MATTER and ENERGY

ELEMENT cannot be broken down or decomposed by chemical means.

Allotropes: different forms of the same element in the same state (different prop.)ex: O2 and O3 , or carbon (graphite) and carbon (diamond)

COMPOUND 2 or more elements combined chemically. Can be broken down by chemical means (e.g., electrolysis or combustion)

SUBSTANCE An element or a compound, (s), (l) & (g)

solids : definite volume and shape (tightly packed and organized)

liquids: definite volume, but not shape (loosely packed and slightly organized)

gases: no definite volume and no definite shape (spread out and disorganized)

MIXTURE: impure substances made from combinations of pure substances. Can be separated by physical means.

Homogeneous mixtures: uniform composition. Also called SOLUTIONS

Aqueous solution (aq): a homogeneous mixture in water Can be separated into parts by: DISTILLATION (using differences in boiling points)

Heterogeneous: unequal distribution of matter

Can be separated into parts by: FILTRATION

ENERGY (Joules) The ability to do work

KINETIC ENERGY: energy of motion. TEMPERATURE: average K.E. of a system

POTENTIAL ENERGY: (chemical) energy stored in bonds (physical) energy of phase changes

ENDOTHERMIC CHANGES:energy is absorbed / energy is a reactantEXOTHERMIC CHANGES: energy is released / energy is a product

PHASE CHANGESsublimation: (s) + heat --> (g)deposition: (g) --> (s) + heat

melting: (s) + heat --> (l)freezing: (l) --> (s) + heat

boiling / vaporizing: (l) + heat --> (g)condensing: (g) --> (l) + heat

ENERGY CHANGES

If there is a change in temperature, use q = m x C x ΔT (C of water is on Table B, 4.18 J/g˚C)

If there is NO change in temperature, it is a PHASE CHANGE

melting / freezing, use q = m x Hf (Hf of water is on Table B, 334 J/g)

boiling / condensing, use q = m x Hv (Hf of water is on Table B, 2260 J/g)

UNIFORM HEATING CURVES

minutes phase physical change energy changes0 - 3 100% (s) temperature increases KE increases3 - 11 mix of (s) and (l) melting at 70 PE increases, KE constant11 - 13 100% (l) temperature increases KE increases13 - 23 mix of (l) and (g) boiling at 160 PE increases, KE constant23 - 25 100% (g) temperature increases KE increases

boiling plateau is longer than the melting plateau, therefore the Hv is greater than Hf

Topic 2: ATOMIC STRUCTURE

Development of atomic theory took the work of different scientists over many years:

1808 1891 1911 1920’shard sphere --> plum pudding model --> mostly empty space -->wave-mechanical model theory electrons evenly distributed electrons outside

orbitals are 3D regions in a positive “jelly” dense, + nucleus of

electron probability

Dalton J.J. Thompson E.Rutherford Bohr and physicists

Experimental Methods

CATHODE-RAY TUBE: cathode ray bends toward positive plate, provides evidence of the existence of electrons.

GOLD-FOIL EXPERIMENT: Positively charged alpha particles shot at thin gold foil.

OBSERVATION CONCLUSIONMost pass through without any deflection. The atom is mostly empty space.Some are deflected. There is a small positive nucleus (which

deflected the positive alpha particles)

SUBATOMICPARTICLE LOCATION MASS CHARGEproton nucleus 1 amu +1neutron nucleus 1 amu 0electron outside nucleus 1/1836 amu -1

Atomic Number = protonsMass Number = protons + neutronsNuclear Charge = # of protonsOverall Charge = protons - electrons

In a neutral atom, protons = electrons

Isotopic Notation: 73Li+1

Mass = 7 Atomic Number = 3 Overall Charge = +1

Protons = atomic # = 3

Neutrons = mass # - protons = 4

Electrons = 2

Isotopes: Atoms with the same number of protons, but different numbers of neutrons

isotope protons nuclear charge neutrons 14

6C 6 +6 812

6C 6 +6 6_____________________________________________________

Weighted Average of Isotopes:

To calculate average isotopic mass: multiply the mass number by its relative abundance. Then, add all products together

Ex: What is the average mass of 75% 35Cl and 25% 37Cl?Answer: (75%) x (35) + (25%) x (37) = 35.5 amu

100%

The Weighted Average of an Element’s Isotopic Mass is equal to the Atomic Mass.electron configuration: the arrangement of electrons around the nucleus.

