Download Equilibrium

Group 3:
Periodic Trends
● Electronegativity: ability of an atom in a molecule to attract shared electrons to itself;
increases across a period and decreases down a family
● Ionization Energy: energy required to remove the valence electron from an atom;
increases across a period and increases up a family
● Atomic Radius: radius of an atom measured from the nucleus to the outermost electron;
decreases across a period and decreases up a family
● Electron Affinity: energy change associated with addition of an electron; increases
across a period and decreases down a family
Ionic Bonds
1. Electrons are transferred
2. Metals react with non-metals
3. Ions paired have lower energy (greater stability) than separated ions
Coulomb’s Law → E= 2.31x10^(-19)Jxnm(Q1Q2/r)
● E= energy in joules
● Q1 and Q2 are numerical ion charges
● r= distance between ion center in nanometers
● negative sign indicates ion attractive force
Covalent Bonds
1. Electrons are shared by nuclei
2. Pure covalent (non-polar, covalent)
-Electrons are shared evenly
3. Polar covalent bonds
-Electrons are shared unequally
-Atoms end up with fractional charges within an atom (negative or positive)
Polarity
● Dipolar molecules:
-Molecules with a somewhat negative and a somewhat positive end (a dipole moment)
-Molecules with a preferential orientation in an electric field
-All diatomic molecules with a polar covalent bond are dipolar
● Molecules with polar bonds but no dipole moment
-Linear, radial, or tetrahedral symmetry of charge distribution
-CO2 = linear
-CCL4 = tetrahedral
●
Kinetic Molecular Theory Postulates:
1. Gases are composed of very small particles (molecules or atoms)
2. Gas particles are tiny compared to the distances between them, so volume of gas is
negligible
3. Gas particles are constantly in motion. The collisions cause pressure of the gas
4. Gas particles neither attract nor repel each other
5. The average kinetic energy of gases are proportional to the kelvin temperature
● Pressure= force/area
● Properties of Gases: They uniformly fill the container, are easily compressed, mix
completely with any other gas, and exert pressure on their surroundings.
● Gases that obey these postulates are ideal gases
● Average velocity of gas particles are called the root mean squared speed
● Gas Law Relationships:
1. Pressure: Boyle’s Law- Pressure is inversely related to K
P1V1=P2V2
2. Volume: Charles's Law- Direct relationship between volume and temperature
V1/T1=V2/T2
3. Combined Gas Law: (P1V1)/T1=(P2V2)/T2
4. Avogadro’s Law: Direct relationship between volume and number of moles
V1/n1=V2/n2
5. Ideal Gas Equation: PV=nRT where R is .0826 (L*atm/mol*K)
● Dalton’s Law of Partial Pressures: Ptotal=ntotal(RT/V)
● Mole fraction: X1=(n1/ntotal)=(P1/Ptotal)
● Graham’s Law of Diffusion and Effusion: The lighter the gas, the faster the rate of
effusion
(r1/r2)=M2/M1
● Non- Ideal Gases:
1. Van Der Waals equation takes into account the volume and attractive forces of real gases.
(P+an2/V2)*(V-nb)=nRT
● Solving for new temperature, volume or pressure: V2=V1(P1/P2)(T2/T1)
● Standard temperature and pressure: O degrees Celsius or 273K; 760 torr or 1 atm
● 1 mole of ideal gas at STP occupies 22.4 L of volume
● density= mass/volume
● n=grams of substance/molar mass
●
A=Central Atom
X=Bonding Atom
E=Electron Pair
● When molecules exhibit resonance, any structures can be used to predict molecular
structure using VSEPR model
● VSEPR works in most cases for non-ionic compounds
Sigma and pi bonds
● Sigma Bonds: Bond in which the electron pair is shared in an area centered on a line
running between the atoms
o
Pi Bonds: Electron pair above and below the s bond
●
Sigma bonds (s bond) are the first bonds i.e. a single bond. Every bond after that
becomes a pi bond (p bonds). For example, a triple bond has one sigma bond and two pi
bonds; a double bond has a sigma bond and one pi bond
Group 5: Chapter 10 (IMF’s/Boiling Point/Melting Point)
Intermolecular Forces (IMF):
Dipole - Dipole
•Polar Molecules
•Also have London Dispersion Forces (LDF’s) b/c they still have electrons
•Constant Dipoles
•Covalent Molecules
To figure out if a molecule is Dipole-Dipole:
1.Look to see if the
2.Draw a Lewis dot diagram
3.Look to see if the molecule is polar
4.If it is look for hydrogen bonding (O-H, F-H, or N-H)
5.If it is not, it has dipole-dipole
Try these examples:
CHCl3
NO
HCl
Hydrogen Bonding: special dipole-dipole
attraction that is stronger than normal dipoledipole attractions.
