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Module 2

  1. chemistry
    the science of the structure and interactions of matter
  2. matter
    anything that occupies space and has mass
  3. mass
    the amount of matter in any object, which does not change
  4. weight
    the force of gravity acting on matter
  5. matter exists in three states:
    solids, liquids, and gas
  6. element
    • a substance which cannot be split into simpler substances by chemical means
    • the building blocks of chemical compounds
    • Scientists now recognize 117; of those 92 occur naturally on Earth
  7. How many different chemical elements normally are present in the human body?
    26; 4 are major, 8 are lesser, and 14 are trace
  8. The major elements
    4 elements which constitute about 96% of the body's mass: Oxygen, Carbon, Hydrogen, and Nitrogen
  9. The lesser elements
    8 elements which constitute 3.8% of the body's mass: Calcium, Phosphorus, Potassium, Sulfur, Sodium, Chlorine, Magnesium, and Iron
  10. The trace elements
    • 14 elements that account for 0.2% of body mass: Aluminum, boron, chromium, cobalt, copper, fluorine, iodine, manganese, molybdenum, selenium, silicon, tin, vanadium, and zinc
    • *All Chrome Colts Born First In Silina Manage Most Sells Via Zest
  11. atom
    the smallest units of matter that retain the properties and characteristics of the element
  12. Oxygen
    • One of the major elements
    • Symbol: O
    • Part of water and many organic molecules
  13. Carbon
    • One of the major elements
    • Symbol: C
    • Forms backbone chains and rings of all organic molecules; carbohydrates, lipids (fats) proteins, nucleic acids
  14. Hydrogen
    • One of the major elements
    • Symbol: H
    • Constituent of water and most organic molecules; ionized form ( H+ ) makes body fluids more acidic

  15. Nitrogen
    • One of the major elements
    • Symbol: N
    • Component of all proteins and nucleic acids
  16. Calcium
    • One of the lessor elements
    • Symbol: Ca
    • Contributes to hardness of bones and teeth; ionized form ( Ca++ ) needed for blood clotting, release of some hormones, and contraction of muscle
  17. Phosphorus
    • P ; one of the lessor elements
    • Component of nucleic acids and ATP (molecule used to store chemical energy); required for normal bone and tooth structure
  18. Potassium
    • K ; one of lessor elements
    • ionized form ( K+ ) is the most plentiful cation (positively-charged particle) inside cells; needed for nerve activity
  19. Sulfur
    • S ; one of lessor elements
    • component of some vitamins and many proteins
  20. Sodium
    • Na ; one of lessor elements
    • Ionized form ( Na+ ) is the most plentiful cation in extracellular fluid; essential for maintaining water balance, needed for nerve activity
  21. Chlorine
    • Cl ; one of lessor elements
    • Ionized form ( Cl- ) is the most plentiful anion (negatively-charged particle) in extracellular fluid; essential for maintaining water balance
  22. Magnesium
    • Mg ; one of lessor elements
    • Ionized form ( Mg++ ) needed for many enzymes (molecules that increase the rate of chemical reactions in organisms)
  23. Iron
    • Fe ; one of lessor elements
    • Ionized forms (Fe++ and Fe+++ ) are part of hemoglobin (oxygen carrying protein in red blood cells) and some enzymes
  24. nucleus
    • the dense central core of an atom
    • within are protons and neutrons
  25. protons
    • positively charged subatomic particle
    • large particles, have mass
  26. neutrons
    • subatomic particle within the nucleus which has no charge
    • large particles, have mass
  27. electrons
    • tiniest, negatively charged subatomic particles which move about in a large space surrounding the nucleus. They do not follow a fixed path or orbit, but instead form a negatively charged "cloud" that envelops the nucleus
    • # of electrons determines the chemical properties of an element
  28. electron shells
    • certain regions around the nucleus which are depicted as simple circles 
    • Each shell holds a specific # of electrons. The first shell can hold 2, second holds 8, third holds 8, the fourth holds 18
  29. The number of electrons in an atom always equals the number of:
    protons *IF IT IS STABLE & HAS NO CHARGE
  30. atomic number
    • the number of protons in the nucleus of an atom
    • gives us the name of th element
  31. mass number
    the sum of an atoms protons and neutrons
  32. isotopes
    • atoms of an element that have different numbers of neutrons and therefore different mass numbers
    • most are stable, meaning their nuclear structure doesn't change over time
  33. the number of _________ of an atom determines it's chemical properties
    electrons
  34. radioactive isotopes
    • isotopes which are unstable, their nuclei decay (spontaneously change) into a stable configuration
    • as they decay, these atoms emit radiation - either subatomic particles or packet of energy - and in the process often transform into a different element
  35. half-life of an isotope
    the time required for HALF of the radioactive atoms in a sample of that isotope to decay into a more stable form
  36. molecule
    • two or more atoms sharing electrons joind by a chemical bond
    • can be two atoms of the same element or two atoms of different elements
  37. compound
    a molecule that contains atoms of different elements
  38. solids
    • substances in which the particles are tightly associated with each other.