Ground State: lowest possible energy level for an atom. Can be found on Periodic Ta-ble

Excited State: higher energy level. An electron from a lower level has made the transi-tion to a higher level (example: from 1st to 2nd)

Configuration Carbon Bromine

Ground state 2 - 4 2 - 8 - 18 - 7

Excited state 2 - 3 - 1 2 - 8 - 18 - 6 - 1

valence electrons = electrons in the outermost energy level of the ground state configuration.Carbon has 4, Bromine has 7

NUCLEAR CHEMISTRY

Nuclear Emanations: alpha (positive), beta (negative), gamma (neutral) (see table O for notations)

Unstable Nuclei decay into more Stable Nuclei

Natural Transmutations:· Alpha Decay : One nucleus decays into two products—one of which is alpha (2

4He)· Beta Decay : One nucleus decays into two products—one of which is beta (-1

0e)

Predicting the Products of Decay14C --> ___________ + ____________

rad.emission new element from Table N

146C --> 0

-1e + 147N

Artificial Transmutations:Transmutation in which there are two reactants.Fission: Heavy nucleus like 235U is bombarded with high-energy particle like 0

1n. The heavy nucleus is split to smaller nuclei

23592U + 1

0n --> 92Kr + 141Ba + 3 10n + ENERGY

Mass Defect:some mass is converted into energy

Fusion: Two light nuclei are crashed together at high speeds to form heavier nuclei. Products are clean and stable (not radioactive).

21H + 2

1H --> 42He + ENERGY

Mass Defect:some mass is converted into energyHalf-Life Formula

half-life events = time half-life

Half-life cannot be changed by temperature, pressure, etc.

Uses of Radioisotopes131I – diagnose problems with thyroid; 99Tc – Diagnose brain tumors; 14C – organic tracer; 238U – geological dating

PERIODIC TABLE

Arranged in order of increasing atomic number

Periods: the horizontal (left to right) rows. The period # indicates the principle energy level oc-cupied by the valence electrons. Period 3 elements have 3 energy levels

Groups: the vertical (up & down) columns. Elements in the same group have similar chemical properties. All Group 17 elements have 7 valence electrons

Metals: BIG LOSERS!!!To the left of the staircase with few exceptionsAll metals are solid at room temp, except for mercury (Hg) which is a liquidMetals are malleable, ductile, and have lusterHave mobile electrons which make them good conductors of heat & electricityTend to lose electrons to form positive ions with smaller radii

Semimetals (metalloids):Touch the staircase and include B, Si, As, Te, At, Ge, and SbHave some properties of the metals and some of the properties of nonmetals

Nonmetals:To the right of the staircase with the few exceptions noted above, and includes HydrogenExist in all states of matter at room temperature:gases = H, N, O, F, Cl, He, Ne, Ar, Kr, Xe, and Rnliquids = Brsolids = all othersTend to be brittle in the solid phasePoor conductors of heat & electricity with LOW MELTING POINTSTend to gain electrons to become negative ions with radii larger than their atoms

Groups of Interest:

1 = alkali metals: very very reactive, never found uncombined in nature.increasing atomic number, increasing reactivity

2 = alkaline earth metals: very reactive, never found uncombined in nature.increasing atomic number, increasing reactivity

3-11 = transition metals: characterized by forming colored solutions & having multiple positive oxidation states

17 = halogens: very very reactive; not found uncombined in natureincreasing atomic number, DECREASING reactivity. F is most reactive.Elements are found in ALL THREE PHASES

18 = noble gases: very stable (unreactive). Have full valence shell of 8 electrons (2 valence electrons only for helium)

Trends Down Groups:

Atomic Radius Increases, Electronegativity Decreases and Ionization Energy DecreasesElements become More Metallic, and Lose Electrons more Readily

Francium at the bottom of Group 1 is the MOST METALLIC METAL

Trends Across Periods:

Atomic Radius Decreases, Electronegativity Increases and Ionization Energy IncreasesElements become Less Metallic, and Gain Electrons more Readily

BONDING

electronegativity: attraction of an atom for a pair of shared electrons

ionization energy: the amount of energy absorbed to remove the most loosely bound electron

H + H --> H2 a bond is FORMED and energy is RELEASED

H2 --> H + H a bond is BROKEN and energy is ABSORBED

(BARF: Break Absorb Release Form)

Covalent is found in MOLECULAR SUBSTANCES, such as NH3, withLOW MELTING POINTSPOOR CONDUCTORSBRITTLE as solids

