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					               Chemistry Honors Review for
               Mid - term Exam 2006

This Study Packet is only a guide.
Your notes, textbook and orange review book should be studied to help
you clarify anything you don’t understand and to fill in information that
is intentionally left out for you to do in this packet.
BRING THIS STUDY PACKET AND YOUR ORANGE BOOK TO
ANY REVIEW CLASSES YOU ATTEND
When doing each topic in the orange book, do the ―review for Regents
―questions at the end of each topic, THE MULTIPLE CHOICE AND THE
FREE RESPONSE. You can also do the questions within the topic for
practice.
Matter and Energy
Properties of solids - definite shape & volume, fixed atoms;
  regular geometric pattern (crystalline) example NaCl (s)
  draw the crystal structure of NaCl (s) here. Show that the elements are ions



Properties of liquids - no definite shape, but definite volume H2O (l). which elements on the periodic
table are liquids at STP

Properties of gases - no definite shape or volume, random particle motion.
  List all the gases on the periodic table here



Elements - all atoms of the same element have the same ATOMIC # and Can NOT
  be broken down chemically . How many elements are listed on your periodic table?

Mixture - 2 or more elements physically combined.
 There are different types of mixtures
  a. Heterogeneous mixtures (uneven - lumpy iced tea) solid granite, mint choco chip
     ice cream, pizza
  b. homogeneous mixtures (evenly mixed SOLUTION - clear tea, salt water, solids such as a
     sugar cookie, pudding

Homogeneous mixtures that are liquids are called solutions

Mixtures that contain water are called aqueous     (aq)
Physical change - no change in the identity of the substance (i.e. gas → liquid → solid)
                                                                                                         1
Chemical change - substance changes into new substance with NEW properties
Reactants combine to form new products

H2 + O2 → 2H2O: Chemical reaction

Pb(NO3 )2(aq) + 2HCl   (aq)   → PbCl2 (s) + 2HNO3 Take note that this is a precipitation reaction.

For energy problems: know the following formula
Q = m∆t c THIS EQUATION IS ON TABLE T

JOULE is a measurement of energy.

USE The above EQUATION FOR ENERGY CHANGE PROBLEMS
This formula is used when there is a change in temperature
This also applies to the diagonal portion of the heating and cooling curve because there is a change in
temperature

 m = mass,
∆t = change in temperature ( final temp – the start temp)
c = specific heat

Specific heat is the amount of heat needed to raise 1 gram of water 1° C

To find the specific heat (c) of water look on Table B
The specific heat of water is 4.18 Joules

Kinetic Energy Is the energy of motion
Temperature is the measure of average kinetic energy


Know how to convert from Celsius to Kelvin (+273) and back (- 273)
   look on table T

Fixed points on a thermometer –

0° C - freezing/melting point of H2O , remember the heating and cooling curve

100° C is the boiling/condensation point of H20 remember the heating and cooling curve

You need 2 points to create a thermometer. 0° C       & 100° C
ABSOLUTE ZERO -273 C° OR 0 K

Density equation Table T
density is how much mass is in a given volume
 D = Mass (g) / volume cm3

Sublimation → a substance turns directly from a solid to a gas
ex. CO2(s) → CO2(g)
 I2(s) purple crystals→ I2(g)purple gas
                                                                                                          2
Deposition a substance goes directly from gas to solid

Heat of fusion – the amount of heat needed to be absorbed to change 1 gram of a solid to 1 gram of a
liquid
Heat of fusion for water 333.6 J/g table B

Heat of vaporization the amount of heat needed to be absorbed to change 1 gram of a liquid to 1 gram
of a gas
Heat of vaporization for water 2260 J /g table B

Endothermic absorbing heat S - L G
Exothermic releasing heat   G L S
Entropy –the amount of disorder, entropy increases SL G
Phase change diagrams

Complete the heating and cooling curve by adding
Exothermic, and endothermic direction of the graph, increasing in entropy, decreasing in entropy, phase change
portion, increase in average kinetic energy, increase in potential energy, melting pt., freezing pt., boiling pt
equations            Q = m∆t c ,     Q= m Hfusion        Q= m Hvaporization




T
E                                                                                         (boiling)                gas

M                                                                                         vaporization

P                                                                                        condensation
                                   melting                         liquid

                                   fusion


                                   freezing
            solid




                                             time

                                                                                                                         3
Boiling point - the temp. at which the vapor pressure of a liquid equals the    atmospheric pressure
Vapor pressure - depends on the Temperature of the liquid
Strength of intermolecular forces determines the vapor pressure
Use the vapor pressure table Table H

Atomic Structure The atom is the smallest unit of matter
Parts of the atom

Proton - (+) charged; 1 atomic mass unit

Neutron - () charged; ~1 atomic mass unit
Electron - (-) charged; 1/1836 atomic mass unit
Nucleon - particles found in the nucleus (protons & neutrons)
Nucleus - contains most of the mass of the atom and has a positive charge;
The # of protons is called the nuclear charge because the neutrons have no charge
1 AMU - the atomic mass

 In a neutral atom the # of protons = the number of electrons.
In an ion the number of electrons don’t = the number of protons.

