Bonding:
1.
Atoms gain, lose or share electrons to form bonds.
2.
When bonds FORM, energy is RELEASED
3.
When bonds BREAK, energy has been absorbed
4.
There are 3 TYPES of bonds: Ionic, Covalent, and
Metallic
5.
Ionic Bonding: a) Involves a transfer of electrons b) from the metal to the nonmetal c) forming CATIONS ( + ions) and ANIONS (- ions)
d) takes place essentially between
a Metal / Nonmetal Ex. Na Cl
a Metal / Pai Ex. K
2
SO
4
a Pai / Pai Ex. (NH
4
)
2
CO
3
6.
Covalent Bonding: a) involves a SHARING of electrons between 2 or more of the SAME or
DIFFERENT nonmetals b) polar covalent bonds: unequal share between 2 DIFF nonmetals Ex. HCl c) nonpolar covalent: equal share
between 2 of SAME nonmetals Ex. H-H nonpolar covalent: equal to zero coordinate covalent:--------------------
d) takes place essentially between
2 or more Nonmetals
Some metalloid/nonmetal combos
7. Metallic Bonding:

a) holds all atoms of a metal together b) “sea of mobile valence electrons” surround
positive metal ions c) accounts for great electric and thermal
conductivity of metals
8. Covalent bonds form molecules (2 or more atoms sharing electrons).
9. Compds with Pai’s have BOTH ionic and covalent bonds Ex. NaOH (ionic holds the Na to the OH, BUT covalent holds the O to the H)
10. Molecular formulas show elements and subscripts making up the formula of a compound, while Lewis dot diagrams show "circles of happiness",and structural formulas replace bonded pairs in a Lewis dot with a dash.
11.
VSEPR theory=electrons repulse each other. This affects bond angles and shape.
Shape etc, helps predict properties.
12. Shapes Chart:
Shape Angle Ex Lone pairs pyramidal 1
Trigonal planar tetrahedral bent linear
Shared pairs
3

13.POLARITY of the WHOLE MOLECULE:
Polar Molecules are ASYMMETRICAL
( uneven charge distribution)
Ex. H----F , NH
3
Nonpolar Molecules are SYMMETRICAL
(even charge distribution)
Ex. H---H, CH
4
14.POLAR MOLECULES must have polar
bonds BUT NONPOLAR molecules COULD
have either nonpolar OR polar BONDS
15.Dipole: molecule with BOTH polar bonds
AND asymmetry
16. Dipole-Dipole attraction: a weak force of attraction between 2 dipole molecules, in which the slightly + end of one, is attracted to the slightly (-) end of the other. Ex. between HCl and HCl
17. Bonds are between atoms WHILE Forces of attraction (not bonds) are BETWEEN molecules. Ex. Dipole-dipole, ion-dipole, and
Hydrogen “Bonding”
18.Hydrogen "bonding": Very strong force of attraction; accounts for the high b.p. of water,
characterized by H---F, H---O, and H---N
bonds [et FON H ome]

19. London dispersion: exists between nonpolar molecules despite no polar ends to attract each other until instantaneous dipole moment forms; increases as molecule distance decreases and as mass increases.
PHASES OF MATTER: SOLIDS & LIQUIDS
20. Solids: definite shape and volume
21. solid changing to liquid= fusion or melting. This requires the input of heat (endothermic change)
22. sublimation: change from the solid directly to the gas phase without showing liquid. Ex. CO
2
(s) (dry ice) and Iodine (s)
23. liquids: definite volume only, take shape of their container
24. liquid changing back to solid = freezing. This releases heat (exothermic change)
25. liquid changing to gas = vaporization (endo)
26. liquid changing to gas at the surface only= evaporation. (endo)