    • have definite shape and definite volume and are not very compressible
  39. liquids
    • contain particles that are interacting with one another, but much more weakly
    • changes shape to match container, so indefinite shape
    • volume is definite, not very compressible
  40. gases
    • widely spaced particles flying around
    • indefinite shape and volume but very compressible
  41. chemical reaction
    occurs when new bonds form or old bonds break btwen atoms
  42. reactants
    the starting substances in a chemical reaction
  43. products
    the ending substances in a chemical reaction
  44. the total mass of the reactants equals:
    the total mass of the products; thus, the number of atoms of each elements is the same before and after the reaction
  45. metabolism
    refers to all the chemical reactions occuring in the body
  46. energy
    • the capacity to do work
    • two forms: potential energy and kinetic energy
  47. potential energy
    • energy stored by matter due to it's position
    • Ex: a coiled spring
  48. kinetic energy
    energy associated with matter in motion
  49. chemical energy
    a form of potential energy that is stored in the bonds of compounds and molecules; the amount of energy at the beginning and end of chemical reaction is the same
  50. law of conservation of energy
    although energy can be neither created nor destroyed, it may be converted from one form to another
  51. the overall reaction in chemical bonds & reactions may:
    either release energy or absorb energy
  52. exergonic reactions
    • (ex- = out)
    • reactions which release more energy than they absorb
  53. endergonic reactions
    • (end- = within)
    • reactions which store energy; hidden within the chemical bonds
    • In biology, they are called anabolic reactions
  54. activation energy of a reaction
    • the collision energy needed to break the chemical bonds of the reactants
    • the initial energy "investment" is needed to start a reaction
  55. anabolic reactrions
    same as endergonic reactions; the energy is hidden within the chemical bonds. reaction that needs energy
  56. exergonic
    • The reaction when chemical bonds are broken and energy is released and used for work
    • also called catabolic reactions
  57. In most chemical reactions that will be studied, the energy created is turned into:
    work; meaning, for example, the breaking of chemical bonds allows muscles to contract
  58. In the biological system, the chemical bond energy that isn't turned into work is turned into:
    • heat
    • heat production is a side effect which has been harnessed by our body
  59. At the microscopic level, heat is:
    the motion of particles
  60. entropy
    • disorganization; law of disorder
    • to overcome this, energy must be added for order
    • Ex: with atrophy, things become more disorganized in our house. If we want to decrease entropy, we have to do work - clean - to increase organization
    • *This means that particles in a container will eventually distribute themselves all over the container
  61. diffusion
    if there are no barrier, substances always move from where they are at high concentration to where they are at low concentration
  62. Specific heat
    the amount of energy it takes to raise a gram of substance one degree of temperature
  63. Metals are better heat _________, which makes them have _____ specific heat
    conductor; lower
  64. Water has specific heat of
    1 cal/g°C
  65. Materials with a high specific heat, like water, are:
    hard to heat up, but once their hot, they stay hot for a while. Good heat insulators
  66. Materials with low specific heat, like metals,
    get hot fast, and cool off fast
  67. Calories
    • one measure of specific heat
    • the amount of energy it takes to raise one gram of water one degree Centigrade
    • official SI unit is a Joule (J)
  68. Food calories
    • often abbreviated Cal
    • one food calorie is equal to 1000 metric calories
    • 4.2 kJ
  69. Symbol for Aluminum
    Al
  70. Symbol for Boron
    B
  71. Symbol for Chromium
    Cr
  72. Symbol for Cobalt
    Co
  73. Symbol for Copper
    Cu
  74. Symbol for Fluorine
    F
  75. Symbol for Iodine
    I
  76. Symbol for Manganese
    Mn
  77. Symbol for Molybdenum
    Mo
  78. Symbol for Selenium
    Se