Metals form positive ions by losing electrons and therefore have more protons then electrons.

Non-metals form negative ions by gaining electrons and therefore have more electrons then protons

Atomic # is equal the # of protons in an atom
The atomic number is used to identify the element
All atoms of the same element have the same number of protons and therefore the same atomic number

Atomic mass = # of protons + # of neutrons

 Isotopes - elements that have the same atomic # but different atomic masses due to a difference in the #
of neutrons in the nucleus.

To calculate the # of neutrons in an element subtract the # of protons FROM the atomic mass.
  14
6C    has 6 protons, 6 electrons and 8 neutrons

Atomic mass on the periodic table is really a weighted average of all of the isotopes that exist in nature
for that element. i.e. Carbons atomic mass = 12.011 because there is 6C12 and 6C14in nature but 6C12
is more abundant and therefore skews the average toward 12.

THE MODELS OF THE ATOM
The Greeks- single units
Dalton Atomic Theory
Thompson – Plum Pudding
Rutherford- Gold Foil Experiment
Bohr- Orbits
Modern- Electron Cloud, Quantum mechanical model

                                                                                                             4
 Empty space concept - states that atoms are made up of mostly empty space and most of the mass is
confined to a very small nucleus. This was proven by the gold foil experiment.
Draw the gold foil experiment




 Bohr's model of the atom - stated that electrons traveled in certain orbits. An absorption of energy will
cause electrons to TEMPORARILY jump to higher levels & when the electrons fall back down to lower
levels they EMIT this energy in the form of light. Lower level electrons are in the ground state, when the
jump they are in the excited state.

Valence electrons - electrons in the outermost energy levels. i.e. 9F19 1s2 2s2 2p5 has 7 valence electrons
SINCE the outer most principle energy level is the 2nd one.

Kernel electrons are the electrons that orbit the nucleus of atom and are NOT considered to be part of
the valence shell.

Electron dot diagram - uses dots for the valence electrons




                                                                                                         5
Electron Configuration
Used to located the electrons in the atom
Principal energy levels – 1,2,3,4,5,6,7
Sublevels s,p,d,f
Orbitals s has 1 orbital
          p has 3 orbitals
          d has 5 orbitals
          f has 7 orbitals

the filling order 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2              4d10 5p6

Orbital diagrams uses boxes to illustrate the orbitals electrons can take around the nucleus.


  S                  p                         d                               f
Arrows represent the electrons & two electrons or arrows can fit into each box or orbital.
The electrons in the same orbital MUST spin in opposite directions.

 Hund's rule - before an orbital can get a second electron each orbital in that sublevel must have at least
one in each.

 Order of filling sublevels: 1s2 2s2 2p6 3s2 3p6 4s2 3d10: WHY? The 4s2 sublevel needs less energy to fill
than the 3d10 sublevel.
Electron configuration tells where the electrons can be found

Pauli’s Exclusion principle. The electrons must have opposite spins in each orbital

Aufbau’s Principle- ―the building up principle‖. Electrons must fill the lowest energy level before
moving to the next energy level

Look on the periodic table for the electron Configuration

                                     C
                                     2-4
                 2 electrons in the 1st principal energy level
                 4 electrons in the 2nd principal energy level, (2 in the s, 2 in the p)

excited state-

ground state -



Bonding –THEN JOINING TOGETHER OF
ATOMS
When a bond is formed energy is released (exothermic); when a bond is broken energy is absorbed
(endothermic)
                                                                                                              6
Atoms bonded together to form OCTETS (eight valence electrons are stable s2 p6 = 8 valence electrons)

Metals tend to lose electrons and form positive ions.
 (Ions formed are smaller than the neutral atoms: Ionic radii < than atomic radii)

Nonmetal tend to gain electrons and from negative ions (ions formed are larger than the neutral atoms:
Ionic radii > atomic radii)

A chemical bond - results from the simultaneous attraction of electrons by two nuclei


Ionic bonds - formed between metal and nonmetal; created by a transfer of electrons;
electronegativity difference > 1.7


Covalent bond - formed by the sharing of electrons; electronegativity difference < 1.7
Electronegativity - the affinity (attraction) for electrons. Highest: Fluorine 4.0 [See Table S]

Exception to 1.7 rule: METAL hydrides are ionic! ex. NaH

Diatomic molecules are considered to have NONPOLAR covalent bonding. i.e. N2 N=N this should be a
triple bond

Helium & Hydrogen need only 2 electrons to fill its outer shell. All the other elements need 8 electrons.

