life can be organized into a hierarchy of structural levels. at each successive level additional...

Life can be organized into a hierarchy of structural levels.

At each successive level additional emergent properties appear.

Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings

Fig. 2.1

Organisms are composed of matter. Matter is anything that takes up space

and has mass. An element is a substance that

cannot be broken down to other substances by chemical reactions. There are 92 naturally-occurring

elements. Each element has a unique symbol,

usually from the first one or two letters of the name, often from Latin or German.

1. Matter consists of chemical elements in pure form and in combinations called compounds

A compound is a substance consisting of two or more elements in a fixed ratio. Table salt (sodium chloride or NaCl) is a

compound with equal numbers of chlorine and sodium atoms.

While pure sodium is a metal and chlorine is a gas, their combination forms an edible compound, an emergent property.

Fig. 2.2

About 25 of the 92 natural elements are known to be essential for life. Four elements - carbon (C), oxygen (O),

hydrogen (H), and nitrogen (N) - make up 96% of living matter.

Most of the remaining 4% of an organism’s weight consists of phosphorus (P), sulfur (S), calcium (Ca), and potassium (K).

2. Life requires about 25 chemicalelements


Trace elements are required by an organism but only in minute quantities. Some trace elements, like iron (Fe), are

required by all organisms. Other trace elements are

required only by some species. For example, a daily intake

of 0.15 milligrams of iodine is required for normal activity of the human thyroid gland.


Fig. 2.4

Atomic number= number of protons Atomic mass= atomic weight= protons +

neutrons – measured in daltons Helium= 4.003 daltons

Atoms with different amounts of neutrons are isotopes

Nonradioactive carbon-12 Nonradioactive carbon-13 Radioactive carbon-14

6 electrons6 protons6 neutrons



Section 2-1

Figure 2-2 Isotopes of Carbon


Fig. 2.6

Radioactive isotopes are also used to diagnose medical disorders. For example, the rate of excretion in the

urine can be measured after injection into the blood of known quantity of radioactive isotope.

Also, radioactive tracers can be used with imaging instruments to monitor chemical processes in the body.


Fig. 2.7

While useful in research and medicine, the energy emitted in radioactive decay is hazardous to life. This energy can destroy cellular

molecules. The severity of damage depends on the

type and amount of energy that an organism absorbs.

Fig. 2.8

Atoms have different energy levels or shells

Moving electrons to outer orbitals increases the atom’s potential energy

The 3D space where an electron is found 90% of the time is called an orbital


Fig. 2.10

The first electron shell can hold only 2 electrons. The two electrons of Helium fill the first

shell. Atoms with more than two electrons

must place the extra electrons in higher shells. For example, Lithium with three electrons

has two in the first shell and one in the second shell.

The second shell can hold up to 8 electrons. Neon, with 10 total electrons, has two in

the first shell and eight in the second, filling both shells.


The chemical behavior of an atom depends mostly on the number of electrons in its outermost shell, the valence shell. Electrons in the valence shell are known

as valence electrons. Atoms with the same number of

valence electrons have similar chemical behavior.

An atom with a completed valence shell is unreactive.

All other atoms are chemically reactive because they have incomplete valence shells.


While the paths of electrons are often visualized as concentric paths, like planets orbiting the sun.

In reality, an electron occupies a more complex three-dimensional space, an orbital. The first shell has room for a single

spherical orbital for its pair of electrons. The second shell can pack pairs of

electrons into a spherical orbital and three p orbitals (dumbbell-shaped).


The reactivity of atoms arises from the presence of unpaired electrons in one or more orbitals of their valence shells. Electrons preferentially occupy separate

orbitals within the valence shell until forced to share orbitals.

The four valence electrons of carbon each occupy separate orbitals, but the five valence electrons of nitrogen are distributed into three unshared orbitals and one shared orbital.

When atoms interact to complete their valence shells, it is the unpaired electrons that are involved.Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings

Covalent Bonds A covalent bond

Is the sharing of a pair of valence electrons

Formation of a covalent bond

Figure 2.10

Hydrogen atoms (2 H)

Hydrogenmolecule (H2)

+ +

+ +

+ +

In each hydrogenatom, the single electronis held in its orbital byits attraction to theproton in the nucleus.

1

When two hydrogenatoms approach eachother, the electron ofeach atom is alsoattracted to the protonin the other nucleus.

2

The two electronsbecome shared in a covalent bond,forming an H2

molecule.

