copyright © 2010 pearson education, inc. biochemistry part a biochemistry- the chemistry of living...

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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)

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