Chapter 2 The Chemistry of LifeChapter 2 The Chemistry of Life

09-04 ekgh from 09-04 ekgh from Biology (Visual Resource Library)– Biology (Visual Resource Library)–

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Biology Concepts & Connections – CampbellBiology Concepts & Connections – CampbellUnity and Diversity of Life – Starr / TaggertUnity and Diversity of Life – Starr / Taggert

10 – 20 BYA the Universe begins.

Physics and Chemistry are basic to the

study of biology

THE CHEMISTRY

OF LIFE

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A. MatterMaterial that takes up space.

1. ElementsPure chemical substances composed of atoms.– Examples?– How many elements exist?– How many of these elements are essential to

life?

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Periodic Table of Elements

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2. AtomThe smallest “piece” of an element that

retains the characteristics of that element.Composed of 3 subatomic particles:

– Protons– Neutrons– Electrons

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Characteristics of Subatomic Particles

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Atomic number# protons in nucleus of an atom (establishes

identity of the atom)

Since most atoms are electrically neutral, atomic number indicates # of electrons as well.

Atomic mass# protons plus # neutrons in nucleus of an atom.Mass is measured in Daltons. 6.02 x 1023 daltons = 1gProtons and Neutrons = 1dElectrons = 1/1840 d

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How can we determine the number of neutrons in an atom?# neutrons = atomic mass - atomic #

• Determine # neutrons in a carbon atom (atomic mass = 12; atomic # = 6).

# neutrons = 12 - 6 = 6Do all carbon atoms have the same number of

protons?Do all carbon atoms have the same number of

neutrons?

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IsotopesAtoms having the same number of protons, but

differing numbers of neutrons.So their mass is different.Neutrons determine nuclear stability.

Ex. Carbon isotopes

carbon 12 (12C) 6 neutrons99% of C isotopes

carbon 13 (13C) 7 neutronscarbon 14 (14C) 8 neutrons– Which is most stable?– Unstable isotopes tend tobreak into more stable formsreleasing radioactivity

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Periodic table information on carbon:

Atomic mass given in table is average mass of all the element’s isotopes.

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3. CompoundA pure substance formed when atoms of

different elements bond.Examples:

– CO2 carbon dioxide

– H2O water

– CH4 methane

– C6H12O6 glucose

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4. MoleculeSmallest piece of a compound that retains

characteristics of that compound.The number of molecules is written as a

coefficient.Examples:– 4CO2 4 molecules of carbon dioxide

– 2C6H12O6 2 molecules of glucose

– 6O2 6 molecules of oxygen

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5. Chemical BondsType of bond formed is determined by the

number of valence electrons in the interacting atoms [octet rule].

a) Covalent bonds - form when atoms share electron pairs.– strongest type of bond– tend to form when atoms have 3, 4 or 5

valence electrons– can be nonpolar or polar

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Nonpolar covalent bonds - electrons are shared equally between atoms.

Ex. methane

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Polar covalent bonds - electrons are drawn more strongly to one atom’s nucleus than the other.

Form when less electronegative atoms bond with more highly electronegative atoms.

Electronegativity is the tendency of an atom to attract electrons

O2 is highly electronegative.

Ex. water

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b) Ionic bonds - form when oppositely charged ions are attracted to each other. – weaker than covalent bonds – atoms with 1, 2 or 3 valence electrons give up

electrons to atoms with 7, 6 or 5 valence electrons

– form salts also found is certain polypeptidesEx. NaCl

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c) Hydrogen bonds - form when opposite charges on two molecules are attracted to each other.– weakest type of bond* but in large numbers they

provide strength/stability – like teeth on zipperEx. DNA Ex. Water drop BETWEEN

molecules. WITHIN water molecule = covalent bonds H2O

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Covalentbond

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1. Water is cohesive and adhesive2. Water has a high Specific Heat3. Water has a high Heat of Vaporization4. Frozen Water has a lower density than

liquid water5. Water is an excellent solvent

THE PROPERTIES OF WATER

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• Cohesion = attraction between like substances

• Insects can walk on water due to surface tension created by cohesive water molecules

• Rocks skip• Result: water flows

freely, molecules don’t separate from each other. Ex. Blood (it’s 80% water)

#1 Property of water: Water is Cohesive and Adhesive

Figure 2.11

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Adhesion• Due to hydrogen

bonding, water molecules can adhere to molecules of other polar substances.

• Movement up glass tube because surface electric charge attracts water = “Capillary Action”

#1 Property of water: Water is Cohesive and Adhesive

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Results of Cohesive - Adhesive Property of Water

• Adhesion at work:Water moves up Xylem

tubes in plants via capillary action

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• Because Water’s hydrogen bonds absorb heat when they break and release heat when they form, they moderate temperature

• It takes a lot of energy to disrupt hydrogen bonds– Result: water is able to absorb a great deal of heat

energy without a large increase in temperature– As water cools, a slight drop in temperature releases a

large amount of heat– Large bodies of water loose heat slowly. The Atlantic

Ocean is at perfect swimming temperature right now.• Living organisms in large bodies of water benefit from this

#2 Property of Water: High Specific Heat (High Heat Capacity)

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• Water has a high boiling point 100 Co

– It takes high energy to break H bonds Sweat

– H bonds absorb heat when they break.

