E_k = (mv²)/2

E_k = kinetic energy

m = mass

v = velocity

E_p = mgh

E_p = potential energy

m = mass

g = gravity of the planet. On Earth it is 9.807 m/s²

h = height above a standard level. The standard level is arbitrary; the difference in height is what matters.

The kinetic and potential energy of the molecules and particles making up a substance. It is made from heat (q) and work (w) added together.

ΔU = q + w

E_tot = E_k + E_p + U

E_tot = total energy

E_k = kinetic energy

E_p = potential energy

U = internal energy

When energy is transferred into or out of the system and moves an object in the surroundings. 

When work is being done on the system, w is positive (add energy).

WHen work is being done by the system, w is negative (take away energy).

A form of internal energy. The molecules of a hot gas have higher kinetic energy since they are moving faster.

When there is a temperature difference between the system and the surroundings, if the system is not insulated, it is transfered. Heat always flows from where there is a high temperature to where there is a low temperature.

When the system absorbs heat q is positive (add energy).

When the system releases heat q is negative (take away energy).

A system may absorb heat energy from heat or work.

Often the change in an amount is more important than the actual values. What happens in the middle is irrelevent. This general formula is used for U, T, h, and many other things.

ΔX = X_f - X_i

ΔX = Difference of X

X_f = final value of X

X_i = initial value of X

Energy is always conserved; it cannot be destroyed or created, only transfered between different forms. The change in internal energy is equal to the sum of energy and work. Remember that q and w may be negative.

ΔU = q + w

ΔU = change in internal energy

q = heat

w = work

Using the equation for pressure-volume work, we can replace w with -PΔV to get

ΔU = q -PΔV.

If temperature and pressure are fixed, q = ΔH, so

ΔU = ΔH - PΔV

The heat (q) that a chemical reaction either evolves or absorbs.

If the reaction absorbs energy, it is positive.

If the reaction evolves energy, it is negative.

When the volume of a vessel decreases, pressures increases. Work is done on the system; w is positive.

When volume of a vessel increases, pressure decreases. Work is done by the system; w is negative.

Pressure-volume depends on outside pressure and the change in volume.

w = -PΔV

w = -PΔV

w = work

P = pressure outside the vessel

ΔV = change in volume

A value dependent on internal energy, pressure, and volume. All of these are state functions, so enthalpy is a state function. It is an extensive propery: it depends on the amount of a substance.

In a reaction where temperatures and pressure are fixed (but volume may change), 

q = ΔH

q = U + PV

H = U + PV

H = enthaply

U = internal energy

P = pressure

V = volume

The change in enthalpy in a chemical reaction at a fixed temperature.

It can be calciulated from BEs if all the reactants are gases.

ΔH_rxn = (energy in broken bonds) - (energy in formed bonds)

QUESTION CARD

A baseball weighing 143g travels 75 miles per hour (33.5 m/s). What is the kinetic energy of the baseball in Joules? In calories?

ANSWER

80.2 J

19.2 cal

A system in which both mass and energy may leave and/or enter freely.

Example: a beaker.

A system where neither mass nor energy can move in or out.

Example: an insulated, closed container.

A type of measurement that is independent on the size of the sample.

Example: temperature.

A type of measurement that depends on the size of the sample.

Example: heat.

A reaction where a substance reacts with oxygen. The molecules are made more stable, evolving heat and light energy. An activation energy is required. All combustion reactions are exothermic. All combustion reactions are redox reactions.

CO₂ and H₂O are often produced, but not always. The products always include one or more oxides.

"Heat measure". The science of measuring quantities of heat generated, consumed, or dissipated by a sample.

First difference in temperature is measured, then it is converted into the amount of heat.

The energy needed to raise 1 mole of a substance 1 degree Celcius (or Kelvin).

Measured in J/°mol

q = CΔT

q = ncΔT

q = msΔT

q = heat

C = heat capacity

ΔT = change in temperature

n = number of moles

c= molar heat capacity

m = mass

s = specific heat

A calorimeter made out of a coffee cup.

It has constant pressure on it (from the atmosphere), so when working with them q = ΔH.

