Chemistry Flashcards
Define molecular weight
Molecular weight is the weight in grams of one mole (6.02 x 10^23) of the substance, or the weight in atomic mass units (which should have the same coefficient)
Define empirical / molecular formula
Empirical formula is the simplest whole number ratio between the numbers of atoms of the different elements making up the compound. E.g. H2O -> 2:1, H2O2 -> 1:1
Molecular formula states the exact number of the different atoms that make up the molecule. E.g. H2O -> 2:1, H2O2 -> 1:1
What are the rules for oxidation numbers?
- In elementary substances, the oxidation number of an element is zero (e.g. O2, Na, S8).
- In monoatomic ions the oxidation number is equal to the charge of the ion.
- In a neutral molecule the sum of all the oxidation numbers that make up the molecule is zero.
- Useful oxidation numbers to memorise (don’t always apply): H = +1, O = -2, Alkali metals = +1, Alkaline Earth metals = +2, Aluminium = +3
Definitions of quantum numbers?
n = principle quantum number (shows period on PT) l = angular momentum quantum number (shape of orbital) ml = magnetic quantum number (which of 3 planes) ms = spin quantum number (rotates which of 2 directions)
Shapes of s, p orbitals?
s orbitals are spherical
p orbitals are “dumbells”
Order for filling atomic orbitals?
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d…
Don’t forget Hund’s rule: each empty ml sub orbital is filled before a second is added to any of them, convention is to add them facing upwards first
What are hybrid orbitals? Shapes?
In bonding, often orbitals combine and become hybridised. This changes their energy and shape. It’s particularly useful for explaining molecular shapes.
- 1 s and 3 p bonding electrons (?) become 4 sp3 hybrids. This molecular shape would be tetrahedral, about the central atom with angles of 109.5°
- 1 s and 2 p bonding electrons (?) become 3 sp2 hybrids. This molecular shape would be trigonal with 120° angles
- 1 s and 1 p bonding electrons (?) become 2 sp hybrids. This would be a linear shape of 180°
Define Lewis structure, acid, base
Lewis structure is a digram which shows each atom and how many valence electrons is has. In addition it may show the bonding as well.
When Lewis structures don’t accurately represent a molecule because its constituents shift around, more than one Lewis structure can be used. This is known as resonance structures
Define octet rule, formal charge
The Octet rule refers to the fact that the configuration of noble elements is the stablest valence electron configuration, due to energy factors other elements will try to achieve this stability of 8 valence electrons through bonding, ionising et cetera
Define temperature (both C and K), gas P and weight
Celsius and Kelvin both have the same unit magnitude. Celsius is defined by the freezing point and vaporising points of water (0° and 100° respectively). Kelvin is an absolute scale, beginning at 0° K (no heat energy). 0°K = -273.15°C.
Gas pressure is the pressure exerted on the walls of the container in which this gas is placed. At constant temperature pressure is inversely related to volume (P1V1 = P2V2) , at constant volume directly related to temperature. (P1V1 / T1 = P2V2 / T2).
The weight of a gas can be calculated by the number (N) of molecules present and vice versa
Define STP, ideal gas, deviation
Standard Temperature and Pressure of a gas is 0°C and 1 atm (760 mmHg)
Define H bonds, dipole forces
Hydrogen bonds are intermolecular forces which occur whenever hydrogen is covalently bonded to an atom which attracts electrons strongly. In particular O, N and F. This molecules become very polar - stronger than usual cases of dipole dipole interaction.
Dipole dipole forces are basically the intermolecular forces between polar molecules.
London forces are intermolecular forces which occur spontaneously in molecules which aren’t “formally” polar, due to random movement of electrons. These three forces are all called Van der Wall forces
Define saturated, supersaturated, nonvolatile
A solution is said to be saturated when under normal conditions no more solute can be dissolved into the solute - adding more solute leads to precipitation.
Supersaturation is where the usual limit of saturation is exceeded. Can be achieved for example by lowering temperature and not disturbing solution.
