SOL Items 1. Density formula D=m/v so for v= m/D2. Volume= L x W x H= units3 3. Area formula: L x W= units2 4. mass: m= D x v5. K=oC+273.166. oC: K-273.16= oC7. Percent Error= (your value-literature value)/literature value x 100 (Units are in %)8. Energy Conversion of calories to joules : 0.2390 calories (cal) = 1 Joule (J) and 1

cal=4.184 J9. Atmosphere to Pascal : 1atm=101,305 Pa10. Physical Properties to remember: mass, length, volume, color, density, malleability,

ductility, and conductivity, crystalline shape, melting point, boiling point, refractive index.

11. 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

12. 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

13. 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.

14. 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 mass15. Average Atomic Mass

Percent (in decimal form) times Atomic Mass for each one and then add the total16. Rutherford's Gold Foil Experiment:

o That the atom is mostly empty spaceo And that the nucleus is positive charged (because of protons) and contains

almost all of the mass of the atom.17. 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)

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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 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 valance electron number

2. 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 atomi. It decreases as it moves across a periodii. 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 cleanConduct heat and electricity wellSolid at room temperatureMost are also ductile and malleable (meaning they can be pounded into thin sheets and drawn into wire.)Chemically reactivePositively chargedThese 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 waterEasily lose a valence electron and form an ion with a +1 charge.Ends in sublevel s1

Alkaline Earth metals: Group 2Reactivity similar to Alkali metals but not as great+2 charge

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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. 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 levelAre silvery metals with relatively high melting pointsUsed 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 sublevelThey 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 Groupo 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 Groupo There are nonmetals (N and P), metalloids (As and Sb), and metals (Bi) o Each has 5 valence electrons and have many different propertieso Charge is -3

Group 16: The Oxygen Group or Chalogenso 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 electronegativityo Halogens make saltso Have 7 valence electrons and often tend to share one electron or gain one.

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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 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 NotesProperties

1. Ionic Bonds are formed by a cation (positive charged metal) bonded to an anion (negative charged nonmetal)

a. Metal loses one or more electronsb. 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 latticeb. Crystal lattice bonds are strong and take lots of energy to break the bonds

3. Usually poor conductors of electricitya. Because they are solid and rigid the ions can’t move freelyb. Only good conductors if dissolved in an aqueous solution where they become

electrolytes4. Don’t consist of molecules

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5. Strongly bonded6. Can form a salt

Ionic Terms

1. Ion : charged particlea. Anion : negatively charged ionb. Cation : Positively charged ion

2. Salt : An ionic compound that forms when a metal atom or a positive radical replaces the H of an acid.

a. Ex. NaClb. Salts are excellent conductors of electricity because they are brittle solids that

can easily be dissolved in an aqueous solution such as water

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 phosphateo 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-IVo 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 iono Ex. NH4Cl is Ammonium chloride

If the polyatomic ion has an oxyanions

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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 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

o The oxyanions with one less oxygen atom is named with the nonmetal and the suffix-ate

SO4-2 is sulfateo 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 hypochloriteo 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

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 electronso 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

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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 Basis of attractionIonic cations and anionsCovalent positive nuclei and shared electronsMetallic 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 forces3 types:

1. Dispersion forcesa. Sometimes called London dispersion forcesb. The force between oxygen molecules

i. Weak forces that result from temporary shifts in the density of electrons in electron clouds:

………

Attraction Temporary attraction Temporary attraction ←|- ←|-

2. Dipole –dipole: Attraction between oppositely charged regions of polar moleculeso Stronger than dispersion forces

The more polar the molecule, the stronger the force3. 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

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δ- δ+ δ+ δδ+

δ+ δ+ δ= δ+

δ- δ+

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 waterNH3 is ammonia

Examples:1. CO2 –Carbon dioxide2. CO-Carbon monoxide3. N2O4-dinitrogen tetroxide4. 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

Acids1. Liquids are tart, sour, or sharp tasting

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2. They conduct electricity (in solutions)-electrolytes3. They produce H2 gas4. Usually in liquid or gas form5. pH is 0-6.9

a. Strong acid -have a low pH and completely ionized in an aqueous solutioni. 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

6. They react to metals-corrosive

Bases1. Commonly found in solid form2. Chemical formula except for NH3 has OH on the end3. 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 baseii. Some are not very soluble in water

b. Weak base -react with water to form the OH- ion and conjugate acid of the base

4. Some are insoluble in water while others are soluble5. Slippery feel because bases react with oils in your skin-soaps and cleaning agents6. Are electrolytes

Three primary theories of acids and bases Theory Acid definition Base Definition

Arrhenius Any substance that releases H+ ions in water solution

Any substance that releases OH- ions in water solution

Bronstead-Lowery Any substance that donates a proton

Any substance that accepts a proton

Lewis Any substance that can accept an electron pair

Any substance that can donate an 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+ ionIII.) 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)

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1.) pH=-log[H+]2.) pOH=-log[OH-]3.) pH + pOH=144.) [H3O+]=10-pH use antilog5.) [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 Formula1st: Calculate the % composition of each element (If not given) and change % into grams2nd: Calculate Molar Mass of each element3rd: Determine simplest whole # ratio4th: Write Empirical Formula

