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SOL Items
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Density formula D=m/v so for v= m/D
Volume= L x W x H= units3
Area formula: L x W= units2
mass: m= D x v
K=oC+273.16
o
C: K-273.16= oC
Percent Error= (your value-literature value)/literature value x 100 (Units are in %)
Energy Conversion of calories to joules: 0.2390 calories (cal) = 1 Joule (J) and 1 cal=4.184 J
Atmosphere to Pascal: 1atm=101,305 Pa
Physical Properties to remember: mass, length, volume, color, density, malleability, ductility, and
conductivity, crystalline shape, melting point, boiling point, refractive index.
Accuracy vs. Precision: Accuracy-Refers to how close the measurement is to the actual value while
Precision refers to how close a set of measurements is together whether or not the measurements are
correct
Separation of Mixtures:
o Distillation-occurs when a liquid is boiled to produce a vapor that is then condensed again to
a liquid. This causes the solid substances that were originally dissolved to stay in the original
container and the water to go into a second receiving container.
o Chromatography-Involves a solid (stationary phase) and a liquid or gas (mobile phase). The
separation occurs because the liquid or gas has a faster rate than the solid.
i. Paper Chromatography-paper (solid) and a liquid are involved. The liquid travels up
the paper and separates according to the heaviness of the individual parts of the
liquid
Antoine Lavoisier a French Chemist (1743-1794)
 Proposed the Law of Conservation of Mass: in ordinary chemical reactions, matter
can be changed in many ways, but it cannot be created or destroyed.
Find on Periodic table: Atomic Number and Atomic Mass, and figure out Neutron #, Electron # and
charge is negative, and Proton # and charge is positive.
o Atomic Mass (Symbol Z) –Atomic number (Symbol A)=neutrons Z-A=Neutrons (neutrons
have no charge and are found in the nucleus with the protons)
o Note: Atomic number + neutrons =Atomic mass
Average Atomic Mass
Percent (in decimal form) times Atomic Mass for each one and then add the total
Rutherford's Gold Foil Experiment:
o That the atom is mostly empty space
o And that the nucleus is positive charged (because of protons) and contains almost all of the
mass of the atom.
Alpha radiation is radiation that was deflected toward the negatively charged plate alpha radiation.

Made up of 2 alpha () particles
 Each alpha particle contains 2 protons and 2 neutrons
 Has a 2+ charge
 Has a mass of 4 amu

Ex.: Ra- 88 (Radium-226) 86 Rn (radon-222) + 4 He (alpha particle)
(Exact model will be shown in class)
18. Beta radiation is radiation deflected toward the positively charged plate beta radiation.

Consist of fast moving electrons known as beta () particles.
 Each beta particle contains an electron with a -1 charge.
 Ex.: C-14/6  14N/7 + 0/+1e
19. Half Life Formula
 Amount Remaining= Original Amount of parent ÷ 2n
1
n=half-life
20. Electron Configuration-The arrangement of electrons in an atom. The order: 1s, 2s, 2p, 3s, 3p, 4s, 3d,
4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s
Rules:
1. Octet law: only a maximum of 8 electrons on the outermost shell (2 in the 1st level)-known as the
2.
valance electron number
The aufbau principle states that each electron occupies the lowest energy orbital available

In order of increasing energy, the sequence of energy sublevels within a principal energy level
is s (2 e-), p (6 e-), d (10 e-), and f (14 e-).
3. The Pauli exclusion principal states that a maximum to 2 electrons may occupy a single atomic
orbital, but only if the electron has opposite spins.
 Represented as ↑↓
21. Periodic Trends:
o Atomic Radius-Deals with the size of an atom
i. It decreases as it moves across a period
ii. Increases as it moves down the Group
o Electronegativity-The attraction it has to bond with other elements- F has the highest
i. It decreases as it moves down a group and increases as it moves across a period
22. Periodic table/Element items:
 Dmitri Mendeleev put together the 1st periodic table-Table not completely correct
 A period in the periodic table is all the elements in a horizontal row.
 A group in the periodic table is all the elements located in the same vertical column, which is assigned
a number from 1-18.
