Avogadro's number
6.022 x 10^23
Moles in a sample
mass/molar mass
Group 1 features
Melting point increases down groupReactivity increases down group (outer electron further from nucleus so easier to remove)Form metal hydroxides with waterForm metal oxides with oxygenForm ionic salts with halogensHave one valence elctron
Group 2 features
Melting point decreases down (excl. Mg)Reactivity increases down groupTwo valence electronsReact with water to form metal hydroxide M(OH)2 except BeForm metal oxides with oxygenForm metal halides with halogens
Transition Metals
Less reactive, more dense and higher melting points than 1&2Tungsten has highest melting point (3422)Three magnetic - iron, cobalt, nickelForm co-ordinate complexesVarious oxidation states; manganese has ten possibleForm coloured compounds
Group 14/4 - Crystallogens
Mix of metals, metalloids and non-metalsVery diverse featuresForm hydrides with hydrogen EH4Form tetrahalides with halogens EX4Four valence electrons
Group 15/5 - Pnictogens
Can form 3 covalent bondsForm pnictides with most elementsFive valence elctrons
Group 16/6 - Chalcogens
Electronegative metals and non-metalsSoft and don't conduct heat wellSix valence electronsForm -2 ions when reacting with electropositive metals
Group 17/7 - Halogens
Only group with elements in all three statesForm diatomic molecules exc. astatineReact with oxygen to form halogen oxidesReact with metal to form metal halidesHalogens are oxidising agents; halide ions are reducing agentsSeven valence electronsReactivity decreases down the group; harder to add electron
Group 18/8 - Noble Gases
Odourless, colourless, monoatomic, unreactiveForm colours when ionisedFull valence shell
Ionisation energy
Energy required to remove an electron from an atomIncreases from left to right and from bottom to top
Electronegativity
Tendency of an atom to attract an electron in a bond it shares with another atomIncreases from left to right and from bottom to top
Covalent bond
two electrons are shared by two nuclei
Hydrogen bond
attractive force between hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule, usually nitrogen, fluorine or oxygen
Ionic bond
electrostatic attraction between two oppositely charged ions
Van der Waals interactions
interaction of electron clouds between molecules
Hydrophobic interactions
occur between non-polar substances
Empirical formula
simplest whole number ratio of atoms of different elements in a compound
Ideal gas characteristics
Gas molecules make completely elastic collisionsGas molecules have zero volumeAverage kinetic energy of gas molecules is directly proportional to the temperatureGas molecules only exhibit forces due to collisions
Ideal gas law
PV = nRTpressure (atm) x volume (L) = number of moles x constant x temperature (K)Charle's Law - volume and temp are directly proportional when pressure is constantBoyle's Law - pressure and volume are inversely proportional when temperature is constant
Partial pressure
amount of pressure contributed by any gas in a mixture; total pressure of the mixture multiplied by the mole fraction of the particular gasP(a) = X(a)*P(total)
Dalton's Law
total pressure exerted by a gaseous mixture is the sum of the partial pressures of each component gas
Diffusion
spreading of one gas into another gas or an empty area
Effusion
spreading of a gas from high pressure to very low pressure through an opening smaller than the average distance between the gas molecules
Collision Theory
in order for a reaction to occur:-kinetic energy of the colliding molecules must pass the threshold energy called the activation energy-colliding molecules must be in the correct spatial orientation
Rate law
rate = k[A]^a[B]^bfor the chemical reaction aA + bB = cC + dD
Reaction order
in reaction aA + bB = cC + dD the superscripts a and b are the order of each respective reactant and the sum of them is the overall order of the reaction
Chemical equilibrium
when the rate of the forward reaction is equal to the rate of the backward reaction
System (in thermodynamics)
the object that experiences a thermodynamic transformation
Surroundings (in thermodynamics)
any part of the universe that is in direct contact with the system
Open systems
exchange both mass and heat/energy with their surroundings
Closed systems
exchange heat/energy but not mass
Isolated systems
do not exchange heat or mass