Non volatile is used to describe solutes which cannot evaporate
Common anions and cations in solution?
Common anions in solution are:
F-, Cl-, Br-, I-, O2-, S2-, N3-, OH-, NO3-, ClO4-, CO3(2-), SO4(2-), PO4(3-), CH3CO2-
Common cations in solution are:
Na+, Li+, K+, NH4+, H3O+, H+, Ca2+, Mg2+, Fe2+, Fe3+
Units of concentration?
M = MolaRity = # moles solute / litRe
m = MolaLity = # moles solute / kiLogram
N = Normality = Equivalence (number of reacting units) / litre
ρ = Density = mass per unit volume
Osm = Osmole = number of moles of particles that contribute to the osmotic pressure of a solution
Osmolarity = osmoles / litres solution
Osmolality = osmoles / kg solution
Define electrolytes with examples
An electrolyte is a substance that ionises when dissolved in suitable ionising solvents such as water
Define Bronsted acid and base, pH
A “Bronsted and Lowry” acid is a proton donar. For example a molecule which would shed a hydrogen in solution. HA + H2O H3O+ A-
A Bronsted base is a proton acceptor. For example in aqueous solution, a base would strip a proton from H2O molecules. B + H2O HB+ + OH-
The pH of a solution is a convenient way of expressing the concentration of hydrogen ions [H+] in solution. It is defined as: pH = -log(sub 10) [H+]. Thus 10^0 which equals 1, would mean maximum hydrogen ion concentration. 10^-7 would mean the same a water. 10^-14 would mean very little hydrogen ions indeed
Examples of strong acids and bases
Perchloric acid HClO4 Chloric HClO3 Nitric HNO3 Hydrochloric HCl Sulfuric H2SO4 Hydrobromic HBr Hydriodic HI Hydronium Ion H3O+
Lithium hydroxide LiOH sodium hydroxide NaOH Potassium hydroxide KOH Rubidium hydroxide RbOH Cesium hydroxide CsOH Calcium hydroxide Ca(OH)2 Strontium hydroxide Sr(OH)2 Barium hydroxide Ba(OH)2
Kw at STP, neutral H2O pH, conjugate acid/base, zwitterions
Two water molecules can ionise into H3O+ and OH-.
At STP, Kw = [H+][OH-] = 1.0 x 10^-14.
Increases with temperature.
In a neutral solution, [H+] = [OH-] = 10^-7 M.
When an acid dissociates in water, it leaves behind a base. This is called the conjugate acid-base pair.
HA H+ + A-.
A zwitterion is a molecule which has a positive charge in one part and a negative charge in another
Ka abd Kb equations (acid and base dissociation constants)
Ka = [H+][A-]/[HA] Kb = [HB+][OH-]/[B]
pKa and pKb equations (logarithmic constants)
pKa = -log10 Ka pKb = -log 10 Kb
Kw, pH and pOH equations
Kw = [H+][OH-] = 1.0 x 10^-14 = ion product constant
pH = -log10 [H+]
pOH = -log10 [OH-]
Equivalence point, indicator, rules of logarithms
Equivalence point in a chemical reaction is the point at which chemically equivalent quantities of acid and base have been mixed. It can be found by means of an indicator, most often phenolphthalein.
Titration is used to determine the concentration of a given sample of acid or base - a solution of known concentration is added until an equivalence point is reached. An indicator changes colour based on the acidity of a solution and is thus used to determine the pH.
log(sub a) a = 1
log(sub a) M^k = k log (sub a)M
log (sub a) (MN) = log(sub a)M + log(sub a)N
log (sub a) (M/N) = log(sub a)M - log(sub a)N
10 log 10^M = M
Examples of common weak acids and bases
Hydrocyanic HCN Hypochlorous HClO Nitrous HNO2 Hydroflouric HF Sulfurous H2SO3 Hydrogen Sulfide H2S Phosphoric H3PO4 Benzioc, acectic and other carboxylic acids
Carbonate ion CO₃²⁻ Methyl amine CH₃NH₂ Ammonia NH₃ Bicarbonate ion HCO₃⁻ Pyridine C₅H₅N Aniline C₆H₅NH₂
Define state function
A state function is also known as a path independent function. It’s a quantity which is viewed independently of the events that lead it being that way, all that is considered is the current state of it
Conversion: thermal to mechanical energy
Historically thermal energy was measured in calories, that same measure was used for other forms of energy such as mechanical.
