1. Chemistry Review
Chapter 2 in Text

2. Life can be organized into a hierarchy of structural levels.
Atoms are organized into molecules, and molecules are organized into cells.
Somewhere in the transition from molecules to cells, we cross the boundary between nonlife and life.
At each successive level, additional emergent properties appear
Scale of the Universe
Chemistry of Life
Living organisms and the world they live in are subject to the basic laws of physics and chemistry.

3. Emergent Properties of Systems
Has properties and functions NOT present in the individual components.
Whole (system) is greater than the sum of the parts.
CANNOT predict properties of system from properties of parts
This is called “Synergy”
“There is nothing in the chemistry of a toenail that predicts the existence of a human being.”
Buckminster Fuller

4. Synergy in Salt Ordinary Table Salt (NaCl, Sodium Chloride)
Na (sodium) is a metal, highly reactive in water, explosive, burns with a yellow flame.
Cl (chlorine) is a deadly greenish gas.
Both are poisonous and deadly BUT

5. SALT When combined, produce something that we cannot live without.
Something unique and new and unpredictable happens when two or more things come together and work together

6. Biological Hierarchy
New properties emerge at each level in the biological hierarchy.
Each level of biological organization has emergent properties.
Biological organization is based on a hierarchy of structural levels, each building on the levels below.
Synergy!

7. Organisms are Composed of Matter
Matter is made up of elements.
About 25 of the 92 naturally occurring elements are known to be essential for life.
Four elements—carbon (C), oxygen (O), hydrogen (H), and nitrogen (N)—make up 96% of living matter.
Most of the remaining 4% of an organism’s weight consists of phosphorus (P), sulfur (S), calcium (Ca), and potassium (K).
Trace elements are required by an organism but only in minute quantities.
Some trace elements, like iron (Fe), are required by all organisms.
Other trace elements are required by only some species.

8. What do we Remember?
Protons
Neutrons
Electrons
Mass
Charge
Location

9. Subatomic particles: proton, neutron, electron
Atomic number - number of protons
Mass number -the sum of the number of protons and the number of neutrons in the nucleus of an atom.
Electron configuration and chemical properties
Valence electrons
An elements properties depend on the structure of its atoms
Hydrogen – forms when hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom. (usually oxygen or nitrogen)
Van der Waals interactions = because electrons are always in motion they are not symmetrically distributed and even a non-polar molecule may develop regions that are slightly more + or -. These only when molecules are very close together or withing a large molecule (protein)

11. Valence Electrons

13. The Energy Levels of Electrons
An atom’s electrons
Vary in the amount of energy they possess
They have “energy levels”
If an atom absorbs energy, then the electrons will be “bumped up” to a “higher energy level”

14. Energy Energy Potential energy
Is defined as the capacity to cause change
Potential energy
Is the energy that matter possesses because of its location or structure
Types?

15. Energy Levels Energy levels Are represented by electron shells
Third energy level (shell)
Second energy level (shell)
First energy level (shell)
Energy
absorbed
lost
An electron can move from one level to another only if the energy
it gains or loses is exactly equal to the difference in energy between
the two levels. Arrows indicate some of the step-wise changes in
potential energy that are possible.
(b)
Atomic
nucleus
Figure 2.7B

16. Chemical Bonds Covalent- two atoms share a pair of valence electrons
The attraction of an atom for the shared electrons of a covalent bond is called its electronegativity.
Ionic - two atoms are so unequal in their attraction for valence electrons that one atom strips an electron completely from the other. “stealing electrons”

17. Covalent Bonds Covalent Bond: Electrons are shared
when the atoms share:
two electrons  single bond
four electrons  double bond
six electrons  triple bond

18. Covalent Bond Example: H2O

19. Ionic Bonds All atoms want 8 valence electrons
They can gain or lose electrons to get 8
Ionic bond: electrons are transferred
gain electron: negative charge
Anion
lose electron: positive charge
Cation
Ions: charged atoms

20. Ionic Bond Example: NaCl

21. Types of Bonds

22. Hydrogen Bonds
Weak bonds that form when the partial positive charge of a hydrogen atom in a polar covalent bond is attracted to the partial negative charge on another polar covalently bonded atom.
 –
 +
Water
(H2O)
Ammonia
(NH3) O H
 +
 – N
A hydrogen
bond results
from the
attraction
between the
partial positive
charge on the
hydrogen atom
of water and
the partial
negative charge
on the nitrogen
atom of
ammonia. + d+
Figure 2.15
 +

23. A Note on Hydrogen
Hydrogen Ion = Proton

24. Molecular shape & function
Molecular shape determines how biological molecules recognize and respond to one another with specificity

25. Figure 2.17 Carbon Nitrogen Hydrogen Sulfur Oxygen Natural endorphin
Morphine
Carbon
Hydrogen
Nitrogen
Sulfur
Oxygen
Natural
endorphin
(a) Structures of endorphin and morphine. The boxed portion of the endorphin molecule (left) binds to receptor molecules on target cells in the brain. The boxed portion of the morphine molecule is a close match.
(b) Binding to endorphin receptors. Endorphin receptors on the surface of a brain cell recognize and can bind to both endorphin and morphine.
Endorphin
receptors
Brain cell
Figure 2.17

26. Chemical Reactions Reactants→ products
Most reactions are reversible leading to chemical equilibrium
The rate of formation of products is the same as the rate of breakdown of products (formation of reactants), and the system is at chemical equilibrium.
At equilibrium, the concentrations of reactants and products are typically not equal, but their concentrations have stabilized at a particular ratio.