Bio for NonMajors college course provided by Lumen Learning

1. The Chemistry
of Life

2. Atoms Make Up All Matter
• Matter
– Takes up space
• Energy
– Ability to do work

3. 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

5. Elements in the Human Body

6. Trace Elements
Trace Element: needed for survival in very small
quantities
FluorideIodineIron

7. Trace Elements
Trace Element: needed for survival in very small
quantities
FluorideIodineIron

8. Atoms
• Smallest possible
“piece” of an element
• Composed of
– Protons – positively
charged particles,
atomic number
– Neutrons –
uncharged particle
– Electron – negatively
charged particle

9. Types of Subatomic Particles
Particle Charge Mass Position
Electron – 0 Around Nucleus
Proton + 1 In Nucleus
Neutron none 1 In Nucleus

10. Atomic Number and Mass Number
• Mass number: the number of protons and
neutrons in the nucleus
• Atomic Number: the number of protons
C
Carbon
Atomic number
Element
Symbol
Atomic mass
6
12.0 112

11. 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

12. Radioisotopes
• Nucleus is unstable and decays (gives of energy)

13. Example Uses of Radioisotopes
Use 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.

14. 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”

15. Tracers
• Radioisotopes can be used to identify biologically active
cells (cancer cells and goiters)

16. Tracers
MRI: isotopes can be used in medical imaging to view
metabolically active cells in the brain

17. Radiation Therapy
• The energy given off by radioisotopes is damaging to
cells and can be used to treat cancers and to treat
goiters.

18. Dangers of Radioactive
Isotopes
FUKUSHIMA, March 11th, 2011

19. 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

20. 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

21. Compound
+ =

22. 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)

23. Electron
“Vacancy” in energy shell
Hydrogen Carbon Nitrogen Oxygen
8p7p6p1p
Electron Distribution Diagrams

24. Electron Distribution Diagrams

25. 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

27. 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

28. Ionic Bonds: Electron Transfer

29. Ionic Bonds

30. Hydrogen Bonds
• Form when partial charges between two
different molecules attract one another

31. 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

32. O
H H
Polar covalent bonds
Hydrogen Bonds
Slightly negative end

33. Water is Essential to Life
• Water Regulates Temperature
– Ability to resist temperature change
• Body temperature
• Coastal climates

34. Water is Essential to Life
• Water Regulates Temperature
– Evaporation
• Body temperature regulation

35. Water is Essential to Life
• Many Substances Dissolve in Water
– Solution = solvent + solute(s)
– Hydrophilic
• “water-loving”
– Hydrophobic
• “water-fearing”

36. Water is Essential to Life
• 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

37. Water is Essential to Life
• 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

38. Figure 2.14
Ice Floats.
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
H2O molecule
Ice
Liquid water

39. Ice Floats
Hydrogen bonds in water Hydrogen bonds in ice

40. Water is Essential to Life
• 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

41. Chemical Reactions
• Chemical Reaction – 2 or more molecules
“swap” atoms to make different molecules
CH4 + 2O2 CO2 + 2H2O
Reactants Products
6CO2 + 6H20 C6H12O6 + 6O2
Reactants Products

42. 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-

43. pH Scale
• Measures amount
of H+ ions
• Ranges from 0 – 14
• 0 – 6 acids
• 7 neutral
• 8 – 14 bases

44. 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

45. Applications of Chemistry to
Biology
• Ocean Acidification

46. Applications of Chemistry to
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.

47. Applications of Chemistry to
Biology

48. Applications of Chemistry to
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.