2. chemical foundation of life, bio 101 fall 2014

The Chemistry of Life

Atoms Make Up All Matter

• Matter

– Takes up space

• Energy

– Ability to do work

Atoms Make Up All Matter

• Elements are fundamental types of matter

– Element cannot be broken down

– Bulk elements

• 25 elements essential to life

• Minerals

• Trace elements

Elements in the Human Body

Trace Elements

Trace Element: needed for survival in very small

quantities

FluorideIodineIron

Trace Elements

Trace Element: needed for survival in very small

quantities

FluorideIodineIron

Atoms

• Smallest possible “piece” of an element

• Composed of

– Protons – positively charged particles, atomic number

– Neutrons –uncharged particle

– Electron – negatively charged particle

Types of Subatomic Particles

Particle Charge Mass Position

Electron – 0 Around Nucleus

Proton + 1 In Nucleus

Neutron none 1 In Nucleus

Atomic Number and Mass Number

• Mass number: the number of protons and

neutrons in the nucleus

• Atomic Number: the number of protons

CCarbon

Atomic number

Element

Symbol

Atomic mass

6

12.0 112

Isotopes

Isotopes: elements with the same atomic number but

different mass number

Isotopes of Carbon

Carbon-12 Carbon-13 Carbon-14

Electrons 6 6 6

Protons 6 6 6

Neutrons 6 7 8

Mass Number

(Protons + Neutrons)

12 13 14

Radioisotopes

• Nucleus is unstable and decays (gives of energy)

Example Uses of RadioisotopesUse Details

Isotopic labeling the use of unusual isotopes as tracers or markers in chemical

reactions. Normally, atoms of a given element are indistinguishable

from each other. However, by using isotopes of different masses,

even different nonradioactive stable isotopes can be distinguished

by mass spectrometry or infrared spectroscopy. For example, in

'stable isotope labeling with amino acids in cell culture (SILAC)'

stable isotopes are used to quantify proteins. If radioactive isotopes

are used, they can be detected by the radiation they emit (this is

called radioisotopic labeling).

Radiometric dating using the known half-life of an unstable element, one can calculate

the amount of time that has elapsed since a known level of isotope

existed. The most widely known example is radiocarbon dating

used to determine the age of carbonaceous materials.

Spectroscopy Several forms of spectroscopy rely on the unique nuclear

properties of specific isotopes, both radioactive and stable. For

example, nuclear magnetic resonance (NMR) spectroscopy can be

used only for isotopes with a nonzero nuclear spin. The most

common isotopes used with NMR spectroscopy are 1H, 2D,15N,

13C, and 31P.

Mössbauer spectroscopy also relies on the nuclear transitions of

specific isotopes, such as 57Fe.

Carbon Dating

• Carbon-14: radioisotope that decays slowly

– Half-life: time for half the original concentration of an isotope to

decay

• C-14 can be used to

“age fossils”

Tracers

• Radioisotopes can be used to identify biologically active

cells (cancer cells and goiters)

Tracers

MRI: isotopes can be used in medical imaging to view

metabolically active cells in the brain

Radiation Therapy

• The energy given off by radioisotopes is damaging to

cells and can be used to treat cancers and to treat

goiters.

Dangers of Radioactive

IsotopesFUKUSHIMA, March 11th, 2011

Summary of Elemental

Chemistry

Term Definition

Element a pure chemical substance consisting of a single type of atom

Atom the smallest unit that defines the chemical elements and their

isotopes

Atomic number the number of protons found in the nucleus of an atom of that

element, and therefore identical to the charge number of the

nucleus

Mass number the total number of protons and neutrons (together known as

nucleons) in an atomic nucleus, also called atomic mass number or

nucleon number

Isotope variants of a particular chemical element such that while all

isotopes of a given element have the same number of protons in

each atom, they differ in neutron number

Atomic mass the mass of an atomic particle, sub-atomic particle, or molecule;

the protons and neutrons account for almost all of the mass of an

atom

Chemical Bonds

• Chemical Bonds – How elements are

hooked together

• Molecule – 2 or more atoms chemically

joined together

– Ex. O2, Cl2, H2

• Compound – Molecule composed of 2 or

more DIFFERENT atoms

– Ex. NaOH, H2O, NaCl, C6H12O6

Compound

+ =

Chemical Bonds

• Its all up to the electrons!

• Electrons live in orbitals – most likely location

of an electron when rotating around nucleus

– Each orbital has 2 electrons - more electrons,

more orbitals

– Orbitals are in shells

– Valence shell – outermost shell, when full, shell is

stable

• Most atoms DO NOT have a full shell, that’s why they

can bond.