METALS


Usually shiny when smooth and clean
Conduct heat and electricity well
Solid at room temperature
Most are also ductile and malleable (meaning they can be pounded into thin sheets and drawn into
wire.)
Chemically reactive
Positively charged
These atoms have only a few electrons in the outer level.
Have a tendency to lose their electrons in the outermost level.
 Alkali metals: Group 1 except H
They react with water
Easily lose a valence electron and form an ion with a +1 charge.
Ends in sublevel s1
 Alkaline Earth metals: Group 2
Reactivity similar to Alkali metals but not as great
+2 charge
Ends in sublevel s2
 Aluminum Group (Sometimes called the Boron Group)
Group 13
+3 charge
Transition Elements for divided into 2 set:
1. Transition metals
 Any element in columns 3-12
 Has 2 electrons in the outer level - 4s2
 Elements with #'s 22-28 also have a 3d sublevel
 In all groups except 12, the d orbitals are only partially filled.
 Share properties such as electrical conductivity, luster, and malleability with other metals.
2
 They have magnetism
2.
Inner transition metals: Lanthanoid Series and Actinoid Series
 They are the highest energy electrons (f electrons) are inside the d sublevel and the outer level.
 Both have outer shells consisting of an s2 sublevel
 Lanthanoid Series- Lanthanum (57) to Ytterbium (70)
Electrons are added to the 4f sublevel instead of the sixth or outer level
Are silvery metals with relatively high melting points
Used extensively as phosphors, substances that emits light when struck by electrons.
 Actinoid Series-Actinium (89) to Nobelium (102)
This series have an increasing # of electrons in the 5f sublevel
They are all radioactive.
NONMETALS







These are usually gases or brittle solids at room temperature.
Dull appearance
Insulators
Outer electrons are held closely by the nucleus
Form negative ions (anions)
Have 5 or more electrons in the outer level than metals.
They often gain electrons or share their electrons in the outermost level.
 Group 14: Carbon Group
o Allotropes are found in this group: forms of an element in the same physical state that have
different structures and properties.
Ex. Carbon in the form of coal, Diamonds and graphite
o Silicates are silicon compounds bound to Oxygen, and each Si atom is surrounded by 4 O
atoms.
 Group 15: The Nitrogen and Phosphorus Group
o There are nonmetals (N and P), metalloids (As and Sb), and metals (Bi)
o Each has 5 valence electrons and have many different properties
o Charge is -3
 Group 16: The Oxygen Group or Chalogens
o Have 2 allotropes: Ozone, and O2
o Some are oxides known as amphoteric: Those that can produce either acidic or basic
solutions. Ex. Sulfur compounds like sulfuric acid (H2SO4)
o Charge is a -2
 The highly reactive Group 17 elements are called Halogens
o Fluorine is the most reactive element- Highly electronegativity
o Halogens make salts
o Have 7 valence electrons and often tend to share one electron or gain one.
o Have a 1- charge so they react with Group 1 the most
 The extremely nonreactive Group 18 and is known as Noble gases.
o All except Helium have 8 electrons in their outer level.
METALLOIDS




These are elements with physical and chemical properties of both metals and nonmetals.
Silicon and germanium are used a lot in making computer chips and solar cells.
Staircase elements between metals and nonmetals
Are often brittle solids
23. Oxidation Number
For metals
3





Positively charged
in Group 1-2: Same as Group number
in Groups 3-12: number varies
Ex. Hydrogen is in Group 1 and the Oxidation # is +1
Ex. Magnesium is in Group 2 and the Oxidation # is +2
For nonmetals
 Negatively charged
 Find valence electron number: 2nd number of group #
 Formula: -8+valence electron number
 Ex. Oxygen is in Group 16, -8 + 6= -2
 Ex. Nitrogen is in Group 15, -8 + 5= -3
 Ex. Chlorine is in Group 17, -8 + 7= -1
For Metalloids
 Positively charged
 Find the number the same way and nonmetals
 Ex. Carbon is in Group #14 and the Oxidation # is +4
24. Ionic Bond Notes
Properties
1. Ionic Bonds are formed by a cation (positive charged metal) bonded to an anion (negative charged
nonmetal)
a. Metal loses one or more electrons
b. Nonmetal gains one or more electrons
Note: Static electrical attraction is the basis for ionic bonds, because the positively charged ion
(cation) is attracted to the negatively charged ion (anion)
2. High Boiling and High Melting point
a. Forms a 3-D crystal lattice
b. Crystal lattice bonds are strong and take lots of energy to break the bonds
3. Usually poor conductors of electricity
a. Because they are solid and rigid the ions can’t move freely
b. Only good conductors if dissolved in an aqueous solution where they become electrolytes
4. Don’t consist of molecules
5. Strongly bonded
6. Can form a salt
Ionic Terms
1.