  79. Symbol for Silicon
    Si
  80. Symbol for Tin
    Sn
  81. Symbol for Vanadium
    V
  82. Symbol for Zinc
    Zn
  83. metalloids
    • elements that share metal and non-metal properties
    • only ones found in the human body are boron and silicon
  84. noble gases
    elements whose atoms that do not combine with other atoms
  85. nuclear fission
    process which causes atoms to break apart
  86. positrons
    • particles which are much like electrons in all respects except they carry a positive charge
    • used in positron-emission tomography (PET)
  87. ions
    • an atom that has a + or - charge because it has unequal numbers of protons and electrons
    • Atoms that have given away an electron become cations
    • Atoms that have taken on an electron become anions
  88. ionization
    • the process of giving up or gaining electrons
    • *an ion of an atom is symbolized by writing it's chemical symbol followed by the # of it's + or - charge. Thus, Ca2+ stands for a calcium ion that has two positive charges because it has lost two electrons
  89. atomic mass
    also called atomic weight; average mass of all naturally-occuring isotopes of an element
  90. In periodic table, the vertical columns represent
    the # of electrons in outer shell
  91. In periodic table, the horizontal rows indicate
    the # of shells an element has
  92. On a block of the periodic table, the _______ is placed in the top center; the ________ is at the bottom
    atomic number; atomic mass
  93. When writing an element with it's atomic # and mass outside of the periodic table...
    • put the atomic number as a subscript and mass number as a superscript before the symbol
    • 16/8O= mass # of 16, Atomic # of 8
  94. 47Ca
    used in nuclear medicine to study bone
  95. 131I
    used in nuclear medicine to destroy thyroid tissue
  96. 133Xe
    Used in nuclear medicine for respiratory studies
  97. dalton
    • the standard unit for measuring the mass of atoms and their subatomic particles
    • also known as an atomic mass unit (amu)
  98. free radical
    • an elecrically charged atom or group of atoms with an unpaired electron in the outermost shell
    • having a free electron makes a free radical unstable, highly reactive and destructive to nearby molecules
  99. deuterium
    • an isotope of hydrogen, hydrogen-2 ( 2H)
    • sometimes represented as D
  100. types of radioactivity
    Alpha, beta, and gamma particles
  101. alpha particles
    • same as helium nuclei (which has an atomic #2 and atomic mass of 4) but in alpha the two electrons are missing
    • α 42He
  102. beta particles
    • electrons expelled at high energy from radioactive atoms
    • β 0-1e
  103. nuclear fission
    the process by with the nucleus of a radioisotope comes apart
  104. gamma particle
    • γ
    • also called gamma ray, or X-ray
    • a high-energy packet of light energy (photon)
    • Photons of lower energy are visible light
    • *comes from nuclei, no charge, no measurable mass
  105. photons
    • particles of light
    • also behave as waves
    • for this class, call them particle waves
  106. electromagnetic waves
    gamma radiation (photons) which are both waves and particles; have both electrical and magnetic properties
  107. How to measure a wave of electromagnetic waves
    *Wavelength (λ) is the distance btwn the peak of each wave crest
  108. The electromagnetic spectrum
    • Electromagnetic wave-particles (photons) vary in energy
    • *Lower energy = longer wavelength ex:radio waves, microwaves

    • *Visible spectrum - color corresponds to wavelength
    • * Higher energy = Shorter wavelength ex: ultraviolet, xrays, gamma rays
  109. Order of wavelength (energy) of electromagnetic spectrum
    From Low energy to high: radio waves, Infrared, Visible light, UVlight, xrays, gamma rays
  110. The discover of light is:
    • ROY G BIV
    • Order of colors: Red, Orange, Yellow, Green Blue, Indigo, Violet
    • before red is infrared, after violet is UV light
  111. cation
    • An-ion
    • An atom which has a deficit of electrons, resulting in an excess of protons, giving it a positive charge
  112. Anion
    • An-ion
    • An atom which has an excess number of electrons, giving it a negative charge
  113. transition metals
    the elements in the center of the periodic table, are not governed by simply filling shells
