between atoms that gain or lose valence

electrons completely

Ions (2)

One atom loses an electron and becomes positively charged• Na+ :11 protons + 10 electrons

One atom gains an electron and becomes negatively charged• Cl– :17 protons + 18 electrons

Ionic Bond

Fig. 2-7, p. 28

Fig. 2-7a, p. 28

a. Ionic bond formation between sodium and chlorine

Cl–

Electron loss Electron gain

Sodiumatom11 e–

11 p+

Na

Sodiumion10 e–

11 p+

Na+

Chlorineatom17 e–

17 p+

Chlorineion18 e–

17 p+

Cl

Fig. 2-7a, p. 28Cl–

Sodiumion10 e–

11 p+

Na+

Chlorineion18 e–

17 p+

Electron loss Electron gain

Sodiumatom11 e–

11 p+

Na

Chlorineatom17 e–

17 p+

Cl

Stepped Art

Fig. 2-7b, p. 28

b. Crystals of sodium chloride (NaCl)

Cl–

Na+

Ionic Bond

Exerts attractive force over greater distance than other bonds

Attractive force extends in all directions

Varies in strength depending on presence of other charged substances

Covalent Bond

Two atoms share a pair of electrons

Shared orbitals occur at discrete angles and directions

Covalent Bonds

Fig. 2-8, p. 30

Fig. 2-8a, p. 30

a. Shared orbitals of methane (CH4)

Fig. 2-8b, p. 30

b. Space-filling model of methane

Fig. 2-8c, p. 30

c. A carbon “building block” usedto make molecular models

Fig. 2-8d, p. 30

d. Cholesterol

Hydrogen

Carbon

Oxygen

Animation: How atoms bond

Unequal Electron Sharing

Electronegativity • Measure of atom’s attractions for electrons

shared in a chemical bond

Nonpolar covalent bond • Electrons shared equally

Polar covalent bond • Electrons shared unequally

Unequal Electron Sharing

Water is a polar molecule

Fig. 2-9, p. 30

Polar Molecules

Tend to associate with other polar molecules and to exclude nonpolar molecules

Polar molecules that associate readily with water are hydrophilic (“water preferring”)

Nonpolar molecules excluded by water are hydrophobic (“water avoiding”)

Hydrogen Bond

Unequal electron sharing between a hydrogen atom and another atom (oxygen, nitrogen, sulfur)• Hydrogen gets partial positive charge, other atom

gets partial negative charge

• Charges attract to form hydrogen bond

Hydrogen Bond

Weak bond, useful in stabilizing large biological molecules such as proteins

Van der Waals Forces

Natural changes in electron density of molecules

Regions of positive and negative charge• Cause molecules to stick together briefly

Weaker than hydrogen bonds• Help stabilize large biological molecules such as

proteins

Van der Waals Forces

Gecko toes:

Fig. 2-11, p. 32

Fig. 2-11, p. 32

d. Pads on a setaa. Gecko inverted on glass b. Gecko toe c. Setae on toe

Chemical Reactions

When molecules form or break chemical bonds Reactants enter into a chemical reaction Products leave a reaction

6 CO2 + 6 H2O → C6H12O6 + 6 O2

carbondioxide

water a sugar molecularoxygen

reactants products

2.4 Hydrogen Bonds and the Properties of Water

Lattice of hydrogen bonds gives water unusual properties

Hydrogen-bond lattice of water forms polar and nonpolar environments in and around cells

The small size and polarity of its molecules makes water a good solvent

The hydrogen-bond lattice gives water other life-sustaining properties as well

Hydrogen-Bond Lattice of Water

Water lattice • Neighboring water molecules form hydrogen

bonds temporarily

• Difficult for nonpolar substances to penetrate the lattice

• Polar or charged substances penetrate easily

Ice Lattice

Ice lattice • A rigid, crystalline structure

• Water molecules in the ice lattice are spread farther apart than in liquid water

• Ice is less dense than water and floats

Hydrogen-Bond Lattice of Water

Fig. 2-12, p. 33

Fig. 2-12, p. 33

b. Hydrogen-bond lattice of icea. Hydrogen-bond lattice of liquid water

Animation: Structure of water

Membrane molecules have one polar end and one nonpolar (lipid) end

In water, lipid molecules are forced into a double layer (bilayer) that forms the membrane• Hydrophilic ends face water

• Hydrophobic ends associate inside

Water, Lipids, and Membranes

Membranes: The Lipid Bilayer

Fig. 2-13, p. 34

Fig. 2-13, p. 34

Membranecovering cellsurface

Watermolecule

Polar end ofmembranemolecule

Nonpolar end of membranemolecule

Polar water solutionoutside cell

Nonpolar regioninside membrane

Polar water solutioninside cell

Hydration Layer

Water forms hydration layer over surfaces of polar and charged biological molecules, particularly proteins