Coordinate covalent bonds - a covalent bond where both of the electrons are donated by one of the
elements. [See diagram #8]. Usually found in polyatomic ions.




 Ions -atoms with a charge. K+ and Cl- have the same # of electron (18) since formation of ions are
caused by the loss or gaining of ELECTRONS.

Ionization energy: the amount of energy required to remove the outermost electron from an element. See
Table S

 Ionic solids: high melting & boiling point; hard; do not conduct electricity UNLESS dissolved in water
-or- in molten (liquid) form.

 Metallic solids: mobile electrons, conductors in solids phase, malleable, ductile, only metal that is a
liquid at room temp is Hg
                                                                                                            7
Molecular solids (intermolecular forces): can be held together by
van der Waals forces; low melting & boiling points; poor conductors; are soft. ex. Sugar C6H12O6, H2 ,
CCl4

 Van deer Waals forces - attractive forces that exist between non polar molecules particles. They
increase when particles increase in mass
Get closer together
It's like GRAVITY!

 Hydrogen bonds - attractive for between molecules that contain hydrogen and atoms of small atomic
radius and HIGH ELECTRONEGATIVITIES. N ,O, F. These bonds result in some compounds having
higher boiling points than expected. Water has a high boiling point.

Polar molecules - molecules in which there is a localization of charge that causes part of the molecule to
have a slightly positive end and a slightly negative. [See diagram #9] These are usually asymmetrical
molecules.
ex. H2O, HF, NH3




Nonpolar molecule - there may still be localization of charge but there is no NET movement of electrons
in any particular direction. This is a tug of war where no one wins. There is an equal distribution of
charge. The molecules show symmetry.




Formula writing and naming compounds remember you
have to know how to use the stock system of naming for
transition metals that have more than one oxidation state.
Formula writing - use the crisscross method. [See diagram #10]

                                                                                                         8
Stock system uses roman numerals to identify how many electrons the transition metal is losing.

Transition metals have multiple oxidation states.

Binary compound- only two elements

Metal and non metal, ending of the name‖ide‖

Polyatomic ions keep the name. table E


Periodic table
Periodic law - states that elements are arranged on the periodic table according to their atomic numbers
and chemical properties.

Elements are classified in 3 categories
Metals - left of stairs
Nonmetals - right of the stairs
Metalloids - touching the stairs

Trends - as you go from left to right across the table in a period
Metallic character decreases
Atomic radius decreases [See Table S]
Ionization energy increases [See Table S]
Electro negativity increases [See Table S]
As you go down a group
Metallic character increases
Atomic radius increases [See Table S]

Atomic radius decreases as you go across a period since there is an increase of nuclear charge (# of
protons) which pulls the electrons in closer thereby shrinking the size of the atom.

Ionization energy decreases [See Table S]

Metalloids - have both metal and nonmetal properties. Contact the "staircase".
                                                                                                           9
Group 1 metal - alkali metals; strongest bases; form +1 ions

Group 2 - alkaline earth metals; form +2 ions

Group 18 NON metal - inert or noble gases; generally non-reactive. Kr and Xe can form some bonds in
the laboratory.

Group 17 -halogens - contain elements in ALL three phases. F & Cl are gases, Br is a liquid and I is a
solid

Elements in the same period fill up the SAME principle energy levels
Elements in the same groups have the same # of valence electrons
The most active metals are in the lower left corner.
The most active nonmetals are in the upper right corner.
The MOST active elements make for the MOST stable compounds! i.e. RbF
Monatomic molecules (one atom) He, Ne, Ar, Kr, Rn
D                                                  2,O2,N2,Cl2,Br2,I2,F2
Transition elements -
Produce COLORED SOLUTIONS.
found in the middle of periodic table (Groups 3 to 12)
emit color in flame test as electrons fall back DOWN from the excited state.

1.. Dmitri Mendeleev (1834-1907) was a Russian chemist who was one of the first persons who tried to
put the known elements (about 70 during Mendeleev's time) in a chart which showed a pattern of
properties. Mendeleev arranged the known elements in vertical columns by increasing atomic mass and
noticed patterns in their properties. From his arrangement Mendeleev could predict the properties of
elements which had not yet been discovered.

Henry Moseley (1887-1915) arranged the elements in order of increasing atomic number (positive
charge). This is the way that the modern periodic table is arranged today.

2 . The elements are arranged in order of increasing atomic number across the periods (horizontal rows).
The periods are arranged so that the elements in the vertical columns (groups or families) have similar
properties. This causes the properties of elements to change as you move horizontally from group to
group across a period.

Note that there are 7 periods - one for each energy level (principle quantum number).

3. The periodic law states that when the elements are arranged according to increasing atomic number
there is a periodic pattern in their physical and chemical properties.

