1. The Chemistry of Life: Biochemistry
Gallagher, Biology 392

2. Put the following terms in size order:
Atom
Cell
Proton
Protein
Carbohydrate
Muscle
Electron

3. The Nature of Matter Matter is defined as anything that takes up space
Is there anything that is not matter?
All living things take up space and are made of matter.
Non-living things that take up space are also made of matter
YOU are made of the same material as your desk!

4. What distinguishes living from non-living?
Protons
Neutrons
Electrons
Atoms
Elements:
Shown in
Periodic Table
Not bonded
bonded
Molecules or
Compounds
Organic
Inorganic
Non-living
matter
Cells
LIFE!

5. ATOMS Basic unit of matter Structure of an atom
Nucleus- protons and neutrons held together by the “strong force”
Protons (+)
Neutrons (o)
Surrounding the nucleus - Electron Cloud (orbitals)
Electrons (-) only contains about 1/200th mass of proton or neutron
Constantly moving within orbital- attracted to the nucleus by the “weak force”
Size: 1,000,000 (million) side by side = 1 cm
Atoms like to be neutral- no charge
Equal number of protons and electrons

6. Electron orbitals 1st orbital can only hold 2 electrons (too
close to nucleus
- not much space)
2nd orbital can
hold up to 8
3rd orbital can
hold up to 8

7. Elements Pure substances A grouping of the same type of atoms
More than 100 elements exist (shown in the periodic table)
Carbon Elements:

8. The Periodic Table of Elements
Shows all the elements that exist naturally or are synthetic
Each element is represented by a single capital letter or a capital letter followed by a lowercase letter
What other information does the periodic table provide for each element?

9. C 6 12.011 Carbon Atomic Number # of protons (and also # of electrons)
Name of
Element
Chemical
symbol
Atomic Mass
The weight
Of carbon
atom or
average
weight of
all isotopes

10. Learning Checkpoint Draw a representation of a lithium atom.
Draw a representation of a Sodium atom.
What are the three subatomic particles and their charges?
What is the only actual difference between gold and mercury?
What is the atomic mass of lead?

11. Isotopes
Atoms of the same element that differ in the number of neutrons they contain
Isotopes are named by their mass number.
Mass number- protons + neutrons
Example: Carbon isotopes- Carbon 12, Carbon 13, Carbon 14
Remember: # of protons does not change
All isotopes of an atom have the same chemical properties.

12. Radioactive Isotopes
Have unstable nuclei that break down at a constant rate over time.
As it breaks down radiation is released
Good things: used to date fossils, cancer treatment, “tracer” to follow a substance in an organism, kill bacteria
Bad things: radiation is dangerous and can cause cancer

13. Compounds and Molecules
Atoms need to bond together to make molecules or compounds
Molecule- 2 or more of the same atom bonded together: H2, O2, O3
Compound- two or more of different atoms bonded together in a specific ratio: C6H12O6
Molecules and compounds are written out in a chemical formula:

15. Bonding
Covalent Bond- atoms share a pair of electrons (sometimes share 2 (double bond) or 3 (triple bond) pairs)
Ionic Bond- One atom (very unstable) gives 1, 2 or 3 electrons away to another atom. The atom that loses electrons becomes positively charged. The atom that gains the electrons becomes negatively charged. The opposite charges cause the atoms to “bond” together (opposites attract).

16. Example of a Covalent Bond
Example of Covalent Bonding- Water
Example of a Covalent Bond

17. Example of Ionic Bonding-NaCl
Na (sodium) is very unstable because it only has one
e- in its outer orbital. Cl’s (chlorine) outer orbital is
almost filled. Na gives its lonely e- to Cl.
Na become Na Cl becomes Cl-
Their opposite charges cause them to be attracted
to one another- This is an ionic bond.

18. Learning Checkpoint What is an isotope?
What are some positive uses of isotopes?
What is the difference between molecules and compounds?
Why is it important that atoms bond?
What causes atoms to bond?
Explain the difference between an ionic bond and a covalent bond.

19. 2-2 WATER! 75% of the Earth is covered with water
60-70% of your body is water
Water can be found in almost everything we eat and drink
Without any water at all you would die in 3 days

20. Properties of Water Less dense when frozen (ice floats) Polarity
Hydrogen bonds
Adhesion
Cohesion
Making Mixtures
Making Acids and Bases

21. Water Density Ice is less dense than liquid water
When water freezes air is trapped within the frozen ice making the cube larger and less dense
Benefits:
Fish and plant life can survive in liquid layers of water under ice

22. Polarity Water molecules are polar
Hydrogen ends
become slightly
positive
Water molecules are polar
Although the molecule is neutral overall there is a shift of charge within the molecule
The much larger molecule, Oxygen,
pulls more on the shared e-
This end of the molecule becomes slightly
more negative.

