Chemistry 1(2)
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Transcript of 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