2.
Ion: charged particle
a. Anion: negatively charged ion
b. Cation: Positively charged ion
Salt: An ionic compound that forms when a metal atom or a positive radical replaces the H of an
acid.
a. Ex. NaCl
b. Salts are excellent conductors of electricity because they are brittle solids that can easily be
dissolved in an aqueous solution such as water
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Ionic Bonding Problems/Diagrams
 Na (+1) + Cl (-1) → NaCl
 Ca (+2) + Cl (-1) +Cl (-1) → CaCl2
25. Naming Ionic Compounds
Many Ionic compounds contain polyatomic ions: ions made up of more than one atom.
 Polyatomic ions exist as a unit, so never change the subscript
o If you have to balance an ionic compound with a polyatomic ion then ( ) and a subscript must
be written
 Ex. Ca (+2) and PO4 (-3) →Ca3(PO4)2 and named Calcium phosphate
o Most polyatomic ions are oxyanions
 Oxyanion is a polyatomic ion composed of an element, usually a nonmetal, bonded
to one or more oxygen atoms
o If a transitional metal and a polyatomic ion is involved
 Ex. Cu (+2) and NO3 (-1) →Cu(NO3)2 and named Copper (II) nitrate
 Note that transitional metals with varying oxidation numbers always have to
state which atom was used in the chemical compound whether or not a
polyatomic ion is used
Rules for Naming Ionic Compounds
 Name the cation (metal) first and the anion (nonmetal) second
 Monatomic cations use element name
 Monatomic anions take their name from the root element name plus the suffix –ide
o Ex. CsBr is Cesium bromide
 Determine oxidation numbers of transitional metals compounds before naming to determine Roman
number I-IV
o Ex. Fe2O3 is Iron (III)oxide
 Some transitional metals only have one charge

Cadmium: Cd+2
 Zinc: Zn+2
 If the compound has a polyatomic ion, simply name the ion
o Ex. NH4Cl is Ammonium chloride
 If the polyatomic ion has an oxyanions
o The ion with more oxygen atoms is named using the root of the nonmetal plus the suffix –ate
 Ex. NO3- is nitrate
 Ex. ClO3- is chlorate
 Ex. CO3-2 is carbonate
o The ion with fewer oxygen atoms is named using the root of the nonmetal plus the suffix –ite
 Ex. NO2- is nitrite
o
o
The oxyanions with the greatest number of oxygen atoms is named using the prefix per-, the
root of the nonmetal, and the suffix-ate
 ClO4- is perchlorate
 IO4- is periodate
 MnO4- is permanganate
The oxyanions with one less oxygen atom is named with the nonmetal and the suffix-ate
 SO4-2 is sulfate
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The oxyanions with two fewer oxygen atoms is named using the root of the nonmetal plus
the suffix –ite
 Ex. ClO2 – is chlorite
o The oxyanions with three fewer oxygen atoms is named using the prefix hypo-, the root of the
nonmetal, and the suffix –ite
 Ex. ClO – is hypochlorite
o Polyatomic ions with 2 transitional metal atoms include a Di-prefix
 Ex. H2PO4- is Dihydrogen phosphate
 Ex. Cr2O7-2 is Dichromate
o Some Hydrogen plus a polyatomic ion are named two ways
 Ex. HSO4- can be named bisulfate or Hydrogen sulfate
 Ex. HCO3- can be named bicarbonate or Hydrogen carbonate
26. Covalent Bonds
o
Properties
 Nonmetal + Nonmetal (usually)
 Most common type of bond
 Covalent bonds form molecules
 Form by sharing electrons
o The sharing of one pair of electrons is a single bond (X-X)
 Another name for single covalent bond is sigma bond symbolized by σ
 Sigma bonds form from the overlap of a s orbital with another s orbital, a s
orbital with a p orbital, or a p orbital with another p orbital
o The sharing of two pairs- double bond (X=X)