Separates ions and molecules from each other so they can enter a solution• Water = solvent

• Dissolved substance = solute

Hydration Layer

Hydration layers around Na+ and Cl– ions keep salt in solution

Fig. 2-14, p. 34

Animation: Spheres of hydration

Solutions

Concentration • Number of ions or molecules per unit volume

Mole• 6.022 X 1023 molecules (Avogadro’s number)• 1 mole of substance = molecular weight in grams• 1 mole NaCl = 23 + 35 = 58g

Molarity (moles per liter)• 1 molar solution = 1 mole/liter = 1M• 1M NaCl = 58 g/L

High Specific Heat

Water temperature increases slowly as heat is added• Must break hydrogen bonds to allow water

molecules to move faster

• Helps moderate and stabilize temperature of organisms and environment

• c (calorie) = heat to raise 1g of water 1°C

• C (Calorie) = 1,000 calories = 1 kilocalorie (kcal)

High Heat of Vaporization

Water absorbs a large amount of heat to break loose from liquid water and form a gas• Some animals sweat (water evaporation cools

skin and underlying blood vessels)

• Plants evaporate water from leaves (cools heat absorbed from sunlight)

Water Molecules Resist Separation

Cohesion • Attraction between water molecules

Adhesion • Attraction of water molecules to surfaces with

charged or polar molecules

Surface Tension

Forms at surface of water in contact with air

Hydrogen bonds resist stretching, giving surface strength

Fig. 2-15, p. 36

Fig. 2-15a, p. 36

H2O

a.

Air Water surface

Fig. 2-15b, p. 36

2.5 Water Ionization and Acids, Bases, and Buffers

Substances act as acids or bases by altering the concentrations of H+ and OH− ions in water

Buffers help keep pH under control

Water Ionization

Water dissociates to form ions:

H2O ↔ H+ + OH–

H+ (protons) = hydrogen ions OH– = hydroxide ions In pure water, concentration of H+ = OH–

Acids and Bases

Acids release H+ as they dissolve in water• Solution becomes acidic

HCl ↔ H+ + Cl–

Bases gather H+ or release OH− in solution• Solution becomes basic

NaOH ↔ Na+ + OH–

pH Scale

Measures relative concentrations of H+ and OH− (acidity) in a water solution on a scale of 0 to 14

pH = –log10[H+]

Pure water: [H+] = [OH−] = 1 X 10-7 M• pH 7 (neutral) = –log10[1 X 10-7]

• pH < 7 is acidic, pH > 7 is basic

pH Scale

Fig. 2-16, p. 37

Fig. 2-16, p. 37

Hydrochloric acid (HCl)

Gastric fluid (1.0–3.0)

Lemon juice, cola drinks, some acid rain

Vinegar, wine, beer, orangesTomatoesBananasBlack coffee

Sodium hydroxide (NaOH)

BreadTypical rainwaterUrine (5.0–7.0)Milk (6.6)Pure water [H+] = [OH–]Blood (7.3–7.5)Egg white (8.0)Seawater (7.8–8.3)Baking sodaPhosphate detergents, bleach, antacids

Soapy solutions,milk of magnesiaHousehold ammonia (10.5–11.9)

Hair remover

Oven cleaner

Animation: The pH scale

Acid Rain

Fig. 2-17, p. 38

Buffers (1)

Regulate pH of living cells

Absorb or release H+ to compensate for changes in H+ concentration

H2CO3 ↔ HCO3– + H+

Buffers (2)

If cell is too acidic• Push reaction to the left

• Remove some H+ ions

If cell is too basic• Push reaction to the right

• Add more H+ ions

Help keep cells close to neutral pH

Hyperventilation

Breathing too fast removes buffering capacity from the blood

Fig. 2-18, p. 39

Fig. 2-18, p. 39

Adverse physiological effects,

such as dizziness, visual impairment, fainting,

seizures, or death

Rapidbreathing

Blood CO2concentration

decreases

Blood carbonic acid level decreases

Blood pHchanges fromnormal levels

Fig. 2-18, p. 39

Adverse physiological effects,such as dizziness, visual

impairment, fainting, seizures, or death

Rapidbreathing

Blood CO2concentration

decreases

Blood carbonic acid level decreases

Blood pHchanges fromnormal levels

Stepped Art

Video: Covalent bonds

Video: The wine of life

Video: Isotopes