Biology

Chapter 3

Composition of Matter: 

“knowing the number of protons, neutrons, and electrons in atoms helps us predict how a substance will behave in cells, multicelled organisms, and in the environment.”

Concepts/Ideas/Notes:

  1. the structure and function of all living things are governed by the laws of chemistry
  2. 90% of the mass of living things is composed of combinations of four elements: oxygen, carbon, hydrogen, and nitrogen
  3. all atoms of a given element have the same number of protons
  4. the net electrical charge of an atom is zero
  5. positive and negative charges attract each other, like charges repel each other
  6. we have about 1 ounce (30 grams) of weight of electrons in our bodies
  7. most naturally occurring isotopes are radioactive
  8. H2O  represents a chemical formula;  H—O—H  represents a structural formula.
  9. amount of energy in the universe remains the same over time but can change in form constantly
  10. the rate at which atoms move determines its state
  11. thermal energy must be added to a substance in order to change its state

History:

  1. J. J. Thomson – 1800’s, “raisons in a plum pudding” model of electrons
  2. Rutherford – 1911, electron orbital model of the atom/electrons

DEFINITIONS:

  1. MATTER – anything that occupies space and has mass
  2. ENERGY – ability to do work
  3. FREE ENERGY – energy in a system that is available to do work
  4. ACTIVATION ENERGY – amount of energy needed to start a reaction
  5. MASS – the quantity of matter an object has
  6. WEIGHT – the pull of gravity on an object
  7. ELEMENTS – pure substances that cannot be broken down chemically to simpler kinds of matter. It is a form of matter that takes up space, has mass, and can’t be degraded into something else.  The are 92 naturally occurring elements. 
  8. ATOM – simplest particle of an element that retains all of the properties of that element
  9. COMPOUND – a pure substance that is made up of atoms of two or more elements.  Molecules with unvarying proportions of two or more elements.
  10. MOLECULE – the simplest part of a substance (compound) that retains all of the properties of the substance and that can exist in a free state (not combined with anything else).  It is an assembly of atoms held together by covalent chemical bonds.  Molecules are made up of two or more atoms, either of the same element or of  two or more different elements. Ionic compounds, such as common salt, are made up not of molecules, but of  ions arranged in a crystalline structure. Unlike ions, molecules carry no electrical charge.
  11. MIXTURE – two or more kinds of molecules that intermingle
  12. PROTON – positively charged subatomic particle (+).  Abbreviated as p+
  13. NEUTRON – an uncharged subatomic particle.  Neutrons can vary in an atom’s nucleus (isotope).  The atomic mass minus the atomic number  (protons) equals the neutrons in an atom.
  14. ISOTOPE – atoms of the same element which have variable neutrons in their nucleus.  Forms of an element whose nuclei have the same atomic number – the number of protons in the nucleus--but different atomic mass because they contain different numbers of neutrons.
  15. RADIOACTIVE – element with an unstable nucleus that emits radiation (energy) as it changes into amore stable form. Heavier elements (84 or more protons) are unstable, and therefore, radioactive.
  16. HALF-LIFE – the rate of decay of a radioactive isotope which is defined as the time for half the atoms in a sample to change into another isotope or into a stable element.  The half-life of an element is constant, therefore it is possible to calculate the fraction of decay over a given isotope over a given period of time.
  17. AUTORADIOGRAPHY – technique in which a sample of material containing a radioactive isotope is placed on a sheet of photographic film and so reveal the exact location of the isotope within the specimen.
  18.  ELECTRON – high energy negatively charged subatomic particle (-) with very little mass.  Abbreviated e-.   the way electrons are arranged in an atom determines the chemical properties of that atom.  Chemical reactions involve changes in the number and the distribution of an atom's electrons.
  19. NUCLEUS – central core of an atom which consists of a proton with a positive charge (+) and a  neutron which has no charge.  Most of the mass in concentrated in the nucleus.
  20.  ATOMIC NUMBER – number of protons in the nucleus of an element
  21.  ATOMIC MASS UNIT (AMU) – the number neutrons plus the number of protons.  To find the number of neutrons in the nucleus just subtract the atomic number from the AMU and you would get the neutrons in the nucleus.  The AMU would be the higher number (and most of the time a number with a decimal) in the element box in the Periodic Table.
  22.  ENERGY LEVELS – a volume of space around the atomic nucleus in which electrons are likely to be at any instant. The first spherical orbital is the lowest energy level.
  23.  SHELL MODEL – illustration/visualization of how atoms are distributed in an atom.
  24. CHEMICAL REACTIONS – the combining or sharing of the electrons of one element with the electrons of another element.  It is the process of breaking or forming of chemical bonds or both.
  25. CHEMICAL EQUATIONS – show how reactions change during chemical reactions.  Reactants are shown on the left side and the products are shown on the right side of an equation.  The same number of atoms must be the same on either side of the equation.   A chemical equation represents the reactants and products in a chemical reaction using their symbols or formulas.
  26. REACTANTS – substances shown on the left side of chemical equation
  27. PRODUCTS – substances shown on the right side of a chemical equation
  28. TYPES of REACTIONS –     A + B ΰ AB               AB ΰ A + B            AB + CD ΰ  AD + C B
  29. CHEMICAL BONDS – union between the electron structures of two or more atoms.  Only electrons in their outer shells (valence shells) are involved.
  30.  ION – atom or molecule with an electrical charge either positive or plus (+) or negative or minus (-)
  31.  CATION – positive charged ion (+)
  32.  ANION – negatively charged ion (-)
  33.  VALENCE – an atoms bonding capacity that equals the number of unpaired electrons in its valence shell.  The valence is determined by the number of electrons that an atom can gain, or lose, or share.  The valence of Na+ is one, as is Cl- one.
  34. VALENCE SHELL – outermost electron shell.  The greater the distance from the nucleus, the higher the energy associated with it, the more electrons it can hold.
  35.  ORBITAL – volume of space around the atomic nucleus in which electrons are likely to be at any instant (90% of time).  The first spherical orbital is the lowest energy level.
  36.  ELECTRONEGATIVITY – the ease with which an atom can attract electrons to itself so forming negative ions, it is the affinity of an atom for electrons.  The measure of the ability of an atom or a molecule to attract a pair of electrons in a covalent bond.  The strength of the attraction depends on the number of protons in the nucleus, the number of electrons, and their proximity to the nucleus.  Electronegativity is expressed on a scale of 0 to 4 with the noble gases having electronegativities of 0.  When an electron moves from an atom that is less electronegative to one that is more electronegative, it moves "downhill" and energy is released.   Electronegativity generally increases from left to right on the periodic table and decreases from top to bottom.  The most electronegative element is fluorine (4).  Electrons tend to spend more time around the more electronegative atom, giving the more electronegative atom in the molecule a slightly negative charge, with the slightly less electronegative atom a slightly positive charge.
  37. ELEMENTS of LIFE – C H O P K I N S Ca Fe
  38. EXERGONIC – Chemical reactions that involve a net release of free energy.  Only exergonic reactions can proceed spontaneously.
  39. ENDERGONIC – Chemical reactions that involve a net input or the absorption of free energy to form bonds.  They do not occur spontaneously.
  40. EXOTHERMIC REACTIONS – release heat.  The energy needed for the reaction to occur is less than the total energy provided. As a result of this, the extra energy is released, usually in the form of heat.
  41. ENDOTHERMIC REACTIONS – absorb heat.
  42. CATALYSTS – chemical substance that can reduce the amount of activation energy that is needed for a reaction without being used in the reaction.
  43. ENZYMES – important class of catalysts in living things
  44. REDUCTIUON-OXIDATION REACTIONS or REDOX REACTIONS –  redox reactions always occur together!  In an oxidation reaction the electron given up by one substance is accepted by the other substance in the reduction reaction!
    1. OXIDATION REACTION – when a reactant loses one or more electrons, thus becoming more positive in charge (+)
    2. REDUCTION REACTION – when a reactant gains one or more electron, thus becoming more negative in charge (-)
  45. INTERMEDIARY METABOLISM – a series of chemical reactions by which energy is released slowly in controlled amounts that will not damage cells.
  46. LAW of CONSERVASTION of MASS – states that in a chemical reaction the total mass of reactants is equal to the total mass of the products.

