Science Framework for California Public School

1
Science Framework for California Public School

• Mr. Gowers
• Chemistry

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Standard 1Atomic and Molecular Structure

• 1.The periodic table displays the elements in
  increasing atomic number and shows how
  periodicity of the physical and chemical
  properties of the elements relates to atomic
  structure.

3
Standard 1

• a. Students know how to relate the position of
  an element in the periodic table to its atomic
  number and atomic mass.
• The number of protons, not electrons or neutrons,
  determines the atomic number.
• Elements are arranged on the periodic table in
  order of increasing atomic mass.
• Isotopes differences in the number of neutrons
  of the same element.

Atomic number (Z) number of protons in nucleus
Mass number (A) number of protons number of
neutrons
atomic number (Z) number of neutrons
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Hydrogen
Deuterium
Tritium
5
Do You Understand Isotopes?
How many protons, neutrons, and electrons?
6 protons, 8 (14 - 6) neutrons, 6 electrons
How many protons, neutrons, and electrons are in?
6 protons, 5 (11 - 6) neutrons, 6 electrons
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Standard 1

1. Students know how to use the periodic table to
   identify alkali metals, alkaline earth metals and
   transition metals, trends in ionization energy,
   electronegativity, and the relative size of ions
   and atoms.

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1. Alkali metals Group 1(ie. Na K) Soft white
   and extremely reactive.
2. Alkaline earth metals Group 2 (ie. Mg Ca).
3. Transition Metals Groups 3 12 (Common metals
   like iron gold) Have electrons in the d
   orbitals.
4. Electronegativity a measure of the ability of an
   atom of an element to attract electrons toward
   itself in a chemical bond (0 4).
5. Ionization energy energy it takes to remove an
   electron from the atom.
6. Both Electronegativity and Ionization energy
   increase from bottom to top and left to right.
7. Atomic ionic sizes increase from top to bottom
   and right to left. (Exception with full or ½ full
   subshells)
8. Cations are smaller than their neutral form and
   Anions are larger than their neutral state.

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Students know how to use the periodic table to
identify metals, semimetals, nonmetals, and
halogens.
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Standard 1

• Students know how to use the periodic table to
  determine the number of electrons available for
  bonding.
• Valence electrons.
• Valence electrons are equal to the Group number
  (ie. Group 1 has 1 valence electron).
• Useful in determining how many electrons are
  involved in bonding (ie. MgCl2).
• e. Students know the nucleus of the atom is much
  smaller than the atom yet very dense (contains
  most of its mass).

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Standard 2

• Students know atoms combine to form molecules by
  sharing electrons to form covalent or metallic
  bonds or by exchanging electrons to form ionic
  bonds.
• Covalent shared electrons electrons overlap
  between the two atomic orbitals.
• Metals valence electrons are not localized
  (delocalized) allowing them to move between
  orbitals of adjacent metals (conduct
  electricity).
• Ionic transfer of electrons.
• Polar covalent unevenly shared electrons.
• Octet Rule

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Valence electrons are the outer shell electrons
of an atom. The valence electrons are the
electrons that participate in chemical bonding.
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The Ionic Bond
He
Ne
1s22s1
1s22s22p5
1s2
1s22s22p6
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A covalent bond is a chemical bond in which two
or more electrons are shared by two atoms.
Lewis structure of F2
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Lewis structure of water


Double bond two atoms share two pairs of
electrons
or
Triple bond two atoms share three pairs of
electrons
or
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Lengths of Covalent Bonds
Bond Type Bond Length (pm)
C-C 154
C?C 133
C?C 120
C-N 143
C?N 138
C?N 116
Bond Lengths Triple bond lt Double Bond lt Single
Bond
Bond Strength Triple bond gt Double Bond gt Single
Bond
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Polar covalent bond or polar bond is a covalent
bond with greater electron density around one of
the two atoms
electron rich region
electron poor region
e- rich
e- poor
d
d-
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Classification of bonds by difference in
electronegativity
Difference
Bond Type
0
Covalent
? 2.0
Ionic
0 lt and lt2.0
Polar Covalent
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Cs 0.7
Cl 3.0
3.0 0.7 2.3
Ionic
H 2.1
S 2.5
2.5 2.1 0.4
Polar Covalent
N 3.0
N 3.0
3.0 3.0 0
Covalent
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Standard 2

• Students know chemical bonds between atoms in
  molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2,
  and many large biological molecules are covalent.
• Organic biological molecules consist primarily
  of carbon, oxygen, hydrogen, nitrogen. These
  elements share valence electrons to form bonds so
  that they have electron configurations like their
  nearest noble gas. (Octet Rule)
• Noble gases inert gases (column 18).