  114. electrolytes
    • ions which are found dissolved in the tissues and fluids of the body. In this state they can (and do) conduct electricity
    • 4 which are usually measured in human medicine are sodium, potassium, chloride, and bicarbonate
  115. Free radicals
    • atoms which have unpaired electron
    • extremely damaging to biological systems and are likely part of the cause of diseases such as cancer
  116. free radical scavengers
    • reducing the amount of free radicals in cells and preventing disease
    • lycopenes, omega-3 fatty acids, and vitamins C and E
  117. ionic bond
    the force of attraction that holds together ions with opposite charges
  118. ionic compounds
    • also called polyatomic ions
    • groups of elements that like to "travel together" as charged groups
    • When writing the chemical formula, the metal goes on the left and the non-metal goes on the right, just like in the periodic table
  119. H+
    • Hydrogen ion
    • Common Ion, is a cation
  120. Na+
    • Sodium ion
    • Common ion, is a cation
  121. K+
    • Potassium ion
    • Common ion, is a cation
  122. NH4+
    • Ammonium ion
    • is a common polyatomic ion
    • is cation
  123. Mg2+
    • Magnesium, common polyatomic ion
    • cation
  124. Ca2+
    • Calcium, common polyatomic ion
    • is cation
  125. Fe2+
    • Iron (II), common polyatomic ion
    • Called ferrous ion
    • important part of hemoglobin, which carries oxygen in human body. 
    • In hemoglobin, ferrous ion (Fe2+) binds oxygen (O2-): "2=2 and o matches o"
  126. Fe3+
    • Iron (III), common polyatomic ion
    • cation
    • is called ferric ion
  127. F-
    • Fluoride, common polyatomic ion
    • anion
  128. Cl-
    • Chloride ion, common polyatomic ion
    • anion
  129. I-
    • Iodide ion, common polyatomic ion
    • anion
  130. OH-
    • Hydroxide ion, common polyatomic ion
    • Also called hydroxyl
    • when found in molecules, such as alcohols and sugars, called a hydroxyl group
    • anion
  131. HCO3-
    • Bicarbonate ion, common polyatomic ion and electrolyte in blood
    • anion
  132. O2-
    • Oxide ion, common polyatomic ion
    • anion
  133. SO42-
    • Sulfate ion, common polyatomic ion
    • anion
  134. PO43-
    • Phosphate ion, common polyatomic ion
    • anion
  135. EXAMPLE: If the hydroxyl anion (OH-) gains an electron, but it's an unpaired electron, then the electron in the free radical is indicated in the chemical equation by:
    a dot (OH⋅)
  136. chemical bond
    the sharing of electrons which produces very stable interactions between atoms
  137. valence of the atom
    • another name for the outer shell of electrons
    • an atoms valence is related to it's position in the periodic table
  138. covalent bonds
    • result when atoms share electrons, nicely
    • covalent single bond usually indicated by a single line (H-H), same with double and triple bonds
  139. covalent single bond
    if two electrons are shared
  140. covalent double bond
    if FOUR electrons are shared
  141. covalent triple bond
    is a total of six electrons are shared
  142. polar covalent bond
    • water is an example
    • the sharing of electrons btwn two atoms is unequal - the nucleus of one atom attracts the shared electrons more strongly than the nucleus of the other
    • The resulting molecule has a partial negative charge near the atom that attracts electrons more stongely
  143. nonpolar covalent bond
    in a covalent bond, where two atoms share the electrons equally-one atom does not attract the shared electrons more strongly than the other
  144. electronegativity
    the power of an atom to attract electrons to itself
  145. Partial charges in polar covalent bond, are indicated by:
    a lowercase Greek delta with a minus or plus sign: δ- or δ+
  146. hydrogen bond
    • a polar covalent bond btwn hydrogen atoms and other atoms
    • results from the attraction of a partial negative and partial positive charge
    • common elements which form hydrogen bonds are oxygen, nitrogen and sulfur
  147. ionic bonds
    • When electrons are completely given away or taken, the atoms "hang out" together because their merely attracted by their opposite charges
    •  typically formed by elements on opposite extremes of periodic table.