• Inert Elements – Have a full outer shell and cannot

bond – Noble gases (Ne, He, Ar, Xe, Kr, Rn)

Electron

“Vacancy” in energy shell

Hydrogen Carbon Nitrogen Oxygen

8p7p6p1p

Electron Distribution Diagrams

Electron Distribution Diagrams

Types of Bonds – Covalent Bonds

• Covalent Bonds – forms when 2 atoms SHARE electrons– Nonpolar Covalent Bond – Equal share of electrons

– Polar Covalent Bond – Unequal share of electrons, one atom pulls electrons more than others.

• Hydrogen bonds – attractions between oppositely charged particles within a single molecule, or between molecules

Types of Bonds – Ionic Bonds

• Ionic Bonds – forms when 1 atom “takes”

an electron from another

– Happens when ions of opposite charge attract

each other and more negative gives up

electron for bond

– Very strong b/c create stability in atoms

Ionic Bonds: Electron Transfer

Ionic Bonds

Hydrogen Bonds

• Form when partial charges between two

different molecules attract one another

Figure 2.10 Hydrogen Bonds in Water.

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

+

+

+O

H H

Hydrogen atoms

slightly positive (δ+)

Oxygen atom

slightly negative (δ−)

Polar

covalent

bonds

a. b. c.

Water molecule

Hydrogen

bond

+

c: © The McGraw-Hill Companies, Inc./Jacques Cornell photographer

Hydrogen Bonds

O

H H

Polar covalent bonds

Hydrogen Bonds

Slightly negative end

Water is Essential to Life

• Water Regulates Temperature

– Ability to resist temperature change

• Body temperature

• Coastal climates


• Water Regulates Temperature

– Evaporation

• Body temperature regulation


• Many Substances Dissolve in Water

– Solution = solvent + solute(s)

– Hydrophilic

• “water-loving”

– Hydrophobic

• “water-fearing”


• Water is Cohesive and Adhesive

– Cohesion – tendency of water molecules to

stick together

• Surface tension

– Adhesion – tendency to form hydrogen bonds

with other substances

• Together responsible for transport in plants


• Water Expands as It Freezes

– Unusual tendency

– Ice less dense than liquid water

• Benefits aquatic life

– Formation of ice crystals deadly

• Adaptations – fur in mammals

Figure 2.14

Ice Floats.

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

H2O molecule

Ice

Liquid water

Ice Floats

Hydrogen bonds in water Hydrogen bonds in ice


• Water Participates in Life’s Chemical

Reactions

– Chemical reaction

• Reactants

• Products

– Reactions happen in water

– Water is either a reactant or product

CH4 + 2O2 CO2 + 2H20

methane + oxygen carbon dioxide + water

Chemical Reactions

• Chemical Reaction – 2 or more molecules

“swap” atoms to make different molecules

CH4 + 2O2 CO2 + 2H2O

Reactants Products

6CO2 + 6H20 C6H12O6 + 6O2Reactants Products

Acids and Bases• Water disassociates into H+ and OH-

• Water = Neutral Solution – H+ = OH-

• Acid – Substance that adds H+ to a solution– Taste sour

– Found in your stomach, orange juice, tomatoes, coffee, coca-cola

– HCl, H2SO4

• Base – Substance that adds OH- to a solution– Taste bitter, feel slippery, soapy

– Found in detergents, soaps, cleaners

– NaOH

• Buffer Systems – Pairs of weak acids and bases that help resist pH changes

H2O H+ + OH-

pH Scale

• Measures amount

of H+ ions

• Ranges from 0 – 14

• 0 – 6 acids

• 7 neutral

• 8 – 14 bases

Buffers

• Buffer systems regulate pH in organisms

– Maintaining correct pH of body fluids critical

– Buffer system

• Pair of weak acid and base that resist pH changes

– Carbonic acid

H2CO3 H+ + HCO3-

carbonic acid bicarbonate

Applications of Chemistry to

Biology

• Ocean Acidification


Biology

• Ocean Acidification

– the ongoing decrease in the

pH of the Earth's oceans,

caused by the uptake CO2

• Effects

– lower metabolic rates and

immune responses of ocean

life

– alter ocean water’s properties

allowing sound to travel

further, affecting prey and

predators

Estimated change in sea pH caused by

human created CO2.


Biology


Biology

Earth formation began

4.6 BYA

Moon formed

4.5 BYA

First solid rock

4.4 BYA

First water

4.3 BYA

First evidence

of life

3.8 BYA

While features of self-organization and self-replication are often considered the

hallmark of living systems, there are many instances of abiotic molecules

exhibiting such characteristics under proper conditions. Palasek showed that

self-assembly of RNA molecules can occur spontaneously due to physical

factors in hydrothermal vents.

It is postulated that this kind of spontaneous generation could have changed

simple inorganic molecules (CO2, H2O, etc.) into organic compounds.

Science

2. chemical foundation of life, bio 101 fall 2014