Nonpolar Covalent EQUAL sharing of electrons between 2 of the SAME NONMETALSPolar Covalent: UNEQUAL sharing of electrons between 2 DIFFERENT NONMETALSTHE BIGGER THE E.N.D. (Electronegativty Difference) the MORE POLAR THE BOND

Ionic: TRANSFER of electrons (from metal to nonmetal). ex: NaClHIGH MELTING POINT due to STRONG ATTRACTIONSPOOR CONDUCTORS AS SOLIDS (s)EXCELLENT CONDUCTORS IN SOLUTION (aq)

THE BIGGER THE E.N.D., the MORE IONIC THE BOND

Metallic: Sea of mobile electrons occurring between metallic elements ex: Fe, HgHIGH MELTING POINTEXCELLENT CONDUCTORS as (s) and (l)

Hydrogen bonding strongest POLAR attractions, found in samples of HF, H2O and NH3

Lewis Dot Diagrams:

Ionic Compounds use square brackets (no dots on the metal, and 8 dots on the nonmetal). Don’t forget ion charges on the outside of brackets, and then determine subscripts by crossing charges and reducing if possible.

Na2O

Covalent Compounds do NOT use brackets. Instead join together with single, double or triple bonds.

N2

H2O

FORMULAS & EQUATIONS

Naming of Ionic Compounds:

Na3P sodium phosphideCu(OH)2 copper (II) hydroxideFeO iron (II) oxide(NH4)3(PO4) ammonium phosphate

Writing Formulas for Ionic Compounds

Nickel (III) oxide Ni+3 O-2 swap, drop, reduct Ni2O3

Magnesium Sulfate Mg+2 (SO4)-2 --> Mg2(SO4)2 --> MgSO4

Classifying Reactions

Synthesis: making larger molecules from smaller ones: Mg + O2 --> MgO

Decomposition: breaking large molecules down into smaller ones: MgO --> Mg + O2

Single Replacement : one element replaces another in a compoundF2 + 2 NaI --> 2 NaF + I2

Double Replacement: elements from different compounds replace each otherNaI + AgNO3 --> NaNO3 + AgI

THE LAW OF THE CONSERVATION OF MASS:

Matter (Atoms) cannot be created or destroyed, only re-arranged.

A BALANCED EQUATION demonstrates CONSERVATION because there are the SAME NUMBER OF ATOMS ON BOTH SIDES OF THE EQUATION

N2 + 3 H2 --> 2 NH3

N H N H 2 6 2 6

MOLES & STOICHIOMETRY

% by Mass Composition

(1) of hydrogen in CH4

(part / whole) x 100% (4 H / CH4 ) x 100%4 / 16 x 100% = 25%

(2) of water in CaSO4 * 2 H2O(part / whole) x 100% (2 H2O / CaSO4 * 2 H2O)

36 / 172 = 20.9%

(3) of water in a hydrate using experiment data.mass of hydrate before heating: 10.0 grammass after heating: -7.5 grammass of water 2.5 grams

(part / whole) x 100% water / hydrate x 100%2.5 / 10.0 x 100% = 25%

EMPIRICAL & MOLECULAR FORMULAS

Molecular Formula: Actual Numbers of Atoms of all Elements in a Compound.Examples: C6H6, H2O, C8H18

Empirical Formula: Lowest Whole-Number Ratio of Atoms of Elements in a CompoundExamples: C6H6 --> CH, H2O --> H2O, C8H18--> C4H9

Question: The empirical formula is CH2 and the molecular mass is 56.0 g/mol. What is the molecular formula?

1. find empirical mass: CH2 12.0 + 2.0 = 14.02. mol.mass / emp.mass: 56.0 / 14.0 = 43. CH2 x 4 = C4H8

CALCULATING MOLES

1. Using the Mole Formula (found on back of Reference Tables—Table T):

moles = given mass gram formula mass

How many moles of molecules are in 5.00 grams of NH3?

moles = 5.00 grams / 17.0 g/mol

= 0.294 moles

2. Using Mole Ratios from a Balanced Equation

Using the balance equation 3H2 + N2 ⇌ 2NH3

How many moles of NH3 can be produced from 0.6 mol H2?