STATES of  MATTER:

1.          SOLID – particles of a solid are linked together in a definite shape. A solid maintains a fixed volume and shape.

2.          LIQUID – particles in liquid are not as tightly linked giving a liquid the ability to flow and to conform to the shape of its container even though it has a fixed volume.

3.          Gas – rapidly moving particles that fill the volume of the container they occupy.  Particles of gas continue to disperse if not contained.

Ways of Modeling Molecules and Compounds:

1.         Ball and stick

2.         Chemical formula

3.         Electron shell model

4.         Structural Formula

5.         Space-filling model

6.         Icon ring models

7.       Lewis Dot

Energy in an Atom: energy goes from the lowest to the highest energy as you move away (outward) from the nucleus.  The first spherical orbital is the lowest energy level.  Energy is needed to move a negatively charged electron farther away from the positively charged nucleus.  For an electron to move from one level to another, the atom must absorb a discrete unit of energy, known as a quantum of energy, which contains precisely the amount of energy needed for the transition.  As an electron moves further from the nucleus it gains additional potential energy.  Energy is release as an electron moves back towards (inward) the nucleus or returns to its previous energy level.   

2n2  = the number of electrons that will fill an orbital where n is the level number or shell  level (eg.  "K" shell would be #1:  thus 1squared =1 and 1 x  2 = 2) .                                                

Names of the valence levels – K,L,M,N,O,P Q                 Names of the electron orbitals – s, p, d, f, g  

"s" orbital - is spherically shaped and contains a maximum of two electrons.

More Vocabulary : Radioisotope, Tracer, Radioactive Decay, Subscript. Superscript, Carbon 14

http://www.chemguide.co.uk/atoms/bonding/electroneg.html

Orbitals: sequence of electrons in orbitals(each orbital carries 2 electrons)                                                                                       1s22s22p63s23p64s23d104p65s24d105p66s24f14

  1. S subshells carry  a maximum of 2 electrons and have 1 orbital.
  2. P subshells carry a maximum of 6 electrons and have 3 orbital.
  3. D subshells carry a maximum of 10 electrons and have 5 orbital.
  4. F subshells carry a maximum of 14 electrons and have 7 orbital

Note Below:  Ordinary lines represent bonds in the plane of the screen or paper. Dotted lines represent bonds going away from you into the screen or paper. Wedged lines represent bonds coming out of the screen or paper towards you.

COVALENT BOND: WATER