Waves

Unit 5BondingChapters 7-8Unit 6Naming and MolesChapter 9-10Unit 7Reactions and StoichiometryChapters 11-12Reference Tables F & JUnit 8Phases and HeatChapters 13 & 17Reference Table HIonsAtoms that have gained or lost electrons, and now have a charge

Must show chargeNa+F-O-2CompoundsTwo Main Types of CompoundsIonicMolecular (Covalent)

Based on type of bonding involvedBondingBondShared or exchanged electrons that hold two atoms together

Three Main TypesIonicCovalentMetallic

BondingIonic BondElectrons are transferred from one atom to another (one gives, one takes)

Metal and nonmetal, NaClLarge electronegativity difference

Polyatomic ion, Mg(NO3)2More than 2 elementsIonic CompoundsIonic compounds are electrically neutral, even though they are composed of charged ionsTotal positive charge equals total negative charge

9Determining FormulasMust be electrically neutralTotal positive charge must equal total negative chargeUse oxidation numbers from Periodic TableGroup 1 +1Group 2 +2Group 13 +3Group 15 -3Group 16 -2Group 17 -1Determining FormulasDetermine number of each ion to balance out chargeUse as subscript for element symbolEx: CaCl2, Na3PO4, Mg(NO3)2Write Positive Ion First

Formula must be smallest whole-number ratio

ExampleCalcium and Fluorine

CaF2Ca+2F-F-Ca1F2Polyatomic IonsGroup of atoms that collectively have gained or lost electrons (Table E)

Sodium and NitrateNa+(NO3)-Na1(NO3)1NaNO3Short-cut (criss-cross method)Magnesium and Carbonate

Mg2(CO3)2Mg+2CO3-2MgCO3Must SimplifyDot StructuresShows valence electronsMust show charge for Ions

NaCl

Na+Cl-Covalent BondsElectrons are shared between two atomsEach atom will try to achieve a full valence shell2 nonmetals

Two types of covalent bondsNon-Polar Covalent Shared equallyPolar Covalent Shared unequallyCovalent BondingH2

O2

N2

HHSingle BondDouble BondOONNTriple BondCovalent BondingH2O

CO2

OHHSingleBondSingleBondCOOMore ExamplesHCl

NH3

CH4

ClHNHHHCHHHHDetermining Bond TypeBond type is based on electronegativity difference (EN) between two bonding atoms

Nonpolar Covalent Bond2 of the same nonmetals

Polar Covalent Bond2 different nonmetals

Ionic BondMetal and a nonmetal

Determining Bond PolarityThe larger the difference in electronegativity, the more polar the bond.Which is more polar?

HIHBrHClHF1.81.00.80.5ENBiggestMost PolarPropertiesIonic CompoundsMost ionic compounds are hard, crystalline solids at room temperature

High melting points

Mostly soluble in water

Can conduct an electric current when melted or dissolved in water(aq).

PropertiesCovalent CompoundsMost molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

PolyatomicsCompounds with polyatomic ions contain BOTH ionic and covalent bondsExample: NaNO3Na+-NOOOAllotropesTwo or more different molecular forms of the same element in the same physical stateDifferent properties because they have different molecular structuresO2 vs O3Diamond, Graphite, Fullerenes (pictured on next slide)

Allotropes

Metallic BondingBonding within metallic samples is due to highly mobile valence electronsFree flowing valence electronsSea of Electrons

Network SolidsAll atoms in a network solid are covalently bonded togethervery high melting and boiling pointsExamplesDiamonds ( C )Graphite ( C )Silicon Dioxide (SiO2)Silicon Carbide (SiC)

Bond EnergyWhen two atoms form a bond, energy is releasedExample: Cl + Cl Cl2 + energy

Energy needs to be added to break a bondExample: Cl2 + energy Cl + ClMolecular PolarityPolar Moleculeone end of a molecule is slightly negative(-) and the other end is slightly positive(+).Asymmetrical charge distribution

Nonpolar MoleculeCan not be separated into different endsSymmetrical charge distributionPolar MoleculesH2O

HCl

NH3

OHH-++HCl+-NH-+++NHHHNonpolar ExamplesCH4

CO2

CHHHH+-+++O=C=O+--PolarityIonic Compounds are Ionic

Nonpolar Covalent Bonds always indicate Nonpolar Molecules

Polar Covalent BondsDetermine SymmetryPolarityNonpolar MoleculesCH4, CO2, H2, N2, O2,

Polar MoleculesH2O, HCl, HBr, NH3, Like Dissolves LikePolar and Ionic substances will dissolve in other Polar Substances