●
●
hydrogen is covalently bonded to a N, O,
or F atom.
Example: Water and DNA
London Dispersion Forces: an instantaneous
dipole; a random movement of electrons that
creates a momentary non-symmetrical
distribution of charge even in non-polar molecules.
●
They exist between all molecules, but they are the weakest
forces of attraction.
Try these Practice Problems:
Identify the interparticle forces present in the solids of each of the following substances.
1. Ar
2. HCl
3. HF
4. CH4
Boiling (liquid -->gas):
● More particles = harder for particles to evaporate = higher boiling point
● The boiling point of a liquid is determined by the strength of the IMF’s in the liquid
● Alcohols have the highest boiling point, then ketones, and alkanes have the lowest
boiling point
● Polarity affects boiling point. Non-polar molecules have the lowest boiling point (vice
versa)
● Mass increase -->boiling point increases
● IMF’s increase -->boiling point increases
Freezing (liquid -->solid):
● More particles = greater energy required to freeze it because more particles get in the
way for the lattice structure to form
Melting (solid --> liquid):
● More particles = less energy required (opposite of freezing)
● Lower melting point -->lower IMF’s (when a solid melt, the structure breaks and the
bonds are weaker in a liquid then in a solid, so there’s less force attraction the particles
together, which results in lower IMF’s
Try these practice problems:
1. Which has the highest boiling point? HBr, Kr, or Cl2
2. Which has the lowest freezing point? N2, CO, or CO2
3. Which has the highest boiling point? HF, HCL, or HBr
Group 6
Metallic bonding: arrangement of metallic atoms in the tightest pattern possible
● Electron sea model: valence electrons circulate freely among the metal cations.
● Band model: electrons are assumed to occupy molecular orbitals.
Covalent Network: network solids contain giant network of atoms covalently bonded together.
● Diamonds: Poor conductors of electricity because of large gaps in energy between occupied and
unoccupied orbitals.
● Graphite: Conductors of electricity. Its structure is based on layers of carbon atoms arranged in
six fused rings
● Silica: Based on interconnected SiO4 tetrahedral.
Triple Point Graph: shows what phase exists at a given temperature and pressure in a closed system.
● Triple point: temperature at which all three phases exist simultaneously
● Critical point: defined by the critical temperature and pressure
○ Critical temperature: the temperature above which the vapor cannot be liquefied no
matter the applied pressure.
○ Critical pressure: the pressure required to produce liquefaction at the critical temperature.
Group 7
Group 8
Chemistry Chapter 11 Study Guide
Freezing Point:
● The temperature at which a substance converts to a solid state from a liquid.
● D
T=Kbmsolute
● Use molality because molarity changes with temperature.
● Freezing point will depress because added molecules interfere with the lattice structure
of ice.
Melting point:
● The temperature at which a substance goes from the solid state to the liquid state.
Vapor Pressure:
● Raoult’s Law: Psoln=XsolventP0solvrent
● Take the mole fraction of the solvent and multiply it by the original pressure of the
solvent.
● The pressure of a solvent in a closed container decreases with the addition of a nonvolatile substance, because less solvent molecules can escape to the vapors to the
gaseous state.
●
N is directly related to pressure; fewer molecules equal less
pressure.