 Another name for multiple bonds is pi bond symbolized by π
 Pi bonds form when parallel orbitals overlap to share electrons
o A double covalent bond has one sigma and one pi bond
o The sharing of three pairs-triple bond (XΞX)
 A triple covalent bond has one sigma and 2 pi bonds
 Bond polarity explains the attraction between the sharing
o Nonpolar electrons are shared equally Ex. F-F (same electronegativity)
o Polar electrons are not shared evenly Ex. H-F (different electronegativity)
Intramolecular Forces in Bonds Table
Force
Ionic
Covalent
Metallic
Basis of attraction
cations and anions
positive nuclei and shared electrons
metal cations and mobile electrons
Intermolecular Forces
Intramolecular forces do not account for all attractions between particles. There are forces of attraction called
intermolecular forces.
o They can hold together identical particles or two different types of particles
o Also called van der Waals forces
3 types:
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1. Dispersion forces
a. Sometimes called London dispersion forces
b. The force between oxygen molecules
i. Weak forces that result from temporary shifts in the density of electrons in electron
clouds:
………
δ- δ+
δ+ δδ+
Temporary attraction
δ+ δ+ δ=
←|δ+
δ-
δ+
Attraction
Temporary attraction
←|-
2. Dipole –dipole: Attraction between oppositely charged regions of polar molecules
o Stronger than dispersion forces
 The more polar the molecule, the stronger the force
3. Hydrogen bonds: One special type of dipole-dipole dealing with hydrogen bonds
 Very strong intermolecular force that is formed with a H end and a F, O, or
N atom on the other dipole
Many physical properties of covalent molecular solids are due to intermolecular forces.
 The melting and boiling points are relatively lower than Ionic (that is why
salt doesn’t burn when you heat it but sugar will)
 Many are gases are vaporized at room temperature
 Hardness is also due to the intermolecular forces so covalent solids are soft
in comparison to ionic solids
Naming Molecular Compounds: Rules for Binary Molecular Compounds are similar to that of naming Ionic
compounds except the names include prefixes indicating the number of atoms in the molecule.
Numerial Prefixes
 Mono-1
 Di-2
 Tri-3
 Tetra-4
 Penta-5
 Hexa-6
 Hepta-7
 Octa-8
 Nona-9
 Deca-10
Exceptions:
H2O is water
NH3 is ammonia
Examples:
1. CO2 –Carbon dioxide
2. CO-Carbon monoxide
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3.
4.
N2O4-dinitrogen tetroxide
SCl6 –Sulfur hexachloride
27. Naming Acids and Bases
Binary acids are acids with only two elements.
 Prefix –hydro, stem of anion, and suffix –ic
 Exception is HN3: Hydroazoic acid, where the root – azo is used for nitrogen.
Ternary acids are acids that contain 3 elements.
 Usually no prefix is used and the suffix is –ic.
Exceptions:
 One less O than the most common : no prefix and suffix used is –ous
 Two less O than the most common: prefix hypo- and suffix –ous
 One more O than the most common: prefix per- and suffix –ic
 Ex. HClO3 is the most common: Chloric acid
 HClO2 has one less O so: Chlorous acid
 HClO has two less O so: Hypochlorous acid
 HClO4 has one more O than most common so: Perchloric acid
Ternary bases
 Arrhenius bases are composed of metallic, or positively charged ions and the negatively charged
hydroxide ion. Therefore, these bases are named by adding the word hydroxide to the name of the positive
ion.
 Ex. Sodium hydroxide is NaOH.
28. Characteristics of Acids and Bases
1.
2.
3.
4.
5.
6.
1.
2.
3.
4.
5.
6.