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Standard 2

• Students know salt crystals (NaCl) are repeating
  patterns of positive and negative ions held
  together by electrostatic attraction.
• Lattice energy the energy that holds ionic
  compounds together.
• Cations positive charge.
• Anions negative charge.

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Standard 2

• Students know the atoms and molecules in liquids
  move in a random pattern relative to one another
  b/c the intermolecular forces are too weak to
  hold the atoms or molecules in a solid form.
• When enough energy is added to the solid, the
  kinetic energy of the atoms molecules increases
  sufficiently to overcome the attractive forces
  between the particles, they break (melting
  point) which forms a liquid (disorder). The
  particles in the liquid are free to move although
  they remain in contact w/ each other.

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Standard 2

• Students know how to draw Lewis dot structures.
• 1) Lewis dot structures show how valence
  electrons and covalent bonds are arranged between
  atoms in a molecule.

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VI. Lewis Structures Representation of
_________ bonding in molecules.
covalent
Group e- dot of bonds lone pairs
1
2
13
14
15
16
17
18

x
1
0

x

2
0


x
3
0




x
4
0



x
3
1






x
2
2






x
1
3








x
0
4




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Writing Lewis Structures
Technique 1. Find the total number of _______
electrons in the molecule. 2. Add __ for each __
charge subtract __ for each __ charge. 3. Write
skeleton structure with _____ bonds (generally
the element which makes the greatest number of
bonds will be the ______ atom). 4. Place
remaining valence electrons around the ____
elements until octets are filled. 5. Place
valence electrons around the ______ atom until
octet is formed. 6. If you run out of electrons
(look at the total you calculated), then ______
or ______ bonds will be formed. 7. The goal is
to draw a structure that uses the correct number
of ________ which has an ____ around each
atom. 8. Hydrogen is an ________ to the octet
rule. It can only form _ bond and have a total
of _ electrons (to become like ______). 9.
Elements will be most _____ if they contain the
number of bonds and the number of lone pairs as
shown in the above table.
valence
1
1

single
central
outer
central
double
triple
electrons
octet
exception
1
2
helium
stable
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Step 1 N is less electronegative than F, put N
in center
Step 2 Count valence electrons N - 5 (2s22p3)
and F - 7 (2s22p5)
5 (3 x 7) 26 valence electrons total
electrons in your budget
Step 3 Draw single bonds between N and F atoms
and complete octets on N and F
atoms.
Step 4 - Check. Did you spend your budget? In
other wards are of e- (electrons) in structure
equal to number of total of valence e- ?
3 single bonds (3x2) 10 lone pairs (10x2) 26
valence electrons
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Standard 3

• Students know how to describe chemical reactions
  by writing balanced equations.
• Nomenclature.
• Tips for balancing equations
• a) Number of atoms of products must equal number
  of atoms of reactants.
• b) Coefficients are whole numbers written at the
  front of the substances.

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Standard 3

• c) All atoms are balanced by the coefficients.
• d) Subscripts are NOT changed.
• e) Keep polyatomic ions together as a unit if not
  changed from reactants to products.
• f) Balance single elements last.
• g) Use the even/odd rule.
• h) If an element is in multiple compounds,
  balance that element last.

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Chemical Nomenclature

• Ionic Compounds
• often a metal nonmetal
• anion (nonmetal), add ide to element name

BaCl2
barium chloride
K2O
potassium oxide
Mg(OH)2
magnesium hydroxide
KNO3
potassium nitrate
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• Transition metal ionic compounds
• indicate charge on metal with Roman numerals

iron(II) chloride
FeCl2
2 Cl- -2 so Fe is 2
FeCl3
3 Cl- -3 so Fe is 3
iron(III) chloride
Cr2S3
3 S-2 -6 so Cr is 3 (6/2)
chromium(III) sulfide
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Molecular Compounds
HI
hydrogen iodide
NF3
nitrogen trifluoride
SO2
sulfur dioxide
N2Cl4
dinitrogen tetrachloride
NO2
nitrogen dioxide
N2O
dinitrogen monoxide
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Balancing Chemical Equations

1. Write the correct formula(s) for the reactants on
   the left side and the correct formula(s) for the
   product(s) on the right side of the equation.

