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Chemistry of Life• All life processes involve chemical
reactions (clinical day)– Ex. Ca++ in muscle contraction
Na+, K+ in nerve impulses
• anything that has mass and takes up space
• can you think of solids, liquids, and gases that might be found in the body?
Matter
• the capacity to do work • No mass & does not
take up space• Types:
– Potential: stored in bonds (chemical)
– Kinetic: doing work (electrical, mechanical, radiant)
• Exergonic & endergonic reactions
Energy
Composition of Matter• 92 naturally occurring elements (112 known, 113-
118 are alleged)• Living organisms require about 26 of these elements
(table 2.1, p.28)• About 96% (by mass) comes from Oxygen (O),
Carbon (C), Hydrogen (H), and Nitrogen (N)
Atoms• smallest complete unit of an element
– Composed of dozens of subatomic particles, but we are only concerned with THREE!
Subatomic Particle
Charge Location What it tells you!
Proton + Nucleus Identity of atom, mass
Neutron 0 Nucleus Isotope, mass
Electron - Surrounds nucleus
Properties of atom, negligible mass
Identifying Elements
• Atomic number
• equal to the number of protons in an atom (& electrons in neutral atom!)
• Atomic mass
• sum of the masses of all the protons & neutrons contained in nucleus
Isotopes• atoms of same element with a different mass
(due to neutrons) but same chemical properties• Ex. C-12 and C-14
• Radioactive isotopes used in many medical tests to tag biological molecules to be followed or traced • i.e. PET scans, I-131 for thyroid activity,
destroying localized cancers (Ra, Co, etc.)
Ions• Charged particles• Form ionic bonds• Cations (+) • Anions (-)
Find the Face (in the Beans)
• Transfer or share electrons in order to fill their valence shell (stability)• All atoms want 8 e- in their valence shell
(except H & He)
• Ionic bond – transfer electrons• Covalent bond – share electrons
– Nonpolar: shares electrons equally
– Polar: shares electron unequally
Chemical Bonds
Chemical Bonding
• Due to electronegativity– How much an atom in
a bond pulls electrons to itself
– Ionic: >1.7– Polar covalent: 0.4-1.7– Covalent: <0.4
Hydrogen Bonding
• Weak bonds
• attraction of H to partial negative charge– Example: polar
covalent bonds between oxygen and hydrogen
• represents the numbers and types of atoms in a molecule– Ex. H2O , C6H12O6
Molecular Formula
Chemical Reactions• Metabolism= sum of all chemical reactions in the
body– Synthesis (anabolism) A + B AB
• Energy absorbing• i.e. growth, repair, protein synthesis
– Decomposition (catabolism) AB A + B• Energy releasing• i.e. digestion of foods, breakdown of glycogen in liver to
produce glucose
– Single replacement AB + C AC + B
– Double replacement AB + CD AD + CB
Rate of Chemical Reactions• Temperature ( temp increases collisions)• Concentration of reactants ( number = faster,
more collisions)• Particle size (smaller = faster, more collisions)• Presence of catalysts
– Affect rate of reaction without being changed by reaction
– Biological catalysts: enzymes (proteins)– Shape matters! Like a puzzle piece
Biochemistry• Inorganic compounds: lack carbon (with few
exceptions)– Small, simple molecules– Water, salts, many acids
& bases
• Organic compounds: carbon-containing compounds– Large covalently bonded molecules– Carbohydrates, lipids, proteins, nucleic acids
Inorganic Compounds• Water
– High heat capacity• Absorbs & releases large amounts of heat• Prevents sudden changes in body temperature (homeostasis!)
– Polarity/solvent properties• “universal solvent”• Chemical reactions depend on solvent• Transport/exchange medium • Lubrication (synovial fluid in joints)
– Chemical reactivity (hydrolysis reactions)– Cushioning
• Protective (CSF, amniotic fluid)
• Oxygen– used to release energy
from glucose
• Carbon dioxide– waste of metabolic
processes
Inorganic Compounds
• Salts– Ionic compound containing cations other than H+
and anions other than OH-– Vital to body functions
• K+ & Na+ essential for nerve function, Fe2+ is essential for hemoglobin, Cl-, Ca++, Mg++, PO4-, CO3-, etc.
– All salts are electrolytes (substances that conduct an electrical current in solvent)
• Release ions when dissolved in water
– Functions in Table 2.1, page 28
Inorganic Compounds
Inorganic Compounds• Acids & Bases
– Electrolytes
• Acids– Release H+ ions in solution– “proton donors”– HCL, acetic acid, carbonic acid
• Bases– Release OH- ions in solution– “proton acceptors”– HCO3- (important base in blood)
• pH scale measures hydrogen ion concentration– pH 7 = neutral– pH >7= basics (more OH-
than H+)– pH <7= acidic (more H+
than OH-)• Normal blood pH for humans is
7.35 to 7.45– If > , alkalosis– If < , acidosis
• Buffers- maintain pH
Organic Compounds
• sugars, starches, glycogen, cellulose– 2-3% body weight– Plants- starches and cellulose (cannot digest)– Animals- source of energy- stored as
glycogen
Carbohydrates
• Monosaccharides: 3 to 7 carbons– Ex. Glucose, fructose, galactose
• Carbohydrateutilized by the
cell
Many C6H12O6
Carbohydrates
• Disaccharides: 2 monosaccharides combine by dehydration synthesis (condensation)– Ex. Sucrose
• Broken apart by hydrolysis (add water)
Carbohydrates
• Polysacchride: 10-100s of monos– Ex. starch
Carbohydrates
• 18-25% in lean adults– Contain C, H, O - neutral– Fats- concentrated energy stored in adipose
tissue
Lipids
• Triglycerides: Glycerol + 3 fatty acids• Monounsaturated- one double bond• Polyunsaturated- more than one double bond
• Saturated- no double bonds
Lipids
• Phospholipids- polar head and 2 non-polar tails (membrane)
Lipids
• Steroids- cholesterol, sex hormones, cortisol, etc.
Lipids
• 12-18% in lean adults– Structural and physiological enzymes– Made of amino acids (20)- held by peptide
bonds– 3D shape held by H-bonds (denatured with
heat)
Proteins
– Base + sugar + phosphate
– DNA and RNA– ATP- provides energy
for the cell
Nucleic Acids
• molecules with the same chemical formula and with the same kinds of bonds between atoms, but in which the atoms are arranged differently.
• share similar if not identical properties in most chemical contexts.
Isomers