23. Hydrogen Bonding
Due to polarity, water molecules attract to one another
Slightly negative oxygen attracts slightly positive hydrogen from another molecule
This attraction between molecules is COHESION.
Water molecules are also attracted to other materials. This is ADHESION.

24. Water molecules attract To glass molecules And form a meniscus
COHESION
ADHESION
Water molecules attract
To glass molecules
And form a meniscus
Water molecules attract
To one another-
causes water to “bead”

25. Mixtures
Two or more elements physically mixed together but not chemically combined
Solutions- a solute (salt) is dissolved into a solvent (water
Distributes evenly
*Due to water’s polarity it can dissolve ionic compounds and other polar molecules making it THE GREATEST SOLVENT ON EARTH!
Suspensions- added substance does not dissolve but breaks into small enough pieces that it remains suspended in the water and does not settle out.

26. Making Acids and Bases
Water molecules can react to form individual ions: H2O H+ + OH-
In pure water this occurs naturally but the amount of H+ is always = to the amount of OH- so water remains neutral
Some solutions made with water become acidic or basic. This is determined by the amount of H+ (hydrogen ions) in the solution
This is measured on the pH scale

27. Acids Any compound that forms H+ ions in solution
H+ ions > OH- ions
Range from just below 7 to 0 on the pH scale
The closer to 0 the more acidic the solution
Examples: stomach acid, lemon juice

28. Bases (Alkaline) Any compound that forms 0H- ions in solution
OH- ions > H+ ions
Range from just above 7 to 14 on the pH scale
The closer to 14 the more basic the solution
Examples: lye, bleach, oven cleaner

29. Buffers Weak acids or bases that can react weith strong acids or bases
Used to regulate pH and prevent sharp sudden changes in pH
There are natural buffers in your blood that keep the pH at 6.5 to7.5

31. Learning Checkpoint Why is ice less dense than water?
What does it mean to say water molecules are polar?
What is the difference between adhesion and cohesion?
What makes a solution acidic or basic?
How is acidity measured?

32. 2-3 Carbon Compounds Why Carbon?
Carbon can from 4 covalent bonds (can create many different compounds)
Carbons can bond to one another forming large chains or rings
Linking of carbons can form very large molecules called Macromolecules
Each individual unit is called a monomer. When they are linked together they are called a polymer.
4 macromolecules necessary for life: carbohydrates, lipids, protein, nucleic acids

33. Nucleic Acid Contain hydrogen, oxygen, nitrogen, carbon and phosphorus
Monomer- nucleotide
Polymer- DNA or RNA
Store or transmit genetic information
*Nucleic Acids will be studied in greater detail when we study genetics

34. Carbohydrates Made of carbon, hydrogen and oxygen (ratio of 1:2:1)
Monomer- monosaccharides (simple sugars): glucose, galactose and fructose
Disaccharides- 2 sugars linked together: sucrose, maltose, lactose
Polymer- polysaccharides: glycogen (animals), starch and cellulose (plants)
Main source of energy

35. Lipids Made mostly of carbon and hydrogen and some oxygen
Not soluble in water: fats, oils and waxes
Monomer: all lipids have an end called glycerol in which fatty acid chains attach
Polymer- lipid
Used to store energy, also for membrane structure

36. Saturated vs. unsaturated fats
Unsaturated- at least
one double bond,
less hydrogen, can bend
Saturated- no double bonds
between carbons, all
possible hydrogens

37. Protein
Contain nitrogen, carbon, hydrogen and oxygen (amino group and carboxyl group)
Monomer- amino acid
Polymer- polypeptide or protein
Control reactions, regulate cell processes, form bones and muscles, transport and help fight disease

38. Identify the Macromolecule

39. Chemical reactions
Process that changes one set of chemicals (reactants) into another set of chemicals (products
Activation energy
-“The starting push” of chemical reactions
-The minimal amount of energy required to get a chemical reaction started

40. Enzymes
proteins that lower the activation energy required and allow reactions to happen at the normal temperature of cells.
Each enzyme is specific (only works on one particular reaction)
Can be used over and over again for that reaction
The reactant that the enzyme helps is called the substrate. The enzyme is usually named after the substrate with the ending –ase added to it.
Coenzymes are non-protein helper molecules that sometimes assist enzymes with their job.

42. Learning Checkpoint
What are the 4 carbon compounds necessary for life?
What is the main function of carbohydrates?
What are some of the functions of protein in the body?
What is the easiest way to distinguish a lipid from a carbohydrate?
What is a monomer and a polymer?