A sealed calorimeter in a tank of water. The temperature of the water is measured to calculate q. The bomb is at constant volume, so pressure may change.

Therefore q ≠ ΔH.

QUESTION CARD

50.0mL of dilute AgNO₃ is added to a solution with OH^- ions in a coffee cup calorimeter. AgNO₃ precipitates and the temperature of the solution rises from 23.78°C to 25.19°C. Assume that the mixture has the same specific heat of water. Calculate q_surr for a mass of 150.0g. Is the reaction exothermic or endothermic?

ANSWER

q_surr = 885 J

The reaction is exothermic.

QUESTION CARD

A 48.9g of a mystery metal is 95.72°C. It is added to 43.58g of water, which is at 23.84°C. The final temperature of the water (with the metal in it) is 28.37°C. What is the mystery metal's specific heat in J/°g?

ANSWER

s = 0.251 J/°g

QUESTION CARD

When ice at 0°C melts to liquid water, it absorbs 334J for every gram of ice. Suppose the heat needed to melt a 35.0g ice cube is absorbed from 0.210kg of water at 21.0°C in a calorimeter. What is the final water temperature?

ANSWER

6.6°C.

*Note this is above freezing point: beware for all party-throwers.

If equation A is the sum of equations B and C, the change in enthalpy of equation A is equal to the sum of the changes of enthalpy in equations B and C. Because enthalpy is a state function, it doesn't matter what path you take.

It is useful for determining the enthalpy changes of reactions that are difficult to do or difficult to measure.

QUESTION CARD

An aqueous solution of hydrogen carbonate (baking soda in water) reacts at constant pressure with hydrochloric acid to produce sodium chloride, water, and carbon dioxide gas. The reaction absorbs 12.7 kJ of heat for each mole of sodim carbonate.

Write the thermochemical equation for this reaction.

ANSWER

NaHCO₃(aq) + HCl (aq) --> NaCl(aq) + H₂O(l) + CO₂(g); ΔH = 127 kJ

QUESTION CARD

Using the following equation, give the thermochemical equation for 1 mole of liquid water dissociating into hydrogen and oxygen gases.

2H₂(g) + O₂(g) --> 2H₂O(l); ΔH = -572 kJ

ANSWER

H₂O(l) --> H₂(g) + (1/2)O₂(g); ΔH = 286 kJ

QUESTION CARD

N2(g) + 3H2(g) --> 2NH3(g); ΔH = -91.8 kJ

How much heat is evolved from 9.07e5 g of ammonia?

The conditions of a reaction when it is said to be standard, e.g. standard molar enthalpy.

1 atm (or 1 bar, or 10⁵ Pa)

1 mole

25°C

The enthalpy change in 1 mole of a substance as it is formed from its elements in their standard states in standard conditions. The value for common compounds may be found in the textbook appendix. We can find the standard enthalpy change of any reaction as long as we have these values for all the products and reactants.

ΔHº = ΣnΔHº_f (products) - ΣnΔHº_f (reactants)

The state an element is most stable in in standard conditions. Each element has one. Most are a single atom of the element. Exceptions are as follows:

H₂, N₂, O₂, P₄, S_8, Cl₂, Br₂, I_2.

Bromine and mercury are the only liquids. Hydrogen, helium, nitrogen, oxygen, fluorine, neon, chlorine, argon, krypton, xenon, and radon are the gasses. Everything else is solid.

Some elements have allotropes. One of the allotropes is the standard state.

QUESTION CARD

What is the enthalpy of reaction (ΔH) for the formaion of tunsten carbide from its lements in standard state?

W(s) + C(s, graphite) --> WC (s)

2W(s) + 3O₂(g) --> WO₃(s); ΔH = -1685.8 kJ

C(s, graphite) + O₂(g) --> CO₂(g); ΔH = -393.5 kJ

2WC(s) + 5O₂(g) --> 2WO₃(s) + 2CO₂(g); ΔH = -239.18 kJ

ANSWER

ΔH = -40.5 kJ

QUESTION CARD

What is the ΔHº_vap of CS₂ in standard conditions?