27. Evaporative cooling: highest KE molecules change to a gas first leaving low KE(cool) molecules.
28. liquid changing to gas throughout the liquid AND matching atmospheric pressure = boiling (endo)
29. vapor pressure of a liquid: the pressure exerted on the walls of a container and against the atmosphere from the vapor molecules above the surface of their own liquid.
30. vapor pressure and temp: the greater the temp of the liquid, the greater the KE of its molecules, the higher the liquid’s vapor pressure.
31. The weaker the intermolecular forces of attraction within a liquid, the higher that liquid’s v.p.,
(the more volatile the liquid is)
32. when vapor pressure(v.p.) of a liquid matches (=) the pressure of the atmosphere above it, boiling occurs
33. When the atmospheric pressure above a liquid is
STANDARD PRESSURE (760 torr, 1atm or 101.3 kPa), the TEMP you must raise the liquid to so the liquid’s

v.p. can match standard pressure is called the liquid’s NORMAL boiling point.
PHASE CHANGES:
33 A.
temp = a measure of average kinetic energy
(energy of movement). Temp (KE) will change on the diagonal lines of a phase change graph where
ONLY one phase is being heated or cooled
33 B.
Potential energy= stored energy or latent heat used to break bonds. Potential energy changes on the horizontal lines ONLY, where two phases exist because one phase is changing to another.
33C.
Zero Celsius = melting/freezing point
H20; 100 celsius= normal boiling point water
33D.
The horizontal line for H20 vaporization is longer (the PE greater) than for other substances, and for its own fusion line, because water is most dense in the liquid phase.
33E.
Going forward on the Heating/Cooling
Curve for water = endothermic (ex. Heating a liquid or changing from liquid to gas phase);
Going backwards =exothermic

PHASES OF MATTER: GASES
34.
Gases have NO definite shape, NO definite volume..they spread out to fill the container. They are highly compressible and therefore are GREATLY influenced by
TEMP and PRESSURE
35.
Pressure= force per unit area. Units of pressure include atmospheres, torrs (mmHg), millibars, pounds per sq.inch, kilopascals
*1 atm=760 torr
*1 atm= 760 mmHg
*1 atm= 101.3 kPa
36.
Barometer: instrument used to measure air pressure.
The higher the altitude, the lower the height of Hg in
the tube.
37.
STP Conditions: Standard conditions for gases include a
Temp of
273 kelvin or 0 celsius and a Pressure of 1 atm or
760 torr.
Sometimes STP = the orig or new conditions for a gas.
38.
Boyles Law: Pressure and volume/ Indirect/ P
1
V
1
= P
2
V
2
39.
Charles Law: Temp and volume (Charlie the Tuna)/
Direct/
V1 = V2
T1 T2 MUST BE IN KELVIN (k= C+273)

40. Gay-Lussac: Temp and Pressure (Need Toilet
Paper in the “Lu”) / Direct/
With an increase in temp but volume constant, the gas
molecules collide MORE FREQUENTLY and with greater force.
P1=P2
T1 T2 MUST BE IN KELVIN
41.
Combined Gas Law: Must use when temp AND pressure
BOTH change. OR Can use in place of any of the 3 gas laws
(cross out the unneeded variable) P
1
V
1
= P
2
V
2
(KELVIN)
T
1
42. Ideal gas law: PV=nRT
P= pressure MUST BE IN ATM
V= volume MUST BE IN LITERS,
n= moles,
R= 0.0821
L-atm/moles-kelvin
T= temp in Kelvin
43. Avogadro:
Equal volumes of gases at the same Temp
and Pressure have the same number of
molecules.
44.
1 mole of any gas at STP conditions occupies
22.4 liters. So… New Moleville box:

1mole=22.4 liters If at STP
45.
Your
4
Best Friends for converting
ONE
substance:
1 mol______
Molar mass
Grams
6.02x10
23
1
________
1
count subscripts
46.
Split Moleville ONLY IF you are comparing two substances in a balanced equation.
RECALL:
Recall: c oeff mol
new
Coeff mol
orig
****** NEED TO SEE A BALANCED EQUATION FOR
STOICHIOMETRY
1 N
2
+ 3 H
2
 2 NH
3
For ex. there are 3 hydrogens to 2 ammonias
47. Dalton’s Law 1: The total pressure of a mixture of gases is equal to the SUM of the partial pressures of each gas: P total
= P
1
+ P
2
+ P
3
48. Kinetic Molecular Theory (KMT): a) constant random motion b) elastic collisions c) because gases spread out they have NO
signif volume
d) because gases spread out they have NO

