The Chemistry of Life - Lecture notes 1 PDF

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This is an overview of the chemistry of life. Topics discussed include atoms and atomic structure, electron shells, atom identification, elements, isotopes, radioactivity, combination of matter, three states of matter, types of mixtures, chemical bonds, molecular formulas, valance electrons, the two...

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1 The Chemistry of Life Atoms and atomic structure  Atoms: smallest, stable unit of matter o Subatomic particles make up atoms  Atomic nucleus o Protons (p +): Positive electrical change o Neutrons (n or no): No electrical change (neutral) 

Electron cloud (shell) outside nucleus o Electrons (e-): Negative electrical charge. Much smaller than protons or neutrons.

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Atoms are electrically neutral: The number of positively charged protons in an atom equals the number of negatively charged electrons.

Electron shells (clouds): Electrons circle nucleus. Circle in orbits (shells) numbered 1-7  1st shell: closest to nucleus, can hold 2 electrons  2nd shell: can hold 8 electrons  3rd shell: can hold 18 electrons but “satisfied” with 8

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Atoms are identified by atomic number and mass number  Atomic number: Number of protons  Mass number: Number of protons and neutrons  Hydrogen atomic number: 1 (only a proton) Elements: Composed only of atoms with same atomics number (protons). Cannot be broken down into simpler substance by chemical means. Represented by a chemical symbol (English or Latin)  Example: Hydrogen [H], Oxygen [O], Potassium [K], Sodium [Na]

Isotopes and radioactivity  Isotopes are different atoms of the same element o Atomic number of isotopes: Same, Same number of protons o Mass number of isotopes: number of subatomic particles in nuclei, different (Different number of neutrons)  Chemical properties of isotopes: Same, some radioactive. Matter combined: Mixtures and chemical bonds o Matter defined: Anything that has mass and occupies space o Solid: Maintain volume and shape at ordinary temperatures and pressures o Liquid: Volume remains constant but shape changes to fit contain o

Gas: Neither constant volume nor fixed shape. Can be compressed or expanded to fill any size container.

Three states of matter o Water: Only substance that exists in all three states of matter at temperatures compatible with life. o Solid: ice

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Liquid: water Gas: water vapor

Matter combined o Mixture: Contains atoms of two or more elements physically intermixed. Chemical nature of individual atoms unchanged o Types of mixtures o Suspension: Contains two or more components with large, unevenly distributed particles. Particles will settle out (separate) when left undisturbed.

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Colloids: Two or more components with small, evenly distributed particles. Particles will not settle out (not separate).

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Solutions: Two or more components with small, evenly distributed particles. One component dissolved in the other. Will not settle out. o Solute: Dissolved substance. (sugar in tea) o Solvent: Substance that dissolves solute. (wax in gasoline)

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Chemical bonds o Defined: An attractive force between atoms. Not a physical structure. o Molecule: Bonding between two or more atoms of same element. o Macromolecule: Composed of many atoms o Compound: Bonding between two or more atoms of different elements. Molecular formulas include o Molecular (element) symbol = which atoms o Number of atoms of each element Valance electrons (outermost) of atoms: Determine which atoms interact to form bonds.

https://www.scienceabc.com/pure-sciences/how-to-find-the-number-of-valence-electrons-in-an-element.html

Two rules of bond formation o Duet rule: atoms with 5 or fewer electrons o Atom most stable when it’s valence electron shell holds 2 electrons.

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Octet rule: Atom most stable when it has 8 electrons in its valence shell (as in CO 2)

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Ions o o o

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Defined: Formed when atoms loose or gain one or more electrons. Losing an electron creates a cation – Fewer electrons (negative) than protons (positive) Cation representation – Ions represented using a symbol followed with a superscript+ o Number of missing electrons indicated by charge number (+1, +2, +3, +4)  +NA for example.

Gaining an electron creates an anion – More electrons (negative) than protons (positive) Anion representation – Ions represented using a symbol followed with a superscript – Number of electrons gained indicated by charge number (-1, -2, -3, -4) o

CA- for example.

