Chapter 2 · · 2015-10-30Campbell Biology: Concepts & Connections, Eighth Edition REECE • TAYLOR • SIMON • DICKEY • HOGAN Chapter 2 Lecture by Edward J. Zalisko ... 2.4

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Campbell Biology: Concepts & Connections, Eighth Edition REECE • TAYLOR • SIMON • DICKEY • HOGAN

Chapter 2

Lecture by Edward J. Zalisko

The Chemical Basis of Life

© 2015 Pearson Education, Inc.


Figure 2.0-1


Figure 2.0-2

Chapter 2: Big Ideas

Elements, Atoms,

and Compounds

Chemical Bonds Water’s Life-

Supporting Properties

O

H H

Introduction

• Chemicals are the raw materials that make up

• our bodies,

• the bodies of other organisms, and

• the physical environment.


Introduction

• Life’s chemistry is tied to water.

• Life first evolved in water.

• All living organisms require water.

• Cells consist of about 75% water.


ELEMENTS, ATOMS,

AND COMPOUNDS


2.1 Organisms are composed of elements, in combinations called compounds

• Living organisms are composed of matter, which is

anything that occupies space and has mass

(weight).

• Matter is composed of chemical elements.

• An element is a substance that cannot be broken

down to other substances by ordinary chemical

means.

• There are 92 elements in nature—only a few exist

in a pure state.



Table 2.1


• A compound is a substance consisting of two or

more different elements in a fixed ratio.

• Compounds are more common than pure

elements.

• Sodium chloride, table salt, is a common

compound of equal parts of sodium (Na) and

chlorine (Cl).



Figure 2.1-0

Sodium

(Na) Chlorine

(Cl)

Sodium chloride

(NaCl)


Figure 2.1-1

Sodium

(Na)


Figure 2.1-2

Chlorine

(Cl)


Figure 2.1-3

Sodium chloride

(NaCl)


• About 25 elements are essential for human life.

• Four elements make up about 96% of the weight of

most living organisms. These are

• oxygen,

• carbon,

• hydrogen, and

• nitrogen.

• Trace elements are essential but are only needed

in minute quantities.


2.2 CONNECTION: Trace elements are common additives to food and water

• Some trace elements are required to prevent

disease.

• Without iron, your body cannot transport oxygen.

• An iodine deficiency prevents production of thyroid

hormones, resulting in goiter.



Figure 2.2a


• Fluoride is usually added to municipal water and

dental products to help reduce tooth decay.



Figure 2.2b


• Several chemicals are added to food to

• help preserve it,

• make it more nutritious, and/or

• make it look better.

• Check out the nutrition facts label on foods and

drinks you purchase.



Figure 2.2c

2.3 Atoms consist of protons, neutrons, and electrons

• Each element consists of one kind of atom.

• An atom is the smallest unit of matter that still

retains the properties of an element.

• Three subatomic particles in atoms are relevant to

our discussion of the properties of elements.

• Protons are positively charged.

• Electrons are negatively charged.

• Neutrons are electrically neutral.



• Neutrons and protons are packed into an atom’s

nucleus.

• Electrons orbit the nucleus.

• The negative charge of electrons and the positive

charge of protons keep electrons near the nucleus.



Figure 2.3

Nucleus

Electron

cloud

Nucleus Protons

Neutrons

Electrons

2

2e−

2

2

−

+

−

+

+

+

+

−


• The number of protons is the atom’s atomic

number.

• An atom’s mass number is the sum of the number

of protons and neutrons in the nucleus.

• The atomic mass is approximately equal to its

mass number.



Figure 2.3

Nucleus

Electron

cloud

Nucleus Protons

Neutrons

Electrons

2

2e−

2

2

−

+

−

+

+

+

+

−


• Although all atoms of an element have the same

atomic number, some differ in mass number.

• Different isotopes of an element have

• the same number of protons

• but different numbers of neutrons.

• Different isotopes of an element behave identically

in chemical reactions.

• In radioactive isotopes, the nucleus decays

spontaneously, giving off particles and energy.



Table 2.3

2.4 CONNECTION: Radioactive isotopes can help or harm us

• Living cells cannot distinguish between isotopes of

the same element.

• Therefore, radioactive compounds in metabolic

processes can act as tracers.

• This radioactivity can be detected by instruments.

• By using these instruments, the fate of radioactive

tracers can be monitored in living organisms.