 4. Although the periodic table is arranged by increasing atomic number (which indicates the number of
protons), the electrons configuration is what really determines the physical and chemical properties of
the elements. The periodic table can be divided into four groups based on electron configuration :

              The Noble gases (Group 0) Group 18 - have their outermost s and p orbitals filled which
               creates a stable and non-reactive (inert) element.
              The representative elements - - have their s and p orbitals being filled. These include :
                   o Group 1 - Li, Na, K etc. - all very reactive with one electron in the outer s orbital
                                                                                                         10
                   o  Group 2 - Be, Mg, Ca etc. - all quite reactive with 2 electrons filling their outer s
                      orbital
                  o Group 13 - Aluminum group - 3 electrons in outer energy level (2s and 1p)
                      properties vary from metallic to metalloid
                  o Group 14 - Carbon group - 4 electrons in outer energy level (2s and 2p) -
                      properties vary from nonmetallic to metalloid to metallic down the group
                  o Group 15 - Nitrogen group - 5 electrons in outer energy level (2s and 3p) -
                      properties vary from nonmetallic to metalloid to metallic
                  o Group 16 - Oxygen group - 6 electrons in outer energy level (2s and 4p) -
                      properties vary from nonmetallic to metalloid
                  o Group 17 - Halogens - all have 7 electrons in the outer energy level (2s and 5p) -
                      properties vary from nonmetallic to metalloid. Very reactive due to the outer
                      energy level being almost filled.
              The transition metals - elements whose d orbitals are being filled - found in the "d-block."
               These are also called the Group B elements
              The Inner transition metals - These are the Lanthanide and Actinide series, element
               whose f orbitals are being filled.

5. The s, p, d, and f groups can be identified on the diagram below. The f block (inner transition metals)
is usually shown separated and below the rest of the table.




6. Periodicity is the property of having periodic properties. The periodic table shows periodicity in the
following properties :

              Atomic size - atoms of elements tend to increase as you go down a group (due to a
               greater number of energy levels) and atoms of elements tend to decrease in size across a
               period (greater positive nuclear charge which draws in electrons -energy levels are
               constant across a period).
           
              Ionization energy - the energy required to remove an electron from the gaseous state of
               the atom. Ionization energy decreases as you go down a group due to the outer electrons
               being further from the positive charge of the nucleus and being shielded from the nucleus'
               positive charge by the inner energy levels. Ionization energy increases as you move
               across a period. This is due to the increase in nuclear charge without the increase in
               number of energy levels.
              Electron affinity - this is the energy change associated with the addition of an electron to
               a gaseous atom. This trend is not as consistent as the others, but in general electron
               affinity decreases down a group and increases across a period.
              Ionic size - The size of ions increases as you go down a group, and decreases as you
               move across a period for the metals and for the nonmetals, for the same reasons as atomic
               size. However, the metallic ions are positive (have lost electrons, which make up the
                                                                                                          11
               space or size of the atom) and are much smaller than the negative nonmetallic ions (have
               gained electrons which create the volume of atoms or ions).
              Electronegativity -the tendency of an atom to gain an electron(s) when combining with
               another element. Electronegativity decreases down a group (due to shielding) and
               increases as you move across a period (due to the increase in nuclear charge).

7. The three major forces affecting periodicity are :

              Nuclear charge - the greater the number of protons in the nucleus, the greater the positive
               charge and the stronger the electrons are held.
              Shielding - the effect of inner energy levels reducing the strength of the nuclear charge on
               the electrons in the outer energy levels.
              Electron configuration - atoms are most stable when their outer orbitals are filled
               (especially the s and p orbitals). This causes the Noble gases to be inert.

8. The representative groups of elements :

              Noble gases - Group 18, helium, neon, argon, krypton, xenon and radon
                  o inert (unreactive) because of stable electron configuration (filled s and p orbitals)
                  o helium is used in weather balloons
                  o helium and neon are used to create artificial, unreactive environments (less
                       soluble than nitrogen and therefore less likely to cause the bends
                  o other noble gases are used to create unreactive environments in flashbulbs or
                       aluminum welding
              Alkali metals - Group 1, lithium, sodium, potassium, rubidium, cesium and francium
                  o very reactive (one electron away from a filled s and p orbital)
                  o low density
                  o low melting point
                  o good electrical conductivity
                  o react with water to form strong bases (sodium hydroxide, lithium hydroxide etc.)
              Alkaline earth elements - Group 2, beryllium, magnesium, calcium, strontium, barium
               and radium
                  o very reactive (2 electrons away from a filled s and p orbital)
                  o react with water to form hydroxides
                  o used to form metal alloys
              Aluminum group - Group 13 - 3 electrons in outer energy level (2s and 1p) properties
               vary from metallic to metalloid
                  o aluminum is the most useful metal of this group being lightweight and strong to
                       make boats, aircraft etc.
              Group 14 - Carbon group - 4 electrons in outer energy level (2s and 2p) - properties vary
               from nonmetallic to metalloid to metallic down the group
                  o diamond and graphite are forms of pure carbon
                  o silicon and germanium are semiconductors used in electronics
                  o tin and lead are useful metals
              Group 15 - Nitrogen group - 5 electrons in outer energy level (2s and 3p) - properties
               vary from nonmetallic to metalloid to metallic
                  o nitrogen and phosphorus are elements necessary to form proteins and nucleic
                       acids in living things
              Group 16 - Oxygen group - 6 electrons in outer energy level (2s and 4p) - properties vary
               from nonmetallic to metalloid
                                                                                                        12
                      o  oxygen is the most abundant element on the earth
                      o  sulfur has many industrial uses (sulfuric acid is the most widely used industrial
                         chemical)
                 Group 17 - Halogens - all have 7 electrons in the outer energy level (2s and 5p) -
                  properties vary from nonmetallic to metalloid. Very reactive due to the outer energy level
                  being almost filled.
                      o iodine is used as an antiseptic
                      o chlorine is a bleaching agent and disinfecting agent
                      o fluorine, as the fluoride ion, is used to maintain the health of our teeth
                      o fluorine is used to make teflon

