The Chemical Basis of Life

Chapter 2

Overview

• Atoms

• Combining Matter– Physically– Chemically

• Water

• Acids, Bases, and pH

• Buffers

Matter and Energy

Matter:– Occupies space– Has mass: liquid, gas, solid

Energy:– Capacity to do work– Measured by effect on matter

Chemistry

Science of the structure of matter

Central to all other sciences

Chemistry is part of all living & non-living things

Life requires ~25 chemical elements

Humans & other living organisms differ from non-living things in elemental composition

96% of body weight made up of C, H, O, N

Other 4%: Ca, P, K, S, Na, Cl, Mg & trace elements essential for life (e.g. Fe)

ElementsBasic units of all matter

Can’t be broken down to simpler substances using ordinary chemical methods

112 known elements → periodic table

Each element is represented by its atomic symbol

(1st letter(s) of element’s name)

e.g. carbon = C

hydrogen = H

oxygen = O

In nature, few elements exist in pure form

(tend to form compounds)

Emergent properties:

e.g. NaCl

Na (metal) + Cl (poisonous gas) = NaCl (table salt)

Atoms

Building blocks of elements

Unique to each elementGive it specific physical and chemical properties

Physical properties:Colour, texture, boiling point, melting point, etc.

Chemical properties:The ways that atoms interact with other atoms

Made up of protons (p+), neutrons (no), electrons (e-)

p+

no

e-

Protons (p+) have positive charge

Neutrons (no) have no overall chargeBoth are heavy particles with approximately

same mass

Electrons (e-) have negative chargeDo not contribute to atomic mass

(1/2000th mass of proton)

In general, # protons = # electronsNo net electrical charge

Generalized Atom

Mass # = p+ + no

Ha

b

Atomic # = p+

Periodic table is ordered by atomic number

The 3 Smallest Atoms

1

H1

1 He4

2 Li7

3

p+

no

e-

2 3

2

1

0

32

4

Atomic Mass

Approximately equal to mass number (# p+ + # no) because e-s weigh so little

In general, atomic weight is about equal to mass # of most abundant isotope

e.g. atomic mass of H = 1.008(indicates that 1H is present in much greater amounts

than 2H or 3H forms)

Isotopes

Different versions of same element

Occur with most natural elements

Differ in # of neutrons(same atomic # but different mass #)

If stable, nucleus remains intact

If unstable, is radioactive

Radioisotopes

Nuclei decompose spontaneously into more stable forms

e.g. 14C: half-life of 5700 years ½ atoms turn into 13N

Used to date rocks and biological remains

Releases particles & energy

(breaks chemical bonds in living organisms)

Damaging to live tissue but used in biological research & medicine

Structure of an AtomNucleus contains protons & neutrons

Electrons move around nucleus

= electron cloud

Atomic orbitals organized into shells

e.g. 11Na

Higher-energy shells hold more e-s (2n2) & are located further from nucleus

Shells fill up in order of increasing energy

e-s can be excited up to higher energy level for brief periods

Spontaneously return to lower level while emitting the energy gained via excitation

e-s in outer (valence) shell dictate chemical behaviour

(these ones interact with those from other atoms)

Regardless of # of e-s in each shell, # that can participate in bonding is 8

= octet rule

The Octet Rule

Atoms want to gain, lose, or share e-s so that have 8 electrons in outer shell

Exception = H

(only has room in 1st energy level for 2 e-s)

Use atomic number to calculate how many e-s are available for bonding

• 1st energy level = 2 e-s / shell

• 2nd and up = 8 e-s / shell

e.g. 6C:

Has 4 e-s in outer shell; wants to gain 4 e-s to fill shell for a total of 8 e-s

7N:

Has 5 e-s in outer shell

Needs 3

8O:

Has 6 e-s in outer shell

Needs 2

11Na:

Has 1 e- in outer shell

Needs 7

17Cl:

Has 7 e-s in outer shell

Needs 1

Combining Matter

Most atoms do not exist in free state

Chemically combine with other atoms to form molecules

If atoms are the same

= molecule of element e.g. O2

If atoms are different

= compound e.g. H2O

Molecular Formulas

A molecule’s chemical composition is written as a formula

Symbols for elements

Subscripts for number of atoms of each element

e.g. H20 = 2 H, 1 O

e.g. 5 H20 = 10 H, 5 O

Ways to Represent Compounds

e.g. methane (CH4)