ΔHº_fCS₂(l) = 89.7 kJ

ΔHº_fCS₂(g) = 116.9 kJ

ANSWER

ΔHº_vap = 27.2 kJ

QUESTION CARD

What is the standard enthalpy for this reaction?

4NH₃(g) + 5O₂(g) --(Pt)--> 4NO(g) + 6H₂O(g)

ΔHº_fNH₃(g) = -45.9 kJ

ΔHº_fNO(g) = 90.3 kJ

ΔHº_fH₂O(g) = -241.8 kJ

ANSWER

ΔHº = -906.0 kJ

A measure of the average strength of a bond between two specific elements. Obtained from enthalpies of reactions. It will be different in different textbooks since it is an average. Determined by breaking the bonds in a molecule and dividing the energy by the number of bonds. The energy to break a bond is the negative of the energy to form the same bond. It is defined as the ΔH of the dissociation of the bond into its atom not ions or molecules.

Double bonds are not twice the BE of single bonds, nor are triple bonds thrice the BE of single bonds. However double bonds have bigger BEs than single bonds and triple bonds bigger BEs than double bonds.

Invented in 1919 by Max Born and Fritz Haber. A way of calculating lattice energy using Hess's law with these 5 reactions:

This example uses NaCl as the ionic solid.

1. Sublimation of the first consituent (Na solid into gas)

2. Ionization of first constituent (Na into Na⁺)

3. Dissociation of second consituent (Cl₂ into Cl)

4. Electron gain of second constituent (Cl to Cl^-)

5. Brining the two constituents together (Na⁺ and Cl^- form NaCl)

The enthalpy of an ionic solid. Depends on the charge of the ions and their bond length. Depends on Coulomb's Law.

V = (k*Q₁*Q₂)/r

k = 8.99e9 J*m/C² (a constant)

Q = charges of ions

r = bond length in meters

The measure of how dispersed the energy of a system is among all the different possible ways that a system can contain energy. Measured in J/K. It is an extensive property (relies on size of system).

q = TΔS

It can be based on the positional (pressure) or thermal (heat) of the situation.

A reaction that occurs without any external intervention. ΔH is negative and ΔS is positive: the reaction is always spontaneous.

ΔS_universe > 0

ΔS_system > q/T

A reaction that requires external intervention in order to occur. Only spontaneous in the reverse direction.

ΔS < 0 this defies the Second Law of Thermodynamics.

ΔH is positive and ΔS is negative: the reaction will never happen.

ΔSsystem > q/T

A reaction that can be undone, e.g. melting water; it can be refrozen. Burning paper is irreversible.

q_rev is the heat supplied reversibly.

ΔS = q_rev/T

The total entropy of the universe is always increasing.

ΔS_total = ΔS_universe = ΔS_system + ΔS_surroundings

*Note that ΔS_system can be negative, as long as ΔS_universe is positive its ok.

The change in entropy of a reaction. If you have the standard entropies of reactans and products it can be calculated.

ΔSrxn = ΣnS°_m(products) - ΣnS°_m(reactants)

The temperatures at which a substance melts or boils, respectively.

At these points the system is at equilibrium.

ΔS° = (ΔH°/T)

Potential energy released from bringing to oppositely charged ions together is found by:

E = (KQ₁Q₂)/r

K = 8.99e9 J*m/C²

Q = charges of the two ions

r = distance between the ions

This gives us E for one pair of atoms.

Ionic bonds that share more electrons have higher potential energy. And ions with smaller diameter (closer to each other) will have higher potential energy.

The number of molecules in a mole.

6.02e23

QUESTION CARD

CCl₄(l) --> CCl₄(g); ΔH = 39.4 kJ

If one mole of liquid carbon tetrachloride at 25°C has an entropy of 216 J/K. What is the entropy of 1 mole of tetrachloride vapour at 25°C.

ANSWER

348 J/mol*K

A criterion for spontaneity of a reaction. ΔG is a state function (independent on path).

G = H - TS

ΔG = ΔH° - TΔS°

ΔG° = ΔH° - TΔS°

As a reaction progresses entropy decreases and Gibb's energy increases. If a reaction is done so that every single bit of energy is used up, no entropy will be created, however this can never happen in real life.