Ethane reacts with oxygen to form carbon dioxide
and water

1. Change the numbers in front of the formulas
   (coefficients) to make the number of atoms of
   each element the same on both sides of the
   equation. Do not change the subscripts.

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Balancing Chemical Equations

1. Start by balancing those elements that appear in
   only one reactant and one product.

start with C or H but not O
multiply CO2 by 2
multiply H2O by 3
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Balancing Chemical Equations

1. Balance those elements that appear in two or more
   reactants or products. (Double your odds)

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Balancing Chemical Equations

1. Check to make sure that you have the same number
   of each type of atom on both sides of the
   equation.

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Standard 3

• Students know the quantity one mole is set by
  defining one mole of carbon-12 atoms to have a
  mass of exactly 12 grams.
• Mole the number of atoms in 12 grams of pure
  carbon-12.
• Students know one mole equals 6.022 x 1023
  particles (atoms or molecules) (Avogadros
  number).

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Standard 3

• d. Students know how to determine the molar
  mass of a molecule from its chemical formula and
  a table of atomic masses and how to convert the
  mass of a molecular substance to moles, number of
  particles, or volume of gas at standard
  temperature and pressure. Road Map!!!!
• e. Students know how to calculate the masses
  of reactants and products in a chemical reaction
  from the mass of one of the reactants or
  products. Road Map!!!!

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C. Road Map (Memorize)

• A B ? C D

Mass (g) A
Mass (g) D
Mol to mol ratio!!!
Particles A Molecules A Atoms A
Particles D Molecules D Atoms D
6.022 x 1023
6.022 x 1023
Moles D
Moles A
Coefficients From Balanced
22.4 L
22.4 L
Chemical Equations
Volume of gas D
Volume of gas A
Coefficients
(Same T P)
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How many H atoms are in 72.5 g of C3H8O ?
1 mol C3H8O (3 x 12) (8 x 1) 16 60 g C3H8O
1 mol C3H8O molecules 8 mol H atoms
1 mol H 6.022 x 1023 atoms H
72.5 g C3H8O
5.82 x 1024 atoms H
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Methanol burns in air according to the equation
If 209 g of methanol are used up in the
combustion, what mass of water is produced?
molar mass CH3OH
molar mass H2O
coefficients chemical equation
209 g CH3OH
235 g H2O
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Theoretical Yield is the amount of product that
would result if all the limiting reagent reacted.
Actual Yield is the amount of product actually
obtained from a reaction.
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Standard 4 Gases and Their Properties

• The kinetic molecular theory describes the motion
  of atoms and molecules and explains the
  properties of gases.
• Students know the random motion of molecules and
  their collisions with a surface create the
  observable pressure on that surface.
• Fluids consist of molecules that freely move, but
  intermolecular forces hold the atoms or molecules
  close to each other.

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Standard 4 Gases and Their Properties

• 2) Gases consist of tiny particles (atoms or
  molecules) spaced far apart from each other and
  move freely at high speeds, near the speed of
  sound.
• 3) Pressure a force per unit area.
• 4) Pressure is caused by the collisions of atoms
  or molecules with the walls of the container.
• 5) Pressure in water increases with depth, and
  pressure in air decreases with altitude.

