MULTIDISCIPLINARY SCIENCE
Biological systems follow the same laws that are dictated by physics and chemistry.
Emergent properties – biological functions are made possible by the interaction
between atoms and molecules that make up an organism.
Discuss:
What is the definition of life?
What is required for something to be considered alive?
At what point do we cross the line between nonliving and living at the molecular
level of cell biology?
Do you consider a virus to be “alive?”
Bombardier Beetle (chemistry) - https://www.youtube.com/watch?v=Pib9qT-pccI
Pistol Shrimp (physics) - https://www.youtube.com/watch?v=QXK2G2AzMTU
THE BASICS - IMPORTANT TERMS
Matter – Anything that has mass and occupies volume.
Element – A substance that cannot be broken down into other substances through
chemical reactions.
Compound – A substance that consists of two or more different elements in a fixed
ratio.
Atom – The smallest unit of an element that maintains the chemical properties of the
element.
Molecule – The smallest unit of a compound that maintains the chemical properties
of the compound,
LE 2-2
Sodium
Chlorine
Sodium chloride
THE BASICS – PERIODIC TABLE
ELEMENTS ESSENTIAL TO LIFE
92 natural known elements
25 elements are essential to life
4 elements make up 96% of living matter
• Carbon (C)
• Oxygen (O)
• Hydrogen (H)
• Nitrogen (N)
• Remaining 4%
• Phosphorous (P)
• Sulfur (S)
• Calcium (Ca)
• Potassium (K)
• Less that 0.01% of living matter is made
up of trace elements
•
•
•
THE STRUCTURE OF ATOMS
•
Subatomic particles
•
Neutrons - ~1 amu (1.7 x10ˆ-24 grams); no charge
•
Protons - ~1 amu; +1e (1.602 x 10ˆ-19 coulombs)
•
Electrons - 5.4858 x 10ˆ-4 amu (or 1/1840 amu); -1e (-1.602 x 10ˆ-19
coulombs)
•
What makes up the atomic nucleus?
•
1 amu=1 dalton
•
•
Atomic Weight - # of Protons + # of Neutrons
• Mass of electrons is negligible
• AKA Atomic Mass; Mass Number
Atomic Number - # of Protons
Atomic Number
Atomic Number
Atomic Weight
Atomic Weight
LE 2-4
Cloud of negative
charge (2 electrons)
Electrons
Nucleus
ISOTOPES
•
Atoms of the same element that have different atomic weights
•
Always have the same number of protons
•
Differ in the number of neutrons
•
Atomic Weight = Average Atomic Weight of all isotopes
•
Some isotopes are stable
•
Most isotopes are unstable (radioactive)
ISOTOPES – APPLICATIONS IN LIFE SCIENCES
• Dating techniques for biological
materials and fossils
• Tracers used to follow metabolic
processes (Rate of DNA synthesis)
• Medical Diagnostics
• PET Scans
• Risk associated with using radioactive
isotopes
• Severity depends on type and amount
of radiation
ISOTOPES - PRACTICE
How many protons, neutrons, and electrons are present in the following isotopes?
•
Technetium-99
•
Flourine-18
•
Iodine-131
•
Cobalt-60
What is the atomic mass of an element, given the following isotopes?
Oxygen-16, oxygen-17, oxygen-18, oxygen-12, oxygen-24, and oxygen-15
ELECTRONS – ENERGY LEVELS
•
Energy is the capacity to cause a change; the capacity to do work (exert a force
on an object).
•
The energy level of an electron is the potential energy of that electron.
•
Average distance of an electron from the nucleus
•
Represented by electron shells
•
Shells further from the nucleus represent higher levels of energy
•
Energy input (light, etc.) required to move an electron to a higher energy level
(further shell)
•
Energy released (usually heat) when electron moves back to original shell (ground
state)
The higher up the ball is,
the more potential energy
is has
ELECTRONS – CONFIGURATION/CHEMICAL
PROPERTIES
•
Chemical properties of an element depend on the number of valence electrons
(outermost shell).
•
Elements with the same number of valence electrons exhibit similar chemical
properties
ELECTRONS – ORBITALS
•
Electron Orbital – The area within an atom that electrons have the highest
probability of being found.