Work
energy transfer that is not heat
Internal energy
average total mechanical energy (kinetic and potential) of the particles that make up the system; for a reaction in a system with constant volume, no internal work is done so the change in internal energy is equal to the heat
Heat
the transfer of energy from a warmer body to a cooler body via conduction, convection or radiation
First Law of Thermodynamics
energy of the system and its surroundings is always conserved, so any change to a system must equal the heat flow in the system plus the work done on the systemΔU = Q + Wwork done on the system = +Wwork done by the system = -Wheat added to the system = +Qheat given off by the system = -Q
Second Law of Thermodynamics
heat cannot be completely changed to work in a cycle-like process
Temperature
in liquids, directly proportional to the translational kinetic energy of its moleculesin gases, the greater the random translational energy per mole of gas, the greater the temp
Kelvin
Celsius + 273.15
Enthalpy
H = U + PVsystem that releases heat (exothermic), has a negative ΔH because the enthalpy of the products is higher than of the reactantssystem that absorbs heat (endothermic), has a positive ΔH
Enthalpy of formation
ΔHf reaction = ΔHf products - ΔHf reactants
Entropy
represents degree of disorder; increases from solid to liquid to gas
Gibbs Free Energy
determines if a reaction is spontaneous or notΔG = ΔH - TΔSH = enthalpy of system (kJ/mol)T = temp (K)S = entropy of system (J/K/mol)if ΔG < 0 at constant pressure, reaction is spontaneousif ΔG > 0, reaction is not spontaneousif ΔG = 0, reaction is in a state of equilibrium
Units of Concentration
Molarity (M) - number of moles of solute divided by volume of solution (mol/L)Molality (m) - number of moles of solute divided by kilograms of solution (mol/kg)Mole fraction - number of moles of a compound divided by the total moles of all species in solution, ratio
Solubility
the limit of solute that can be dissolved in a given amount of solvent at equilibrium; when the rate of dissolution and precipitation are equal, the solution is saturated
Solubility product (Ksp)
equal to the aqueous products over reactants raised to the power of their coefficients
Bronsted and Lowry definitions
acid - anything that donates a protonbase - anything that accepts a proton
Arrhenius definition
acid - produces hydrogen ions in aqueous solutionbas - produces hydroxide ions in aqueous solution
Lewis definition
acid - accepts a pair of electronsbase - donates a pair of electrons
pH
measure of hydrogen ion concentrationpH = -log[H+]each point of the pH scale represents a tenfold difference in ion concentration
Conjugate acids and bases
when an acid and base react, they produce their corresponding conjugate base and conjugate acid respectivelystronger acid; weaker conjugate basestronger base; weaker conjugate acidstrong acids and bases dissociate completely in water
Polyprotic acids
acids that can donate more than one proton
Acid-base equilibrium equations
HA(aq) + H2O <-> A-(aq) + H3O+(aq)Ka is used in the Henderson-Hasselbach equation when determining pH of an acid-base reaction at equilibriumKw = [H3O+][OH-] = 1.0x10-14Ka = [H3O+][A-]/[HA]KaKb = KwpKa + pKb = pKw = 14pKa = -log(Ka)pKb = -log(Kb)pH = pKa + log([base]/[acid])
Logarithm rules
log(x*y) = log(x) + log(y)log(x/y) = log(x) - log(y)log(x^y) = y*log(x)logb(b) = 110^log10(M) = M
Buffers
solutions that resist changes in pH when a small amount of base or acid is added; composed of a weak base and its salt or a weak acid and its salt
OILRIG
oxidation is loss of electronsreduction is gain of electrons
Reduction half reaction
shows the half reactions all as reduction potential, to find oxidation potential simply reverse half equation and swap sign of potential
Galvanic cells
turns chemical energy into electrical energy using spontaneous redox reactions e.g.Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s)in this example, oxidation occurs at the Zn electrode (anode) and reduction at the Cu electrode (cathode)
Potential of cell
Ecell = Ecathode - EanodeZn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s)standard reduction potential of Zn = -0.76standard reduction potential of Cu = +0.34so Ecell = 0.34 - (-0.76) = 1.10V
Electrolytic cell
A current is used to drive a non-spontaneous redox reaction