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2014-03-14 08:50:33

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  1. quantum number
    a number which occurs in the theoretical expression for the value of some quantized property of a subatomic particle, atom, or molecule and can only have certain integral or half-integral values.
  2. Integer
    a number which is not a fraction; a whole number. A thing complete in itself.
  3. molecule
    a group of atoms bonded together, representing the smallest fundamental unit of a chemical compound that can take part in a chemical reaction.
  4. Substance
    a chemical substance is a form of matter that has constant chemical composition and characteristic properties. It cannot be separated into components by physical separation methods, i.e. without breaking chemical bonds. It can be S,L,G
  5. solute
    The substance present in smaller amount in a solution.

    The thing which is being dissolved in a solution.

    All solutes that dissolve in water fit into two categories –electrolytes and nonelectrolytes. All ionic solids are strong electrolytes but they are not equally soluble. Sodium Chloride is a strong electrolyte which means that, when it dissolves in water, it is completely dissociated into ions.
  6. Solvent
    A liquid in which substances (or solutes) are dissolved forming a solution.
  7. True or False, electrons can be arranged in Atoms and in Molecules
    True. Arranged in Molecules simply means electrons are bonding, therefore, we need to understand how they are bonding, and what the electrons are doing in the Bonding state.
  8. Valence electrons
    Those which lie in the outermost orbitals, the ones with the largest Principal Quantum Number.
  9. Ionic Bonds
    a chemical bond in which one atom loses an electron to form a positive ion and the other atom gains an electron to form a negative ion.
  10. Covalent Bond
    a bond between two or more atoms that is provided by electrons that travel between the atoms nuclei, holding them together but keeping them a stable distance apart. sharing electrons between two or more atoms. Covalent usually having no Charge.
  11. Double Bond
    a chemical bond between two chemical elements involving four bonding electrons instead of the usual two. The most common double bond, that is between two carbon atoms, can be found in alkenes.
  12. Octet Rule
    Atoms of low (<20) atomic number tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. Note heavier elements can accomodate 10, 12 or even more electrons e.g. Phosphorus (10) and sulphur (12).
  13. Lewis Structure (Steps 1)
    1. Count valence electrons for each atom
  14. Lewis Structure (Step 2)
    Assemble bonding framework using single covalent bonds (pairs of electrons). Note more Electronegative atoms or H on the outside.
  15. Lewis Structure (Step 3)
    Place three nonbonding pairs of electrons on each outer atom except H
  16. Lewis Structure (Step 4)
    Assign remaining valence electrons to inner atoms
  17. Lewis Structure (Step 5)
    Minimise formal charges on all atoms.

    Note: usually by adding double bonds. Note you can only add tripple bonds if the atom's V.electron is > 6 or if Atom can hold tripple bonds. Can't just add tripple bond simply because there's available 'lone pairs'. Use Formal Charge equation to work out if additional bond can be added.

    Note - Orbitals usually only hold 8e, that's how you know how many double bonds or tripples you should use. Even though sharing is occuring, you still count the shared ones towards the atom you are counting the electrons for.
  18. Formal Charge
    Formal Charge = Valence electrons on 'free atom' - Electrons assigned in Lewis Structure.

    V.E on free atom, is the number of V.E the atom can have (mostly 6).
  19. Anion
    An ionic species having a Negative charge
  20. nonpolar covalent bond
    A covalent bond where the electrons are shared equally between the two atoms. Usually only between atoms of the same element.
  21. polar covalent bond
    A polar bond is a covalent bond between two atoms where the electrons forming the bond are unequally distributed. This causes the molecule to have a slight electrical dipole moment where one end is slightly positive and the other is slightly negative. Polar bonds are the dividing line between pure covalent bonding and pure ionic bonding.
  22. dipole
    a pair of equal and oppositely charged or magnetized poles separated by a distance.
  23. Cation
    An ionic species having a POSATIVE charge
  24. Polyatomic Ions
    A polyatomic ion, also known as a molecular ion, is a charged chemical species (ion) composed of two or more atoms covalently bonded or of a metal complex that can be considered to be acting as a single unit. The prefix "poly-" means "many," in Greek, but even ions of two atoms are commonly referred to as polyatomic.
  25. What do S Orbitals, P Orbitals, and D Orbitals look like?
  26. What is Valence Shell Electron Pair Repulsion Theory? VSEPR
    A model in chemistry used to predict the shape of individual molecules based upon the extent of electron-pair electrostatic repulsion. So the shape of the moecule adopts will be the one where these electron pair repulsions are best minimised.

