Chemistry 1(2)

CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g)

Coefficients: relative number of moles1 mole of CH4 + 2 moles of O2 1 mole CO2 + 2 moles H2O

Reaction between methane and oxygen to produce carbon dioxide and water

3Fe (s) + 4H2O (g) Fe3O4 (s) + 4H2 (g)

States of matter: solid/liquid/gasSolid iron + gas water solid iron oxide + gas hydrogen

CH EM I S T R Y 1 – S EM E S T E R 2

REVIEW GUIDEChemical ReactionsChemical Reaction: process by which one or more substances are changed into different substances.Total mass of reactants must equal total mass of productsEvidence of reactions:

- Heat- Light- Gas- Precipitate- Color change

Chemical Equation: a representation of a chemical reaction using symbols and formulas

Word Equations: chemical reactions represented in words without formulas or symbols (or coefficients) Sodium oxide + water sodium hydroxide Na2O + H2O NaOH Unbalanced

Balanced Equations:- Correct formulas- Correct number of moles

Unbalanced: C2H6 + O2 CO2 + H2O Balanced: 2C2H6 + 7O2 4CO2 + 6H2OUnbalanced: H2SO4 + NaOH Na2SO4 + H2O Balanced: H2SO4 + 2NaOH Na2SO4 + 2H2O

Reactants: substances being reactedProducts: substances produced by reaction

Never change subscriptsPolyatomic ions are single unitsCount all atoms at the end to make sure

+ →

Ionic Equations:Overall Ionic Equation: balance equation write out all ions w/ charges except product which forms precipitate

Balanced Equation: Pb(NO3)2 (aq) + 2KI (aq) 2KNO3 (aq) + PbI2 (s)Pb2+ (aq) + 2(NO3)1- (aq) + 2K1+ (aq) + 2I1- 2K1+ + (aq) + 2NO3

1- (aq) + PbI2 (s)Net Ionic Equation: use overall ionic equation use product which forms precipitate write out what reactions formed that product

Pb2+ (aq) + 2I1- (aq) PbI2 (s)

Relative Masses: 4 Li + O2 2Li2O4 moles 1 mole 2 moles

4 mol Li x (6.94g/1mol) = 27.76g of Li1 mol O2 x (16.0g/1 mol) = 16.0g of O2

2 mol Li2O x (29.9g/1mol) = 59.7g of Li2O

Types of ReactionsI. Synthesis

2 or more substances react to make 1 new product

2Mg + O2 2MgO

II. DecompositionOne compound broken to produce two or more simpler substances

2H2O 2H2 + O2

III. Single ReplacementOne element replaces a similar element in a compound

+→

A + BX → AX + B

+ → +

+ → +

2Al + 3ZnCl2 3Zn + 2AlCl3

Double Replacement

Ions of two compounds exchange places to form two new compounds

FeS + 2HCl H2S + FeCl2

IV. CombustionReaction involving oxygen which usually releases energy as heat

C3H8 (g) + 5O2 (g) → 3CO2 (g) 4H2O (g)

Stoichiometric calculationsStoichiometry: mass relationship between reactants and products in a chemical reaction

Mole ratio: conversion factor in moles of any two substances in a reaction

Has to follow ACTIVITY SERIES

Elements higher up the activity series will replace elements lower down the activity series

eg: Mg replace Zn Li replace Fe

AX + BY → AY + BX

One of the 2 new compounds formed HAS to form

PRECIPITATEFound using “solubility

guidelines”

hydrocarbon + oxygen carbon dioxide + water

Molar Mass: the mass of 1 mole of a substance

Mole to mole calculation:

Write balanced equation and find mole:mole ratio using coefficients in equation

Lithium and oxygen react and form LiO2. 2 moles of LiO2 form, how many moles of O2 did we start with?

Li + O2 LiO2 1 – 1 mole ratio.

2 moles of LiO2 x (1mol O2/1mol LiO2) = 2 moles of O2

Mole to Mass calculation:

Write balanced equation - use mole:mole ratio and molar mass from periodic table

If 0.500 mol of NaN3 react, what mass in grams of nitrogen would result?