• Radioactive tracers are frequently used in medical

diagnosis.

• Sophisticated imaging instruments are used to

detect them.

• An imaging instrument that uses positron-emission

tomography (PET) detects the location of injected

radioactive materials.

• PET is useful for diagnosing heart disorders and

cancer and in brain research.



Figure 2.4a


Figure 2.4b

Healthy person Alzheimer’s patient


• In addition to benefits, there are also dangers

associated with using radioactive substances.

• Uncontrolled exposure can cause damage to some

molecules in a living cell, especially DNA.

• Chemical bonds are broken by the emitted energy,

which causes abnormal bonds to form.


CHEMICAL BONDS


2.5 The distribution of electrons determines an atom’s chemical properties

• Of the three subatomic particles—protons,

neutrons, and electrons—only electrons are

directly involved in the chemical activity of an

atom.

• Electrons can be located in different electron

shells, each with a characteristic distance from the

nucleus.

• An atom may have one, two, or more electron

shells.

• Information about the distribution of electrons is

found in the periodic table of the elements. © 2015 Pearson Education, Inc.


Figure 2.5a


Figure 2.5b-0

Hydrogen Helium

Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon

First

shell

Second

shell

Third

shell

Sodium Magnesium Aluminum Silicon Phosphorus Sulfur Chlorine Argon


Figure 2.5b-1

Hydrogen

Lithium Beryllium Boron

Sodium Magnesium Aluminum

Carbon

Silicon


Figure 2.5b-2

Helium

Nitrogen Oxygen Fluorine Neon

Phosphorus Sulfur Chlorine Argon


• The number of electrons in the outermost shell,

called the valence shell, determines the chemical

properties of the atom.

• Atoms whose outer shells are not full tend to

interact with other atoms in ways that enable them

to complete or fill their valence shells.



• When two atoms with incomplete outer shells

react, each atom will share, donate, or receive

electrons, so that both partners end up with

completed outer shells.

• These interactions usually result in atoms staying

close together, held by attractions called chemical

bonds.


2.6 Covalent bonds join atoms into molecules through electron sharing

• In a covalent bond, two atoms, each with an

unpaired electron in its outer shell, actually share a

pair of electrons.

• Two or more atoms held together by covalent

bonds form a molecule.

• A covalent bond connects two hydrogen atoms in a

molecule of the gas H2.


Animation: Covalent Bonds



• There are four alternative ways to represent

common molecules.

• The hydrogen atoms in H2 are held together by a

pair of shared electrons.

• In an oxygen molecule (O2), the two oxygen atoms

share two pairs of electrons, forming a double

bond, indicated in a structural formula by a pair of

lines.



Figure 2.6-0

Molecular

Formula

Electron Distribution

Diagram

Structural

Formula

Space-Filling

Model

O2

Oxygen

CH4

Methane

H2O

Water

Polar covalent bonds

in a water molecule

Single bond

Double bond

Nonpolar covalent

bonds

Polar covalent

bonds

(slightly −)

(slightly +) (slightly +)

H H

H

H

H

H

H H

O O

O

C

O

H H

H2

Hydrogen


Figure 2.6-1

Molecular

Formula


Diagram

Structural

Formula Space-Filling

Model

O2

Oxygen

H H

O O

H2

Hydrogen Single bond

Double bond


Figure 2.6-2

Molecular

Formula


Diagram

Structural

Formula

Space-Filling

Model

CH4

Methane

H2O

Water

Nonpolar covalent

bonds

Polar covalent

bonds

H

H

H

H H

O

C H


Figure 2.6-3

Polar covalent bonds

in a water molecule

(slightly +) (slightly +)

O

H H

(slightly −)


• H2 and O2 are molecules composed of only one

element.

• Methane (CH4) and water (H2O) are compounds,

substances composed of two or more different

elements.



• Atoms in a covalently bonded molecule continually

compete for shared electrons.

• The attraction (pull) for shared electrons is called

electronegativity.

• More electronegative atoms pull harder.



• In molecules of only one element, the pull toward

each atom is equal, because each atom has the

same electronegativity.

• The bonds formed are called nonpolar covalent

bonds.



• Water has atoms with different electronegativities.

• Oxygen attracts the shared electrons more strongly

than hydrogen.

• So the shared electrons spend more time near

oxygen.

• The oxygen atom has a slightly negative charge

and the hydrogen atoms have a slightly positive

charge.

• The result is a polar covalent bond.