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

(a)

(b)

Name(molecularformula)

Electron-shell

diagram

Structuralformula

Space-fillingmodel

Hydrogen (H2). Two hydrogen atoms can form a single bond.

Oxygen (O2). Two oxygen atoms share two pairs of electrons to form a double bond.

H H

O O

Figure 2.11 A, B

Single and double covalent bonds

Name(molecularformula)

Electron-shell

diagram

Structuralformula

Space-fillingmodel

(c)

Methane (CH4). Four hydrogen atoms can satisfy the valence ofone carbonatom, formingmethane.

Water (H2O). Two hydrogenatoms and one oxygen atom arejoined by covalent bonds to produce a molecule of water.

(d)

HO

H

H H

H

H

C

Figure 2.11 C, D

Covalent bonding in compounds

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

In a nonpolar covalent bond The atoms have similar

electronegativities Share the electron equally

Figure 2.12

This results in a partial negative charge on theoxygen and apartial positivecharge onthe hydrogens.

H2O

–

O

H H+ +

Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen.

In a polar covalent bond The atoms have differing

electronegativities Share the electrons unequally

Structural formula

Molecular formula

Ionic Bond

After the transfer, both atoms are no longer neutral, but have charges and are called ions.

Sodium has one more proton than electrons and has a net positive charge. Atoms with positive charges are cations.

Chlorine has one more electron than protons and has a net negative charge. Atoms with negative charges are anions.


Fig. 2.14

An anion Is negatively charged ions

A cation Is positively charged

Weak Chemical Bonds Several types of weak chemical

bonds are important in living systems

Hydrogen Bonds

– +

+

Water(H2O)

Ammonia(NH3)

OH

H

+

–

N

HH H

A hydrogenbond results from the attraction between thepartial positive charge on the hydrogen atom of water and the partial negative charge on the nitrogen atom of ammonia.+ +

Figure 2.15

A hydrogen bond Forms when a

hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom

Van der Waals Forces

Electrons may gather on one side of an atom momentarily by chance causing one side to become slightly positive and the other negative

Weak chemical bonds Reinforce the shapes of large

molecules Help molecules adhere to each other

Molecular Shape and Function The precise shape of a molecule

Is usually very important to its function in the living cell

Is determined by the positions of its atoms’ valence orbitals

s orbital

ZThree p orbitals

X

Y

Four hybrid orbitals

(a) Hybridization of orbitals. The single s and three p orbitals of a valence shell involved in covalent bonding combine to form four teardrop-shaped hybrid orbitals. These orbitals extend to the four corners of an imaginary tetrahedron (outlined in pink).

Tetrahedron

Figure 2.16 (a)

In a covalent bond The s and p orbitals may hybridize,

creating specific molecular shapes

Space-fillingmodel

Hybrid-orbital model(with ball-and-stick

model superimposed)UnbondedElectron pair

104.5°

O

HWater (H2O)

Methane (CH4)

H

H H

H

C

O

H

H

H

C

Ball-and-stickmodel

H H

H

H

(b) Molecular shape models. Three models representing molecular shape are shown for two examples; water and methane. The positions of the hybrid orbital determine the shapes of the moleculesFigure 2.16 (b)

Molecular shape Determines how biological molecules

recognize and respond to one another with specificity

Molecules such as neural transmitters must have a specific shape to fit in the receptor molecule

Morphine

Carbon

Hydrogen

Nitrogen

Sulfur

OxygenNaturalendorphin

(a) Structures of endorphin and morphine. The boxed portion of the endorphin molecule (left) binds toreceptor molecules on target cells in the brain. The boxed portion of the morphine molecule is a close match.

(b) Binding to endorphin receptors. Endorphin receptors on the surface of a brain cell recognize and can bind to both endorphin and morphine.

Naturalendorphin

Endorphinreceptors

Morphine

Brain cell

Figure 2.17

Within a cell, weak, brief bonds between molecules are important to a variety of processes. For example, signal molecules from one

neuron use weak bonds to bind briefly to receptor molecules on the surface of a receiving neuron.

This triggers a momentary response by the recipient.

Weak interactions include ionic bonds (weak in water), hydrogen bonds, and van der Waals interactions.

