IB CHEMISTRY HL

Boyle's Law

P1V1=P2V2 (inversely proportional)

Charle's Law

V1/T1=V2/T2

Gay-Lussac's Law

P1/T1=P2/T2

Combined Gas Law

P1V1/T1=P2V2/T2

Ideal Gas Law

PV=nRT

conditions of PV=nRT

Pascals

ideal gas

low molar mass and weak to no IMF

effective nuclear charge formula

of protons - # of protons of previous noble gas

effective nuclear charge

extent to which the nucleus is attracting the valence electrons (more difference

effective nuclear charge increase across periods

linear

atomic radius

half the difference between the nuclei of an element (depende on what it is bonded with)

atomic radius trend

increases down a group because more shells

more effective nuclear charge means

smaller atomic radius

ionic radius

half the distance between the nucleus of the cation and the anion (Average)

parent ion has _______ ionic radius than cation

higher (because loses electron and less shells)

ionic radius trend

decreases across a period and increases down a group

ionization energy

minimum amount of energy required to remove one mol of electrons from a neutral gaseous atom in its ground state

ionization energy trend

decreases from top to bottom in a group

electron \ trend

increases from left to right in a period

electron affinity

energy released when 1 mol of electrons is attached to 1 mol of neutral gaseous atoms or molecules

melting point trend

1. decreases down group 1

2. increase down 17

3. increase across a period and reach a maximum at group 14

alkali metal properties (group 1)

-very reactive -form ionic compounds with non-metals -react with H2O to produce hydroxide and gas -intensity increases down group (valence further away)

Halogen properties (group 7)

-some gas/liquid/solid -colored -very reactive -diatomic -Fluorine

sulfuric acid reaction

SO3 + H2O = H2SO4

sulfurous acid reaction

SO2 + H2O = H2SO4

ionic bond electronegativity difference

more than 1.8

coordination number

how many anions around cation in a lattice

lattice energy

how much energy needed to break ionic bond

ionic compound physical properties

brittle

volatility

ability of a liquid/solid to become a gas at room temperature

pure covalent electronegativity difference

0

polar covalent electronegativity difference

between 0 and 1.8

pure covalent properties

symmetric electron distribution

polar covalent properties

asymmetric electron distribution

dative bond

pair of electrons donated from one atom to another

dipole-dipole

attraction between polar molecules

hydrogen bonds

strong dipole-dipole

London dispersion

temporary dipole

what is responsible for partial bonds and resonance structures

delocalized pi electrons

formal charge formula

valence electrons-available electrons

formal charge characteristics

-closer to 0 more stable -used to see which suitable Lewis Structure more appropriate -most electronegative atoms will be further from ideal formal charges

linear BA

180

bent/v-shape/angular BA

118

trigonal planar bA

120

tetrahedral BA

109

bent with 2 lone pairs BA

105

trigonal pyramidal BA

107

trigonal bipyramidal BA

120 and 90

square pyramidal BA

180 and 90

square planar BA

180

ammonium

NH4+

Hydroxide

OH-

Nitrate

NO3-

Hydrocarbonate

HCO3-

carbonate

CO3 2-

Sulfate

SO4 2-

phosphate

PO4 3-

metallic bonding definition

lattice of metal cations in a sea of delocalized electrons

malleability

ability to be shaped

ductility

ability to form threads

alloys

homogeneous mixture of two metals or metal with another substance. metals combine to improve the individual qualities of the metals

hybridization

overlap of atomic orbitals to create hybrid orbitals that can form covalent bonds

radical

substance or species with one or two unpaired electrons

conditions when measuring heat change

298K

specific heat capacity

amount of energy needed to increase temperature of a substance by 1C or 1K

enthalpy change

amount of energy released or absorbed PER MOLE of substance

Hess' Law

enthalpy of any reaction is independent of the route you take

standard enthalpy change of formation

enthalpy change when one mole of a substance is formed from its elements in standard states

bond enthalpy

energy required to break one mole of bonds in gaseous molecules under standard conditions

wavelength formula

(planck's x speed of light) / (enthalpy/avogadro)

Lattice Enthalpy

enthalpy change when one mole of a solid ionic compound is separated into its respective gaseuous ions (BORN HABER CYCLE)

factors affecting lattice enthalpy

-ionic radius (assume spherical shape) -ionic charge (Coulomb's Law) -STP

Enthalpy change of solution

enthalpy change when 1 mole of an ionic substance dissolves in water to give a solution of infinite dilution (100% of sample diluted)

steps in forming liquid solution (particles)

-separate solute -separate solvent -allow solute and solvent to interact

enthalpy change of hydration

Enthalpy change when 1 mole of aqueous ions is formed from gaseous ions

entropy

the tendency of a system to become disorganized

negative ΔS

less disorder

positive ΔS

more disorder

absolute entropy

ΔS reaction= ΔSproducts-ΔSreactants ΔS surrounding = -ΔS system/ T absolute

gibbs formula

G = H - TS

at low temps (exo)

ΔG is < 0

at low temps (endo)

ΔG is > 0

at high temps (endo)

ΔG is < 0

at high temps (exo)

ΔG is > 0

chemical equilibrium characteristics

rate of forward/reverse reactions equal amount of reactants/products constant

Q meanings

if Q > Kc, left

if Q = Kc, equilibrium

if Q > Kc, right

Kc formula

Le Chatelier's principle

if a system is at equilibrium and a change is made that disturbs the equilibrium, then the system responds in such a way as to counteract the change and eventually a new equilibrium is established

concentration effect on Kc

adding more reactants will favour the products side

pressure effect on Kc

added pressure, shift to side with less moles

temp effect on Kc

if temp increases, shift in endo direction

how to get overall Kc

multiply individual ones

Gibbs free energy

the energy of a system that is available to do work at a constant temperature and pressure

rate of reaction can be measured through

water displacement and gas formation

collision theory

not every collision forms a reaction. successful collisions are chemical reactions and need activation energy and certain orientation

factors affecting rate of reaction

-catalyst -concentration -surface area -temperature -pressure (gases)

Maxwell-Boltzmann distribution curve

factors that change K in kinetics

catalyst and temp

winkler method

used to measure Biological Oxygen Demand (a measure of the dissolved oxygen required to decompose organic matter in water over a set time period)

lewis acids

accept pairs of electrons