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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.

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The Chemical Basis of Life


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overview
Overview
  • Atoms
  • Combining Matter
    • Physically
    • Chemically
  • Water
  • Acids, Bases, and pH
  • Buffers
matter and energy
Matter and Energy

Matter:

  • Occupies space
  • Has mass: liquid, gas, solid

Energy:

  • Capacity to do work
  • Measured by effect on matter
chemistry
Chemistry

Science of the structure of matter

Central to all other sciences

Chemistry is part of all living & non-living things

slide5
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)

elements
Elements

Basic units of all matter

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

112 known elements → periodic table

slide7
Each element is represented by its atomic symbol

(1st letter(s) of element’s name)

e.g. carbon = C

hydrogen = H

oxygen = O

slide8
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
Atoms

Building blocks of elements

Unique to each element

Give 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

slide11
Protons (p+) have positive charge

Neutrons (no) have no overall charge

Both are heavy particles with approximately same mass

Electrons (e-) have negative charge

Do not contribute to atomic mass

(1/2000th mass of proton)

In general, # protons = # electrons

No net electrical charge

generalized atom

Mass # = p+ + no

H

a

b

Atomic # = p+

Generalized Atom

Periodic table is ordered by atomic number

the 3 smallest atoms

H

He

Li

1

4

7

p+

1

2

3

2

3

1

no

0

2

4

e-

1

2

3

The 3 Smallest Atoms
atomic mass
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
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
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 atom
Structure of an Atom

Nucleus contains protons & neutrons

Electrons move around nucleus

= electron cloud

Atomic orbitals organized into shells

e.g. 11Na

slide18
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

slide19
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
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)

slide21
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

slide22

8O:

Has 6 e-s in outer shell

Needs 2

7N:

Has 5 e-s in outer shell

Needs 3

slide23

11Na:

Has 1 e- in outer shell

Needs 7

17Cl:

Has 7 e-s in outer shell

Needs 1

combining matter
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
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
Ways to Represent Compounds

e.g. methane (CH4)

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

special structure carbon ring
Special Structure: Carbon Ring

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

Same goes for 5-carbon rings

=

mixtures
Mixtures

2 or more substances

No chemical bonding

= physical intermixing

Living material contains 3 types:

Solutions

Colloids

Suspensions

mixture 1 solution
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
Mixture #2: Colloid

Heterogeneous

Translucent or milky

Does not settle out

Can undergo sol-gel transformation

e.g. cytosol in cells

mixture 3 suspension
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

Chemical Bonds

Attractive forces between atoms

ionic bonds
Ionic Bonds

Transfer of e- s from one atom to another

Become ions (charged particles)

Gain e- →negative charge = anion

Lose e- →positive charge = cation

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

salts
Salts

Release 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 bonds
Covalent Bonds

E- sharing

Each atom fills outer shell part of the time

Can be single, double, or triple bonds

e.g. H2: H-H; O2: O=O; N2: NN

Can be polar or non-polar bonds

slide36
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

O

C

O

Non-polar Covalent Bonds

Electrically balanced

Equal sharing of e-s

polar covalent bonds
Polar Covalent Bonds

Unequal 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 bonds
Hydrogen Bonds

Not 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

water s life giving properties
Water’s Life-Giving Properties

The 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
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

-

+

+

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
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

water heat
Water & Heat

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

heat of vaporization
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)

slide47
When water cools:

H-bonds reform

H-bonds release heat energy as temperature drops

specific heat capacity
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

slide49
Water’s high specific heat capacity:
  • Helps regulate Earth’s climate by buffering large changes in temperature
  • Helps moderate internal temperature
density of water
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

slide51

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

    • Causes molecules to be further apart than normal

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

Insulates lakes & other bodies of water in the winter

cohesion and adhesion
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
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
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
Acids and Bases

Acid:

Dissociates in H2O

Releases H+ ions = proton donor

Concentration of protons determines acidity of a solution

acids and bases56
Acids and Bases

Base:

Takes up H+ ions = proton accepter

Dissociates in H2O

Releases hydroxyl (OH-) ions

These bind to protons in solution, produce water, & lower acidity of solution

slide57
Neutral:

Acid and base form H2O and salt

e.g. HCl + NaOH = H2O + NaCl

strong acids
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
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
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

more on acids and bases
More on Acids and Bases

Strong acids and bases can cause severe chemical burns

e.g. battery acid (pH ~ 1.0)

In high concentrations, can kill organisms in an ecosystem

acid precipitation
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 us
Acid Precipitation in the US

Eastern US: pH 2-3 (rain)

Los Angeles: pH 1.7 (fog)

effects on terrestrial systems
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
Effects on Aquatic Systems

Kills aquatic life

Especially prevalent in spring:

Combo of snow melt & breeding season

buffers
Buffers

Buffers resist changes in pH by:

  • Acting as acids (releasing H+) when pH 
  • Acting as bases (binding H+) when pH 
buffer systems
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
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

slide70
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
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