Example #1The 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 element48.64 g of C X 1 mol of C/12.01g of C (atomic mass) =4.050 mol of C8.16 g of H X 1 mol of H/1.008g of H (atomic mass) =8.10 mol of H43.20 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 14.050/2.7 = 1.5 mol of C 8.10/2.7 = 3 mol of H2.7/2.7 = 1 mol of O

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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 C8.10/2.7 = 3 mol of H X 2=6 mol of H2.7/2.7 = 1 mol of O X2=2 mol of O

Step 3: Create Empirical Formula from moles in Step twoC3H6O2

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 C5.08 g of H X 1mol H/1.008g (atomic mass) of H= 5.04 mol of H54.24g 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 13.390/3.390 = 1 mol of C 5.04/3.390 = 1.5 mol of H3.390/3.390 = 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 C5.04/3.390 = 1.5 mol of H X 2=3 mol of H3.390/3.390 = 1 mol of O X 2 =2 mol of OStep 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 #1Find the percent composition by mass of Hydrogen and Oxygen in water. (Formula: H2O)Step one: Find individual mass (if not already stated in problem) of elements

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# 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 H1X16=16 g of O 18g/mol of H2O (just add individual amounts together)

Step three: Use Percent Composition formula to solve problem2.0 g of H/18.0g of H2O X 100= 11% of H16g of O/18g of H2O X 100= 89% of O

Example #2Find 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 1X23=23 g of Na# of Atoms of H: 1 (also called # of moles) 1 atom of H X Atomic mass of H 1X1=1 g of H# of Atoms of C: 1 (also called # of moles) 1 atom of C X Atomic mass of C 1X12=12 g of C# of Atoms of O: 3 (also called # of moles) 3 atoms of O X Atomic mass of O 3X16=48 g of O

Step three: Find Mass of Compound (if not already stated in the problem)1X23=23 g of Na1X1= 1 g of H1X12=12 g of C3X16=48 g of O 84g/mol o NaHCO3 (just add individual amounts together)

Step four: Use Percent Composition formula to solve problem23 g of Na/84.0g of NaHCO3 X 100= 27.3% of Na1.0 g of H/84.0g of NaHCO3 X 100= 1.190% of H12 g of C/84.0g of NaHCO3 X 100= 14.28% of C48g of O/84.0g of NaHCO3 X 100= 57% of O

31. Lewis Dot-1-8 dots according to valance electron #32. Molecular Molecules (drawn chemical structures)-VSEPR33. Naming types and Balancing Chemical Equations 34. Stoichiometry Formulas: Also include Limiting Reagent and Percent Yield

I. Moles to Moles: Given Moles x mole ratio

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II. 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 1023 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)

35. Thermochemistry and Chemical KineticsLaw 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

Entropy (S) is a measure of randomness or disorder in a system and is a thermodynamic property

Measuring Heat

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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 releasedii. C (sometimes seen as Cp) =the specific heat (also called Molar heat

capacity)iii. m=mass in giv. ΔT = change in temperature in oC ΔT

Lots of variationsi. n (number of moles)= q ÷ C ΔTii. n= mass (m) ÷ Molar mass (M) n=m ÷ Miii. c (calorie) =C ÷ Miv. ΔT =Tf (Final temp) – Ti (Initial Temp)v. ΔH (Change in Heat) = C x ΔTvi. q= n x C x ΔTvii. C = q ÷ n x ΔTviii. 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 Curves37. Phase Diagrams38. Gas Laws

P = Pressure T= Temperature V= Volume

1atm=101.3 kPa760 mm Hg =101.3 kPa so 1atm=760 mm Hg 760 torr=1 atm

Manometers are used to measure the pressure in a closed containerSTP: T =273.16 K and 1 atm of pressure or 101.3 kPa

I. Boyle’s Law: P1V1=P2V2 II. Charles’ Law: V1/T1=V2/T2III. Gay Lussac’s Law: P1/T1=P2/T2

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IV. Combined Gas Equation: P1V1/T1=P2V2/T2V. Dalton’s Law of Partial Pressures: PT (total) = P1 + P2 +etcVI. Ideal Gas Law: PV = nRT

o R = 8.314L x kPa / mol x K or R = 0.0821 atm x L / mol x Ko Amount of gas (n) = moleo M=Molar mass g/molo m=moles

39. Boiling point depression Formula: iKbm 5 steps1. Determine the moles of solute2. Find Molality

a. Add grams of solution together if neededb. convert g into kgc. 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 CaCl22nd: m=moles of solute/kg of solution 3.47moles/1.230kg of H2O=2.82 m3rd: Ionic so add subscripts CaCl2 1+2=34th: 3 x 0.51 x 2.82=4.3146 oC5th: 100 + 4.3146=104.3146 oC is ∆Kb

40. Freezing point depression Formula: iKfm 5 stepsa. Determine the moles of soluteb. Find Molality

i. Add grams of solution together if neededii. convert g into kgiii. 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 solution42. m (Molality)= mole of solute ٪ Kilograms of solution43. 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 5 5

44. Keq and Ksp problems45. Scientific Notation46. Significant Digits47. Conversion of Metric Units

Add: Chemical and Physical change

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Chemical Properties 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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