Transcript Biochemistry Part A PPT

Biochemistry Part A
• Biochemistry- the chemistry of living things
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Matter
•
Anything that has mass and occupies space
•
States of matter:
1. Solid—definite shape and volume
2. Liquid—definite volume, changeable shape
3. Gas—changeable shape and volume
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Organic/Inorganic
• Inorganic matter- mostly non living, but
essential to living organism
• *in general does not contain “C”- Carbon
• Exceptions: CO, CO2
• Abundant, and represent raw materials
needed to build life
• Organic matter- Is living, was living, came
from a living thing
• *in general contains “C”- carbon
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Composition of Matter
• Elements
• Cannot be broken down by ordinary chemical means
• Each has unique properties:
• Physical properties
• Are detectable with our senses, or are
measurable
• Chemical properties
• How atoms interact (bond) with one another
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Composition of Matter
• Atoms
• Unique building blocks for each element
• Atomic symbol: one- or two-letter chemical
shorthand for each element
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Major Elements of the Human Body
• Oxygen (O)
• Carbon (C)
• Hydrogen (H)
• Nitrogen (N)
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About 96% of body mass
Lesser Elements of the Human Body
• About 3.9% of body mass:
• Calcium (Ca), phosphorus (P), potassium (K),
sulfur (S), sodium (Na), chlorine (Cl),
magnesium (Mg), iodine (I), and iron (Fe)
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Trace Elements of the Human Body
• < 0.01% of body mass:
• Part of enzymes, e.g., chromium (Cr),
manganese (Mn), and zinc (Zn)
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Atomic Structure
• Determined by numbers of subatomic
particles
• Nucleus consists of neutrons and protons
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Atomic Structure
• Neutrons
• No charge
• Mass = 1 atomic mass unit (amu)
• Protons
• Positive charge
• Mass = 1 amu
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Atomic Structure
• Electrons
• Orbit nucleus
• Equal in number to protons in atom
• Negative charge
• 1/2000 the mass of a proton (0 amu)
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Models of the Atom
• Orbital model: current model used by
chemists
• Depicts probable regions of greatest electron
density (an electron cloud)
• Useful for predicting chemical behavior of
atoms
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Models of the Atom
• Planetary model—oversimplified, outdated
model
• Incorrectly depicts fixed circular electron paths
• Useful for illustrations (as in the text)
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Nucleus
Nucleus
Helium atom
Helium atom
2 protons (p+)
2 neutrons (n0)
2 electrons (e–)
2 protons (p+)
2 neutrons (n0)
2 electrons (e–)
(a) Planetary model
Proton
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Neutron
(b) Orbital model
Electron
Electron
cloud
Figure 2.1
Identifying Elements
• Atoms of different elements contain different
numbers of subatomic particles
• Compare hydrogen, helium and lithium (next
slide)
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Proton
Neutron
Electron
Hydrogen (H)
(1p+; 0n0; 1e–)
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Helium (He)
(2p+; 2n0; 2e–)
Lithium (Li)
(3p+; 4n0; 3e–)
Figure 2.2
Identifying Elements
• Atomic number = number of protons in
nucleus
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Identifying Elements
• Mass number = mass of the protons and
neutrons
• Mass numbers of atoms of an element are not
all identical
• Isotopes are structural variations of elements
that differ in the number of neutrons they
contain
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Identifying Elements
• Atomic weight = average of mass numbers of
all isotopes
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Proton
Neutron
Electron
Hydrogen (1H)
(1p+; 0n0; 1e–)
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Deuterium (2H)
(1p+; 1n0; 1e–)
Tritium (3H)
(1p+; 2n0; 1e–)
Figure 2.3
Radioisotopes
• Spontaneous decay (radioactivity)
• Similar chemistry to stable isotopes
• Can be detected with scanners
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Radioisotopes
• Valuable tools for biological research and
medicine
• Cause damage to living tissue:
• Useful against localized cancers
• Radon from uranium decay causes lung
cancer
• Other Values of Radatiosotopes…
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Molecules and Compounds
• Most atoms combine chemically with other
atoms to form molecules and compounds
• Molecule—two or more atoms bonded
together (e.g., H2 or C6H12O6)
• Compound—two or more different kinds of
atoms bonded together (e.g., C6H12O6)
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Chemically Inert Elements
• Stable and unreactive
• Outermost energy level fully occupied or
contains eight electrons
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(a)
Chemically inert elements
Outermost energy level (valence shell) complete
8e
2e
Helium (He)
(2p+; 2n0; 2e–)
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2e
Neon (Ne)
(10p+; 10n0; 10e–)
Figure 2.5a
Chemically Reactive Elements
• Outermost energy level not fully occupied by
electrons
• Tend to gain, lose, or share electrons (form
bonds) with other atoms to achieve stability
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(b)
Chemically reactive elements
Outermost energy level (valence shell) incomplete
1e
Hydrogen (H)
(1p+; 0n0; 1e–)
6e
2e
Oxygen (O)
(8p+; 8n0; 8e–)
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4e
2e
Carbon (C)
(6p+; 6n0; 6e–)
1e
8e
2e
Sodium (Na)
(11p+; 12n0; 11e–)
Figure 2.5b
Types of Chemical Bonds
• Ionic
• Covalent
• Hydrogen
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Ionic Bonds
• Ions are formed by transfer of valence shell
electrons between atoms
• Anions (– charge) have gained one or more
electrons
• Cations (+ charge) have lost one or more
electrons
• Attraction of opposite charges results in an
ionic bond
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Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
+
–
Sodium ion (Na+)
Chloride ion (Cl–)
Sodium chloride (NaCl)
(a) Sodium gains stability by losing one electron, and
chlorine becomes stable by gaining one electron.