   8. The accepted theory which explains the origin of the elements and all matter in the universe is the Big
   Bang theory. It states that the universe began with an explosion of tremendous energy. This energy was
   converted into matter, according to Einstein's equation E=mc2. At first all matter was in the form of
   quarks. As the universe expanded it cooled allowing matter to condense and form the lightest elements
   first and then the heavier elements.

   VALENCE ELECTRONS\
   1. Valence electrons are the electrons in the highest energy level of an atom. For example, in the calcium
   atom (electron configuration 1s22s22p63s23p64s2) the 4s2 electrons are the valence electrons. In the
   titanium atom (electron configuration 1s22s22p63s23p64s23d2) The 4s2 electrons are still the valence
   electrons -they are in the highest energy level. In the phosphorus atom (electron configuration
   1s22s22p63s23p3) the 3s23p3 are the valence electrons.

   The valence electron number of the representative elements are :

                 Group 1 - 1 valence electron
                 Group 2 - 2 valence electrons
                 Group 13 - 3 valence electrons
                 Group 14 - 4 valence electrons
                 Group 15 - 5 valence electrons
                 Group 16 - 6 valence electrons
                 Group 17 - 7 valence electrons
                 Group 18 - 8 valence electrons

2. Dot formulas use dots surrounding the symbol of the element to represent the valence electrons. The dot
   formulas for period 2 and 3 would appear as follows.




   Note electrons are usually shown as far apart as possible -they have the same charge and therefore repel
   each other.
                                                                                                          13
3. The Noble gases (Group 18) have a stable electron configuration (s2p6) with 8 electrons filling the outer s
   and p orbitals. This stability comes from the low energy state of this configuration and also accounts for
   the low reactivity of these elements (most elements react with other elements to get to a lower, more
   stable energy state).

For example the halogens (Group 17) have 7 valence electrons (s2p5) and want to gain one electron to get
  the low energy, stable electron configuration of the noble gases.

The elements in group 16 (s2p4) want to gain 2 electrons to get the low energy, stable electron configuration
  of the noble gases.

The Group 1 elements (s1) want to lose their outer electron to empty their outer shell and get a stable
  electron configuration. For example if sodium (1s22s22p63s1) loses its 3s1 electron it will have filled s
  and p orbitals in its outer energy level.

4. Gilbert Lewis, in 1916, proposed the octet rule : Atoms react by changing their number of electrons so as
   to acquire the stable electron configuration of a noble gas (s2p6).

5. An exception to the octet rule is the electron configuration of helium. Helium(1s2) is a noble gas, only it
   has only one orbital, the s orbital. It is filled and therefore stable and elements close to it (lithium,
   beryllium and sometimes hydrogen) try to acquire its electron configuration by losing or gaining
   electrons).

6. The pseudo noble-gas electron configuration has the outer three orbitals filled, the s, p and d- s2p6d10 (18
   electrons total) and so is fairly stable. Elements that attain this electron configuration are at the right side
   of the transition metals (d-block).

7. An ion is a charged atom that is formed by the gaining or losing of electrons. When an atom gains or
   loses electrons it is no longer neutral because the number of electrons (negative charges) and protons
   (positive charges) are not equal. For example when a sodium atom loses an electron to get the noble gas
   electron configuration of neon it gets a charge of +1and becomes the sodium ion, because it now has 11
   protons (+) and only 10 electrons(-). When a chlorine atom gains one electron to get the noble gas
   electron configuration of argon it gets a charge of -1 and becomes the chloride ion, because it has 17
   protons (+) and 18 electrons (-).

   Cations are positive ions and are formed by the losing of electrons.

