Chapter 2jfalabella.weebly.com/uploads/2/5/2/4/25249222/chapter_2.pdf · 2014-08-26 · 2.4...

PowerPoint Lectures 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




• 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

Electroncloud

NucleusProtons

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

Electroncloud

NucleusProtons

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.


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


Structural Formula

Space-Filling Model

O2 Oxygen

H H

O O

H2 Hydrogen

Single bond

Double bond


Figure 2.6-2

Molecular Formula


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

ClChlorine atom

Na Cl


Figure 2.7a-2

Na Sodium atom

ClChlorine atom

Na+Sodium ion

Cl−Chloride ion

Na+Na Cl−Cl

−+



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

(−)

(+)

(−)(+)

(−)(+)(−)

(+)

Hydrogenbond

Polar covalentbonds


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

Hydrogenbond

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 attractedto negative chloride ion

Negative oxygen ends ofwater molecules attractedto 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

Acidicsolution

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−

Neutralsolution

Basicsolution

NEUTRAL[H+] = [OH−]

Incr

easi

ngly

AC

IDIC

(Hig

her H

+ con

cent

ratio

n)In

crea

sing

ly B

ASI

C

(Hig

her O

H− c

once

ntra

tion)

pH scale

Battery acid

Lemon juice, gastric juiceVinegar, cola

Tomato juice

RainwaterHuman urineSalivaPure water Human blood, tearsSeawater

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


Incr

easi

ngly

AC

IDIC

(Hig

her H

+ con

cent

ratio

n)

pH scale

Battery acid

Lemon juice, gastric juiceVinegar, cola

Tomato juice

RainwaterHuman urineSalivaPure water

0

1

2

3

4

5

6

7


Figure 2.14-2


Incr

easi

ngly

BA

SIC

(Hig

her O

H− c

once

ntra

tion)

Pure water Human blood, tearsSeawater

Milk of magnesia

Household ammonia

Household bleach

Oven cleaner

7

8

9

10

11

12

13

14

pH scale


Figure 2.14-3

Acidicsolution

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−

Neutralsolution

Basicsolution


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.


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


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

• As seawater acidifies, the extra hydrogen ions (H+) combine with carbonate ions (CO3

2–) to form bicarbonate ions (HCO3

–).

• This reaction reduces the carbonate ion concentration available to corals and other shell-building animals.


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

• 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)

Cal

cific

atio

n ra

te

(mm

ol C

aCO

3/m2 ×

day

)

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.


Figure 2.UN01

Protons (+ charge) determine element

−+

−+ Electrons (− charge)

form negative cloudand determine chemical behavior

Neutrons (no charge) determine isotope Atom

Nucleus


Figure 2.UN02

Liquid water: Hydrogen bondsconstantly break and re-form

Ice: Stable hydrogen bonds hold molecules apart


Figure 2.UN03-0Atoms

atomic number of each element

(a)

ChemicalBonds

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 transferbetween atoms

creates

electron sharingbetween atoms

creates

attraction betweenions creates

unequal sharing creates

equalsharing 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

−

ChemicalBonds


Figure 2.UN03-2

ChemicalBonds

ions

nonpolar covalent bonds

water

(e)

(f) (g)

(h)

Na H H

−

Cl

H H

H

O

O

(−)

(−)

(+) (+)

(+)

electron transferbetween atoms

creates

electron sharingbetween atoms

creates

attraction betweenions creates

unequal sharing creates

equalsharing 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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