2.7 Ionic bonds are attractions between ions of opposite charge

• An ion is an atom or molecule with an electrical

charge resulting from gain or loss of one or more

electrons.

• When an electron is lost, a positive charge results.

• When an electron is gained, a negative charge

results.

• Two ions with opposite charges attract each other.

• When the attraction holds the ions together, it is

called an ionic bond.

• Salt is a synonym for an ionic compound.


Animation: Ionic Bonds



Figure 2.7a-1

Na

Sodium atom

Cl

Chlorine atom

Na Cl


Figure 2.7a-2

Na

Sodium atom

Cl

Chlorine atom

Na+

Sodium ion

Cl−

Chloride ion

Na+ Na Cl− Cl

− +

Sodium chloride (NaCl)


Figure 2.7b-0

Na+

Cl−


Figure 2.7b-1

2.8 Hydrogen bonds are weak bonds important in the chemistry of life

• In living organisms, most of the strong chemical

bonds are covalent, linking atoms to form a cell’s

molecules.

• Crucial to the functioning of a cell are weaker

bonds within and between molecules.

• One of the most important types of weak bonds is

the hydrogen bond, which is best illustrated with

water molecules.



• The hydrogen atoms of a water molecule are

attached to oxygen by polar covalent bonds.

• Because of these polar bonds and the wide V

shape of the molecule, water is a polar

molecule—that is, it has an unequal distribution of

charges.

• This partial positive charge allows each hydrogen

to be attracted to a nearby atom that has a partial

negative charge.



• Weak hydrogen bonds form between water

molecules.

• Each hydrogen atom of a water molecule can form

a hydrogen bond with a nearby partially negative

oxygen atom of another water molecule.

• The negative (oxygen) pole of a water molecule can

form hydrogen bonds to two hydrogen atoms.

• Thus, each H2O molecule can hydrogen-bond to as

many as four partners.


Animation: Water Structure



Figure 2.8

(−)

(+)

(−) (+)

(−)

(+) (−)

(+)

Hydrogen

bond

Polar covalent

bonds

2.9 Chemical reactions make and break chemical bonds

• Remember that the structure of atoms and

molecules determines the way they behave.

• Atoms combine to form molecules.

• Hydrogen and oxygen can react to form water:

2 H2 + O2 2 H2O



• The formation of water from hydrogen and oxygen

is an example of a chemical reaction.

• The reactants (H2 and O2) are converted to H2O,

the product.

• Chemical reactions do not create or destroy

matter.

• Chemical reactions only rearrange matter.



Figure 2.9

+

Reactants Products

2 H2 O2 2 H2O


• Photosynthesis is a chemical reaction that is

essential to life on Earth.

• Carbon dioxide (from the air) reacts with water.

• Sunlight powers the conversion of these reactants

to produce the products glucose and oxygen.


WATER’S LIFE-SUPPORTING

PROPERTIES


2.10 Hydrogen bonds make liquid water cohesive

• The tendency of molecules of the same kind to

stick together is cohesion.

• Cohesion is much stronger for water than for other

liquids.

• Most plants depend upon cohesion to help

transport water and nutrients from their roots to

their leaves.

• The tendency of two kinds of molecules to stick

together is adhesion.


2.10 Hydrogen bonds make liquid water cohesive

• Cohesion is related to surface tension—a

measure of how difficult it is to break the surface of

a liquid.

• Hydrogen bonds give water high surface tension,

making it behave as if it were coated with an

invisible film.

• Water striders stand on water without breaking the

water surface.


Animation: Water Transport



Figure 2.10

2.11 Water’s hydrogen bonds moderate temperature

• Thermal energy is the energy associated with the

random movement of atoms and molecules.

• Thermal energy in transfer from a warmer to a

cooler body of matter is defined as heat.

• Temperature measures the intensity of heat—that

is, the average speed of molecules in a body of

matter.



• Heat must be absorbed to break hydrogen bonds.

• Heat is released when hydrogen bonds form.

• To raise the temperature of water, hydrogen bonds

between water molecules must be broken before

the molecules can move faster. Thus,

• when warming up, water absorbs a large amount of

heat and

• when water cools, water molecules slow down,

more hydrogen bonds form, and a considerable

amount of heat is released.



• Earth’s giant water supply moderates

temperatures, helping to keep temperatures within

limits that permit life.

• Water’s resistance to temperature change also

stabilizes ocean temperatures, creating a favorable

environment for marine life.