Acids
Liquids are tart, sour, or sharp tasting
They conduct electricity (in solutions)-electrolytes
They produce H2 gas
Usually in liquid or gas form
pH is 0-6.9
a. Strong acid-have a low pH and completely ionized in an aqueous solution
i. The closer the substance’s pH is zero the stronger the acid
b. Weak Acid-have pH closer to 6.9 and are only slightly ionized in an aqueous solution
They react to metals-corrosive
Bases
Commonly found in solid form
Chemical formula except for NH3 has OH on the end
pH range is 7.1-14
a. Strong base-dissociates completely into metal ions and OH- ions in aqueous solution
i. The closer the substance’s pH is 14 the stronger the base
ii. Some are not very soluble in water
b. Weak base-react with water to form the OH- ion and conjugate acid of the base
Some are insoluble in water while others are soluble
Slippery feel because bases react with oils in your skin-soaps and cleaning agents
Are electrolytes
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Theory
Arrhenius
Bronstead-Lowery
Lewis
Three primary theories of acids and bases
Acid definition
Base Definition
Any substance that releases H+
Any substance that releases OHions in water solution
ions in water solution
Any substance that donates a
Any substance that accepts a
proton
proton
Any substance that can accept an
Any substance that can donate an
electron pair
electron pair
Examples:
I.)
Arrhenius acid: HCl (g)→H+(aq) +Cl-(aq) Arrhenius base: NaOH (cr) →Na+ (aq) + OH-(aq)
II.)
Bronstead-Lowery: HCl (g) + H2O → H3O+(aq)
+
Cl-(aq)
Acid +base→
conjugate acid + conjugate base
 Conjugate acid-is the particle formed when a base gains a H+ ion
 Conjugate base-is the particle that remains when an acid has donated a H+ ion
 Conjugate acid-base pair-consists of 2 substances related by the gain or loss of a single H+ ion
III.)
Lewis: H3N: (Lewis base) + BF3 (Lewis acid) → H3N: BF3 (Product)
pH
 pH is a measurement of the H3O+ ion concentration of an acid or a base.
Problem Formulas: (Actual problem examples will be stated in class)
1.) pH=-log[H+]
2.) pOH=-log[OH-]
3.) pH + pOH=14
4.) [H3O+]=10-pH use antilog
5.) [OH-]=10-pH
6.) [OH-]=antilog (-pOH)
7.) Kw=[OH-] x [H+] which equals 1 x 10-14M
Titration
Titration-is a procedure used to bring a solution of a known concentration into a reaction with a solution of an
unknown concentration in order to determine the unknown concentration or the quantity of the solute in the
unknown.
 The point in the titration at which stoichiometrical equivalent quantities of reactants are brought
together –equivalence point
 An indicator can be used to show the end point of the titration, (at equivalence point)
o In acid-base titrations, the dyes used are colorless and only change to pink (basic) or blue
(acidic) at the end point
 Example dyes are phenolphthalein colorless for an acid and pink for a base, and
NaOH blue with acids and pink for base
 The neutralization reaction occurs between an acid and a metal OH (base) and
produces water and a salt (is a crystalline compound composed of the negative ion of
an acid and the positive ion of a base.)
29. Empirical and Molecular Formulas
Empirical Formula: The smallest whole number mole ratio of elements in a compound
 Assume that each percent by mass represents the mass of the element in a 100.00-g sample
How to Calculate Empirical Formula
1st: Calculate the % composition of each element (If not given) and change % into grams
2nd: Calculate Molar Mass of each element
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3rd: Determine simplest whole # ratio
4th: Write Empirical Formula
Example #1
The mass of C is 48.64g, the mass of H is 8.16g, and the mass of O is 43.20g. Find the Empirical Formula
(EF).
Step 1: Find Molar mass of each element
48.64 g of C X 1 mol of C/12.01g of C (atomic mass) =4.050 mol of C
8.16 g of H X 1 mol of H/1.008g of H (atomic mass) =8.10 mol of H
43.20 g of O X 1 mol of O/16.00g of O (atomic mass) =2.700 mol of O
Step 2: Determine simplest ratio by dividing the lowest amount of moles determined in step 1
4.050/2.7 = 1.5 mol of C
8.10/2.7 = 3 mol of H
2.7/2.7 = 1 mol of O
Then look at the three numbers of moles and determine the lowest number they can be multiplied by to get all
whole numbers. In this case the number is 2.