Structural formula Ball-and-stick model Space-filling model

Special Structure: Carbon Ring

If icon for ring shows no atoms, assume that C occupies each corner

Same goes for 5-carbon rings

=

Mixtures

2 or more substances

No chemical bonding

= physical intermixing

Living material contains 3 types:Solutions

Colloids

Suspensions

Mixture #1: Solution

Homogeneous

Transparent

Does not settle out

Solvent– Present in largest quantity– Usually liquid– Water is body’s principle solvent

Solute– Present in smaller quantity

Mixture #2: Colloid

Heterogeneous

Translucent or milky

Does not settle out

Can undergo sol-gel transformation

e.g. cytosol in cells

Mixture #3: Suspension

Heterogenous

Settles out

e.g. blood settles out into plasma & cells

Chemical Bonds

Inert if outer e- valence shell is filled

– Do not tend to form bonds

e.g. He

Reactive if outer shell is not filled

– React with other atoms to gain / lose / share e-s

to fill shells

e.g. O

Attractive forces between atoms

Ionic BondsTransfer of e- s from one atom to another

Become ions (charged particles)Gain e- → negative charge = anionLose e- → positive charge = cation

Both become stable & combine to form ionic compound (a.k.a. salt)

SaltsRelease ions other than H+ and OH-

Usually form when acids and bases mix

Dissociate in water into component ions(electrolytes that can conduct electricity)

Important in living organisms:e.g. Na+, K+, Ca2+ used in nerve transmission, muscle

contraction

e.g. plant cells use salts to take up water from soil

Covalent BondsE- sharing

Each atom fills outer shell part of the time

Can be single, double, or triple bondse.g. H2: H-H; O2: O=O; N2: NN

Can be polar or non-polar bonds

Atoms can be:

Electropositive:

1-2 valence shell e-s

Tend to lose e-s

Electronegative:

6-7 valence shell e-s

Tend to attract e-s strongly

Electrically balanced

Non-polar Covalent Bonds

Electrically balanced

Equal sharing of e-s

CO O

Polar Covalent BondsUnequal e- sharing

One element has more protons

= stronger pull on e-s

= has e-s more of the time

= slightly electronegative

Results in molecule with + & - charges at either end

Often occurs when atoms are of different sizes

O

H H

+ +

-

Hydrogen BondsNot a true bond

= can’t form molecules

Attraction between covalently-bound H atom & electronegative atom

(can be different molecule or different area of same molecule)

e.g. between water molecules, between complementary bases in

DNA

TYPE Mixture Compound

“BOND” Physical mixing Chemical

SEPARATION BY: Physical means

Chemical means

COMPOSITIONHomogeneous

or heterogeneous

Homogeneous

Mixtures vs. Compounds

Water’s Life-Giving PropertiesThe universal solvent

Water is important because:

• Life originated in it

• All known living things depend on water

(metabolic processes, respiration, photosynthesis)

• Maintains cell structure/shape

Characteristics of Water• Polar molecules• Specific heat capacity• Heat of vaporization• Density of water• Cohesion• Adhesion• Surface tension• Good solvent

All result from H-bonding

Polarity of the Water Molecule

One end slightly positive, other slightly negative

= no net charge

Attracts other water molecules (cohesion)

Attracts sugar & other polar (hydrophilic) molecules

Repels oil & other non-polar (hydrophobic) molecules

-

+ +

Why is Polarity Important?

If water were linear (non-polar), not bent (polar):

– It would not liquify except at high pressures

– It would probably not remain liquid over more than about a 20°C. temperature range

• Polarity helps water stay liquid because molecules so strongly attracted to each other

– It would dissolve very few other substances• Polarity of water molecules can cause temporary polarity in non-polar molecules; virtually everything

will dissolve to a small extent in water

In consequence, life could not exist anywhere

H-bonds make it difficult to separate molecules

H-bonds are constantly forming & breaking

When temperature is stable, H bonds form at the same rate that they break

Water & Heat

Heat of Vaporization

When temperature increases:

H bonds break & stay broken

Individual molecules escape into air

= evaporation

Heat energy changes liquid H2O into gaseous form

High boiling point (100°C)

When water cools:

H-bonds reform

H-bonds release heat energy as temperature drops

Specific Heat Capacity= energy required to raise given amount of

substance by 1°C

Water has high specific heat capacity:

At high temperatures, water absorbs heat as H-bonds break

(can absorb a lot before temperature measurably rises)

As water cools, heat released from formation of H-bonds slows down cooling

Water’s high specific heat capacity:

• Helps regulate Earth’s climate by buffering large changes in

temperature

• Helps moderate internal temperature

Density of Water

Water reaches max. density at 4°C

(becomes less dense at lower temps)

When temp decreases below 0°C:

Molecules don’t move enough to break H-bonds so become locked

= ice

Lower density causes ice to “float” or form sheets at top of

water column

Insulates lakes & other bodies of water in the winter

Water expands as freezes due to hexagonal configuration of molecules caused by H-bonds

Causes molecules to be further apart than normal

Cohesion and Adhesion

Cohesion:– Water sticks to itself– H-bonds cause attraction

between water molecules

Adhesion:– Water sticks to other things– Due to electrostatic forces of

molecules/H-bonds

e.g. transpiration in plants:– Adhesion = water sticks to xylem– Cohesion = holds water column

together

Surface Tension

How hard it is to break a liquid’s surface

Causes liquid to act as elastic sheet

Caused by H-bonds between water molecules

Liquid compresses to have smallest surface area possible

e.g. water beading

Water as a Solvent

Ions & other polar molecules dissolve readily in water

H2O molecules cluster around ions / molecules in sphere

of hydration

Acids and Bases

Acid:

Dissociates in H2O

Releases H+ ions = proton donor

Concentration of protons determines acidity of a

solution

Base:

Takes up H+ ions = proton accepter

Dissociates in H2O

Releases hydroxyl (OH-) ionsThese bind to protons in solution, produce

water, & lower acidity of solution

Acids and Bases

Neutral:

Acid and base form H2O and salt

e.g. HCl + NaOH = H2O + NaCl

Strong Acids

Dissociate completely & irreversibly in water

e.g. 100 HCl molecules in H2O becomes 100 H+ and 100 Cl-

(reaction occurs in one direction only)

Dramatically affect pH

Weak Acids

Dissociate partially in water

e.g. HAc H+ + Ac-

(molecules of intact acid are in dynamic equilibrium with dissociated ions )

Do not affect pH as much as strong acids

Important in body’s chemical buffer systems

pH (potential of hydrogen)Relative concentration of H+ ions in a solution

pH scale 0-14

Each pH unit is 10-fold change in [H+]

At pH = 7, [H+] = [OH-]

= neutral

Body’s internal environment

= pH 7.3-7.5

pH Scale

More on Acids and BasesStrong acids and bases can cause severe

chemical burnse.g. battery acid (pH ~ 1.0)

In high concentrations, can kill organisms in an ecosystem

Acid Precipitation

Rain, snow, or fog with pH < 5.6

Caused by S oxides & N oxides in air(from N-containing fertilizers & burning

of fossil fuels)

Oxides react with water vapour in air to form H2SO4 & HNO3

Acid Precipitation in the USEastern US: pH 2-3 (rain)

Los Angeles: pH 1.7 (fog)

Effects on Terrestrial Systems

Has damaged / destroyed forests in US, Canada,

Europe

Physical damage from acid contact

Essential minerals in soil washed away

Effects on Aquatic Systems

Kills aquatic life

Especially prevalent in spring:

Combo of snow melt & breeding season

Buffers

Buffers resist changes in pH by:

• Acting as acids (releasing H+) when pH

• Acting as bases (binding H+) when pH

Buffer Systems

It is imperative for cells to respond to changes in pH

Changes disrupt cellular processes & functioning of biological molecules

Buffer systems help resist large and abrupt swings in pH

Bicarbonate Buffer System

Maintains blood pH (7.3 - 7.5)

If pH increases, carbonic acid releases H+ to neutralize excess OH-

H+ combines with OH- to form water

OH- + H2CO3 → HCO3- + H2O

When pH begins to drop, bicarbonate consumes excess H+ to shift reaction

back towards acid

HCO3- + H+ → H2CO3

System is constantly buffering pH changes

A Final Word on Buffers

Buffer systems work within narrow range

When range is exceeded, extremely severe effects

If blood pH drops to 7.0:

= respiratory acidosis, coma

If blood pH rises to 7.8:

= alkalosis, tetany