    The premise of VSEPR is that the valence electron pairs surrounding an atom mutually repel each other, and will therefore adopt an arrangement that minimizes this repulsion, thus determining the molecular geometry.

    VSEPR theory is usually compared with valence bond theory, which addresses molecular shape through orbitals that are energetically accessible for bonding.
  27. What is the Geometry of the sets of Electron Pairs, Molecular Structure, and Angle?

    Number of Pairs: 2
    • Geometry: Linear
    • Structure: O-O-O
    • Angel: 180

    • Note: H2O is:
    • Geometry: Bent
    • Structure: Tetrahedrally arranged
    • Angel: 109.5
    • (NOTE: H2O has two "Lone Pairs" and Lone Pairs are more repelling than Shared Pairs, hence the BENT STRUCTURE.
  28. What is the Geometry of the sets of Electron Pairs, Molecular Structure, and Angle?

    Number of Pairs: 3
    • Geometry: Trigonal Planar
    • Structure:

    Angel: 120
  29. What is the Geometry of the sets of Electron Pairs, Molecular Structure, and Angle?

    Number of Pairs: 4
    • Geometry: Tetrahedral
    • Structure:

    • Angel: 109
  30. What is the Geometry of the sets of Electron Pairs, Molecular Structure, and Angle?

    Number of Pairs: 6
    • Geometry: Octahedral
    • Structure: (above)
    • Angel: 90
  31. True or False:
    VSEPR makes a distinction between electron pairs in single, double, and/or tripple bonds?
    FALSE: The types of bonds are irrelavent. The important part is HOW MANY BONDS there are.
  32. What is Valence Bond Theory?

    Says that in making a bond between two atoms (Chemical Bonding), the atomic orbitals, in which the valence electrons from each atom lie, overlap to make a new orbital in which the pair of electrons exist.
  33. What is the Molecular Formula for Methane?

    It is the simplest alkane and the main component of natural gas. The relative abundance of methane makes it an attractive fuel. However, because it is a gas at normal conditions, it is difficult to store it.
  34. What is the Molecular Formula for Ammonia?

    • is a compound of nitrogen and hydrogen with the formula NH3. It is a colourless gas with a characteristic pungent smell.
    • Ammonia contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to food and fertilizers. Ammonia, either directly or indirectly, is also a building-block for the synthesis of many pharmaceuticals and is used in many commercial cleaning products.
  35. What is the Molecular Formula for Carbon Dioxide?

    Carbon dioxide (chemical formula CO2) is a naturally occurring chemical compound composed of 2 oxygen atoms each covalently double bonded to a single carbon atom. It is a gas at standard temperature and pressure and exists in Earth's atmosphere in this state, as a trace gas at a concentration of 0.039 per cent by volume.[1]
  36. What is the Molecular Formula for Hydrogen Peroxide?

    is the simplest peroxide (a compound with an oxygen-oxygen single bond). It is also a strong oxidizer. Hydrogen peroxide is a colorless liquid, slightly more viscous than water. In dilute solution, it appears colorless. Due to its oxidizing properties, hydrogen peroxide is often used as a bleach or cleaning agent. The oxidizing capacity of hydrogen peroxide is so strong that it is considered a highly reactive oxygen species. Concentrated hydrogen peroxide, or 'high-test peroxide', is therefore used as a propellant in rocketry.
  37. What is a Sigma Orbital?
    an orbital that is spherically symmetrical about the internuclear axis. When you turn it side on the 'p orbital' and look down it's internuclear axis, you will see the atoms will look superimposed.
  38. Orbital Diagrams (arrows pointing up and down, representing electron placing/bonding)
    for 1s and 2s, arrows fill up systematically i.e. up, down, up, down. After the 2s orbitals are filled, we move to p orbitals which have three orbitals. Instead of the up, down, up down, up down sequence, paring up is more difficult for P orbitals therefore, each individual P orbital is first filled up e.g. up, up, up, then up down, up down, up down.
  39. What are hybrid orbitals?
    In chemistry, hybridisation is the concept of mixing atomic orbitals into new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory. Hybrid orbitals are very useful in the explanation of molecular geometry and atomic bonding properties. E.g. in CH4 (Methane) a hybridise orbital of sp is required, the hybrid of the sp raises the energy level of the s orbital and decreases the energy of the p orbital to get a weighted average of energy.