ΔG is always negative in spontaneous reactions; free energy decreases. A reaction is at equilibrium when ΔG is as small as possible.

In the the reaction

aA + bB --> cC + dD

Q_c = ([C]^c[D]^d)/([A]^a[B]^b)

Q_p = (P_C^cP_D^d)/(P_A^aP_B^b)

At equilibrium Q = K. If Q > K, the reaction will proceed towards the reactants. If Q < K, the reaction will proceed towards the products. This number is dimensionless.

Expresses the partial pressures (in atm) or the concentrations (in M) of things in equilibrium. Q is this value when not at equilibrium.

aA + bB -> cC + dD

K = ([A]^a[B]^b)/([C]^c[D]^d)

K = (P_A^aP_B^b)/(P_C^cP_D^d)

The hypothetical charge of an atom if it were on its own. Rules for assigning oxidation numbers:

1. Lone atoms or atoms in single-element molecules have oxidation number of 0.

2. Oxidation number of an ion is equal to the charge.

3. Except in peroxides (H and O only in the compound), oxygen has an oxidation number of -2.

4. Except in binary compounds with metals (when it is -1), hydrogen has an oxidation number of +1.

5. The sum of the oxidation numbers should equal the overall charge of the molecule.

QUESTION CARD

Write the cell reactions for 

a)  Tl(s)|Tl⁺(aq)||Sn²⁺(aq)|Sn(s)

b)  Zn(s)|Zn²⁺(aq)||Fe³⁺(aq),Fe²⁺(aq)|Pt 

ANSWER

a)  2Tl(s) + Sn²⁺(aq) --> 2Tl⁺(aq) + Sn(s)

b)  Zn(s) + 2 Fe³⁺(aq) --> Zn²⁺(aq) + 2Fe²⁺(aq)

The difference in electrical potential of to half-reactions. The energy produced by the cell is based on it. It is measured in volts.

Electrical work = Charge*potential difference

The magnitude of a charge in one mole of electrons. Equal to 9.6485e4 C/moles of e-.

The difference between the maximum voltages of the two half-reactions, or oxidation and reduction potentials. Measured with an electric digital voltometer. If the number is negative, the reaction is non-spontaneous.

w_max = -nFE_cell

E_cell = E_cathode - E_anode

n = number of electrons transfered

F = Faraday constant

QUESTION CARD

The cell potential of this voltaic cell is 0.650 V. Calculate the maximum electrical work of this cell when 0.500 g of H₂ is consumed.

Hg²⁺(aq) + H₂(g) --> 2Hg(l) + 2H⁺(aq)

ANSWER

-3.09e4 J

The potential of the oxidation half-reaction.

Oxidation potenital = -Reduction potential of reverse reaction.

The potential of the reduction half-reaction. The ability of a species to act as an oxidizing agent. There is a list of reduction potentials, very useful when doing calculations. An intensive property (doesn't depend on the amount of a species).

Oxidation potenital = -Reduction potential of reverse reaction.

E_cell = E°_cell - (RT/nF) ln Q

or

E_cell = Eº_cell - (0.0592/n) log Q

Note: the second one only works at 25ºC.

A battery used in flashlights, radios, and other portable appliances. Voltage is aound 1.5V, but decreases over time. Not good in cold weather. The anode is zinc. The cathode is a graphite rod surrounded by a paste made of manganese dioxide, ammonium, zinc chloride, and black carbon.

Zn(s) --> Zn²⁺(aq) + 2e^-

2NH₄⁺(aq) + 2MnO₂(s) + 2e^- --> Mn₂O₃(s) + H₂O(l) + 2NH₃(aq)

Similar to the Leclanché, but potassium hydroxide replaces ammonium chloride. This makes it perform beter in cold weather.