   Anions are negative ions and are formed by the gaining of electrons.

8. The charge on an ion depends on the number of electrons gained or lost. An element that loses two
   electrons becomes an ion with a +2 charge. An element which gains three electrons becomes an ion with
   a -3 charge.




                                                                                                                14
Stoichiometry and mathematics in chemistry

   1. 1 Mole = 22.4 liters of a gas at S.T.P.
                    = 6.02 X 1023 PARTICLES
          =THE FORMULA MASS IN GRAMS, this is the
molar mass
use table T
calculate formula mass
Mass of a hydrate
% composition
% error
number of moles
ppm
% comp volume

mole conversions
2 step math problems

Use the mole conversion chart Draw it here




                                       moles




Gas Laws (all temperatures must be in Kelvin)
STP table A
densities table S

Avogadro's Law - equal volumes of gases contain equal # of molecules

Boyles give formula and draw graph




                                                                       15
Charles give formula and draw graph




Gay Lussac give formula and draw graph




Combined gas Law




Vapor pressure- the pressure exerted by a vapor above a liquid. Liquids that evaporate easily , have high
vapor pressures. Table H




SOLUTIONS use table F and G
Molarity = # of moles of solute
           Liters of solvent


Solution - homogeneous mixture (evenly mixed)

Unsaturated solution - holds less solute than the maximum

Saturated - holds the exact amount of solute the solvent can hold

Super-saturated - holds more than the maximum amount of solute

Concentrated solution - holds a large amount of solute

Dilute solution - holds a little amount of solute



                                                                                                      16
Solubility of a solid- (ability to dissolve) generally increases as temperature increases.
Use solubility curve and rules on the reference tables

Solubility of a gas increase as temperature decreases and pressure increases. Think of when soda goes
flat (CO2 escapes)

Boiling point elevation - for every mole of substance dissolved in solution the boiling point increase by
0.520 °C

Freezing point depression - for every mole of substance dissolved in solution the freezing point
decreases by 1.86 °C




When figuring out boiling point elevation and freezing point depression keep in mind that electrolytes
(molecules that split into ions) create more moles in solution than the
lose both s & d electrons & therefore have multiple oxidation states

Van der Waals forces increase as you go down a group since the size of the atom increase. This causes
the boiling and melting points to increases as well. Remember this when you get to ORGANIC
chemistry.
# of moles = given mass (grams)
                Gram molecular mass (add up masses from periodic table)
This formula
                Is on table T




                                                                                                         17
CHEMICAL EQUATIONS
Balancing equations
Informal rules. Balance the metals , non-metal, carbon , hydrogen, oxygen

Types of reactions

Synthesis


Decomposition


Single replacement


Double replacement



Unit VI - Kinetics and equilibrium
Table I gives the Delta H . look at the bottom of the table. It says Minus sign indicates an
EXOTHERMIC reaction.

   1. Heat of reaction (H)- the difference between the potential energy of the reactants and the
      products
      (does NOT change with the addition of a catalyst)
   2. Diagrams of exothermic and endothermic reactions. [See diagrams # 13 & 14]




                                                                                                    18
   3. Exothermic reactions  release energy, (H) = -, products formed are MORE stable
      compounds than the reactants
   4. Endothermic reactions  absorb energy, (H) = +, products formed are LESS stable
      compounds than the reactants
   5. If the heat is listed on the right side (with the products) the reactions is exothermic.
   6. If the heat is listed on the left side (with the products) the reactions is endothermic.
   7. Factors effecting the reactions rate
           a. Catalyst - speeds up the reaction by reducing the activation energy needed to start a
              reaction. A catalyst does NOT effect the heat of reaction or the potential energy of the
              products or the reactants.
           b. Increasing the concentration of one of the substances  shifts the equilibrium away
              from the increase to the other side of the reaction while decreasing the concentration of
              ALL of the other compounds on the side of the increase.
           c. increase in temperature  shifts the equilibrium away from the heat. Favors the
              endothermic reaction.
           d. Increase in pressure  shifts the equilibrium to the side with the least number of moles.
           e. Increase in surface area  increases the reaction rate in both directions {like pounding
              it into a powder]
   8. Entropy --the randomness of a system. If + then there is an increase in entropy or Randomness
      and if - then there is a decrease.
   9. Order of increasing entropy: solids liquids gas

Equilibrium – the rate of the forward reaction =’s the rate of the reverse reaction
Le Chatelier’s Principle
Factors affecting shift
       1. concentration
       2. temperature
       3. pressure


-Acids and Bases And salts (all salts are ionic
compounds) Use tables K , L , M
Properties of Salts -----ionic compounds
   1. metal bonded to a non-metal
                                                                                                     19
   2. conductivity—only conducts electricity when in the aqueous or liquid form
      why does it only conduct electricity in these forms?
      Ionic solids have crystalline structure, the ions are held together by
      the ionic bond which is the transfer of electron(s) from the metal to the
      non- metal.
      What happens when the ionic solid is placed in water?
      The ions dissociate in the water because of the polarity of the water.