Ionic bonds o Involved electron transfer. Electrical attraction occurs between and anion o Example: Sodium chloride  Salt

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Salts are inorganic compounds: Inorganic compound of any cation (except hydrogen) and any anion (except hydroxide) o Example: NaCl, NaHCO [baking soda] 3

Covalent bonds o Electrons shared between atoms (result is a molecule)

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Number of covalent bonds o Single covalent bond: Share 1 pair of electrons (1 each) o Double covalent bond: Share 2 pairs of electrons (2 each) o Triple covalent bond: Share 3 pairs of electrons (3 each) o Quadruple covalent bond: Involves 4 electrons Types of covalent bonds o

Nonpolar covalent bond: Equal sharing of electrons between atoms. No electrical charge on final molecule.

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Polar covalent bond: Unequal sharing of electrons (more time around 1 atom). One atom becomes partially negative (8-). One atom becomes partially positive (6+) o Polar covalent bonds form a dipole: Polar molecule with partially positive and partially negative ends.

Hydrogen bonds o Small appositive charges on hydrogen atoms of a polar molecule attracted to negative charges on atoms in other polar molecules.

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Hydrogen bonds affect molecules – Can change shape of molecules or pull molecules together.

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Hydrogen bonds affect properties of water o Surface tension: Water molecules on surface interact more with each other creating a surface tension.

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Hydrogen bonds slow rate of evaporation – Water becomes vapor (gas) when all hydrogen bonds broken (boiling water) o Frozen water (ice): Crystal lattice is formed. Bonds hold water molecules in place. Accounts for expansion of water when freezing.

Chemical reactions and chemical notation o When do chemical reactions occur? Every time a chemical bond is formed, broken, or rearranged. When electrons are transferred between two or more atoms (or molecules) Chemical notation: series of symbols and abbreviations. Show what happens in chemical reactions. Chemical equation (uses chemical notation) has two parts o Reactants: Left side of the equation. What will undergo reaction. o Products: Right side of the equation. Show results of chemical reaction. Direction of chemical reactions o Reversible reactions: Can proceed in either direction. Indicated by two arrows pointing in opposite directions.  CO2+H2O -> 2 MgO Chemical reactions should be balanced o Number of atoms must be same on both sides of the equation.  Balanced equation 2H +O →2H O 2 2 2 o Unbalanced equation H +O →H O 2 2 2 o

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Energy and chemical reactions o Energy: Capacity to do work. Put matter into motion. Fuel chemical reactions. o Two general forms of energy o Potential energy: Stored energy can be released to do work at some later time. o Kinetic energy: Potential energy that has been released or set in motion to perform work.

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Energy found in 3 forms in human body o All three forms may be potential or kinetic depending on location or process.  Chemical energy: Found in bonds between atoms. Drives nearly all chemical processes.  Electrical energy: Generated by movement of charged particles or ions.  Mechanical energy: Energy transferred from one object to another (motion that does the work) Chemical reactions may require or release energy o Endergonic reactions: Require input of energy (converting water to water vapor). Products contain more energy than reactants.

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Exergonic reactions: Energy is released. Products have less energy than reactants.

Homeostasis and types of chemical reactions o Anabolic reactions: Individual subunits and united by chemical bonds yielding larger units. o General notion – A+B -> AB Endergonic

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Catabolic reactions: Large substance is broken down into smaller substances. Bonds are broken.

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General Notation – AB -> A+B Usually exergonic

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Exchange reactions: Atoms in reactants are exchanged with one another. o General Notation - AB + CD  AD + BC / HCL + NaOH -> H2O + NaCL

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Oxidation-reduction reactions (redox reactions): Type of exchange reaction. Electrons and energy are exchanged instead of atoms. Usually exergonic (can released large amounts of energy). o Oxidation: Reactant that loses electrons is oxidized. o Reduction: Reactant that gains electrons is reduced.

Reaction rates and enzymes o Activation energy (Ea): Energy required for a reaction to take place. Atoms must collide with enough energy to overcome the repulsive force of their electrons.

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Reducing activation energy needed for chemical reactions increases the reaction rate

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Ways to reduce activation energy: o Reactant concentration: Affects number of collisions (interactions) between reactants. Greater concentration more collisions. Lower concentration fewer collisions. o Temperature: Affects movement (kinetic energy) of molecules. Increasing temperature get more forceful and effective collisions between reactants. Particle properties influence reaction rates o Smaller particles - Move faster with more energy than larger particles o Particle phase: Gaseous phase particles have highest kinetic energy. Liquid phase has higher kinetic energy than solid phase. Catalyst: Lower activation energy. Increases reaction rate. Does not alter reactants or products. Not consumed or changed in reaction. Enzymes: Biological catalysts. Most are proteins. Highly specific for individual substrates Substance that can bind to the enzyme’s active site.