•
Each shell has a specific number of orbitals.
•
Each orbital can contain up to 2 electrons.
•
Different orbitals represented by different shapes
•
Types of orbitals –
• s orbitals (spherical) – 1 per shell
• p orbitals (dumbbell-shaped) – 3 per shell
• d orbitals – 5 per shell
GROUND STATE ELECTRON CONFIGURATION
The ground state electron configuration for oxygen is…
1s22s22p4
For aluminum…
1s22s22p63s23p1
What is the ground state electron configuration for…
•
Sodium (Na)
•
Sulfur (S)
•
Potassium (K)
•
Scandium (Sc)
CHEMICAL BONDING
• Most stable atoms are those with completely filled
valence shells (usually eight valence electrons).
• Atoms interact with each other to form bonds in order
to fill their valence shells.
• Types of chemical bonding
• Covalent
• Ionic
• Metallic
COVALENT BONDS
Atoms fill their valence shells by sharing electrons with other atoms
Occur between nonmetals
Generally have a low to intermediate difference in electronegativity
Polar covalent bonds – electrons are shared unequally; one atom has a more
negative charge (δ-) and the other has a more positive charge (δ+).
Nonpolar covalent bonds – electrons are shared equally and result in a net charge of
zero. No difference in electronegativity.
Bond order – ½(# of bonding electrons); single, double, or triple bonds.
Bonding capacity = atom’s valence (number of unpaired valence electrons).
The structure of a molecule depends on the type and number of covalent bonds
present, as well as the number of paired electrons.
Tetrahedral
Bent
IONIC BONDS
Atoms that have a high difference in electronegativity
Atom with higher electronegativity strips electron away
from atom with lower electronegativity to create two
ions.
Cation – positively charged ion
Anion – negatively charged ion
Resulting ions are attracted to each other, forming an
ionic bond.
Metals form cations; nonmetals form anions
IONIC COMPOUNDS
Also referred to as salts
Most commonly found in nature as crystals
Formula only represents ratio of atoms (NaCl); individual molecules do not form
Ionic compounds form lattice structures
Cl–
Na+
WEAK CHEMICAL BONDS
Hydrogen bonds – Hydrogen atoms in a polar covalent bond are attracted to another
electronegative atom
Van der Waals Interactions – Due to delocalized electrons, molecules form temporary
“hotspots” of charge that enable molecules to stick together.
MOLECULAR SHAPE AND FUNCTION
Molecular shape depends on the positions of its atoms’ valence orbitals.
In covalent bonds, overlapping of orbitals results in orbital hybridization.
Specific shapes are associated with specific hybrid orbitals.
MOLECULAR SHAPE AND FUNCTION
Molecular function depends on molecular shape
Molecules interact with each other depending on shape
Enzymes – Fit into specific receptor cites based on shape
CHEMICAL REACTIONS
Molecular bonds break Atoms rearrange themselves  New molecular bonds form
CHEMICAL REACTIONS
Chemical Equilibrium – The rate of a chemical reaction is the same as the rate of the
reverse reaction
Concentration of reactants and products is NOT necessarily equal when chemical
equilibrium is reached.
Reactions still occur after equilibrium is reached, but there is no net change in the
concentration of either reactants or products.
EXAM PRACTICE
The atomic number of sulfur is 16. Sulfur
combines with hydrogen by covalent bonding
to form a compound, hydrogen sulfide.
Based on the electron configuration of sulfur,
we can predict that the molecular formula of
the compound will be;
A.) HS
b.) HS2
c.) H2S
d.)H3S2
E.) H4S
EXAM PRACTICE
Draw the Lewis Structure for the following molecules…
HCN
H2S
CH2Br2
EXAM PRACTICE
Why is it important that we understand chemistry when we
are studying biology?
What coefficients must be placed in the blanks so that all
atoms are accounted for in the products?
___C6H12O6 → ___C2H6O + ___CO2
___S8 + ___O2 → ___SO2
___Al2(SO4)3 + ___Ca(OH)2 → ___Al(OH)3 + ___CaSO4