    Note it's about taking ATOMIC ORBITALS and turning them into MOLECULAR ORBITALS. And if you start with 4 atomic, must end with 4 molecular orbitals. No creating or destroying orbitals. See Lecture 3 Slide #18 (Marked A). So sp3 orbital means 1 s orbital and 3 p orbitals hybrid. Note 'Hybrid Orbitals always lead to Sigma Bonds.
  40. Sodium Chloride

    - Sodium Chloride is a strong electrolyte which means that, when it dissolves in water, it is completely dissociated into ions i.e.

    NaCl(s) + H2O -> Na+(aq) + Cl-(aq)

    This equation tells us that all the solid NaCl (solute) that enters the water (solvent) dissociates to give Na+(aq) cations and Cl-(aq) anions.
  41. Spectator Ion
    An ion that, although present in the solution, takes no part in the reaction.
  42. Precipitate
    An insoluble solid which separates from a solution.
  43. Insoluble Solid
    incapable of being dissolved
  44. Limiting Reagent (Reactant)
    The reagent (reactant) which is used up first in a reaction.
  45. Electrolyte
    A substance that dissolves to produce ions in a solution.
  46. lead nitrate
  47. Lead(II) chloride
  48. Soluble
    the maximum quantity of a substance that may be dissolved in another. The maximum amount of solute that may be dissolved in a solvent.
  49. Solubility Rules for Common Ionic Compounds(in Water at 25oC)
    1. All Group 1 compounds are soluble.

    2. All ammonium (NH4+) compounds are soluble.

    3. Most hydroxides (OH–) are insoluble; the exceptions are the Group 1 hydroxides and barium hydroxide(Ba(OH)2). Calcium hydroxide, (Ca(OH)2), is only slightly soluble.

    4. Most compounds containing nitrate (NO3–), bicarbonate (HCO3–) and bisulfate (HSO4–) are soluble.

    5. Most compounds containing chlorides (Cl–), bromides (Br–), or iodides (I–) are soluble. Exceptions arethose containing Ag+ and Pb2+.

    6. All carbonates (CO32–), phosphates (PO43–) and sulfides (S2–) are insoluble. The exceptions are those ofthe Group 1 metals and the ammonium ion. See 1 and 2 above.

    7. Most sulfates (SO42–) are soluble. Calcium sulfate (CaSO4) and silver sulfate (Ag2SO4) are slightlysoluble. Barium sulfate (BaSO4) and lead sulfate (PbSO4) are insoluble. Calcium sulfate (CaSO4) and silver sulfate (Ag2SO4) are slightly soluble.
  50. Substances can be defined as?
    pure substances or mixtures.Mixtures can be classified as homogeneous (uniform throughout) or heterogeneous.

    Solutions are an important examples of homogeneous mixtures
  51. Solutions in which water is the solvent are?
    aqueous solutions
  52. What are electrolytes?
    Substances that give rise to charged particles when they dissolve in water are classed as electrolytes.

    Ionic solids are electrolytes because to the extent that they dissolve they dissociate into their ions.

    NaCl(s)  Na+(aq) + Cl–(aq) Electrical current is movement of charge, and these ions can move in the solution to complete an electrical circuit.
  53. Solubility Rules for Ionic Solids (per bestchoice)
    1. All Group 1 compounds (containing Na+, K+, Li+ ....) are soluble.

    2. All ammonium (NH4+) compounds are soluble.

    3. All compounds containing nitrate (NO3–), hydrogen carbonate (HCO3–) and hydrogen sulfate (HSO4–) ions are soluble.

    4. Most compounds containing chloride (Cl–), bromide (Br–), or iodide (I–) ions are soluble.Exceptions: Halides of Ag+ and Pb2+

    5. Most hydroxides (OH–) are insoluble.Exceptions: Group 1 hydroxides (see 1 above) and barium hydroxide (Ba(OH)2).Calcium hydroxide, (Ca(OH)2), is slightly soluble.

    6. All carbonates (CO32–), phosphates (PO43–) and sulfides (S2–) are insoluble.Exceptions: Compounds containing Group 1 metal cations and/or ammonium ion. See 1 and 2 above.