3. Weak chemical bonds play important roles in the chemistry of life


Strength of Bonds Ionic (dry) -110 kcals/mole Covalent – 50 kcals/mole Hydrogen – 5 kcals/mole Ionic (aqueous) – 5 kcals/mole Vander waals – 1-2 kcals/mole Weak bonds help to shape the

molecule

Concept 2.4: Chemical reactions make and break chemical bonds

A Chemical reaction Is the making and breaking of

chemical bonds Leads to changes in the composition

of matter

Reactants Reaction Product

2 H2 O2 2 H2O

+

+

Chemical reactions Convert reactants to products

Photosynthesis Is an example of a chemical reaction

Figure 2.18

Chemical equilibrium Is reached when the forward and

reverse reaction rates are equal

Chapter 3

Water and the Fitness of the Environment

Because water tends to form hydrogen bonds with four other water molecules, it is very cohesive

& 38Page 32 & 38

Cohesion among water molecules plays a key role in the transport of water against gravity in plants. Water that evaporates from a leaf is

replaced by water from vessels in the leaf.

Hydrogen bonds cause water molecules leaving the veins to tug on molecules further down.

This upward pull is transmitted to the roots.

Adhesion, clinging of one substance to another, contributes too, as water adheres to the wall of the vessels.

Copyright © 2002 Pearson Education, Inc., publishing as Benjamin CummingsFig. 3.2

The cohesion of water creates surface tension which can be strong enough to support light animals on water

Ever try to float a paper clip on the surface of water?

Page 38 Water is adhesive to the charges found on the surface of glass and this causes capillary action.

Why does the water move up less in the larger glass tube and not at all in the plastic tube?

Atoms and molecules have kinetic energy, the energy of motion, because they are always moving. The faster that a molecule moves, the more

kinetic energy that it has. Heat is a measure of the total quantity of

kinetic energy due to molecular motion in a body of matter.

Temperature measures the intensity of heat due to the average kinetic energy of molecules. As the average speed of molecules increases, a

thermometer will record an increase in temperature.

Heat and temperature are related, but not identical.


Page 39

Water has a high heat capacity and a high vaporization temperature

These two factors causes water to regulate temperature fluctuations of the surrounding environment.

Water resists changes in temperature because it takes a lot of energy to speed up its molecules. Viewed from a different perspective, it absorbs or

releases a relatively large quantity of heat for each degree of change.

Water’s high specific heat is due to hydrogen bonding. Heat must be absorbed to break hydrogen bonds

and is released when hydrogen bonds form. Investment of one calorie of heat causes

relatively little change to the temperature of water because much of the energy is used to disrupt hydrogen bonds, not move molecules faster.


When 1 cal of heat energy is transferred to one gram (1 g) of pure liquid water, the temperature of that sample of water is raised by one degree Celsius (1º C)

1000 cal= 1 kilocalorie = 1 Calorie

The transformation of a molecule from a liquid to a gas is called vaporization or evaporation. This occurs when the molecule moves

fast enough that it can overcome the attraction (hydrogen bonds) of other molecules in the liquid.

Even in a low temperature liquid (low average kinetic energy), some molecules are moving fast enough to evaporate.

Heating a liquid increases the average kinetic energy and increases the rate of evaporation.


Heat of vaporization is the quantity of heat that a liquid must absorb for 1 g of it to be converted from the liquid to the gaseous state. Water has a relatively high heat of vaporization,

requiring about 580 cal of heat is to evaporate 1g of water at room temperature.

This is double the heat required to vaporize the same quantity of alcohol or ammonia.

This is because hydrogen bonds must be broken before a water molecule can evaporate from the liquid.

Water’s high heat of vaporization moderates climate by absorbing heat in the tropics via evaporation and releasing it at higher latitudes as rain.Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings

As a liquid evaporates, the surface of the liquid that remains behind cools - evaporative cooling. This occurs because the most energetic

molecules are the most likely to evaporate, leaving the lower kinetic energy molecules behind.

Evaporative cooling moderates temperature in lakes and ponds and prevents terrestrial organisms from overheating. Evaporation of water from the leaves of

plants or the skin of humans removes excess heat.Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings

When water reaches 0oC, water becomes locked into a crystalline lattice with each molecule bonded to to the maximum of four partners.

As ice starts to melt, some of the hydrogen bonds break and some water molecules can slip closer together than they can while in the ice state.

Ice is about 10% less dense than water at 4oC.

Fig. 3.5

Page 41

Nonpolar

Polar

A liquid that is a completely homogeneous mixture of two or more substances is called a solution. A sugar cube in a glass of water will eventually

dissolve to form a uniform mixture of sugar and water.

The dissolving agent is the solvent and the substance that is dissolved is the solute. In our example, water is the solvent and sugar the

solute. In an aqueous solution, water is the

solvent. Water is not a universal solvent, but it is very

versatile because of the polarity of water molecules.