4.050/2.7 = 1.5 mol of C x2=3 mol of C
8.10/2.7 = 3 mol of H X 2=6 mol of H
2.7/2.7 = 1 mol of O X2=2 mol of O
Step 3: Create Empirical Formula from moles in Step two
C3H6O2
Example #2
Succinic acid is a substance produced by lichens. Chemical analysis indicates it is composed of 40.68% C,
5.08% H, and 54.24% oxygen and has a molar mass of 118.1g/mol. Determine the empirical formula for
succinic acid.
Step 1: Determine molar mass.
1st: Convert percentages into grams of elements.
2nd: Use molar mass formula to find moles.
40.68 g of C X 1mol C/12.01g (atomic mass) of C=
3.390 mol of C
5.08 g of H X 1mol H/1.008g (atomic mass) of H= 5.04 mol of H
54.24g of O X 1mol O/16.00g (atomic mass) of O=
3.390 mol of O
Step 2: Determine simplest ratio by dividing the lowest amount of moles determined in step 1
3.390/3.390 = 1 mol of C
5.04/3.390 = 1.5 mol of H
3.390/3.390 = 1 mol of O
Then look at the three numbers of moles and determine the lowest number they can be multiplied by to get all
whole numbers. In this case the number is 2.
3.390/3.390 = 1 mol of C X 2 =2 mol of C
10
5.04/3.390 = 1.5 mol of H X 2=3 mol of H
3.390/3.390 = 1 mol of O X 2 =2 mol of O
Step 3: Create Empirical Formula from moles in Step two: C2H3O2
30. Percent Composition
% Composition formula: Mass of element/ Mass of Compound X 100 = % of mass in grams
Note: You may only be given the name of the compound and not the formula. If so, you will have to use
your rules you learned from Ionic and Covalent Bonds.
Example #1
Find the percent composition by mass of Hydrogen and Oxygen in water. (Formula: H 2O)
Step one: Find individual mass (if not already stated in problem) of elements
# of Atoms of H: 2 (also called # of moles)
2 atoms of H X Atomic mass of H
2X1=2 g of H
# of Atoms of O: 1 (also called # of moles)
1 atom of O X Atomic mass of O
1X16=16 g of O
Step two: Find Mass of Compound (if not already stated in the problem)
2X1= 2 g of H
1X16=16 g of O
18g/mol of H2O (just add individual amounts together)
Step three: Use Percent Composition formula to solve problem
2.0 g of H/18.0g of H2O X 100= 11% of H
16g of O/18g of H2O
X 100= 89% of O
Example #2
Find the percent composition of each element in Sodium Hydrogen Carbonate.
Step one: Figure out chemical formula: NaHCO3
Step two: Find individual mass (if not already stated in problem) of elements
# of Atoms of Na: 1 (also called # of moles)
1 atom of Na X Atomic mass of Na
Na
# of Atoms of H: 1 (also called # of moles)
1 atom of H X Atomic mass of H
# of Atoms of C: 1 (also called # of moles)
1 atom of C X Atomic mass of C
# of Atoms of O: 3 (also called # of moles)
3 atoms of O X Atomic mass of O
Step three: Find Mass of Compound (if not already stated in the problem)
1X23=23 g of Na
1X1= 1 g of H
1X12=12 g of C
3X16=48 g of O
84g/mol o NaHCO3 (just add individual amounts together)
Step four: Use Percent Composition formula to solve problem
23 g of Na/84.0g of NaHCO3 X 100= 27.3% of Na
1.0 g of H/84.0g of NaHCO3 X 100= 1.190% of H
12 g of C/84.0g of NaHCO3 X 100= 14.28% of C
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1X23=23 g of
1X1=1 g of H
1X12=12 g of C
3X16=48 g of O
48g of O/84.0g of NaHCO3 X 100=
31.
32.
33.
34.