    7. Most sulfates (SO42–) are soluble. Exceptions:  Barium sulfate (BaSO4) and lead sulfate (PbSO4) are insoluble.Calcium sulfate (CaSO4) and silver sulfate (Ag2SO4) are slightly soluble.

    Despite their wide range of solubilities, all ionic solids are electrolytes; that is, to the extent that they dissolve they exist as hydrated ions.

    • It may be helpful to you to make the general observation from this set of rules that compounds where the charge on the ions is higher tend to be insoluble.  
    • Why is this so?  One of the factors that determines whether a substance is soluble in a given solvent is the relative strengths of the possible interactions:
    • solute-solute (in pure solute) 
    • solvent-solvent (in pure solvent)
    • solute-solvent (in a solution) 

    For soluble substances the solute-solvent attractive forces are stronger than the attractive forces between solute particles or solvent particles.For ionic solids where the ions are highly charged, the attractive forces between the ions in the lattice are stronger than for ions with lower charges; thus the difference between the solute-solute attractive forces and solvent-solute attractive forces are greater.

    • Summary of rules:
    • Look at cation. Compounds containing Na+, K+ or NH4+ are soluble.

    If the cation is not one of the above, look at anion.

    • If anion has –1 charge and is not OH–, the compound is probably soluble.
    • Exceptions: Chlorides, bromides and iodides of Ag+, Pb2+ are insoluble.

    • If the anion has –2 or –3 charge or is OH–, the compound is probably insoluble.
    •  Exceptions: Sulfates other than those of Ca2+ , Ba2+ , Pb2+ are soluble
  54. What is a 'precipitate'?
    An insoluble solid

    As shown below, the formula for the precipitate can be predicted from knowledge of the solubility rules and the formulae of the ionic solids mixed.

    • Consider mixing solutions of two soluble ionic solids AX and BY.
    • - AX exists as A+ and X– ions in solution.  - BY exists as B+ and Y– ions in solution.
    • - The mixed solution contains two cations (A+ and B+) and two anions (X– and Y–). 

    • What are the possible precipitates?
    • Precipitates are insoluble ionic compounds and have no overall charge. The two combinations that can form a precipitate are AY and BX.
    • Recall that AX and BY are soluble compounds!
    • These each have the cation of one of the solids in solution and the anion of the other. 

    Precipitation occurs if one of these pairs has the cation and the anion of an insoluble solid.

    No precipitate forms if both of the possible products are soluble!

    • What happens to the other pair of ions?
    • Because these ions are not changed due to reaction, they are referred to as spectator ions.These remain in solution.The precipitate can be separated from the solution by filtration.  The soluble substance can be recovered by evaporation of the filtrate.
  55. Electronegativity
    is a measure of the tendency of a bonded atom to attract the bonding electrons to itself.
  56. Electronegativity Facts
    electronegativities are higher for elements to the right and top of the periodic table

    metals are less electronegative than non-metals

    Metals are referred to as electropositive elements. hydrogen has the lowest

    electronegativity of the elements commonly appearing in organic compounds.
  57. Electronegativity for bonded atoms = >2
    For bonded atoms where the difference in electronegativity is greater than 2.0, the bond is ionic.

    The atom of lower electronegativity exists as the cation (usually a metal such as Na+ in NaCl).

    The atom of higher electronegativity exists as an anion (usually a non-metal such as Cl– in NaCl).
  58. Electronegativity for bonded atoms = 0
    If the difference in electronegativity between the bonded atoms is zero, the bond is covalent and the electron pair is shared equally as shown in the electron density map for chlorine.
  59. Electronegativity for bonded atoms = >0 but <2
    If the difference in electronegativity is more than 0 and less than 2.0, the bond is polar covalent.

    Both atoms have a partial charge, the one of higher electronegativity being more negative

    The relative electron densities at the bond atoms depends on the difference in electronegativities.
  60. s orbitals and p orbitals
    two sp - linear - two p unhybridized

    three sp2 - trigonal planar - one p unhybridized

    four sp3 - tetrahedral - none unhybridized
  61. Molar Mass is given in units of?
  62. mass is given in units of?
  63. Concentration is given in units of?
  64. Diatomic Elements
    • Hydrogen (H2)
    • Nitrogen (N2)
    • Oxygen (O2)
    • Fluorine (F2)
    • Chlorine (Cl2)
    • Iodine (I2)
    • Bromine (Br2)