Group 9
Chapter 12 - Chemical Kinetics
Chemical Kinetics - Study of the factors that control the rate (speed) of a chemical reaction.
Rate:Change in concentration of a reaction over time.
Rate Laws
● Differential Rate Law: describes the rate as a function of concentration.
●
●
●
●
Rate Law = =k
k=rate constant
n=order (not related to coefficient in balanced equation)
Integrated Rate Law: describes concentration as a function of time.
●
Rate Law = k
○
○
○
○
○
If n=0, then [A]=-kt +
=
If n=1, then ln[A]=-kt + ln
=
If n=2, then
=kt +
○
=
● K is equal to the absolute value of the slope concentration v. time plot.
Reaction Mechanisms: Elementary steps that lead to an overall reaction.
● An elementary step is a rate law for the step that can be written from the molecularity of
the reaction.
● Two requirements for an acceptable mechanism:
○ the elementary steps sum to give the correct overall balanced equations
○ mechanism agrees with experimentally determined rate law
● Slowest step is the rate determining step
Kinetic Models: The simplest model to account for reaction kinetics is the collision model.
● Molecules must collide to react
● Collision kinetic energy supplies potential energy that the reactants need to form
products.
● Activation energy (Ea): is a threshold energy necessary for the reaction to occur.
● Collision orientation of the reactants is a reaction rate determining factor.
● Arrhenius equation: k=A
, R= 8.314Jmol*K T= Temperature in Kelvin
Catalyst:
● Speeds up reaction without being consumed.
● Lowers the Ea of the reaction.
● Enzymes are biological catalysts.
○ Homogeneous exists in the same phase as the reactants.
○ Heterogeneous exists in a different phase than the reactants.
Group 10
Chapter 13: Reactions and the Equilibrium Constant
Equilibrium
●
●
Dynamic (forward and reverse reactions are always occurring at equilibrium)
Rate of forward reaction = rate of reverse reaction
Equilibrium Expression & Constant
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●
●
●
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K = ratio of products to reactants at equilibrium
αA + βB <==> σS + τT…
K = [S]σ[T]τ…/[A]α[B]β…
Does not change unless temperature changes (NOT concentration or pressure)
If K<1, equilibrium lies to the left (more reactants than products at equilibrium)
●
●
If K>1, equilibrium lies to the right (more products than reactants at equilibrium)
Solids and liquids are not included in the equilibrium expression
Manipulating the Equilibrium Constant
●
●
●
Equation is flipped: K becomes K-1
Equation is doubled: K becomes K2
Equation is halved: K becomes ÖK
Le Châtelier’s Principle
● Ex: 2A(g) + B(g) <==> C(g) + D(g) + energy
● Add A, reaction shifts right, no change in K
● Remove C, reaction shifts right, no change in K
● Increased temperature shifts reaction left, decreases K (more reactants/less products)
● Decreased temperature shifts reaction right, increases K (more products/less reactants)
● Increased volume, decreases pressure, shifts reaction left (side with more moles to increase
pressure back to normal), no change in K
● Decreased volume, increases pressure, shifts reaction right (side with less moles to
decrease pressure back to normal), no change in K
Example:
At 25°C, K=0.090 for the reaction
H2O(g) + Cl2O(g) <==> 2HOCl(g)
Calculate the concentrations of all species at equilibrium for: 1.0 mol pure HOCl is placed in a 2.0-L
flask.
[HOCl]0 = 1.0 mol / 2.0 L = 0.50 M
H2O(g)
+
Cl2O(g)
<==>
2HOCl(g)
R
I
0M
0M
0.50 M
C
+x
+x
-2x
E
x
x
0.50 – 2x
K = [HOCl]2/[H2O][Cl2O]
0.090 = (0.50 – 2x)2/(x)(x)
0.30 = (0.50 – 2x)/x
0.30x = 0.50 – 2x
2.30x = 0.50
x = 0.22
[H2O] = [Cl2O] = x = 0.22 M
[HOCl] = 0.50 – 2x = 0.06 M