the chemistry of life chapter 2
TRANSCRIPT
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 numberElement
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 Isotopes
FUKUSHIMA, 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 atomsslightly positive (δ+)
Oxygen atomslightly negative (δ−)
Polarcovalentbonds
a. b. c.
Water molecule
Hydrogenbond
+
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 is Essential to Life
• Water Regulates Temperature– Evaporation
• Body temperature regulation
Water is Essential to Life
• Many Substances Dissolve in Water– Solution = solvent + solute(s)– Hydrophilic
• “water-loving”
– Hydrophobic• “water-fearing”
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
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
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 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
Chemical Reactions
• Chemical Reaction – 2 or more molecules “swap” atoms to make different molecules
CH4 + 2O2 CO2 + 2H2OReactants 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
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.
Applications of Chemistry to Biology
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.