2NaN3 (s) → 2Na (s) + 3N2 (g)

0.5 mol of NaN3 x (3 mol N2/2 mol NaN3) x (28.01gN2/1mol) = 21.01g of N2

Mass to mole calculations:

Balance equation - start with grams and use molar mass and mole:mole ratio

How many moles of NH3 are needed to produce 30.0 g of PtCl2(NH3)2?

K2PtCl4 + 2NH3 → 2KCL + PtCl2(NH3)2

30g x (1 mol/283.02) x (2mol/1mol) = 0.212moles of NH3

Mass to Mass calculations:

Balance equation – start with grams molar mass mole:mole molar mass

What mass in grams of MgCl2 will be produced if 3.00 g of Mg(OH)2 reacts?

Mg(OH)2 + 2HCl → 2H2O + MgCl2

3.0g x (1 mol/58.32g) x (1 mol/1mol) x (95.21g/1mol) = 4.90g of MgCl2

Limiting reactants: reactant that limits the amount of other reactants which can be used

The reactant with the lower number of moles relative to the mole:mole ratio

Balanced equation number of moles present of each reactant mole:mole ratio

If 2.00 mol of HF is exposed to 4.5 mol SiO2, which is the limiting reactant?

SiO2 (s) + 4HF (g) → SiF4 (g) + 2H2O (l)

2mol HF x (1mol SiO2/4mol HF) = 0.5mol SiO2

4.5 moles of SiO2 present so HF is limiting reactant

Theoretical yield: amount of product that is predicted to be produced from mole:mole ratio

Balance equation mole:mole ratio moles of limiting reactant moles of product

Energy required to break bondsEnergy required to form bonds

What is theoretical yield of H2O if there are 6 moles of C6H8O7 and 4 moles of ZnCO3?

3ZnCO3 + 2C6H8O7 → Zn3(C6H5O7)2 + 3H2O + 3CO2

Limiting reactant is – ZnCO3

4mol ZnCO3 x (3mol H2O/3mol ZnCO3) = 4mol H2O

Actual Yield: measured amount of product obtained from reaction

Percent Yield: (actual yield / theoretical yield) x 100

If 50.0 g of CO reacts to produce 55.4 g of CH3OH, what is the percent yield of CH3OH?

CO + 2H2 → CH3OH

Theoretical yield: 50g x (1mol/28g) x (1mol/1mol) x (32g/1mol) = 57.1g

Percent yield: (55.4g/57.1g) x 100 = 97.0%

Molarity: the number of moles of a solute in 1 liter of a solution

M = mol/L

Given mass and volume - calculate molarity: 3.50L of solution that has 90g of NaCl, what is molarity?

Grams moles molarity

90g of NaCl = 1.54 moles

1.54mol NaCl/3.50L = .440mol/1L M = .440mol/L

Given concentration and volume – calculate grams: 0.8L of 0.5M HCl solution. How many moles of HCl?

M = mol/L

0.5M = moles/0.8L 0.4moles of HCl

BondingChemical Bond: mutual electrical attraction between nuclei and valance electrons of different atoms

VALANCE ELECTRONS REDISTRIBUTE THEMSELVES TO BECOME STATBLE

Energy kJ/mol

>1.7 Ionic0.3 to 1.7 Polar Covalent< 0.3 CovalentO = C = O

One shared pair of electrons

Ionic Bond: electrical attraction between cations and anions

Covalent Bond: sharing of electrons between two atoms

Coordinate Covalent Bond: both electrons in the bond are provided by one atom

Metallic Bond: attraction between metal atom and surround electrons

Bond Length: average distance between two bonded atoms

HIGHER # OF BONDS = SHORTER BONDS = STRONGER (MORE ENERGY)

SMALLER ATOM = SMALLER LENGTH = MORE ENRGY

Electronegativity: atom’s ability to attract electrons

Electronegativity Difference:

Exceptions: CO2

Octet Rule: chemical compounds want to gain/lose/share electrons to reach stability – outer s and p orbitals filled [exception: Boron]

Electronegativity difference more than 0.3 but oxygen on both sides of C -- nonpolar

C = C

H Cl

Cl HC = C

Cl Cl

H H

Na N NeTwo shared pairs of electrons

Each H atom forms 1 bond – 1 valence electronCarbon forms 4 bonds – 4 valence electronsOxygen forms 2 bonds – 6 valence electrons

Structural formula: representation of number of bonds within compounds C = O H H

4 valence electrons

1 valence electron

Covalent Bond:

Isomers:different orientations for the same compound

Ionic Compounds:

Lewis Dot Structures:

Ionic Bond:

Positive and negative charges must balance

One valence electron gives 1 electron (forms 1 bond)

5 valence electrons – forms 3 bonds to reach stability

8 valence electrons – forms no bonds – stable

+

AB2180o bond angle

Pulling electronsEqual pulls balance outNON-polar

Crystal Lattice formation: repeating pattern of 3D units – electrostatic attraction

Metallic Bonding:

- Electrons are free to move in sea of electrons- High electrical conductivity- Malleable – bonding same in all directions

Alloys: homogenous mixtures

- Mixture of metals (eg: gold + silver)- Harder than pure metals

Allotropes of Carbon:crystal lattice structures

All set up in different formations

- Diamond – tetrahedral- Graphite – hexagonal in layers- Buckminister fullerene – giodixic domes

Molecular GeometryProperties of molecules depend on the bonding and geometry

Linear Molecules:

Trigonal Planar: Trigonal Pyramidal:

Bonds between ions are stronger than bonds between molecules – Ionic Bonds are harder to

break except when in water

VSEPR – Valence Shell Electron Pair Repulsion

Repulsion between the sets of valence electrons causes electron pairs to orient far apart

Unshared PairPull towards NPolar

Bent Angular: Tetrahedron:

Intermolecular ForcesIntermolecular forces: Bonds between molecules, not within

Polar molecules with Polar molecules – “dipole-dipole”

- Strong force because each molecule is a dipole

Dipole: having both a positively charged region and a negatively charged region

AB3120o Bond Angles

BF3 – exception to OCTET ruleNo unshared pairsNonpolar

AB3ELone electron pair

distorts shape107o Bond Angles

2 unshared pairsPull towards OPolar

AB3E22 Lone electron

pairs distorts shape even more

AB4109.5o bond angles

No unshared pairsEqual pull in all

directions balance outNon polar

Partially negatively charged region because

hydrogen atoms’ electrons move towards

nitrogen

Partially positively charged region because nitrogen attracts electrons

towards itself

Polar molecule causes dipole in a non-polar molecule by temporarily

attracting its electrons

Dipole-Dipole attraction: negative regions attracted to positive regions

Hydrogen Bonds: a type of dipole-

Hydrogen atom bonded to highly electronegative atom is attracted to unshared pair of electrons of an electronegative atom

Water has unique properties

- Surface tension - High heat of vaporization- High specific heat - Less dense in solid state

Attraction between molecules: Vanderwaals Forces

Positive hydrogens attracted to

negative oxygens

Electrons push away electrons

Particles of matter are always in motion above absolute zero

temperature – 0 Kelvin = 273oCMovement stops at 0 Kelvin

(Another unit for temperature)

London dispersal forces: constant movement of electrons create instantaneous dipoles

Kinetic Molecular Theory - Gases

Theory for Ideal Gases

- Lots of particles far apart- Fast, constantly moving particles- No forces of attraction or repulsion- All collisions are elastic – no energy lost- Speed of particles depend on temperature

o Increase temp – increase speed

Gases:

- Expand to fill container- Compress easily - Low density- Flow- Diffuse from high to low

Rate of diffusion:

2 different gases: ½MA • vA2 = ½ MB • vB

2

vA2 = MB

vB2 MA

Rate of Diffusion of A = √MB

Rate of Diffusion of B √MA

Creating partial charges in helium

Kinetic EnergyKE = ½ mv2

m – mass of particles v – speed of particles (proportional to temperature) Different gases at the same temp – SAME KE – so velocity varies

Graham’s Law

k is a constant that depends on the quantities of the other two variables

P1 • V1 = P2 • V2 T1 T2

Ideal gases vs. Real gases.