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Standard 4

• Students know the random motion of molecules
  explains the diffusion of gases.
• Heavier gases have a slower rate of diffusion.
• Students know how to apply the gas laws to
  relations between the pressure, temperature, and
  volume of any amount of an ideal gas or any
  mixture of ideal gases.
• Boyles Law P1V1 P2V2 (Inverse)
• Charle's Law V1/T1 V2/T2 (Direct)
• Gay-Lussacs Law P1/T1 P2/T2 (Direct)
• Combined Gas Law P1V1/T1 P2V2/T2

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Standard 4

• Students know the values and meanings of STP (0
  oC/ 273.15 K and 1 atm or 760 mmHg).
• Students know how to convert between Celsius and
  Kelvin. (K oC 273.15)
• Students know there is no temperature lower than
  0 Kelvin.
• The greater the atomic and molecular motion, the
  greater the observed temperature of a substance.
• 0 Kelvin or -273.15 oC all motion stops.
• g. Ideal gas law PV nRT

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Physical Characteristics of Gases

• Gases assume the volume and shape of their
  containers.
• Gases are the most compressible state of matter.
• Gases will mix evenly and completely when
  confined to the same container.
• Gases have much lower densities than liquids and
  solids.

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A sample of chlorine gas occupies a volume of 946
mL at a pressure of 726 mmHg. What is the
pressure of the gas (in mmHg) if the volume is
reduced at constant temperature to 154 mL?
P1 x V1 P2 x V2
P1 726 mmHg
P2 ?
V1 946 mL
V2 154 mL
P2
4460 mmHg
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A sample of carbon monoxide gas occupies 3.20 L
at 125 0C. At what temperature will the gas
occupy a volume of 1.54 L if the pressure remains
constant?
V1/T1 V2/T2
V1 3.20 L
V2 1.54 L
T1 398.15 K
T2 ?
T1 125 (0C) 273.15 (K) 398.15 K
T2
192 K
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Standard 5 Acids Bases

• Acids, bases, salts are three classes of
  compounds that form ions in water solutions.
• Students know the observable properties of acids,
  bases, salt solutions.
• Acids taste sour, change color of litmus paper
  from blue to red, react w/ certain metals to
  produce hydrogen gas, react w/ strong bases to
  produce water salt.
• Bases taste bitter, slippery, change litmus
  paper from red to blue, react w/ strong acids
  to produce water salt.

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Standard 5

• Students know acids are hydrogen-ion-donating and
  bases are hydrogen-ion-accepting substance.
• Bronsted-Lowry acid-base definition acids donate
  hydrogen ions, bases accept hydrogen ions.
• Nonmetals in the 1st 2nd rows of the periodic
  table easily dissociate to produce hydrogen ions
  b/c they have high electronegativities compared
  to hydrogen.

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A Brønsted acid is a proton donor A Brønsted base
is a proton acceptor
acid
base
acid
base
conjugate base
conjugate acid
acid
base
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Standard 5

• Students know strong acids bases fully
  dissociate and weak acids and bases partially
  dissociate.
• 1) Acids dissociate (ionize) by donating hydrogen
  ions.
• 2) Bases dissociate (ionize) by donating
  hydroxide ions.
• 3) Nearly complete dissociation is strong
  partial dissociation is weak.
• 4) The strength of an acid or base can vary
  depending on temperature concentration.

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Standard 5

• c. Students know how to use the pH scale to
  characterize acid and base solutions.
• 1) The pH scale measures the concentration of
  hydrogen ions (H) in solution.
• 2) The scale is logarithmic at pH 2, the
  concentration of H is 10 xs greater than it is
  at pH 3.
• 3) The pH scale below 0 (very acidic) to above
  14 (very basic). (0-14)
• 4) pH values less than 7 are acidic, greater than
  7 are basic, and 7 is neutral.

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pH A Measure of Acidity
pH -log H
Solution Is
At 250C
H OH-
neutral
H 1 x 10-7
pH 7
H gt OH-
acidic
H gt 1 x 10-7
pH lt 7
H lt OH-
basic
H lt 1 x 10-7
pH gt 7
pH 6
0.000001
1 x 10-6
pH 13
0.0000000000001
1 x 10-13
0.01
1 x 10-2
pH 2
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Standard 6 Solutions

• Solutions are homogeneous mixtures of two or more
  substances.
• Students know the definitions of solute and
  solvent.
• 1. Solute Substance present in smaller
  amounts.
• 2. Solvent Substance present in larger amounts
• Students know how to describe the dissolving
  process at the molecular level by using the
  concept of random molecular motion. LIKE
  DISSOLVES LIKE
• When a solid is in contact w/ a liquid, at least
  some dissolution occurs.
• When salts dissolve in water, positive and
  negative ions are separate and surrounded by
  polar water molecules.