Any substance that can conduct electricity is called an electrolyte.

   3. Electrolyte - a compound that breaks into ions in solution or when melted. Usually ionically
      bonded.
   4. Non-electrolyte - a compound that does not break into ions in solution or when melted.
      Covalently bonded
   5. Arhennius theory of Acids
         a. Acid  gives of a H+ ion, as the ONLY positive ion

                    HCl     H+      +   Cl -

                    HNO3      H+ +      NO3-

                    H2SO4     2H+ + SO42-

   6. Arhennius theory of Bases

           a. Base  gives off an OH- ion as the only negative ion

                    NaOH         Na+ + OH-

                    Ba(OH)2  Ba2+ +        2OH-

   7. Bronsted-Lowry Theory of acids and bases

          b. Acid = proton donor (losses H+ )
          c. Base = proton acceptor (gains H+)



   8. Salt - a metal combined with a nonmetal [ex. NaCl, Na is the metal & Cl is the nonmetal]

   9 . Organic compounds- begins with C. i.e. C6H12O6 - usually NOT electrolytes.       Except
   organic acids [functional group –COOH]

   10. Properties of Acids

          a.   Turns blue litmus red
          b.   pH less than 7.0
          c.   Reacts with metals (below H on chart N) to form salt and H2 gas
          d.   Taste sour

                                                                                                 20
       e. Reacts with base to form salt and water (neutralization)
       f. The more they ionize, the better they conduct electricity
       g. They contain more H+ (H3O+) than (OH-)

11. Properties of bases

       a.   Turns red litmus blue, pink in phenolthalein
       b.   pH greater than 7.0
       c.   Reacts with acids - neutralization
       d.   Taste bitter
       e.   Feel slippery
       f.   The more they ionize, the better they conduct electricity
       g.   They contain more OH- than H+

12. Ionization Constant of water (Chart M) = Kw [H+] x [OH-] = 1 x 10-14. Use this to figure
out pH. [See diagram #16]




13.   pH scale - a measure of the concentration of Hydrogen ions (H+)

      the only concern is the number of hydrogen ions compared to the number

      hydroxide ions.

      When the concentration of hydrogen ions is greater then the number of
                                                                                            21
         Hydroxide ions the solution is acidic and has a low pH.

         When the concentration of the hydroxide ions is greater then the

          Concentration of the hydrogen ions the pH is high and the solution is

         Basic.    pH 7 is neutral, the concentration of H+ and OH- is equal

           0 ---1---2---3---4---5---6---7---8---9---10---11---12---13---14

               strong acids                     strong base    

  14.Neutralization

          a. Acid + Base  salt + water
          b. H+ + OH-  H2O (net reaction)

In neutralization, moles of acid and moles of base must be equal.

Formula for titration Table T

Indicators - color changes tell if acidic or basic Table M

Common acids and bases Table K and L


It should be noted that Group IA and IIA are strong bases when combined with OH; Bases [OH
combined with a metal] get weaker as you move across the periodic table from left to right.


- Redox       and electrochemistry (Table J)
   1. Know the rules for determining the oxidation states.
   2. Sum of the oxidation states in a neutral atom must always equal ZERO.
   3. Oxidation - loss of electrons causes the oxidation # to increase (LEO)
   4. Reduction - gaining of electrons causes the oxidation # to decrease.(GER)
   5. To have a Redox reaction there must be a change in oxidation # and you CANNOT have
      oxidation without having reduction.
   6. Spectator ions - ions that are not involved in being reduced or oxidized.
   7. Hydrogen is used as the standard on which the entire table is based.
   8. Electrochemical cell - [See diagram # 19], spontaneous, electrons flow to better reducer, salt
      bridge allows for the migration of ions in BOTH directions to sustain the reaction. Cathode
      is (+) electrode & the anode is the (-) electrode.




                                                                                                 22
9. Electrolytic cell - need a battery to get going, Anode is (+) electrode & the cathode is the (-)
    electrode.
10. Electroplating - plating occurs at the reduction or negative electrode. Car bumpers can be
    coated with protective metal in this manner. Mass increases at the site of plating and
    decreases at the oxidation or positive electrode. [See diagram # 20]




11. Balancing Redox equations - balance with respect to charge and mass. [See Diagram # 21]




                                                                                                 23
 12. The substance being reduced is considered to be the oxidizing agent.
 13. The substance being oxidized is considered to be the reducing agent.
 14. RED CAT; AN OX


Unit IX - Organic Chemistry
 1.   Hydrocarbons - contain only hydrogen and carbon
 2.   Homologous series - successive members differ by -CH2 groups.
 3.   Alkanes - CnH2n+2 contain ALL single bonds - saturated compounds, ending is - ane.
 4.   Alkenes - CnH2n contain one double bond - unsaturated compounds ending are - ene.
 5.   Alkynes - CnH2n-2 contain one triple bond - unsaturated compounds ending is - yne.