5. Water is the solvent of life


Water is an effective solvent because it so readily forms hydrogen bonds with charged and polar covalent molecules. For example, when a crystal of salt

(NaCl) is placed in water, the Na+ cations form hydrogen bonds with partial negative oxygen regions of water molecules.

The Cl- anions form hydrogen bonds with the partial positive hydrogen regions of water molecules.

Fig. 3.7

Page 42

Ionic and charged molecules form hydration spheres and dissolve readily in water- nonpolar molecules do not dissolve in water.

Each dissolved ion is surrounded by a sphere of water molecules, a hydration shell.

Eventually, water dissolves all the ions, resulting in a solution with two solutes, sodium and chloride.

Polar molecules are also soluble in water because they can also form hydrogen bonds with water.

Even large molecules, like proteins, can dissolve in water if they have ionic and polar regions.


Fig. 3.8

Any substance that has an affinity for water is hydrophilic. These substances are dominated by ionic

or polar bonds. This term includes substances that do

not dissolve because their molecules are too large and too tightly held together. For example, cotton is hydrophilic

because it has numerous polar covalent bonds in cellulose, its major constituent.

Water molecules form hydrogen bonds in these areas.


Substances that have no affinity for water are hydrophobic. These substances are dominated by non-

ionic and nonpolar covalent bonds. Because there are no consistent regions

with partial or full charges, water molecules cannot form hydrogen bonds with these molecules.

Oils, such as vegetable oil, are hydrophobic because the dominant bonds, carbon-carbon and carbon-hydrogen, exhibit equal or near equal sharing of electrons.

Hydrophobic molecules are major ingredients of cell membranes.Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings

Hydrophobic molecules get clumped together by water molecules

Page 42

The advantage of using moles as a measurement is that a mole of one substance has the same number of molecules as a mole of any other substance. If substance A has a molecular weight of 10

daltons and substance B has a molecular weight of 100 daltons, then we know that 10 g of A has the same number of molecules as 100 g of substance B.

The actual number of molecules in a mole is called Avogadro’s number, 6.02 x 1023.

A mole of sucrose contains 6.02 x 1023 molecules and weighs 342g, while a mole of ethyl alcohol (C2H6O) also contains 6.02 x 1023 molecules but weighs only 46g because the molecules are smaller.


In “wet” chemistry, we are typically combining solutions or measuring the quantities of materials in aqueous solutions. The concentration of a material in

solution is called its molarity. A one molar solution has one mole of a

substance dissolved in one liter of solvent, typically water.

To make a 1 molar (1 M) solution of sucrose we would slowly add water to 342 g of sucrose until the total volume was 1 liter and all the sugar was dissolved.Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings

Animation 3.2 Dissociation o.MOV

Occasionally, a hydrogen atom shared by two water molecules shifts from one molecule to the other. The hydrogen atom leaves its electron

behind and is transferred as a single proton - a hydrogen ion (H+).

The water molecule that lost a proton is now a hydroxide ion (OH-).

The water molecule with the extra proton is a hydronium ion (H3O+).

Introduction


Unnumbered Fig. 3.47

A simpler way to view this process is that a water molecule dissociates into a hydrogen ion and a hydroxide ion: H2O <=> H+ + OH-

This reaction is reversible. At equilibrium the concentration of

water molecules greatly exceeds that of H+ and OH-.

In pure water only one water molecule in every 554 million is dissociated. At equilibrium the concentration of H+ or

OH- is 10-7M (25°C) .

Because hydrogen and hydroxide ions are very reactive, changes in their concentrations can drastically affect the proteins and other molecules of a cell.

Adding certain solutes, called acids and bases, disrupts the equilibrium and modifies the concentrations of hydrogen and hydroxide ions.

The pH scale is used to describe how acidic or basic (the opposite of acidic) a solution is.


An acid is a substance that increases the hydrogen ion concentration in a solution. When hydrochloric acid is added to water,

hydrogen ions dissociate from chloride ions:

HCl -> H+ + Cl-

Addition of an acid makes a solution more acidic.

1. Organisms are sensitive to changes in pH


Any substance that reduces the hydrogen ion concentration in a solution is a base.