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Standard 6

• Students know temperature, pressure, and surface
  area affect the dissolving process.
• In a liquid solvent, solubility of gases and
  solids is a function of temperature.
• Increasing temperature usually increases
  solubility of solid solutes but always decreases
  the solubility of gaseous solutes.
• The solubility of a gas in a liquid is directly
  proportional to pressure.
• Solubility describes only how much solute will
  dissolve at equilibrium, not how quickly this
  process occurs.

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Standard 6

• Students know how to calculate the concentration
  of a solute in terms of grams per liter,
  molarity, parts per million, and percent
  composition.
• Grams per liter represents the mass of solute
  divided by the volume of solution.
• Molarity describes moles of solute divided by
  liters of solution.
• Parts per million is a ratio of one part of
  solute to one million parts of solvent (dilute
  solutions). (Units mg/L) or (x 1 million (106))
• Percent composition is the ratio of one part of
  solute to one hundred parts of solvent. (x 100)

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Standard 7 Chemical Thermodynamics

• Energy is exchanged or transformed in all
  chemical reactions and physical changes of
  matter.
• Students know how to describe temperature and
  heat flow in terms of the motion of molecules (or
  atoms).
• Temperature a measure of the average kinetic
  energy of molecular motion in a sample.
• Heat energy transferred from a sample at higher
  temperature to one at lower temperature.
• System (rxn boundaries) vs Surroundings (outside
  boundaries).

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Standard 7

• b. Students know chemical processes can either
  release (exothermic) or absorb (endothermic)
  thermal energy.
• Breaking bonds requires energy(absorbs energy)
• Making bonds releases energy.
• c. Students know energy is released when a
  material condenses or freezes and is absorbed
  when a material evaporates or melts.
• Evaporation/Melting absorb energy
• Condensation/Freezing release energy

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Exothermic process is any process that gives off
heat transfers thermal energy from the system
to the surroundings.
Endothermic process is any process in which heat
has to be supplied to the system from the
surroundings.
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Standard 7

• d. Student know how to solve problems involving
  heat flow and temperature changes, using known
  values of specific heat and latent heat of phase
  change.
• Specific heat energy needed to change the
  temperature of one gram of substance one degree
  Celsius.
• Latent heat of fusion (Enthalpy of fusion) Hfus
• Latent heat of vaporization (Enthalpy of
  vaporization) Hvap

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• 2.      Determine the amount of heat needed (kJ)
  to completely vaporize 25.4 g of solid ice at
  26.0?C and heat the vapor to 108?C.

108
(5)
(4)
100.0
(3)
(2)
0
(1)
-26.0
Q
(mc?T)(1)
(mHfus)(2)
(mc?T)(3)
(mHvap)(4)
(mc?T)(5)
Q
(25.4 g)(2.06 J/goC)(26.0 oC)
(25.4 g)(334 J/g)
(25.4 g)(4.184 J/goC)(100.0 oC)
(25.4 g)(2260 J/g)
(25.4 g)(2.02 J/goC)(8 oC)
Q
78286 J
80 kJ
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Standard 8 Reaction Rates

• Chemical reaction rates depend on factors that
  influence the frequency of collision of reactant
  molecules.
• Students know the rate of reaction is the
  decrease in concentration of reactants or the
  increase in concentration of products with time.
• Reaction rate the rate of decrease in
  concentration of reactants or the rate increase
  in concentration of products.
• A balanced equation expresses that the
  concentration of reactants must decrease, the
  concentration of products must increase in
  proportion to their mole ratios.

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• F. Reaction rate A positive quantity that
  expresses how the concentration (_________) of a
  reactant or product changes with ______. X
  concentration of X (M)
• Units of Rate _____________
• 1. In general
• rate
• 2. Example
• rate

Molarity
time
mol/Ls
aA bB ? cC dD
N2O5 ? NO2 O2
2
4
Reactants decrease (?) over time, products
increase () over time.
coefficients
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Standard 8

• b. Students know how reaction rates depend on
  such factors as concentration, temperature, and
  pressure.
• They increase the number of collisions in turn
  increasing the rate of reaction.
• Pressure only increases reaction rate for gases.
• c. Students know the role a catalyst plays in
  increasing the reaction rate.
• Lowers activation energyincreasing the rate of
  reaction without being consumed.
• Enzymes biological catalyst.
• Catalysts are used in automobile exhaust systems
  to reduce the emission of smog-producing unburned
  hydrocarbons.