Zn(s) + 2OH^-(aq) --> Zn(OH)₂(s) + 2e^-

2MnO₂(s) + H₂O(l) + 2e^- --> Mn₂O₃(s) + 2OH^-(aq)

Electrodes are lead alloy grids. The anode is packed with spongy lead. The cathode is packed with lead (IV) oxide. They are soaked in aqueous sulfuric acid. Delivers 2V (adding more electrodes increases this number in factors of 2). Can be recharged by an external current. More water needs to be added, unless it is maintenance-free.

Pb(s) + HSO₄^-(aq) --> PbSO₄(s) + H⁺(aq) + 2e^-

PbO₂(s) + 3H⁺(aq) + HSO₄^-(aq) + 2e^- --> PbSO₄(s) + 2H₂O(l)

A common storage battery. Anode is cadmium. Cathode is hydrated nickel oxide on nickel. The electrode is potassium hydroxide. Used in calculators, portable power tools, shavers, and tooth-brushes.

Cd(s) + 2OH^-(aq) --> Cd(OH)₂(s) + 2e^-

NiOOH(s) + H₂O(l) + e^- --> Ni(OH)₂(s) + OH-(aq)

A common replacement for a nicad cell. Use metal hydrides as anodes instead of cadmium; less toxic. Can undergo more discharge-recharge cucles.

Used in electrid hybrid cars.

A battery, but has a continuous supply of fuel. 

Hydrogen gas is the anode. The H⁺ ions enter the PEM and electrons enter the external circuit. The H+ ions react with oxygen at the cathode, which is taking up electrons from the external circuit. Water is formed. Potential of 0.7. It can run at low temperatures. Used in space, emergency systems, cars, and buses.

H₂(g) --> 2H⁺(aq) + 2e^-

O₂(g) + 4H⁺(aq) + 4e^- --> 2H₂O(l)

Other versions can use hydrocarbons or methanol as fuel.

A drop of water sitting on iron is a tiny voltaic cell. Water and air are the the cathode. The iron is the anode.

O₂(g) + 2H₂O(l) + 4e^- --> 4OH^-(aq)

Fe(s) --> Fe²⁺(aq) + 2e^-

To prevent rusting of iron pipes, a magnesium anode can attach to the pipe, and then the iron can't act as an anode as well, greatly reducing rust.

QUESTION CARD

Calculate standard free energy change for this reaction (need to use standard reduction table)

Zn(s) + 2Ag⁺(aq) --> Zn²⁺(aq) + 2Ag(s)

ANSWER

-301 kJ

QUESTION CARD

Calculate standard potential for this reaction (will need to use tables)

Zn(s) + Cl₂(g) --> Zn²⁺(aq) + 2Cl^-(aq)

ANSWER

E°_cell = 2.12 V

QUESTION CARD

The standard cell potenital for this reaction is 1.10 V. Calculate the equilibrium constant.

Zn(s) + Cu²⁺(aq) <--> Zn²⁺(aq) + Cu(s)

ANSWER

K_c = 2e37

QUESTION CARD

The standard cell potential of this cell is 1.10 V. What is the cell potential at 25°C?

ANSWER

E_cell = 1.22 V

1. Concentration of reactants

2. Concentration of catalyst

3. Temperature

4. Surface area, if a solid reactants or catalyst is involved.

The increase in molar concentration of products (or decrease in reactants) per unit of time. Measured in mol/L*s. Determined by removing samples at different points in a reaction, or monitoring a physical property of the reaction (volume, temp, colour).

Average rate = Δ[A]/Δt

QUESTION CARD

How is the rate of formation of NO₂F related to the formation of F2 in this reaction?

2NO₂(g) + F₂(g) --> 2NO₂F(g)

ANSWER CARD

(1/2)(Δ[NO₂F]/Δt) = -(Δ[F₂]/Δt)

An equation that tells the rate of reaction in comparison to concentration of reactants and catalysts, raised to various powers.

Rate = k[A]^m[B]ⁿ

A factor in the rate law. A constant at fixed temperature. It depends on 

1. Z: the collision frequency

2. f: fraction of collisions having energy higher than E_a.

3. p: the fraction of collisions with the molecules properly oriented.

Expresses the dependence of rate constant on temperature. Written by Svante Arrhenius.

k = Ae ^-E_d/RT

Where A is the frequency factor.