– A water molecule takes a large amount of energy with it when it evaporates– This leads to evaporative cooling

Figure 2.12

#3 Property of Water: Water has a high heat of vaporization

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• Molecules in ice are farther apart than those in liquid water

#4 Property of water: Ice is less dense than liquid water [Ice Floats]

Figure 2.13

Hydrogen bond

ICEHydrogen bonds are stable

LIQUID WATERHydrogen bonds constantly

break and re-form

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Ice is therefore less dense than liquid water, which causes it to float

Frozen water Frozen benzene

– If ice sank, it would seldom have a chance to thaw

– Ponds, lakes, and oceans would eventually freeze solid

– And why would this be bad……

– What benefit does floating ice provide…

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• Water is a strong solvent because it separates charged atoms or molecules

• Solutes whose charges or polarity allow them to stick to water molecules dissolve

• in it– They form

aqueous solutions

#5 Property of Water: Excellent Solvent

Figure 2.14

Ions in solution

Salt crystal

Cl–

Na+

Cl–

–

– –

–

–Na+

+

+

+

+

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• The nature of the water molecule, specifically, its polarity, makes it an excellent solvent.

• if the Atoms in a covalently bonded molecule share electrons equally, creating a nonpolar molecule – Fats. Insoluable in polar molecules

• If electrons are shared unequally, a polar molecule is created

# 5 Water is an excellent solvent

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– This makes the oxygen end of the molecule slightly negatively charged

– The hydrogen end of the molecule is slightly positively charged

– Water is therefore a polar molecule

• In a water molecule, oxygen exerts a stronger pull on the shared electrons than hydrogen

Figure 2.9

(–)

O

(–)

(+)(+)

H H

– This makes the oxygen end of the molecule slightly negatively charged

– The hydrogen end of the molecule is slightly positively charged

– Water is therefore a polar molecule

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• The charged regions on water molecules are attracted to the oppositely charged regions on nearby molecules– This attraction

forms weak bonds called hydrogen bonds

Result: Water’s polarity leads to hydrogen bonding and other unusual properties

Figure 2.10A

Hydrogen bond

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• Like no other common substance, water exists in nature in all three physical states:

Campbell Figure 2.10B

– as a solid– as a liquid– as a gas

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• A compound that releases H+ ions in solution is an acid. Proton donor.

• one that accepts H+ ions in solution is a base• Acidity is measured on the pH scale: • pH = -log[H]

– 0-7 is acidic – 8-14 is basic – Pure water and solutions that are neither basic

nor acidic are neutral, with a pH of 7 Because – when water dissociates it forms equal amounts of H+ and OH-

The chemistry of life is sensitive to acidic and basic conditions

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• The pH scale

Figure 2.15

pH scale

Acidic solution

Neutral solution

Basic solution

Incr

easi

ng

ly A

CID

IC(H

igh

er c

on

cen

trat

ion

of

H+)

Incr

easi

ng

ly B

AS

IC(L

ow

er

con

cen

trat

ion

of

H+)

NEUTRAL[H+] = [OH–]

Lemon juice; gastric juice

Grapefruit juice

Tomato juice

Urine

PURE WATER

Seawater

Milk of magnesia

Household ammonia

Household bleach

Oven cleaner

Human blood

H+

OH–

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The math behind pH

substance [H] log[H+] pH

Gastric juice .01 10 -2 2Pure water .0000001 10-7 7Ammonia .00000000001 10-11 11

The pH scale is logarithmic: a change in 1 on the scale means there’s a 10 fold change in H+

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Acids & Bases• Acids - substances that add

H+ to a solution.• Neutral • Bases - substances that

remove H+ from solution by combining with them.

pH scale is measure of acidity/alkalinity based on H+ concentration.

Each unit represents 10 fold change in H+ concentration

So… tomato juice = ____x greater [H+] than coffee

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• Cells are kept close to pH 7 by buffers

• Buffers are substances that resist pH change. “reservoirs for H+– They accept H+ ions when they are in

excess and donate H+ ions when they are depleted

– Buffers are not foolproof

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• Some ecosystems are threatened by acid precipitation• Acid precipitation is formed when air pollutants from burning

fossil fuels combine with water vapor in the air to form sulfuric and nitric acids

Connection: Acid precipitation threatens the environment

Figure 2.16A

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– These acids can kill fish, damage buildings, and injure trees

– Regulations, new technology, and energy conservation may help us reduce acid precipitation

Figure 2.16B

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Fluctuations in pH kept to minimum by buffers

Buffers- Donate H+ ions to

basic solutions and removes them from acid solutions

- Buffers = pairs of compounds, an acid and a base

- Blood buffers: H2CO3 = carbonic acid

and HCO3- = bicarbonate

ion

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The End