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Standard 9 Chemical Equilibrium

• Equilibrium is a dynamic process.
• Concentrations remain constant.
• Stresses cause chemical equations (rxns) to
  shift.
• Chemical equilibrium is a dynamic process at the
  molecular level.
• Students know how to use Le Chateliers principle
  to predict the effect of changes in
  concentration, temperature, pressure.
• If an equilibrium system is stressed it will
  respond to partially undo the stress.
• b. Students know equilibrium is established when
  forward and reverse reaction rates are equal.
• Overall concentrations of each reactant and
  product remain constant over time.

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4.63 x 10-3
Equilibrium Expression
Equilibrium Will
K gtgt 1
Lie to the right
Favor products
K ltlt 1
Lie to the left
Favor reactants
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• Le Chateliers Principle When a system at
  equilibrium is disturbed by applying a ______, a
  new ____________ position is attained to _______
  the stress.
• 1. Temperature effects on equilibrium
• Example
• (a) STRESS Raise temperature (addition of
  ______)
• Equilibrium is shifted ____.
• (b) STRESS Lower temperature (removal of
  ______)
• Equilibrium is shifted ______.

stress
equilibrium
relieve
(exothermic)
NO2 ? N2O4 58.8 kJ
2
heat
left
heat
right
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• 2. Pressure effects on equilibrium When
  pressure is __________, the stress is relieved by
  favoring the reaction with _______ gas molecules
  (fewer gas molecules ______ pressure).
• Example
• STRESS Increase pressure
• Equilibrium is shifted _______.
• (b) STRESS Decrease pressure
• Equilibrium is shifted ______.

increased
fewer
lower
NO2 (g) ? N2O4 (g)
2
2 molecules on left, 1 molecule on the right
(lower pressure)
right
(fewer gas molecules)
left
(greater of gas molecules)
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• 3. Concentration effects on equilibrium
• Example
• STRESS Add NH3
• Equilibrium is shifted _____________.
• STRESS Add H2O
• Equilibrium is shifted _____________.

?Acid/Base indicator
Ni(H2O)62 6 NH3 ? Ni(NH3)62 6 H2O
blue
green
Need to remove NH3
Right (blue)
Need to remove H2O
Left (green)
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Standard 10 Organic and Biochemistry

• The bonding characteristics of carbon allow the
  formation of many different organic molecules of
  varied sizes, shapes, and chemical properties and
  provide the biochemical basis of life.
• Students know large molecules (polymers), such as
  proteins, nucleic acids, and starch, are formed
  by repetitive combinations of simple subunits
  called monomers.
• Ex. Starch is made from a number of simple sugar
  molecules (glucose) joined together.

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Standard 10

• b. Students know the bonding characteristics of
  carbon that result in the formation of a large
  variety of structures ranging from simple
  hydrocarbons to complex polymers and biological
  molecules.
• Carbon can form 4 bonds.
• 2) Carbon can form single, double, triple
  bonds which determine the geometry of the
  molecules.

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Standard 10

• 3) Hydrocarbons (Carbon Hydrogen only) ie.
  methane (CH4) ethane (C2H6)
• 4) Biological molecules protein
• Manufactured polymers polyester, nylon,
  polyethylene.
• Polymers A polymer is a compound with a
  repeating unit, called a monomers and contains a
  high molar mass.

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Standard 10

• c. Students know amino acids are the building
  blocks of proteins.
• Proteins large single-stranded polymers often
  made up of thousands of relatively small subunits
  called amino acids.
• Peptide bonds attach amino acids.
• Amino acids vary in composition giving them
  different shapes and functions (R-group).
• DNA is the blueprint for building proteins.