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Enzymes facilitate chemical reactions o Active site: Region of an enzyme where substances must bind. Site shape determined by structure or enzyme. Provides specificity. o Substrates: Reactants in enzymatic reactions. Form products.

Enzyme substrate interaction  Enzyme-substance complex forms: Substance binds to active site on enzyme.  Substrate binding results in a temporary, reversible change in shape of enzyme.  Product detached from active site.  Enzyme is able to repeat the process.

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Saturation can occur – Too much reactant, enzyme cannot process the amount.

INORGANIC COMPOUNDS: WATER , ACIDS, BASES, AND SALTS BONDS  Biochemistry: The chemistry of life  Inorganic compounds: Generally, do not contain carbon bonded to hydrogen. o Examples: H O, acids, bases, salts, CO, CO 2 2  Organic compounds: contain carbon bonded to hydrogen. o Examples: sugars, fats  Water makes up 60-80% of human body mass o Water has several vital key properties  High heat capacity: Can absorb heat significantly changing temperature itself.  Evaporates readily from liquid to gas – Carries significant heat with it when it…  Cushions: Cerebral Spinal Fluid helps protect the brain (“floats”)  Lubricates (reduces friction): Between two adjacent surfaces (joints)  Universal solvent: Many solutes will dissolve in it entirely or to some degree  Forms hydrogen spheres: Water molecules surround ions or polar molecule in solution. o Hydrophilic (water loving) molecules: Have positive or negative charge (pole). Interact readily with water (dissolve). Like interacts with like (ionic and polar molecules). Hydration sphere around molecule.

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 Examples: Salt (NaCl), glucose Hydrophobic (water fearing) molecules: Lack of positive or negative charge (pole). Uncharged nonpolar covalent molecules. No Hydrogen sphere around molecule.

Acids and bases  Acid: Solute that dissociates and release hydrogen ions. o A proton donor  After losing an electron, a hydrogen ion consists solely of a proton. o More acid in solution means more hydrogen ion (H+) in solution. o Strong acids dissociate completely split from other ions.

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Base (alkali): solute that removes hydrogen ions from solution. Proton acceptor [H+ in solution decrease]. May release a hydroxide ion. o Strong bases dissociate completely  Example: Sodium hydroxide (NAOH) NaOH -> Na+ OH-

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Weak acids and weak bases do not dissociate completely – Have less of an impact on pH than strong acids and bases. o Example: carbonic acid (H2CO3) H2CO3 H+ +CO3

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Salts are inorganic compounds: A cation [+] (not hydrogen) and an anion [-] (not hydroxide). Held together by iconic bonds. Many dissociate completely in water, releasing cations and anions

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Buffers: Stabilize the pH of a solution by removing or replacing hydrogen ions. Helps maintain normal pH of body fluids (most body fluids slightly basic) o Buffer systems: Help maintain pH within normal limits.  Examples: Carbonic acid (H CO ), Sodium bicarbonate (NaHCO ) 2 3 3 pH scale: Measure of H+ concentration. Ranges from 0-14. o Note: each unit represents a 10 filed change not a linear change.

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ORGANIC COMPOUNDS: CARBOHYDRATES, LIPIDS, PROTEINS, AND NUCLEOTIDES o Monomers and polymers o o

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Monomer: Single subunit (glucose) Polymer: Group of subunits linked together. Formed by dehydration (water releasing) reaction (anabolic reaction)

Hydrolysis reactions: Breakdown of polymer to monomers (Catabolic reaction). Uses water molecule.

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Carbohydrates o

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Lipids

Contain carbon, hydrogen, and oxygen in ration near 1:2:1 Important as energy sources (sugars and starches)  Types of carbohydrates  Monosaccharide: single sugar (simple sugar). Contains three to seven carbon atoms  Triose: three carbon atoms  Tetrode: four carbon atoms  Pentose: five carbon atoms  Hexose: six carbon atoms  Heptodes: seven carbon atoms  Glucose is a monosaccharide - Hexose, most important  Disaccharide: Two monosaccharides joined. Very soluble in water.  Polysaccharides (complex carbohydrates): Formed from multiple disaccharides and/or monosaccharides Polysaccharides formed from glucose  Starches (plants): Broken down into monosaccharides by digestive system. Major dietary energy source (potatoes and grains)  Glycogen (animals): Muscle and liver cells make up glycogen.  Polysaccharides can bind to proteins and lipids  Glycoproteins: Class of proteins that have carbohydrate groups attached to the polypeptide chain  Glycolipids: Structure