Real gases do not behave like ideal gases

- Intermolecular forceso Hydrogen bondso London dispersal

- High pressures/Low temperatureso Increasing pressure on a gases moves particles closero Reducing temperature causes particles to move slowly

Pressure: force/area

Force: accelerating an object of given mass

Units: - pascals

mmHg

atmospheres

Measured by barometers

Gas LawsBoyle’s Law: pressure volume relationship

P • V = k or P = k V

Inversely proportional

P1 • V1 = k = P2 • V2

P1 • V1 = P2 • V2

Gay Lussac’s Law: pressure temperature relationship

P = k or P = ktTDirectly proportional

P1 = P2 T1 T2

Combined Gas Laws: Gases often change pressure, temperature, and volume all at the same time

P • V = k T

Charles’ Law: volume temperature relationship

V = k or V = ktTDirectly proportionalV1 = V2 T1 T2

Pressure = 1.5 N/cm2

Area of contact: 325cm2

PT = P1 + P2 + P3 + …….

V α 1 P

Boyle’s

V α T Charles’

V α nAvogadro’s

V α 1 x T x n PV = R 1 x T x n PP • V = n • R • T

Substitute m/M for n

Derivations from PV=nRTn = m M

P • V = mRT M

M = mRT P • V

D = m v

M = DRT P

Rearrange

Substitute D for m/v

Dalton’s Law of Partial Pressures: total pressure of a mixture of gases, like air, is equal to sum of partial pressures of each component gas

Vapor pressure: subtract from atmospheric pressure when collecting gas over water to find the pressure of the gas

Avogadro (what on earth were his parents thinking naming him this?): 1 mole of any gas at STP is equal to 22.4L

Further proportionalities:

V = kn - volume is directly proportional to moles

T = kn – temperature is directly proportional to moles

P = kn – pressure is directly proportional to moles

R – gas constant

Less compressible than gasesHigher densityLower KE – due to lower temperatureSlower diffusion rateSurface tension – evidence of I.M forces

Exothermic Endothermic

Liquids/Solids

Vaporization: process by which liquid or solid turns to gas – inevitable

Evaporation: process by which particles escape the surface of a liquid and become gas

Increase EvaporationIncrease temp.Increase surface areaDecrease humidity

Using a Lid: only certain amount of gas can escape bottle – particles moving slower will fall back down

Lid creates a really HIGH vapor pressure so the equilibrium pressure is higher – causes boiling point to be higher

Equilibrium: equal rates of evaporation and condensation in a stoppered container

Equilibrium Vapor Pressure:

Pressure exerted by vapor in equilibrium with corresponding liquid

Boiling temperature: conversion of a liquid to vapor

where vapor pressure = atmospheric pressure

Rate of evaporation: time it takes to reach atmospheric pressure

at given temperature

Solid:

Definite shape and volume

More closely packed – very structured

More dense

Lower KE due to lower temperature

Ether evaporates fastest as it reaches 1atm at lowest temperature

Liquid: 4.184 J g • OC

Solid (ice): 2.06 J g • OC

Vapor: 2.02 J g • OC

Phase Diagrams: diagrams which show all three phases of a substance with lines of state change

Specific Heats of Water: energy required to change the temp. of one gram without changing phase

Triple Point: temperature and pressure at which

solid, liquid and gas phases exist in

equilibrium and substance is

constantly changing states

Critical Point:

Critical temp + Critical pressure

Critical temperature: temperature above which substance is always in gas phase, no matter the pressure

Critical pressure: lowest pressure at which substance can be in liquid form at critical temperature

Phase Change lines: lines over which

substance changes phases at given

pressure and