• When a substance evaporates, the surface of the

liquid that remains behind cools down; this is the

process of evaporative cooling.

• This cooling occurs because the molecules with

the greatest energy leave the surface.



Figure 2.11

2.12 Ice floats because it is less dense than liquid water

• Water can exist as a gas, liquid, or solid.

• Water is less dense as a solid than a liquid

because of hydrogen bonding.

• When water freezes, each molecule forms a stable

hydrogen bond with its neighbors.

• As ice crystals form, the molecules are less densely

packed than in liquid water.

• Because ice is less dense than water, it floats.



Figure 2.12-0

Hydrogen

bond

Ice

Hydrogen bonds are stable.

Liquid water

Hydrogen bonds constantly break and re-form.


Figure 2.12-1

2.13 Water is the solvent of life

• A solution is a liquid consisting of a uniform

mixture of two or more substances.

• The dissolving agent is the solvent.

• The substance that is dissolved is the solute.

• An aqueous solution is one in which water is the

solvent.


2.13 Water is the solvent of life

• Water’s versatility as a solvent results from the

polarity of its molecules.

• Polar or charged solutes dissolve when water

molecules surround them, forming aqueous

solutions.

• Table salt is an example of a solute that will go into

solution in water.



Figure 2.13

Salt crystal

Na+ Cl−

− +

Na+ Cl−

−

+

+ +

+

+

+

+

−

−

−

−

−

− −

Positive hydrogen ends of

water molecules attracted

to negative chloride ion

Negative oxygen ends of

water molecules attracted

to positive sodium ion

2.14 The chemistry of life is sensitive to acidic and basic conditions

• In liquid water, a small percentage of water

molecules break apart into ions.

• Some are hydrogen ions (H+).

• Some are hydroxide ions (OH–).

• Both types are very reactive.



• A substance that donates hydrogen ions to

solutions is called an acid.

• A base is a substance that reduces the hydrogen

ion concentration of a solution.

• The pH scale describes how acidic or basic a

solution is.

• The pH scale ranges from 0 to 14, with 0 the most

acidic and 14 the most basic.

• Each pH unit represents a 10-fold change in the

concentration of H+ in a solution.



• A buffer is a substance that minimizes changes in

pH. Buffers

• accept H+ when it is in excess and

• donate H+ when it is depleted.



Figure 2.14-0

Acidic

solution

H+

OH−

H+ H+

H+

H+

OH− OH−

OH−

OH−

OH−

OH− H+

H+ H+

H+

H+ H+

H+

H+

OH−

H+

H+

OH− OH− OH−

OH−

OH− OH−

Neutral

solution

Basic

solution

NEUTRAL

[H+] = [OH−]

Inc

rea

sin

gly

AC

IDIC

(Hig

he

r H

+ c

on

ce

ntr

ati

on

)

Inc

rea

sin

gly

BA

SIC

(Hig

he

r O

H− c

on

ce

ntr

ati

on

)

pH scale

Battery acid

Lemon juice,

gastric juice

Vinegar, cola

Tomato juice

Rainwater

Human urine

Saliva

Pure water

Human blood, tears Seawater

Milk of magnesia

Household ammonia

Household bleach

Oven cleaner

0

1

2

3

4

5

6

7

8

9

10

11

12

13

14


Figure 2.14-1

NEUTRAL

[H+] = [OH−]

Incre

asin

gly

AC

IDIC

(Hig

her

H+ c

on

cen

trati

on

)

pH scale

Battery acid

Lemon juice,

gastric juice

Vinegar, cola

Tomato juice

Rainwater

Human urine

Saliva

Pure water

0

1

2

3

4

5

6

7


Figure 2.14-2

NEUTRAL

[H+] = [OH−]

Inc

reasin

gly

BA

SIC

(Hig

her

OH

− c

on

cen

trati

on

)

Pure water

Human blood, tears

Seawater

Milk of magnesia

Household ammonia

Household bleach

Oven cleaner

7

8

9

10

11

12

13

14

pH scale


Figure 2.14-3

Acidic

solution

H+

OH−

H+ H+ H+

H+

OH− OH−

OH−

OH−

OH−

OH− H+

H+ H+

H+

H+

H+

H+

H+ OH−

H+

H+

OH− OH− OH−

OH−

OH− OH−

Neutral

solution

Basic

solution

2.15 SCIENTIFIC THINKING: Scientists study the effects of rising atmospheric CO2 on coral reef ecosystems

• Carbon dioxide is

• the main product of fossil fuel combustion,

• increasing in the atmosphere, and

• linked to global climate change.