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Standard 11 Nuclear Processes

• Nuclear processes are those in which an atomic
  nucleus changes, including radioactive decay of
  naturally occurring and human-made isotopes,
  nuclear fission, and nuclear fusion.
• Students know protons and neutrons in the nucleus
  are held together by nuclear forces that overcome
  the electromagnetic repulsion between the
  protons. (Nuclear binding energy)

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Standard 11 Nuclear Processes

• b. Students know the energy release per gram of
  material is much larger in nuclear fusion or
  fission reactions than in chemical reactions.
  The change in mass (E mc2) is small but
  significant in nuclear reactions.
• Fusion 2 nuclei come together to form a heavier
  nucleus.
• Fission a heavy nucleus splits to form two
  lighter nuclei.
• Nucleon term for a proton or neutron.
• Fusion/Fission produce one million times more
  energy than chemical reactions.

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Atomic number (Z) number of protons in nucleus
Mass number (A) number of protons number of
neutrons
atomic number (Z) number of neutrons
A
1
1
0
0
4
Z
1
0
-1
1
2
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Standard 11

• c. Students know some naturally occurring
  isotopes of elements are radioactive, as are
  isotopes formed in nuclear reactions.
• Isotopes atoms with the same number of protons
  but a different number of neutrons.
• (C-12, C-13, C-14)
• 2) Parent isotopes less stable isotopes of one
  element that undergo radioactive decay,
  transforming to more stable isotopes called
  daughter products.

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Standard 11

• d. Students know the three most common forms of
  radioactive decay (alpha, beta, gamma) and know
  how the nucleus changes in each type of decay.
• Radioactive isotopes transform to more stable
  isotopes, emitting particles from the nucleus.
• Alpha emit helium-4, beta emit electrons or
  positrons, gamma emits high-energy
  electromagnetic.
• Alpha decay forms isotopes w/ 2 less protons and
  2 less neutrons.
• Beta decay forms elements w/ the same of
  nucleons but one less or more protons.
• Gamma decay does not change the number of
  nucleons in the nucleus but lowers the energy
  state.

84
212Po decays by alpha emission. Write the
balanced nuclear equation for the decay of 212Po.
212 4 A
A 208
84 2 Z
Z 82
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Nuclear Stability and Radioactive Decay
Beta decay
Decrease of neutrons by 1
Increase of protons by 1
Gamma decay
86
n/p too large
beta decay
n/p too small
positron decay or electron capture
87
Standard 11

• e. Student know alpha, beta, gamma radiation
  produce different amounts and kinds of damage in
  matter and have different penetrations.
• They are ionizing radiation which can ionize as
  many as ½ million atoms.
• Alpha particles shortest range and can penetrate
  a few millimeters of paper.
• Beta particles longer ranges and can penetrate
  several centimeters of aluminum.
• Gamma rays can penetrate matter up to several
  meters of lead.
• These 3 types of radiation interact w/ matter by
  losing energy and ionizing surrounding atoms.

88
Observations
Hypothesis
Summarizes the results of many observations and
experiments.
89
Important Points KNOW THE FOLLOWING!

• Ionization energy is lowest in group 1.
• Nuclear Binding energy Nuclear forces holding
  the nucleus
• PV nRT
• Temperature Average Kinetic energy
• pH scale 0 (acid) to 14 (base)
• Gas molecules move randomly diffusion
• Radioactive gives of alpha, beta gamma
• 1 mole 6.02 x 1023
• Equilibrium forward and reverse same rate
• Le Chatliers Principle Temp/Pressure/Concentrat
  ion
• Ionic bond Cation anion electrostatic
• Road Map
• Periodic table Lft to Rt increase atomic or
  of protons.

90
Contd

• Lewis structure full octet
• ? H -- strong acid
• Nucleus is small but dense
• Gas Laws PV V/T P/T
• Atom becomes an ion by gaining or losing
  electrons
• Electronegativity High difference is ionic Low
  diff is covlaent
• Endothermic rxn absorbs energy from surroundings
  (feels cold)
• Exothermic rxn release energy into the
  surroundings (feels hot)
• Acids are sour and react w/most metals to produce
  H2(g).
• Increased collisions lead to increased pressure.
• Catalyst lows activation energy
• Parts Per Million (ppm) mg/L ( is parts per
  100).
• Acids turn litmus paper from blue to red Blue to
  Red Acid

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