Nonpolar substance will dissolve in other nonpolar substancesIntermolecular ForcesIntermolecular Forces of Attractionattraction between two molecules or ions that hold them together (not a bond)

Determines melting and boiling points of compoundsStronger intermolecular forces, higher melting and boiling pointsIntermolecular ForcesVan der WaalsDispersionDipole-DipoleMolecule-IonHydrogen BondingWeakestStrongestHydrogen BondingHydrogen bonded to N, O, or F, is attracted to the N, O, or F of another molecule.Not actual bond, just attractionHFHFHydrogen BondNaming IonsPositive Ions, cations, simply retain their name.Na+ Sodium IonMg2+ Magnesium Ion

Negative Ions, anions, change ending of element to ideCl- Chloride IonBr- Bromide Ion

Polyatomic IonsSelected polyatomic ions are on Table E in the Reference Tables.

Polyatomic ions keep their namesExcept in AcidsNaming SystemsIonic SystemMetals and Nonmetal, more than 2 elements

Stock System (Roman Numerals)Use when the metal element has more than one positive oxidation numberRoman Numeral is the charge of the metal ion

Binary Covalent System (Prefixes)2 nonmetals (including metalloids)Second element ends in ide

Naming Ionic CompoundsName positive ion first, then negative ion.NaCl Sodium chlorideMg(OH)2 Magnesium hydroxideStock System ExampleSnCl4

Tin(IV) ChlorideSnClClClCl-1-1-1-1+4+4-4Roman NumeralsCation ChargeRoman Numeral+1I+2II+3III+4IV+5V+6VI+7VII+8VIIIBinary Covalent ExampleN2Cl3Dinitrogen Trichloride

CO2Carbon Dioxide

PCl5Phosphorus PentachloridePrefixesNumber of atomsPrefix1mono-2di-3tri-4tetra-5penta-6hexa-7hepta-8octa-Avogadros Number6.02 x 1023Number of representative particles in a mole1 mol He = 6.02 x 1023 atoms1 mol H2 = 6.02 x 1023 moleculesGram Formula MassMass of the formula in g/molSimply add the atomic masses of each element in the formula together

H2O = 1 + 1 + 16 = 18 g/molAlso known as gram atomic mass, gram molecular mass, molar massMole - Mass ConversionExample: 96 g of Oxygen gas = ? mol

PracticeHow many moles are there in 506g of ethanol, C2H6O?

What is the mass of 8 moles of CCl4?

1232g CCl411 mol C2H6OMolar VolumeAt STP, 1 mol of any gas occupies 22.4L of spaceMole Road Map

Percent Composition

Percent CompositionWhat is the percent composition of oxygen in H2SO3?

58.5%

HydratesCompounds that have a specific number of water molecules attached

Dot means plus (+)CuSO45H2OEmpirical FormulaSimplest Whole-Number ratio of atoms in a compoundExamplesC6H12O6CH2O

Molecular Formula is a multiple of the Empirical Formula

Empirical FormulaA molecular formula has an empirical formula of CH2 and a molecular mass of 28 g/mol.

A molecular formula has an empirical formula of CH2 and a molecular mass of 42 g/mol.

C2H4C3H6ReactionsReactantsProductsBalancing ReactionsReactions must maintain conservation of mass, charge, and energy

Reactants and Products must have the same number of atoms of each element2H2 + O2 2H2O

Reactants must have the same total charge as ProductsCu+1 + Fe+3 Cu+2 + Fe+2

Balancing ReactionsTo balance a reaction:Do NOT change chemistry (compounds, subscripts)Only change coefficients (big numbers in front of chemicals)Coefficients can only be whole numbers2H2 + O2 2H2O

Balancing Reactions4Na + O2 2Na2O

2Al + 3Br2 2AlBr3

4Ni + 3O2 2Ni2O3

2HNO3 + Ca(OH)2 Ca(NO3)2 + 2H2OReaction TypesSynthesisDecompositionSingle ReplacementDouble ReplacementCombustionSynthesisChemical change in which two or more substances react to form a single new substance (1 product)Also called Combination

2Mg + O2 2MgODecompositionChemical change in which a single compound breaks down into two or more simpler products1 reactant

2NaCl 2Na + Cl2Single ReplacementChemical change in which one element replaces a second element in a compoundMetal replaces metal (hydrogen included)Nonmetal replaces nonmetal

Zn + 2HCl ZnCl2 + H2Double ReplacementChemical change involving an exchange of positive ions between compounds

AgNO3 + NaCl NaNO3 + AgClCombustionChemical change in which an element or a compound reacts with oxygen, often producing energy in the form of heat and light

