copyright © 2010 pearson education, inc. biochemistry part a biochemistry- the chemistry of living...
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Copyright © 2010 Pearson Education, Inc.
Biochemistry Part A
• Biochemistry- the chemistry of living things
Copyright © 2010 Pearson Education, Inc.
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)
About 96% of body mass
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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)
Copyright © 2010 Pearson Education, Inc. Figure 2.1
(a) Planetary model (b) Orbital model
Helium atom
2 protons (p+)2 neutrons (n0)2 electrons (e–)
Helium atom
2 protons (p+)2 neutrons (n0)2 electrons (e–)
Nucleus Nucleus
Proton Neutron Electroncloud
Electron
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Identifying Elements
• Atoms of different elements contain different numbers of subatomic particles
• Compare hydrogen, helium and lithium (next slide)
Copyright © 2010 Pearson Education, Inc. Figure 2.2
Proton
Neutron
Electron
Helium (He)(2p+; 2n0; 2e–)
Lithium (Li)(3p+; 4n0; 3e–)
Hydrogen (H)(1p+; 0n0; 1e–)
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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
Copyright © 2010 Pearson Education, Inc. Figure 2.3
Proton
Neutron
Electron
Deuterium (2H)(1p+; 1n0; 1e–)
Tritium (3H)(1p+; 2n0; 1e–)
Hydrogen (1H)(1p+; 0n0; 1e–)
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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
Copyright © 2010 Pearson Education, Inc. Figure 2.5a
Helium (He)(2p+; 2n0; 2e–)
Neon (Ne)(10p+; 10n0; 10e–)
2e 2e8e
(a) Chemically inert elements
Outermost energy level (valence shell) complete
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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
Copyright © 2010 Pearson Education, Inc. Figure 2.5b
2e4e
2e8e
1e
(b) Chemically reactive elementsOutermost energy level (valence shell) incomplete
Hydrogen (H)(1p+; 0n0; 1e–)
Carbon (C)(6p+; 6n0; 6e–)
1e
Oxygen (O)(8p+; 8n0; 8e–) Sodium (Na)
(11p+; 12n0; 11e–)
2e6e
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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
Copyright © 2010 Pearson Education, Inc. Figure 2.6a-b
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.
(b) After electron transfer, the oppositely charged ions formed attract each other.
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Formation of an Ionic Bond
• Ionic compounds form crystals instead of individual molecules
• NaCl (sodium chloride)
Copyright © 2010 Pearson Education, Inc. Figure 2.6c
CI–
Na+
(c) Large numbers of Na+ and Cl– ions associate to form salt (NaCl) crystals.
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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
Copyright © 2010 Pearson Education, Inc. Figure 2.7a
+
Hydrogenatoms
Carbonatom
Molecule ofmethane gas (CH4)
Structuralformulashows singlebonds.
(a) Formation of four single covalent bonds: carbon shares four electron pairs with four hydrogen atoms.
or
Resulting moleculesReacting atoms
Copyright © 2010 Pearson Education, Inc. Figure 2.7b
or
Oxygenatom
Oxygenatom
Molecule ofoxygen gas (O2)
Structuralformulashowsdouble bond.(b) Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
Resulting moleculesReacting atoms
+
Copyright © 2010 Pearson Education, Inc. Figure 2.7c
+ or
Nitrogenatom
Nitrogenatom
Molecule ofnitrogen gas (N2)
Structuralformulashowstriple bond.(c) Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
Resulting moleculesReacting atoms
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Covalent Bonds
• Sharing of electrons may be equal or unequal
• Equal sharing produces electrically balanced nonpolar molecules
• CO2
Copyright © 2010 Pearson Education, Inc. Figure 2.8a
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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
Copyright © 2010 Pearson Education, Inc. Figure 2.8b
Copyright © 2010 Pearson Education, Inc. Figure 2.9
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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 PLAY Animation: Hydrogen Bonds
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(a) The slightly positive ends (+) of the watermolecules become aligned with the slightlynegative ends (–) of other water molecules.
+
–
–
–– –
+
+
+
+
+
Hydrogen bond(indicated bydotted line)
Figure 2.10a
Copyright © 2010 Pearson Education, Inc. Figure 2.10b
(b) A water strider can walk on a pond because of the highsurface tension of water, a result of the combinedstrength of its hydrogen bonds.
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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)
4H + C CH4 (methane)
(reactants) (product)
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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
Copyright © 2010 Pearson Education, Inc. Figure 2.11a
ExampleAmino acids are joined together toform a protein molecule.
(a) Synthesis reactions
Smaller particles are bondedtogether to form larger,
more complex molecules.
Amino acidmolecules
Proteinmolecule
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Decomposition Reactions
• AB A + B
• Reverse synthesis reactions
• Involve breaking of bonds
• Catabolic
Copyright © 2010 Pearson Education, Inc. Figure 2.11b
ExampleGlycogen is broken down to releaseglucose units.
Bonds are broken in largermolecules, resulting in smaller,
less complex molecules.
(b) Decomposition reactions
Glucosemolecules
Glycogen
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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