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Chapter 2 Matter Any substance that has mass and takes up space Composed of 1 or more elements Found in 1 of 3 states Gas – no definite shape or volume Liquid – shape conforms to container Solid – definite shape Human Elements Substances that can’t be broken down by ordinary chemical processes 92 occur naturally on Earth Oxygen Carbon Hydrogen Nitrogen Calcium Phosphorus Potassium Sulfur 61.0% 23.0 10.0 2.6 1.4 1.1 0.2 0.2 About 25 are essential to life 96% of human body, as well as other living organisms, from 4 elements (CHON) Composed of atoms 2+ of same atoms is a molecule (O2) 2+ of different atoms is a compound (NaCl, H2O) Metal + poisonous gas = edible product Gas + gas = liquid 96% Atoms Smallest particles that retain the properties of an element Made up of subatomic particles: Protons (+) in nucleus Electrons (-) orbits nucleus Neutrons (no charge) in nucleus Protons and neutrons Mass of about 1 dalton (atomic mass unit, amu) Electrons Mass is negligible (1/2000 amu) Reading A Periodic Table Elements differ depending on the number of subatomic particles Atomic symbol 1st letter or 2 (usually) Atomic number Determined by number of protons Neutral atoms contain equal # of electrons Element specific Mass number Determined by number of protons + neutrons Atomic weight: average of relative weights of all isotopes, versions of the element The Periodic Table Most Common Elements in Living Organisms Chemical Bonds Energy relationship between electrons Energy is the capacity to do work, cause change, or move matter Kinetic: energy of motion Potential: energy due to location or structure; capability Electrons repel one another (magnets), but attracted to protons (opposites attract) Determine chemical properties, reactivity, of atoms Orbitals are key Closer to the nucleus = lower energy and filled first Octet rule: hold up to 8 electrons (not 1st = 2) Outermost is valence shell Chemical reactions are making and breaking bonds Electron Shell Models Inert versus reactive elements Valence electrons Ionic Bonds One atom loses electron(s) to become _____ charge? Another atom gains these electrons to become _____ charge? Charge difference attracts the two ions to each other Clinically called electrolytes Very weak bond E.g. Salts cation anion Covalent Bonds Atoms share a pair(s) of electrons to fill valence shell Form single, double, or triple covalent bond, based on number of electrons shared Strong bonds What’s Mine is Yours or Just Mine Nonpolar molecules Electrons shared equally Example: carbon dioxide(CO2) Polar Molecules Electrons spend more time near the nucleus with the most protons (electronegativity) Example: water (H20) Hydrogen Bonds Special case of covalent bonds Attractions between molecules Polar hydrogen (slightly ‘+’ charge) attracted to another polar molecule (slightly ‘-’ charge) Hold large molecules in a specific 3D shape Chemical Reactions Chemical equation: reactant(s) + reactant(s) = product(s) May be reversible Move to equilibrium Types Synthesis: (A + B AB) usually anabolic and endergonic Decomposition: (AB A + B) usually catabolic and exergonic Exchange: (AB + CD AD + BC) may or may not be endergonic/exergonic Redox: may gain or lose electrons Oxidized – loses electrons (LEO) Reduced – gains electrons (GER) Affected by temperature, concentration, catalysts, etc. Biologically Important Compounds Inorganics lack carbon (generally) E.g. salts, water, acids, and bases Organics contain carbon, are covalently bonded, and generally large E.g. carbohydrates, lipids, proteins, and nucleic acids Properties of Water Temperature-stabilizing effects Absorbs and releases large amounts of heat w/o significant change Sun, wind, muscle activity High heat of evaporation Absorption of heat breaks bonds; liquid gas Sweating Polar molecule Hydrophobic and hydrophilic Solvent/solute relationship Blood circulation, waste disposal, and cleaning Reactivity Hydrolysis and dehydration reactions Electrolytes Ions able to conduct electrical current Kidneys regulate Salts contain ions other than H+ or OH- E.g. NaCl, CaCO3, and calcium phosphates Acids are hydrogen ion (H+) donors Concentration determines acidity of a solution E.g. pH < 7; HCl, H2CO3 Bases are hydrogen acceptors Form water upon disassociation E.g. pH > 7; Mg(OH) 2, HCO3-, and NH3 Buffers release H + with increasing pH and accept H + when decreasing H2CO3 HCO3- + H+ Building Organic Molecules Monomers: small repeating units Universal, similar in all forms of life Polymers: chains of monomers, functional components of cells (macromolecules) DNA is composed of 4 monomers (nucleotides) Variation based on arrangement Proteins from 20 different amino acids (AA’s) Variation distinguishes within and between species Making and Breaking Polymers Dehydration reaction Hydrolysis reaction Links monomers Breaks polymers Loss of water for each Addition of water for each monomer added Forms a covalent bond 1 2 broken bond 4 3 1 1 2 3 2 3 4 4 1 2 3 4 Carbohydrates General (CH2O)n ratio, end in ‘ose’ Fuel source for cells Glycosidic bonds Dehydration vs. hydrolysis Monosaccharides (simple) Pentoses: ribose and deoxyribose Hexoses: glucose, fructose, & lactose (energy production) Disaccharides (simple) Maltose, lactose (glu + gal), & sucrose (glu + fru) Polysaccharides (complex) Glycogen (animal storage) Starch (plant storage) Lipids Composed of fatty acids (long carbon chains) and a glycerol (3 carbons) Triglycerides 3 FA’s Most usable form of energy Fats (animal) and oils (plants) Saturated or unsaturated (mono- or poly-) Phospholipids 2 FA’s and a phosphate group Amphipathic molecule Cell membranes Steroids Hydrocarbon rings Cholesterol and sex hormones Eicosanoids In all cell membranes Prostaglandins role in blood clotting, inflammation, and labor contractions Proteins Chains of amino acids joined by peptide bonds 20 different types (alphabet) Peptide, polypeptides, and proteins (words) are all slightly different Structural levels Primary (1°) – sequence of amino acids Secondary (2°) – primary level folds to form alpha (α) – helixes and beta (β) - pleated sheets Tertiary (3°) – folding of secondary structures on each other Quaternary (4°) – 2+ polypeptides interact to form a protein Denaturation destroys structure which alters or inhibits function Changes in pH and temperature Reversible or permanent depending on extend of change (fevers) Protein Types Fibrous (structural proteins) Building materials of the body Keratin, elastin, and collagen Movement Globular (functional proteins) Enzymes Transport Immunity Enzymes Globular proteins acting as catalysts to speed a reaction Lower energy of activation (EA) End in ‘ase’ and named for substrate Mechanism of enzyme action: Enzyme binds substrate at its active site on the enzyme. Enzyme-substrate complex undergoes an internal rearrangement that forms a product. Product released and now catalyzes another reaction Nucleic Acids DNA and RNA Composed of nucleotides with 3 components Pentose sugar Phosphate group (PO4) Nitrogenous base form complementray pairs Purines (2 rings): adenine (A) and guanine (G) Pyrimidines (1 ring): thymine (T), cytosine (C), and uracil (U) How DNA and RNA Differ DNA (deoxyribonucleic acid) RNA (ribonucleic acid) Directs protein synthesis; Carries out protein synthesis replicates self; genetic material Sugar is deoxyribose Sugar is ribose Has -OH Bases are A, C, G, and U Has –H Bases are A,C, G, and T Single-stranded Double-stranded helix Not confined to nucleus Only in nucleus 3 major types 1 type Adenosine Triphosphate (ATP) RNA nucleotide with 3 phosphate groups Stores energy from break down of glucose Transfers phosphate groups to release energy, phosphorylation Controls energy release