2Mg + O2 2MgOCombustionComplete combustion of a hydrocarbon produces carbon dioxide and water

2C2H6 + 7O2 4CO2 + 6H2OSpontaneous ReactionsA single replacement reaction will only occur if:The single element in the reactants is more active than the element it replaces in the compoundUse Reference Table JThe more active element:Does not want to be aloneWants to be combined with someone else

Spontaneous ReactionsA double replacement reaction will only occur if:A precipitate (solid) is producedUse Reference Table FA liquid is producedH2O(l)Stoichiometry ProcessBalance reactionCopy coefficients from reaction into ratio rowPlace numbers from question in row above ratio row (Setting up proportion)Solve for X using a proportionChemical ExampleHow many moles of NH3 can be made with 6 moles H2 and excess N2?

ratio1326XamountX = 4 mol NH3=Conservation of EnergyReactions must maintain conservation of energyEnergy term written in reaction

Endo/ExothermicEndothermic Energy is absorbedEnergy term is on the left side

Exothermic Energy is releasedEnergy term is on the right side

Treat just like a coefficientExample

How much energy is produced when 6 moles H2 reacts with excess N2?ratio1326Xamount91.8X =183.6 kJSolidsform of matter that has a definite shape and definite volume.

Use (s) to denote solids in chemical reactions

In most solids the atoms, ions, or molecules are packed tightly together

Liquidsform of matter that has a definite volume, indefinite shape, and flows.

Use (l) to denote liquids in chemical reactions

In liquids the atoms or molecules are able to slide past each other.

In liquids there are intermolecular attractions between the atoms or molecules, which determine the liquids physical properties.Gasesform of matter that takes both the shape and volume of its container

Use (g) to denote gases in chemical reactionsPhase ChangesSix ChangesSolid LiquidMeltingLiquid SolidFreezingLiquid GasVaporizationGas LiquidCondensationSolid GasSublimationGas SolidDepositionPhase ChangesDuring any given phase change, both phases can exist together in equilibrium

ExampleAt 0C, water can exist in both the liquid and solid phases in equilibriumTypes of EnergyKinetic EnergyEnergy of motionRelated to the speed (velocity) and mass of molecules

Potential EnergyStored energyTemperatureRelationship between energy, particle speed (velocity), and temperature

Temperature DefinitionAverage Kinetic Energy

ConversionK = C + 273

PressureNumber of collisions between particles and container walls

Vapor PressurePressure exerted by vapor that has evaporated and remains above a liquidAs temperature increases, vapor pressure increases (Table H)

Boiling vs. EvaporationBoilingVapor pressure equals external, or atmospheric pressureNormal means at Standard Pressure

EvaporationSome molecules gain enough energy to escape the liquid phaseAt temp. less than boiling pointHeatEnergy transferred from one object to another, usually because of a temperature difference

Measured in Joules (J)

Heat flows from hot to cold

Specific Heat CapacityAmount of heat needed to raise the temperature of 1 g of a substance by 1CUnique for each phase of each substance

4.18 J/(gK) for liquid waterListed in Table B of Reference TablesHeat EquationHeat, qMass, mSpecific Heat Capacity, CChange in Temperature, T

q=mCTExampleHow much energy is required to raise the temperature of 50g of water from 5C to 50C?

q = mCTq= (50g) (4.18J/gK) (45C)q= 9405 J

Phase ChangeWhat happens to temperature during phase changes?Temperature remains constantPotential Energy changes

Temperature remains CONSTANT during a phase change

Heat of Vaporization, HvAmount of energy needed to vaporize 1g of a substance

Water = 2260 J/g

q=mHvUse for Liquid Gas orGas LiquidHeat of Fusion, HfAmount of energy needed to melt 1g of a substance

Water = 334 J/g

q=mHf

Use for Solid Liquid orLiquid SolidExampleHow much energy is needed to vaporize 10g of water at 100C?q = mHvq = (10g) (2260J/g)q = 22600 J

Heating (Cooling) CurvesShows relationship between temperature and time during constant heating or cooling.

Also shows phases, and the phase changes between them.Heating CurvesDiagonal lines are phasesHorizontal lines are phase changesTime (s)Temp (C)GasLiquidSolidHeating CurvesDiagonal lines are phasesHorizontal lines are phase changesTime (s)Temp (C)VaporizationCondensationMeltingFreezingConservation of EnergyEnergy can not be created or destroyed, only transferred or converted from one form to another.

Energy lost by one object must be gained by another object or the environmentqlost = qgainedExampleA chunk of iron at 80C is dropped into a bucket of water at 20C.

What direction will heat flow?From the iron to the waterHot to cold

Final Temperature between 20C and 80C