    • easily disrupted by water
  148. dipoles
    in a bond, where one end of bond is more negative, the other is more positive
  149. polar
    name for a bond in which electrons are not shared equally
  150. Bond strengths from strongest to weakest & examples of each in the human body:
    • Covalent (triple, double, single): gasses, in blood, single strands of DNS,RNA
    • ionic: teeth and bones, dissolved ions in electrolytes
    • hydrogen: holding protein structure, holding two strands of DNA together
  151. Inorganic compounds
    • usually lack carbon and are structurally simple
    • have either ionic or covalent bonds
    • include water and many salts, acids, and bases
  152. organic compounds
    always contain carbon, usually contain hydrogen, and always have covalent bonds
  153. Examples of inorganic compounds that HAVE carbon
    carbon dioxide, bicarbonate ion, and carbonic acid
  154. important property of water
    • it's polarity- Each water molecule is a dipole~the uneven sharing of valence electrons.
    • *makes it an excellent solvent for other ionic or polar substances
    • *gives water molecules cohesion (the tendency to stick together
    • *allows water to resist temp changes
  155. solution
    • made up of two components:solute and solvent
    • the size of dissolved particles is very small (usually the size of an atom or molecule)
  156. solvent vs. solute in a solution
    • the solution is made up of 2 components, solute and solvent
    • the solvent is what's doing the dissolving, the solute is what's being dissolved
  157. hydrophilic
    • hydro- = water; -philic=loving
    • What a solute is called if it is charged or contain polar covalent bonds making it dissolve easily in water
  158. hydrophobic
    • -phobic = fearing
    • What a molecule is called if it contains mainly nonpolar covalent bonds which make it not very water soluble
  159. because water is a dipole, it means 2 things:
    • 1. Water forms hydrogen bonds
    • 2. Water dissolves stuff~ think of mickey mouse hat & example of dissolved NaCl where the different charges came toether
  160. hydrogen shell
    when water molecules dissolve a substance, such as salt (NaCl), the circle of H2O around the element being dissolve.
  161. surface tension
    • a property that results from the cohesion of molecules to each other
    • *water has strong surface tension (because of the strong hydrogen bonding), which makes rounded surfaces as the cohesion btwn water molecules tries to pull the water into a ball. This is what causes a "dome" rising over the rim of a clean glass
  162. surfactant
    • same thing as soap, has the ability to destroy surface tension by interfering with hydrogen bonds btwn water molecules
    • also used to trap grease and dirt, and make a hydration shell around them
  163. capillary action
    • results from the cohesive forces within liquid and the adhesive forces btwn the liquid and walls
    • if the adhesive force btwn the wall of the container and water is greater than the adhesive force btwn molecules, then water will "climb" the walls of the tube (the meniscus of a test tube)
  164. 100 grams of water = ? mL water
    100
  165. 1000 mL = ? L
    1 L
  166. Mole
    • convenient way of counting large numbers of small things
    • 1 mole = 6.02 x 1023 = Avogadro's number
    • One mole of something is equal to the same number of grams as the atomic mass
    • Ex: for oxygen, w an atomic mass of 15.999 grams, 15.999 grams of oxygen atoms is = to 6.02 x 1023 oxygen atoms
  167. colloids
    • like solutions, with dissolved particles, except the solid particles are large enough to scatter light  (larger than an atom or molecule) are translucent, particles are still small enough to remain dispersed, even over long periods of time
    • Ex: the water droplets in fog scatter light from a car's headlights
  168. emulsion
    • like a colloid but the particles are even bigger & all components ARE LIQUID
    • Ex: milk
  169. suspension
    • the particles in a solution are larger than an emulsion. Material may mix w the liquid or suspending medium but they will settle to the bottom of container given time. (muddy water)
    • blood is a suspension of cells in a colloid of proteins and a solution of salts in water
  170. acid
    • a substance which dissolves in water to form one or more hydrogen ions (H+) as a cation (positive ion) and one or more anions  (negative ion)
    • HCl → H+ + Cl-
    • Because H+ is a single proton w 1 positive charge, an acid is also referred to as a proton donor
    • When mixed with a base, will form a salt
  171. base
    • a substance that dissolves in water to form one or more OH- ions (negative ions) and one or more cations (positive ions)
    • KOH → K+ + OH- 
    • Because it removes H+ from a solution, also called proton acceptor.