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(b) After electron transfer, the oppositely
charged ions formed attract each other.
Figure 2.6a-b
Formation of an Ionic Bond
• Ionic compounds form crystals instead of
individual molecules
• NaCl (sodium chloride)
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CI–
Na+
(c) Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
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Figure 2.6c
Covalent Bonds
• Formed by sharing of two or more valence
shell electrons
• Allows each atom to fill its valence shell at
least part of the time
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Reacting atoms
Resulting molecules
+
Molecule of
Hydrogen
Carbon
methane gas (CH4)
atoms
atom
(a) Formation of four single covalent bonds:
carbon shares four electron pairs with four
hydrogen atoms.
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or
Structural
formula
shows
single
bonds.
Figure 2.7a
Reacting atoms
Resulting molecules
+
Oxygen
atom
or
Oxygen
atom
Molecule of
oxygen gas (O2)
(b) Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
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Structural
formula
shows
double
bond.
Figure 2.7b
Reacting atoms
Resulting molecules
+
Nitrogen
atom
or
Nitrogen
atom
Molecule of
nitrogen gas (N2)
(c) Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
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Structural
formula
shows
triple
bond.
Figure 2.7c
Covalent Bonds
• Sharing of electrons may be equal or unequal
• Equal sharing produces electrically balanced
nonpolar molecules
• CO2
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Figure 2.8a
Covalent Bonds
• Unequal sharing by atoms with different
electron-attracting abilities produces polar
molecules
• H2O
• Atoms with six or seven valence shell
electrons are electronegative, e.g., oxygen
• Atoms with one or two valence shell
electrons are electropositive, e.g., sodium
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Figure 2.8b
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Figure 2.9
Hydrogen Bonds
• Attractive force between electropositive
hydrogen of one molecule and an
electronegative atom of another molecule
• Common between dipoles such as water
• Also act as intramolecular bonds, holding a
large molecule in a three-dimensional shape
PLAY
Animation: Hydrogen Bonds
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+
–
Hydrogen bond
(indicated by
dotted line)
+
+
–
–
–
+
+
+
–
(a) The slightly positive ends (+) of the water
molecules become aligned with the slightly
negative ends (–) of other water molecules.
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Figure 2.10a
(b) A water strider can walk on a pond because of the high
surface tension of water, a result of the combined
strength of its hydrogen bonds.
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Figure 2.10b
Chemical Reactions
• Occur when chemical bonds are formed,
rearranged, or broken
• Represented as chemical equations
• Chemical equations contain:
• Molecular formula for each reactant and
product
• Relative amounts of reactants and products,
which should balance
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Examples of Chemical Equations
H + H  H2 (hydrogen gas)
(reactants)
(product)
4H + C  CH4 (methane)
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Patterns of Chemical Reactions
• Synthesis (combination) reactions
• Decomposition reactions
• Exchange reactions
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Synthesis Reactions
• A + B  AB
• Always involve bond formation
• Anabolic
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(a) Synthesis reactions
Smaller particles are bonded
together to form larger,
more complex molecules.
Example
Amino acids are joined together to
form a protein molecule.
Amino acid
molecules
Protein
molecule
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Figure 2.11a
Decomposition Reactions
• AB  A + B
• Reverse synthesis reactions
• Involve breaking of bonds
• Catabolic
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(b) Decomposition reactions
Bonds are broken in larger
molecules, resulting in smaller,
less complex molecules.
Example
Glycogen is broken down to release
glucose units.
Glycogen
Glucose
molecules
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Figure 2.11b
Chemical Reactions
• All chemical reactions are either exergonic or
endergonic
• Exergonic reactions—release energy
• Catabolic reactions
• Endergonic reactions—products contain more
potential energy than did reactants
• Anabolic reactions
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Chemical Reactions
• All chemical reactions are theoretically reversible
• A + B  AB
• AB  A + B
• Chemical equilibrium occurs if neither a forward nor
reverse reaction is dominant
• Many biological reactions are essentially irreversible
due to
• Energy requirements
• Removal of products
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Rate of Chemical Reactions
• Rate of reaction is influenced by:
•  temperature   rate
•  particle size   rate
•  concentration of reactant   rate
• Catalysts:  rate without being chemically
changed
• Enzymes are biological catalysts
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