57% of O
Lewis Dot-1-8 dots according to valance electron #
Molecular Molecules (drawn chemical structures)-VSEPR
Naming types and Balancing Chemical Equations
Stoichiometry Formulas: Also include Limiting Reagent and Percent Yield
Moles to Moles: Given Moles x mole ratio
Moles to Mass: Given Moles x mole ratio x (Molar mass of unknown ÷ 1 mol of unknown)
o Note – Any Stoichiometry problem dealing with Mass must be in grams. Sometimes you
have to convert into Grams (In Mass to Mass problems)
III. Moles to Volume (using STP):
o Given Moles X Mole ratio X (22.4 L of unknown ÷1 mol of unknown)
IV. Mass to Moles: Given Mass x (1 mol of known ÷ Molar Mass of known) x Mole ratio
V. Mass to Mass: Given Mass x (1 mol of known ÷ Molar Mass of known) x Mole ratio x (Molar mass
of unknown ÷ 1 mol of unknown)
VI. Mass to STP Volume: Given Mass x (1 mol of known ÷ Molar Mass of known) x Mole ratio x (22.4
L of unknown ÷1 mol of unknown)
VII. Volume to Volume-Density:
o Given Volume x Density of given (g/L) x (1 mol of given÷ Molar mass of given ) x mole
ratio x (unknown Molar mass ÷ 1 mol of unknown) x Density of unknown (but one liter of
unknown ÷ grams)
VIII. Volume to Volume using just STP: Given volume x (1 mol of known ÷22.4 L) x mole ratio x (22.4 L
of unknown ÷ 1 mol of unknown)
IX. Molecules to Molecules: Given molecules x (1 mol of given ÷ 6.022 x 10 23 of given molecules) x mole
ratio x (6.022 x 1023 of unknown molecules ÷ 1 mol of unknown)
X. Molecules to Grams: Given molecules x (1 mol of given ÷ 6.022 x 1023 of given molecules) x mole
ratio x (molar mass of unknown ÷ 1 mol of unknown)
I.
II.
35. Thermochemistry and Chemical Kinetics
Law of Conservation of Energy: In any chemical reaction or physical process, energy can be converted from
one form to another, but neither can be created or destroyed
Heat-The energy transferred between objects that are at different temperatures
 It is an extensive property, which means that the amount of the energy transferred as heat by
a sample depends on the amount of the sample
Temperature-a measure of how hot (or cold) something is, specifically it is a measure of the average kinetic
energy in the particles of an object
 It is an intensive property, which means that the temperature of a sample does not depend on
the amount of the sample
Enthalpy- represented as H, is the total energy content of a sample.
 If pressure remains constant the enthalpy increases in a sample of matter equal to the energy
as heat that is received.
Molar Heat Capacity- (C) in a pure substance it is the energy as heat is needed to increase the temperature of
one mole of a substance by 1 Kelvin.
Calorimetry-the measurement of heat related constants, such as specific heat
Calorimeter-a device used to measure the heat absorbed or released in a chemical or physical change
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Entropy (S) is a measure of randomness or disorder in a system and is a thermodynamic property
Measuring Heat
1.
Calories- The amount of heat required to raise the temperature of one gram of pure water by one
degree Celsius.
 kcal = 1000 calories
 SI units of heat and energy is joules (J)
 One J=0.2390 cal
 One cal=4.184 J
 KJ=1000 Joules
2.
Specific Heat is the amount of heat required to raise the temperature of on g of that substance by one
degree C.