The change in energy of a system is given by the equation:
ΔE = Q + W
Where Q is the heat absorbed by the system ( > 0), or released by the system (Q < 0). W is work done by the system on its surroundings (W > 0), or by the surroundings on the system (W < 0)
Define endothermic and exothermic
Exothermic is a reaction where heat is released (H < 0).
Endothermic is a reaction in which heat is required (H > 0)
Gibbs free energy?
A state function used as a criterion for spontaneity.
ΔG = ΔH - TΔS
H is the enthalpy of the system in a given state.
T is the temperature.
S is the entropy of the system.
Note that if TΔS is more negative than ΔH, ΔG will be negative. TΔS positive and greater than ΔH will be negative too.
If both are negative and ΔH is greater it ΔG will be positive. If TΔS is positive and less than positive ΔH, it will be positive.
Reaction is spontaneous if ΔG < 0.
Not spontaneous if ΔG > 0.
State of equilibrium is ΔG = 0
Reaction order
The rate of a reaction can be expressed as a function of the concentration of the reactants.
Rate = k [A]^m [B]^n
Where [ ] is the concentration of the corresponding reactant in moles per litre.
k is the reaction constant.
m is the order of reaction with respect to A.
n is the order of reaction with respect to B.
According to the above rate law, the reaction is said to be an (m+n)th order reaction
Define rate determining step
The rate of the overall reaction is limited by the slowest step, therefore the rate determining step in the mechanism of a reaction is the slowest step
Generalised potential energy diagrams
Potential energy diagrams show the change in enthalpy of a reaction over time with two axis: potential energy and reaction progress.
The energy of activation can be seen as the hump that is required for the change to a different state.
Going from a lower energy state to a higher one would be endothermic.
Starting out at a higher energy state than finishing would me energy was released, therefore it was an exothermic reaction
Define activation energy, catalysis
Activation energy is the energy required for a particular reaction to get started.
Catalysis is the increase in the rate of a chemical reaction due to the presence of a catalyst. These lower the activation energy required without losing their form
Define saturation kinetics, substrate
When the concentration of a substrate is great enough to occupy all available active sites on a catalyst, adding further substrate will impact not effect the rate of reaction. This is called saturation kinetics.
Catalysts tend to work on very specific reactants. These reactants are known as substrates and they typically interact with the active site of a catalyst
Define anode, cathode, anion, cation
An anode is the electrode that the electrons flow out of. In an electrolytic cell electrons flow out of the anode to the battery.
A cathode is the electrode that receives the electrons. In an electrolytic cell, electrons flow out of the battery into the cathode.
An anion is a negatively charged ion.
A cation is a positively charged ion.
*LEO is A GERC (Lose Electrons Oxidation is Anode, Gain Electrons Reduction at Cathode)
Define standard half-cell potentials
The oxidation/reduction capabilities of substances are measured by their standard half-cell potentials “Eo”. They are typically tabulated with standard conditions.
The more positive the Eo value, the more likely the reaction will occur spontaneously as written. The strongest oxidising agents will have the highest values, the strongest reducing agents will have the most negative values.
Remember the oxidising agent is reduced and the reducing agent is oxidised
Define strong/weak oxidising/reducing agents
The strongest oxidising agents have high Eo values.
The strongest reducing agents have negative Eo values.
Remember electrons are always on the left hand side of a reduction half reaction. And always on the right hand side in an oxidising half reaction.
This is electrochemistry, where redox reaction are used to transfer electrons from one ionic species to another. For example in a galvanic cell where a current is created by two “half reactions” connected by a conductor to create the circuit for the current to flow