 6. Naming compounds - [See diagram #23]




 7. Alkyl radials (side groups) regular prefixes but they end in -yl.
 8. As the molecular mass of each of these homologous series increase so to do their boiling
     points and melting points due to an increase in the van der Waals forces.
 9. Properties of organic molecules - non-electrolytes, low boiling points. & melting points .,
     insoluble in polar solvents (like water), react slowly & are molecular in structure.
 10. Isomers - have the same chemical formula but a different structural formula, which means
     that they behave differently
 11. General formulas:

      a.Alcohol's  R –OH

      b.Organic acids  R-COOH

      c.Ester  R1 - COO - R2


                                                                                             24
       d.ketones  R1 – CO – R2 [the oxygen is double bonded to the carbon]

       e.ethers  > R1 – O – R2 [the oxygen connects two carbon chains]

       f.aldehydes R – COH [the oxygen is double bonded to the carbon]

   12. Types of alcohol's
          a. Monohydroxl - one OH group
          b. Dihydroxl - two OH groups
          c. Trihydroxyl - three OH groups. (Glycerol) [See diagram #24]




   13. Primary Alcohol's - OH group bonded to a carbon atom that is in turn bonded to ONE (or
       none) other carbon atom.
   14. Secondary Alcohol's - OH group bonded to a carbon atom that is in turn bonded to TWO
       other carbon atoms.
   15. Tertiary Alcohol's - OH group bonded to a carbon atom that is in turn bonded to THREE
       other carbon atoms.
   16. Organic reactions to know
          a. Addition - adds a pair of halogens to an unsaturated hydrocarbon. One product.
          b. Substitution - adds a halogen to a saturated hydrocarbon. Two Products.
          c. Esterfication - acid + alcohol  ester + water
          d. Fermentation - C6H12O6 alcohol + carbon dioxide
          e. Saponification - fat + base  soap + glycerol
          f. Polymerization - n (C2H4)  (C2H4)n
          g. Combustion - hydrocarbon + oxygen  carbon dioxide + water
          h. Cracking --> the separation of a polymer.


Nuclear chemistry
TRANSMUTATION
Natural Radioactivity is the spontaneous disintegration of the nucleus of certain atoms with the emission
of particles and energy to form more stable atoms.
This type of decay occurs among atoms with atomic numbers from 1 to 83 (those that have stable
isotopes)
Any element above 83 has no stable isotopes
Types of decay and decay products Table O
ALPHA decay produces ALPHA particles
Draw it here

                                                                                                      25
Look on Table O for symbols
Alpha particles have the same configuration of a helium nucleus
Alpha particles appear as products in a nuclear equation
Alpha particles have a +2 charge and a mass of 4
Alpha particles will be attracted to the negative end of a magnet


BETA decay produces a BETA particle
Look on table O for the symbols
Beta particles have the properties of a high speed electron
Beta particles have a negative charge and no mass
Beta particle will be attracted to the positive end of a magnet
In this process a NEUTRON is transformed into a PROTON and high speed ELECTRON




POSITRON EMISSION in this process an proton is transformed into a NEUTRON and a high speed
electron.


GAMMA radiation are rays and not particles.
Gamma rays do not have charge or mass
Gamma rays will not be attracted to any en d of a magnet
HALF LIFE            know how to use table N for half life
Half life is the amount of time required for half the sample to disintegrate


ARTIFICAL TRANSMUTATION
THE BOMBARBMENT OF A NUCLEUS WITH A PARTICLES
Fission Reactions—The splitting of heavy nuclei (large unstable one) into lighter nuclei (smaller more
stable )
Fusion Reactions- Is the energy releasing process in which two light nuclei FUSE together to form a
heavier one.

USES OF RADIOISOTOPES
1. Dating material C-14 15 dpm (disintegrations per minute)
                          half life 5730 years
                         good for dating material no older than 11,000 years
                         U- 238 used to date rocks and material older than 11,000 years

2. Chemical tracers – follow the path of a material
                      P-31 found in fertilizer
                      C-14 path of carbon in metabolic processes

3. Industrial Applications test thickness of materials, gamma rays

4. Medical application – they have short half lives and quickly emitted for the body
                                                                                                     26
                      I -131 thyroid
                      Co -60 gamma rays that are targeted at tumors
                      Gamma rays – irradiate food to kill bacteria
                      Co-60 and Cs 137 destroy anthrax
                      Tc-99 kills tumors
RADIATION RISKS

Can kill normal cells
Nuclear Power plants – the spent fuel rods need to be disposed of
                        Release of radioactivity in to air and water




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