• About 25% of this human-generated CO2 is

absorbed by the vast oceans.

• CO2 dissolved in seawater lowers the pH of the

ocean in a process known as ocean acidification.



• As seawater acidifies, the extra hydrogen ions (H+)

combine with carbonate ions (CO32–) to form

bicarbonate ions (HCO3–).

• This reaction reduces the carbonate ion

concentration available to corals and other shell-

building animals.



• In a controlled experiment, scientists looked at the

effect of decreasing carbonate ion concentration

on the rate of calcium deposition by reef

organisms.

• The lower the concentration of carbonate ions, the

lower the rate of calcification, and thus the slower

the growth of coral animals.

• The results from experimental and observational

field studies of sites where pH naturally varies have

dire implications for the health of coral reefs and the

diversity of organisms they support.



Figure 2.15a

220 240 260 280

20

10

0

[CO32−] (μmol/kg of seawater)

Ca

lcif

ica

tio

n r

ate

(mm

ol C

aC

O3/m

2 ×

da

y)

Source: Adaption of figure 5 from “Effect of Calcium Carbonate Saturation State

on the Calcification Rate of an Experimental Coral Reef” by C. Langdon, et al., from

Global Biogeochemical Cycles, June 2000, Volume 14(2). Copyright © 2000 by

American Geophysical Union. Reprinted with permission of Wiley Inc.


Figure 2.15b

Rising CO2 bubbles

lower the pH of the

water

2.16 EVOLUTION CONNECTION: The search for extraterrestrial life centers on the search for water

• The emergent properties of water support life on

Earth.

• When astrobiologists search for signs of

extraterrestrial life on distant planets, they look for

evidence of water.

• The National Aeronautics and Space

Administration (NASA) has found evidence that

water was once abundant on Mars.


You should now be able to

1. Describe the importance of chemical elements to living

organisms.

2. Explain the formation of compounds.

3. Describe the structure of an atom.

4. Distinguish between ionic, hydrogen, and covalent

bonds.

5. Define a chemical reaction and explain how it changes

the composition of matter.

6. List and define the life-supporting properties of water.

7. Explain the pH scale and the formation of acid and base

solutions. © 2015 Pearson Education, Inc.


Figure 2.UN01

Protons (+ charge)

determine element −

+ −

+ Electrons (− charge)

form negative cloud

and determine

chemical behavior

Neutrons (no charge)

determine isotope Atom

Nucleus


Figure 2.UN02

Liquid water:

Hydrogen bonds

constantly break

and re-form

Ice: Stable

hydrogen bonds

hold molecules

apart


Figure 2.UN03-0

Atoms

atomic number

of each element

(a)

Chemical

Bonds

ions

nonpolar

covalent bonds

water

(b) (c)

(d)

(e)

(f) (g)

(h)

Na

have positively

charged have neutral have negatively

charged

number present

equals

number may

differ in

number in outer

shell determines

formation of

H

H H

−

−

Cl

H H

H

O

O

(−)

(−)

(+) (+)

(+)

electron transfer

between atoms

creates

electron sharing

between atoms

creates

attraction between

ions creates

unequal

sharing creates

equal

sharing creates

example is can lead to

has important

qualities due

to polarity and


Figure 2.UN03-1

Atoms

atomic number

of each element

(a) (b) (c)

(d)

have positively

charged have neutral have negatively

charged

number present

equals

number may

differ in

number in outer

shell determines

formation of

H

−

Chemical

Bonds


Figure 2.UN03-2

Chemical

Bonds

ions

nonpolar

covalent bonds

water

(e)

(f) (g)

(h)

Na H H

−

Cl

H H

H

O

O

(−)

(−)

(+) (+)

(+)

electron transfer

between atoms

creates

electron sharing

between atoms

creates

attraction between

ions creates unequal

sharing creates

equal

sharing creates

example is can lead to

has important

qualities due

to polarity and


Figure 2.UN04

Fluorine

atom

F K

Potassium

atom

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Chapter 2 · · 2015-10-30Campbell Biology: Concepts & Connections, Eighth Edition REECE • TAYLOR • SIMON • DICKEY • HOGAN Chapter 2 Lecture by Edward J. Zalisko ... 2.4