Some bases reduce H+ directly by accepting hydrogen ions. Ammonia (NH3) acts as a base when the

nitrogen’s unshared electron pair attracts a hydrogen ion from the solution, creating an ammonium in (NH4

+). NH3 + H+ <=> NH4

+

Other bases reduce H+ indirectly by dissociating to OH- that combines with H+ to form water. NaOH -> Na+ + OH- OH- + H+ -> H2O

Solutions with more OH- than H+are basic solutions.


Some acids and bases (HCl and NaOH) are strong acids or bases. These molecules dissociate completely in

water. Other acids and bases (NH3) are weak

acids or bases. For these molecules, the binding and

release of hydrogen ions are reversible. At equilibrium there will be a fixed ratio of

products to reactants. Carbonic acid (H2CO3) is a weak acid:

H2CO3 <=> HCO3- + H+

At equilibrium, 1% of the molecules will be dissociated.


In any solution the product of their H+ and OH- concentrations is constant at 10-14. [H+] [OH-] = 10-14

In a neutral solution, [H+] = 10-7 M and [OH-] = 10-

7 M Adding acid to a solution shifts the balance

between H+ and OH- toward H+ and leads to a decline in OH-. If [H+] = 10-5 M, then [OH-] = 10-9 M Hydroxide concentrations decline because some

of additional acid combines with hydroxide to form water.

Adding a base does the opposite, increasing OH- concentration and dropping H+ concentration.


The H+ and OH- concentrations of solutions can vary by a factor of 100 trillion or more.

To express this variation more conveniently, the H+ and OH- concentrations are typically expressed via the pH scale. The pH scale, ranging from 1 to 14,

compresses the range of concentrations by employing logarithms.

pH = - log [H+] or [H+] = 10-pH


Page 44

In a neutral solution [H+] = 10-7 M, and the pH = 7.

Values for pH decline as [H+] increase.

While the pH scale is based on [H+], values for [OH-] can be easily calculated from the product relationship.

Fig. 3.9

The pH of a neutral solution is 7. Acidic solutions have pH values less

than 7 and basic solutions have pH values more than 7.

Most biological fluids have pH values in the range of 6 to 8. However, pH values in the human stomach

can reach 2. Each pH unit represents a tenfold

difference in H+ and OH- concentrations. A small change in pH actually indicates a

substantial change in H+ and OH- concentrations.


The chemical processes in the cell can be disrupted by changes to the H+ and OH- concentrations away from their normal values near pH 7.

To maintain cellular pH values at a constant level, biological fluids have buffers.

Buffers resist changes to the pH of a solution when H+ or OH- is added to the solution. Buffers accept hydrogen ions from the

solution when they are in excess and donate hydrogen ions when they have been depleted.


Buffers typically consist of a weak acid and its corresponding base. One important buffer in human blood and

other biological solutions is carbonic acid. The chemical equilibrium between

carbonic acid and bicarbonate acts at a pH regulator.

The equilibrium shifts left or right as other metabolic processes add or remove H+ from the solution.


Acid precipitation is a serious assault on water quality and therefore the environment for all life where this problem occurs. Uncontaminated rain has a slightly acidic

pH of 5.6. The acid is a product of the formation of

carbonic acid from carbon dioxide and water.

Acid precipitation occurs when rain, snow, or fog has a pH that is more acidic than 5.6.

2. Acid precipitation threatens the fitness of the environment


Acid precipitation is caused primarily by sulfur oxides and nitrogen oxides in the atmosphere. These molecules react with water and

sunlight to form strong acids. These fall to the surface with rain or snow.

The major source of these oxides is the burning of fossil fuels (coal, oil, and gas) in factories and automobiles.

The presence of tall smokestacks allows this pollution to spread from its site of origin to contaminate relatively pristine areas. Rain in the Adirondack Mountains of

upstate New York averages a pH of 4.2Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings

The effects of acids in lakes and streams is more pronounced in the spring during snowmelt. As the surface snows melt and drain

down through the snow field, the meltwater accumulates acid and brings it into lakes and streams all at once.

The pH of early meltwater may be as low as 3.

Acid precipitation has a great impact on eggs and early developmental stages of aquatic organisms which are abundant in the spring.

Thus, strong acidity can alter the structure of molecules and impact ecological communities.


Direct impacts of acid precipitation on forests and terrestrial life are more controversial.

However, acid precipitation can impact soils by affecting the solubility of soil minerals. Acid precipitation can wash away key soil

buffers and plant nutrients (calcium and magnesium).

It can also increase the solubility of compounds like aluminum to toxic levels.

This has done major damage to forests in Europe and substantial damage of forests in North America. Fig. 3.10

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