 Basic Equation q=C x m x ΔT
i. q= the heat absorbed or released
ii. C (sometimes seen as Cp) =the specific heat (also called Molar heat capacity)
iii. m=mass in g
iv. ΔT = change in temperature in oC ΔT
 Lots of variations
i. n (number of moles)= q ÷ C ΔT
ii. n= mass (m) ÷ Molar mass (M) n=m ÷ M
iii. c (calorie) =C ÷ M
iv. ΔT =Tf (Final temp) – Ti (Initial Temp)
v. ΔH (Change in Heat) = C x ΔT
vi. q= n x C x ΔT
vii. C = q ÷ n x ΔT
viii. ΔT = q ÷ n x C
Chemical Kinetics-The Study of Reaction Rate (Also see Unit 10 booklet and textbook-Chapter 16)
 Activation Energy-The minimum amount of energy required to start a chemical reaction
 Exothermic reactions- the products are lower energy level than the reactants (makes chemical
reactions rise in temperature) so the ΔH is negative
 Endothermic reactions- The energy of the products is greater than the reactants (chemical
reaction lowers in temperature) so the ΔH is positive
 Catalyst –speeds up a reaction by providing the reactants with an alternate pathway that lowers the
activation energy
 Inhibitor-slows down and can stop a reaction
36. Solubility Curves
37. Phase Diagrams
38. Gas Laws
P = Pressure T= Temperature V= Volume
1atm=101.3 kPa
760 mm Hg =101.3 kPa
 so 1atm=760 mm Hg
 760 torr=1 atm
Manometers are used to measure the pressure in a closed container
STP: T =273.16 K and 1 atm of pressure or 101.3 kPa
I.
Boyle’s Law:
P1V1=P2V2
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II.
III.
IV.
V.
VI.
Charles’ Law: V1/T1=V2/T2
Gay Lussac’s Law: P1/T1=P2/T2
Combined Gas Equation: P1V1/T1=P2V2/T2
Dalton’s Law of Partial Pressures: PT (total) = P1 + P2 +etc
Ideal Gas Law: PV = nRT
o R = 8.314L x kPa / mol x K or R = 0.0821 atm x L / mol x K
o Amount of gas (n) = mole
o M=Molar mass g/mol
o m=moles
39. Boiling point depression
Formula: iKbm 5 steps
1. Determine the moles of solute
2. Find Molality
a. Add grams of solution together if needed
b. convert g into kg
c. then find molality
3. Add up the # of ions present (add up subscripts if Ionic or put one for Covalent compound)
4. i (# of ions) x Kb of solvent x Molality = Kb
5. Change in Kb (∆Kb) = Boiling point of solvent + answer in step 4
Example: What is the expected boiling point of CaCl2 solution containing 385g of CaCl2 in 1230g of water?
(Kb of water is 0.51oC and the Boiling point of water is 100oC)
1st: 385g of CaCl2 x 1mole of CaCl2/110.8g of CaCl2 =3.47 moles of CaCl2
2nd: m=moles of solute/kg of solution 3.47moles/1.230kg of H2O=2.82 m
3rd: Ionic so add subscripts CaCl2 1+2=3
4th: 3 x 0.51 x 2.82=4.3146 oC
5th: 100 + 4.3146=104.3146 oC is ∆Kb
40. Freezing point depression Formula: iKfm 5 steps
a. Determine the moles of solute
b. Find Molality
i. Add grams of solution together if needed
ii. convert g into kg
iii. then find molality
c. Add up the # of ions present (add up subscripts if Ionic or put one for Covalent compound)
d. i (# of ions) x Kb of solvent x Molality = Kb
e. Change in Kf (∆Kb) = Freezing of solvent + answer in step 4
41. M (Molarity)=mole of solute ٪ Liter of solution
42. m (Molality)= mole of solute ٪ Kilograms of solution
43. Molarity with Dilution Formula
Formula: M1V1=M2V2 M=Molarity and V=Volume
 Convert Volume to Liters
Example: How much 16M HCl is needed to prepare 200mL of 5M solution?
Solve: 16M x .2L =5M x V2 (?)
3.2=5V2
so V2= .64L needed
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44.
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46.
47.
Keq and Ksp problems
Scientific Notation
Significant Digits
Conversion of Metric Units
Add:
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Chemical and Physical change
Chemical Properties
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Lab safety
Scientific method
Valence electrons
Actual pH scale (pH mentioned but not the scale itself)
Properties of gases
Heat of fusion ∆Hf x mass
Heat of vaporation ∆Hv x mass
Parts of solutions: solute (what is being dissolved) solvent (what is dissolving the solute)
Kinetic molecular theory
Average kinetic theory of water increases as the temperature increases energy
Gases have the highest entropy
Exothermic vs. endothermic Graph
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