Transcript The Chemistry of Life

The Chemistry of Life:
Biochemistry
Gallagher, Biology 392
Put the following terms in size order:
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Atom
Cell
Proton
Protein
Carbohydrate
Muscle
Electron
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!
What distinguishes living from non-living?
Protons
Neutrons
Electrons
Molecules or
Compounds
Organic
Cells
LIFE!
Atoms
bonded
Inorganic
Non-living
matter
Not bonded
Elements:
Shown in
Periodic Table
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
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
Elements
• Pure substances
• A grouping of the same type of atoms
• More than 100 elements exist (shown in the periodic
table)
• Carbon Elements:
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?
Atomic Number
# of protons
(and also # of
electrons)
Chemical
symbol
6
C
Name of
Element
Carbon
12.011
Atomic Mass
The weight
Of carbon
atom or
average
weight of
all isotopes
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?
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.
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
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:
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).
Example of Covalent Bonding- Water
Example of a Covalent Bond
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.
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.
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
Properties of Water
• Less dense when frozen (ice floats)
• Polarity
• Hydrogen bonds
– Adhesion
– Cohesion
• Making Mixtures
• Making Acids and Bases
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
Polarity
• Water
molecules are
polar
• Although the
molecule is
neutral overall
there is a shift
of charge
within the
molecule
Hydrogen ends
become slightly
positive
The much larger molecule, Oxygen,
pulls more on the shared eThis end of the molecule becomes slightly
more negative.
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.
COHESION
Water molecules attract
To one anothercauses water to “bead”
ADHESION
Water molecules attract
To glass molecules
And form a meniscus
Mixtures
•
Two or more elements physically mixed
together but not chemically combined
1. 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!
2. Suspensions- added substance does not
dissolve but breaks into small enough pieces
that it remains suspended in the water and
does not settle out.
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 OHso 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
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
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
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
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?
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
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
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
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
Saturated vs. unsaturated fats
Saturated- no double bonds
between carbons, all
possible hydrogens
Unsaturated- at least
one double bond,
less hydrogen, can bend
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
Identify the Macromolecule
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
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.
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?