forces of attraction between molecules
49. Ideal gases such as Hydrogen and Helium,
Behave as the theory predicts (meaning
Boyles, Charles, Ideal, etc. match real life
50. All other gases are called “Real gases” (any gas other than H
2 and He) ONLY behave as theory predicts IF you
SPREAD THEM OUT= INCREASE TEMP AND DECREASE
PRESSURE
51. Real gases DEVIATE (move away) from the theory when they move close together = LOWER THE TEMP AND
INCREASE THE PRESSURE
SOLUTIONS:
52.
SOLUTIONS: Solute= material BEING dissolved. Ex. Sugar or coffee grinds into coffee. Solvent= the material DOING the dissolving. Ex. The hot water for the coffee
53.
Factors that affect RATE of solubility (How FAST a solute dissolves) include stirring, temp, and surface area
54. Rate(stirring): For SOLID and LIQUID solutes= as stirring increases so does how fast the solute dissolves. (Stirring sugar makes it dissolve faster)
*** For Gases: stirring decreases degree and increases rate of solubility
55. Rate(temp): increasing the temp of a solvent increases the rate of dissolving for a solid or liquid solute. (Hot coffee dissolves sugar faster).
*** Increasing the temp of a solvent decreases the rate at which a gas dissolves (hot soda=flat soda)
56. Rate(surface Area): the greater the surface area of a solid solute, the faster the solute dissolves
57. Degree of Solubility= HOW MUCH can be dissolved. Factors
Affecting include: Nature of solute/solvent (bond types), temp, how much already dissolved and pressure (gases only)
58.
Degree(temp): the higher the temp of the solvent the more
SOLID/LIQUID solute dissolves
**** For Gases: the higher the temp of solvent the
LESS gas dissolves (warm soda=flat soda)

59.
Degree(Pressure) GASES ONLY: More gas dissolves as
the pressure of the gas over its solvent increases
(Henry’s Law) Ex. CO
2
(carbonation) added to soda under high pressure;
Nitrogen in diver’s bloodstream at great pressure (bends)
2
2
2
2
61. Dilute(“weak”): a solution containing a relatively small amount of solute
62. Concentrated(“strong): a solution containing a relatively large amount of solute
63. unsaturated: holds less than the max amount of solute the solution could hold at a given temp; One more added crystal will dissolve
64. saturated: holds the max amount of solute the solution can hold at a given temp; one more crystal will not dissolve
65. supersaturated: holds more than the max amount of solute the solution can hold at a given temp accomplished through heating the solvent; upon cooling the excess can “grow” out of solution if a nucleus exists OR the excess can come out all at once after cooling IF system is agitated
66. Solubility curves: hit the line=saturated, falling below the line= unsaturated, above the line= supersaturated

67. Solubility curves: make a proportion IF the amount of water is more or less than 100 ml. Read the value FOR 100 ml DIRECTLY from the graph, then set up a ratio.
68. Solubility curves: IF BOTH the amount of water AND the temp is changed, repeat steps in #64 above 2x, once for the first temp, a second time for the new temp (higher or lower than the first), then take the difference from your 2 ratios. NEVER place the temps in the ratio you set up.
69. Molarity (M); a way to express concentration; the
# of moles of solute per liter of SOLUTION; solute
AND solvent ADD together to get to the volume
70. For (M): convert grams to moles and milliliters to liters; Formula is M=n/L. (1000ml=1 liter)
71. molality(m) : a way to express concentration; the number of moles of solute per kilogram of SOLVENT; solvent is full amount before solute is added.
72. For (m) m=n/Kg
H20
; convert grams of SOLUTE to moles; convert grams of SOLVENT(water) to kilograms (liters). (1000 g= 1kg)
73. Dilution: M
1
V
1
=M
2
V
2
M
1
=original molarity/
V
1
=original amount of stock or concentrated solution withdrawn/ M
2
=concentration of solution you want/
V
2
= final volume of diluted solution you want

74. Colligative properties: characteristics such as boiling point that change depending on the NUMBER of solute particles added to solution: the more solute added, the HIGHER the b.p. of a soln, the LOWER the f.p.
75. For water the b.p. elevation is 0.52 degrees per molal solution and 1.86 degrees for f.p. depression per molal solution.
76. Formula to determine CHANGE in b.p or f.p=
 T = k x m
(x ions if needed)
Change in temp constant molality for IONIC SOLUTES
77. If solute is covalent (molecular) the formula ends at molality . If solute is IONIC you must multiply by the # of ions that dissociated
78. For FINAL b.p. water use: 100 (+)  T
FINAL f.p. water use: 0 (-)  T