0.60 moles H2 x coefficient NH3 = ________ moles NH3

coefficient H2

0.60 moles H2 x 2 NH3 / 3 H2 = 0.40 moles NH3

GAS LAWS

MOLECULES of IDEAL GASES ...have no forces of attraction...travel in straight-lines randomly...occupy no volume themselves...collide elastically...are separated by large distances

Gases behave ideally at HIGH TEMPERATURE and LOW PRESSURES

Hydrogen and Helium behave MOST IDEALLY

At constant temperature P & V are inversely proportional (one goes up the other goes down)

At constant Volume, P & T are directly proportional At constant Pressure, V & T are directly proportional

THE COMBINED GAS LAW

P1V1 = P2V2 *If a variable (P,V, or T) is held constant, remove that variable T1 T2 from the formula or substitute “1”

*Temperature MUST be in KELVIN!

SOLUTIONS

solute: substance being dissolved (present in smaller amount)solvent: substance doing the dissolving (present in greater amount)

Characteristics of Solutions:1) are homogeneous mixtures2) do not settle on standing3) do not scatter light4) can have color5) have higher boiling points and lower freezing points than pure solvents

NaCl (aq) has a BP above 100˚C and a FP below 0˚C

Solubility of Solids and Liquid: more soluble at higher temperatures pressure has no effect on their solubility

Solubility of Gases: more soluble at lower temperaturesmore soluble at higher pressures

Use Table F to determine if a sample is soluble

CaSO4 Although SO42- is in the SOLUBLE COLUMN, Ca is an EXCEPTION

CaSO4 is not SOLUBLE.

NaOH Although OH- is in the INSOLUBLE COLUMN, Na is an EXCEPTIONNaOH is SOLUBLE

Table G to determine how many grams of a sample can be dissolved in 100 g H2O at different temperatures Table G:

Points on the line represent SATURATED SOLUTIONS which are at equilibrium

Points above the line represent SUPERSATURATED SOLUTIONS

Points below the line represent UNSATURATED SOLUTIONS

To calculate how many grams can be dissolved in an amount of H2O other than 100 grams:

mass that can be dissolved according to graph = 100 g H2O X Different amount of H2O

To calculate the mass that crystallizes when a saturated solution is cooled:

mass in saturated solution at higher temp - mass in saturated solution at lower temp

Molarity = moles of solute / Liters of solution (on table T)

ppm = (grams of solute/grams of solution) x 1,000,000 (on table T)

KINETICS & EQUILIBRIUM

COLLISION THEORY: Particles must have effective collisions (proper energy and orien-tation) in order for reaction to occur.

Based on collision theory – temperature, concentration, pressure, surface area, nature of reactants (ionic = fast…covalent/organic = slow), and presence of a catalyst affect rate of reaction

Using Table I(1) to determine whether a reaction is exothermic or endothermic

NaCl (s) --> Na+ (aq) + Cl- (aq) ∆H = +3.88 kJ/mol

This reaction is endothermic because ∆H = +

(2) to calculate the amount of energy absorbed or released in a reaction

How much energy is released when 4 moles of NH3 (g) is produced?

3 H2 (g) + N2 (g) --> 2 NH3 ∆H = -92 kJ/mol

Because the production of 2 moles of NH3 releases 92 kJ, then the production of 4 moles of NH3 releases twice the energy, or 184 kJ

Equilibrium: Rate of forward reaction equals rate of reverse reaction; concentrations are constant, but not necessarily equal

Catalysts increase the rate of the forward and the reverse reaction by lowering the activation com-plex

catalysts do NOT affect: EQUILIBRIUMcatalysts do NOT affect: heat of reaction

Le Chatelier’s Principle: Equilibrium will shift to counteract a stress applied to the system.

Example: 4NH3(g) + 5O2(g) ⇌ 4NO(g) + 6H2O(g) + heat

Temperature: Decrease temperature: down arrow at heat, equilibrium shifts to the right towards “heat”The forward reaction is favored. It points towards NO and H2O, which increase in concentrationIt points aways from NH3 and O2, whose concentrations decrease.

Concentration: Increase concentration of NO: put an up arrow on NO, equilibrium shifts to the opposite sidein this case, to the reverse.The reverse reaction point towards the NH3 and O2, so their concentrations increaseThe reverse reactions points away from “heat”, so the temperature decreases

Pressure: Increase pressure: equilibrium shifts to fewer moles of gas. Right side 4 + 6 = 10 moles of gas; left side 4 + 5 = 9. Equilibrium shifts toward the left, so the reverse is favored.

ACIDS, BASES and SALTS

Arrhenius Acid: Substance that produces H+ (H3O+) as only positive ion in aq. solution; pH be-low 7; [H+] > [OH-] (examples: see Table K).

