The Chemical Basis of Life Chapter 2. Overview Atoms Combining Matter –Physically –Chemically...
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Transcript of The Chemical Basis of Life Chapter 2. Overview Atoms Combining Matter –Physically –Chemically...
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
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
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
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
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
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)
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
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
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
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