    • When mixed with an acid, forms salt
  172. salt
    • a substance that dissolves in water and releases ions, neither of which is H+ or
    • OH-
    • In the body, salts are electrolytes that are important for carrying electrical currents
  173. pH scale
    • used to express a solution's acidity or alkalinity (base)
    • extends from 0 to 14, based on the concentration of H+ in moles per liter
    • Midpoint of pH scale is 7, where concentrations of H+ and OH- are
    • equal
  174. acidic solution
    • a solution that has more H+ than OH- and
    • has a pH below 7
  175. alkaline solution
    (basic) a solution that has more OH- than H- and has a pH above 7
  176. strong verses weak acids
    • A strong acid is one in which the hydrogen-containing compound almost completely 
    • dissolves in water. 
    • In a weak acid, some of the acid molecules hold onto H+. 
    • So the stronger the acid, the higher the proportion of H+ released
  177. strong versus weak bases
    • A strong base is one in which almost all the OH- containing molecules break apart when dissolved in water.
    • In a weak base, not all the available OH- ions are released
  178. pH balance of blood
    • perfect is at 7.40
    • 7.35 - 7.45; slightly more basic than pure water
  179. buffer
    • a function that can convert strong acids or bases into weak ones by removing or adding protons (H+)
    • Acts as a H+ and/or OH- "sponge" so that pH is kept relatively constant
  180. The most important buffer system in human biology
    • carbonic acid-bicarbonate buffer system
    • carbonic acid = H2CO3  
    • bicarbonate ion (weak base so they act as the sponge) = HCO3-
  181. concentration of a substance is expressed in shorthand as:
    [substance] - where the square brackets mean "concentration of"
  182. pH
    • "negative logarithm of the hydrogen ion concentration"
    • scale goes from 0 to 14
    • Ex: a [H+] of 10-3 mol/L is a pH of 3.
  183. acidic solutions
    • pH < 7.00,  btwn 0 and 6.99
    • More H+ than OH-
  184. neutral solutions
    • have [H+] = 10-7, so pH equals EXACTLY 7
    • the mid-point on scale between acids and bases, so it's neutral
  185. basic solutions
    • pH > 7.0,  btwn 7.01 and 14
    • more OH- than H+
  186. acids in the human body
    gastric juice, vaginal fluid, urine and saliva
  187. bases in the human body
    blood, semen, cerebrospinal fluid, pancreatic juice and bile
  188. In pH when H+ is abundant
    (acidic conditions) excess H+ is "sponged up" by HCO3- (bicarbonate ions) to form H2CO3 (carbonic acid) which can continue on to make H2O and CO2 (carbon dioxide)
  189. In pH, when H+ is scarce (alkaline conditions):
    making it alkaline conditions, H+ is released by H2CO3 (carbonic acid) to form HCO3- (bicarbonate ion) and H+ 
Author
jskunz
ID
235838
Card Set
Module 2
Description
Bio Core
Updated

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