Life’s Chemical Basis -...
Transcript of Life’s Chemical Basis -...
Life’s Chemical Basis
Life’s Chemical Basis
Ø Atoms and Elements
Ø Why Electrons Matter
Ø Atomic Bonds
Ø Water molecule properties
Ø Hydrogen Power (pH)
Matter & Elements
Ø Matter is anything that occupies space and has mass
Ø Matter is composed of chemical elements
Ø An element is a substance that cannot be broken down into other substances
Elements and Atoms
Ø The three subatomic particles are: • Protons are positively charged • Neutrons are electrically neutral • Electrons are negatively charged
Ø Each element consists of one kind of atoms
Ø An atom is the smallest unit of matter that still retains the properties of an element
Terminology/ Elements Ø Atomic Number: The number of protons in an atom
atomic number
mass number
element symbol
Ø Mass Number: The sum of the numbers of protons and neutrons in the nucleus
atomic number
element symbol
mass number
elemental substance element name
carbon
Periodic Table of Elements
Elements Required for Life Ø O, C, H, N make up
about 96% of the human body weight
Ø Ca, P, K, S, Na, Cl, Mg make up most of the rest
Ø Trace elements make less than 0.01% of human body weight
Required in minute amounts but they are crucial!
Isotopes Isotopes of an element have the same number of protons but different numbers of neutrons (different mass numbers)
Stable isotope: enough binding energy to hold the nucleus together permanently 12C and 13C
Isotopes
the nucleus decays spontaneously, giving off particles and energy
Unstable isotope, Radioactive isotopes, Radioisotopes Isotope with an unstable nucleus/ 14C
nucleus of 14C, with
6 protons, 8 neutrons
nucleus of 14N, with
7 protons, 7 neutrons
Ø Under natural conditions, elements occur as a mixture of isotopes
Ø Different isotopes of an element behave identically in chemical reactions Ø Living cells cannot distinguish between isotopes of the same element
Ø Isotopes have fixed half-life/ radioisotopes decay at a predictable rate into predictable products
14C 14N
Biological Applications of Radioactive Isotopes
Basic Research/ radioactive tracers/ detectable component/can be tracked after delivery into the body or system Archaeological Dating/ based on decay rate of radioisotopes (half life) Medical Treatments and Diagnosis
β-amyloid/ PiB
Why do Atoms Interact?
Ø Of the three subatomic particles, only electrons are directly involved in chemical activity
Ø Electrons occupy different orbitals (volumes of space around an atom’s nucleus)
Ø Orbitals are filled from lower to higher energy Ø The farther an electron is from the nucleus, the
greater its energy
Arrangement of electrons around the nucleus
shell model
first shell hydrogen (H) helium (He)
second shell carbon (C) oxygen (O)
third shell sodium (Na) chlorine (Cl)
neon (Ne)
argon (Ar)
one proton
one electron
Vacancies/ Can the outer shell hold more electrons?