Arrhenius Base: Substance that produces OH – as only negative ion in aqueous solution; pH above 7; [H+] < [OH-] (examples: see Table L).

Bronsted-Lowry Acid: Substance that can donate protons (H+).

Bronsted-Lowry Base: Substance that can accept protons (H+).

Neutralizaion reaction: Acid + Base → Salt + Water (The pH of products is 7)HNO3 + KOH --> KNO3 + H2O

Acids: Taste sour: turn litmus red; pH below 7; react with metals to produce H2; conduct electric current (electrolyte); the higher [H3O+] the lower the pH and the more acidic

Bases:Taste bitter; turn litmus blue; feel slippery; pH above 7; conduct electric current (elec-trolyte); the higher [OH -] the higher the pH and the more basi

Using the TITRATION / NEUTRALIZATION equation

Table T: (#H) x (MA) x (VA) = (MB) x (VB) x (#OH)

where #H is the # of leading hydrogens in an acidMA is the molarity of the acidVA is the volume of the acid#OH is the # of hydroxides in the baseMB is the molarity of the baseVB is the volume of the base

Table M: a list of indicators and the pH range of their color changes

REDOXWhat reactions are redox?

If a reaction contains a free element, it must be redoxDouble Replacement, therefore, is never redox

What redox reactions are spontaneous?

Redox reactions in which the free element is higher on Table J than its corresponding element in a compound

spontaneous: F2 + NaI --> I2 + NaFspontaneous: Zn + HCl --> ZnCl2 + H2non-spontaneous Ag + FeCl2 --> AgCl + Fe(Ag is at the bottom of Table J, and is thus not terribly reactive)

Redox reactions transfer electrons and can be recognized by a change in oxidation #.LEO says GER: LOSE electrons OXIDATION / GAIN electrons REDUCTIONOxidation: Oxidation # goes up; e- placed on right side of reactionReduction: Oxidation # goes down; e- placed on left side of reactionOxidizing Agent: The species that’s reduced. Reducing Agent: The species that’s oxidized.

In a voltaic cell, which metal is the ANODE? Metal that is higher on table “J” which is oxidized

An Ox (anode = oxidation) and the Fat Red Cat (cathode = reduction)Electrons flow through the external wire from anode to cathode

Voltaic Cell = chemical reaction spontaneously produces electricity; anode = negative side; has two half cells and a salt bridge

Writing Half-Reactions

Half-Reactions must show Conservation of Mass and Charge

Fe + Cu+2 --> Fe+2 + Cu

Reduction: Cu2+ + 2e- --> Cu(+2) + (-2) = (0)

Oxidation: Fe --> Fe+2 + 2e-(0) = (+2) + (-2)

Electrolytic = nonspontaneous redox; anode = positive side; one cell; power supply needed

ORGANIC CHEMISTRY

Carbon can form up to four bonds and make a tetrahedronLarge variety of organic molecules b/c carbon easily forms covalent bonds with other carbon atomsFor homologous series see Table “Q”Saturated = alkanes = single bonds.Unsaturated = alkenes/alkynes = double/triple bondsWhen naming, place side chains in alphabetical order with lowest possible numberFunctional groups…see Table “R”

-- aldehydes, acids, & amides MUST be at end of chain-- ethers, ketones, & amines MUST be in “middle” of chain-- alcohols & halides can be anywhere in chain-- alcohol prefixes go in middle of parent name (e.g., 2,2-pentanediol)

ORGANIC REACTIONS

combustion: hydrocarbons + oxygen → CO2 + water

addition: adding atoms on to a molecule at double or triple bonds

unsaturated hydrocarbon + halogen --> halocarbonC2H4 + Cl2 --> C2H4Cl2

substitution:

saturated hydrocarbon + halogen --> halocarbon + acidC2H6 + Cl2 --> C2H5Cl + HCl

esterification: acid + alcohol → ester + water

fermentation: sugar → alcohol + CO2 (often the catalyst zymase is used)

saponification: fat + base → glycerol + soap

polymerization: joining of monomers into a long chain of repeating units

· Isomers = same # of each element, but different arrangements and different properties

C4H10 can be drawn two ways:(1) as a single chain of 4 carbons (butane) BUTANE(2) as a chain of 3 carbons (propane) with a one carbon branch from the middle (methyl) METHYL PROPANE

· Characteristics of organic molecules: low melting point, don’t conduct electricity, decom-pose easily, slow reaction rates, usually not very polar