15 o Contain carbon hydrogen, and oxygen. Carbon to hydrogen ratio is near 1:2. May contain small quantities of phosphorus, nitrogen, or sulfur. Most are insoluble in water. Special transport mechanisms for them in the blood.  Examples: fats, oils, waxes o Fatty acids are long carbon chains with attached hydrogen atoms  Head: Carboxyl group (-COOH). Hydrophilic  Tail: Carbon and hydrogen. Hydrophobic. The longer the tail, the lower the solubility.

 Fatty acids can be saturated or unsaturated o Saturated fatty acid (carbon with 4 single bonds): Each carbon in tail attached to two hydrogens and two carbons. Solid at room temperature.

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Unsaturated fatty acid (carbon with less than 4 single bonds):  1 double bond = monounsaturated (maybe liquid)  >1 double bond = polyunsaturated (liquid)  Double bonds change shape of tail.  Changes the way the body mobilize the acid.

Glycerides (glycerol molecule with attached fatty acid chains) o Neutral fats  Monoglyceride: one fatty acid attached.  Diglyceride: two fatty acids attached.  Triglyceride: three fatty acids attached.

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Cholesterol: Functions to maintain plasma membranes. Needed for cell growth and division. Structural base for all steroids. Sex hormones (estrogen and testosterone): Regulation of sexual and other metabolic functions.

Proteins o Proteins are macromolecules o Enzymes are proteins o

Play structural roles [e.g., membranes]

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Involved in movement

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Involved in the body’s defenses

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Can be used as fuel

o Proteins are formed from amino acids o 20 amino acids [monomeric unit of proteins] o Molecule has both + and – chargers, but net charge of zero (0). R group gives amino its reactive properties

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Covalent bond (peptide bond) joins amino acids forming peptides – connects carboxylic acid group of one amino acid to amino group of another o Dipeptide: Two amino acids linked together o Polypeptide: Three or more amino acids linked together

17 o Proteins: more than 100 amino acids; one or more peptide chains o Proteins may be classified as fibrous or globular o Fibrous proteins  Long rope-like strands  Mostly nonpolar amino acids  Link things together and add strength and durability to structures

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Globular proteins o Spherical or globe-like o Mostly of polar amino acids o Function as enzymes, hormones, and other cell messengers

 Protein structure  Primary structure: sequence of amino acids

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Secondary structure: Bonds form between atoms at different parts of polypeptide chain o Example: hydrogen bonds o Secondary protein structure results in specific shapes – simple spiral (alpha helix) and Flat pleated sheet (beta sheet)

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Tertiary structure: coiling and folding gives protein a final 3D shape. Interactions of proteins and surrounding water molecules. Interactions between R groups of amino acids in protein.

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Quaternary structure: interaction between multiple polypeptide chains (More than one chain bonding)

Proteins may denature o Change in protein tertiary or quaternary structure o Protein shape changes and function deteriorates o Occurs under extreme conditions o Body temperature above 43 C or 110 F o Fatal due to denaturation of structural proteins and enzymes

Nucleotides and nucleic acids o Nucleotides are the monomers of nucleic acids o

Nucleotide components

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Nitrogenous base [hydrocarbon ring structure]

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Pentose (five-carbon) sugar

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deoxyribose - DNA

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ribose – RNA and ATP

Phosphate group

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Deoxyribonucleic acid (DNA) is one of two classes of nucleic acids  Location: In nuclei of eukaryote cells o Components: Sugar is deoxidize (pentose ribose minus oxygen atom)  Phosphate bound to sugar

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DNA bases: adenine, guanine, cytosine, and thymine DNA structure: Two nucleotide chains twist around each other forming a double helix. Sugar and phosphate group form “backbone” Base faces center of helix.

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DNA structure (A-T, G-C) o Purine A [adenine] bind pyrimidine T [thymine] o

Purine G [guanine] binds pyrimidine C [cytosine]

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Strands run in opposite directions [3’ and 5’]

DNA function o Pieces of DNA form chromosomes (23 and me) o Gene is region of chromosomal DNA o Genes provide code for protein syn...