Ø No vacancies: an atom’s outermost shell is filled with electrons Ø Most stable state
Ø Vacancy: an atom’s outermost shell has room for other electrons Ø Chemically active; atoms
interact with one another Ø Example: oxygen atom has
six electrons in its outer (second) shell, which can hold eight
Oxygen (O) Neon (Ne)
Chemical Bonds An atom can get rid of vacancies by participating in chemical bonds with other atoms Atoms with incomplete outer shells share, donate or receive electrons
Chemical bond is the attractive force that keeps atoms together in a molecule
electron loss
Sodium atom Sodium ion 11p+
11e-
charge: 0
11p+
10e-
charge: +1
Ø Atoms with an unequal number of protons and electrons are called ions Ø Carry a net (overall) charge
Compounds
Ø A compound is a substance consisting of two or more different elements in a fixed ratio
Ø Compounds are more common than pure elements
Ø Sodium chloride, table salt, is a common compound of equal parts of sodium (Na) and chlorine (Cl)
Sodium Chlorine
Sodium Chloride
+
Types of chemical bonds
Ø Covalent Bonds
Ø Ionic Bonds
Ø Hydrogen Bonds
Ionic Bonds
Ø An ion is an atom or group of atoms with an electrical charge resulting from gain or loss of electrons
Ø Two ions with opposite charges attract each other
Ø When the attraction holds the ions together ionic bond
Example of ionic bond: Sodium Chloride
Transfer of electron
Sodium atom Chlorine atom Na+
Sodium ion Cl-
Chloride ion
Sodium chloride (NaCl)
Ions retain their respective charges when participating in an ionic bond
Electronegativity measure of the ability of an atom to pull electrons away from other atoms
Polarity separation of charge into positive and negative regions
Covalent Bonds
Ø Two atoms with incomplete outer shells share one or more electrons
Types of Covalent Bonds Non-polar:
Ø Bond between atoms with the same electronegativity
Polar: Ø Bond between atoms of
different electronegativity
Methane Water
Hydrogen Bonds Ø Hydrogen, as part of a polar
covalent bond, has a partial positive charge
Ø Hydrogen atoms that are covalently bonded in a molecule can be attracted to atoms with slight negative charges in other molecules such as oxygen and nitrogen
Ø Because the positively charged region is always a hydrogen atom, the bond is called a hydrogen bond
Ø H-bonds help define the properties of water, create the shape of proteins, hold DNA strands together, etc
Water’s Life-supporting Properties
Ø 70% of Earth is covered with water
Ø All living organisms require water more than any other substance
Ø Most cells are surrounded by water, and cells themselves are about 70-95% water
Polar Molecule Hydrogen Bonds
Universal Solvent
Solution/ homogeneous mixture Solvent + Solute Aqueous solution
hydrophilic (“water-loving”) molecules/ polar hydrophobic (“water-hating”) molecules/ non-polar
Universal Solvent
Ø Water is an effective solvent because of its polarity
Ø When an ionic compound is dissolved in water, each ion is surrounded by a sphere of water molecules, a hydration shell
Universal Solvent
Non ionic compounds such as sugar dissolve as water molecules surround them and form hydrogen bonds with their polar regions
Cohesion Ø The tendency of molecules of the same kind to stick together
• Hydrogen bonds make liquid water cohesive
• Surface tension
Ø Cohesion in water is much stronger than other liquids
https://youtu.be/iAT_iBgn1QI
Ø Because of hydrogen bonds between water molecules, it takes more heat to raise the temperature of water compared with other liquids (Temperature: measure of molecular motion)
Ø Below 0°C (32°F), water molecules become locked in the bonding pattern of ice (lattice pattern) Ø Sheets of ice that form on the surface of ponds,
lakes, and streams insulate the water Ø Protects aquatic organisms during cold winters
Water stabilizes temperature
Ice is less dense than liquid water
Ice floats in liquid water because hydrogen bonds in ice are more “ordered,” making ice less dense
Ø pH: measure of the number of hydrogen ions in a fluid
Ø Base: substance that accepts hydrogen ions in water Ø pH above 7
Ø Acid: substance that releases hydrogen ions in water Ø pH below 7
Ø Most biological systems can function properly only within a narrow range of pH
Ø The fluids inside cells stay within a consistent range of pH because they are buffered
Ø Buffer: set of chemicals that can keep the pH
of a solution stable by alternately donating and accepting ions that contribute to pH
household ammonia
gastric fluid
acid rain lemon juice
cola vinegar
tomatoes, wine bananas
urine, tea, typical rain
corn
milk pure water
blood, tears
baking soda Tums
oven cleaner
bleach
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HCl is a strong acid (the gastric fluid) Carbonic acid is a weak acid: CO2 + H2O = H2CO3 H2CO3 H+ + HCO3
- H2CO3 Carbonic Acid Bicarbonate
Ø Acid rain (pH < 5.2)
Changes in Acidity can Have Environmental Consequences
Ø Ocean acidification
CO2 + H2O H2CO3