DP chem SL everything

mixture

a combination of pure substances that retain their individual properties

example of a heterogenous mixture

soil, blood, sand and water, oil and salt etc

example of a homogenous mixture

rain, air, saltwater, steel, brass, coffee etc.

crystallisation

• used to separate a dissolved solid from a solution, when the solid is more soluble in hot solvent than in cold (e.g. copper sulphate from a solution of copper (II) sulphate in water)

• based on their ease of evaporation

recrystallisation

used to purify impure solids

simple distillation

used to separate a liquid and soluble solid from a solution (e.g. water from a solution of saltwater) or a pure liquid from a mixture of liquids

fractional distillation

separate two or more liquids that are miscible with one another (e.g. ethanol and water from a mixture of the two)

paper chromotography

separate substances that have different solubilities in a given solvent

solids

• fixed volume and shape

• they have a high density

• atoms vibrate in position but can’t change location

• particles are packed very closely together in a fixed and regular pattern

liquids

• fixed volume but adopt the shape of the container

• generally less dense than solids, but much denser than gases

• particles move and slide past each other which is why liquids adopt the shape of the container and also why they are able to flow freely

gases

• do not have a fixed volume

• take up the shape of the container (fill the whole container)

• very low density

• can be compressed into a much smaller volume

• particles are far apart and move randomly and quickly in all directions

• particles collide with each other and with the sides of the container

change of states graph

• 1&2 = particles are vibrating and gaining kinetic energy + temperature rises

• 2&3 = all the energy goes into breaking bonds; no increase in kinetic energy or temperature

• 3&4 = particles are moving around and gaining kinetic energy

• 4&5 = substance is boiling; bonds are breaking; there is no increase in kinetic energy or temperature

• 5&6 = particles are moving around rapidly and increasing in kinetic energy

isotopes

different atoms of the same element with the same number of protons + electrons but a different number of neutrons

relative atomic mass

average mass of one atom of an element compared to 1/12 of the mass of an atom of carbon-12

total mass of 100 atoms = (% abundance_A x mass_A) + (% abundance_B x mass_B) 

electromagnetic spectrum

• range of frequencies that covers all electromagnetic radiation and their respective wavelengths and energy

• shows the relationship between frequency, wavelength, and energy

• frequency = how many waves pass per second

• wavelength = distance between two consecutive peaks on the wave

frequency is … to wavelength

inversely proportional

emission spectra

• produced when an atom moves from a higher to a lower energy level

• energy emitted is a mixture of different frequencies

principal quantum numbers (n)

• used to number the energy levels or quantum shells

• the lower the principal quantum number, the closer the shell is to the nucleus

• mathematical relationship between the number of electrons and the principal energy level is 2n² 

ground state

• most stable electronic configuration of an atom which has the lowest amount of energy

• achieved by filling the subshells of energy with the lowest energy first

atomic orbital

• region around an atomic nucleus where there is a 90% chance of finding the electron

• shape depends on electron’s energy

avogadro constant

• 6.02 x 10²³ mol^-1

• mole (mol)

• SI unit of amount of substance

moles/particles triangle

moles/mass triangle

molecular formula

• shows the actual number and type of each atom in a molecule

• (mass of empirical formula)x

empirical formula

simplest whole number ratio of the elements in a compound

molar concentration

amount of solute per volume, given in mol dm^-3

moles/concentration triangle

standard solution

solution of known concentration

ultraviolet-visible spectroscopy

• uses direct relationship between concentration of solution and its absorbance

• absorbance is measured → results plotted in a calibration curve

avogadro’s law

equal volumes of all gases, when measured at the same temperature and pressure, contain an equal number of particles

ideal gases

• negligible volume

• no intermolecular forces between particles

• gas particles have range of speeds and move randomly; average kinetic energy is proportional to temperature

• collisions are elastic; kinetic energy is conserved

gases in a container exert … as the gas molecules are constantly … with the walls of the container

pressure … colliding

pressure is … to volume

inversely proportional

volume is … to temperature

directly proportional

temperature is … to pressure

directly proportional

constant relationships in gases

• PV = a constant

• V / T = a constant

• P / T = a constant

ideal gas equations

PV = nRT

• P = pressure (pascals, Pa)

• V = volume (m³)

• n = number of moles of gas (mol)

• R = gas constant (8.31 J K^-1 mol^-1)

• T = temperature (Kelvin, K)

real gases

• volume is not negligible

  • collisions more frequent than predicted

  • pressure is greater

• forces between particles are present

  • reduces speed of colliding particles

  • pressure is lower

density

mass/volume

if volume of gas particles is not negligible…

PV/nRT > 1

if attractive forces between particles are present…

PV/nRT < 1

ionic bonds

• transfer of electrons (between metal + non metal)

• pulled together by electrostatic attraction

• attraction is very strong and requires a lot of energy to overcome

metals have … ionization energies

low; positive ions

non-metals have … ionization energies

high; negative ions

polyatomic ions

• more than one ion which together lost or gained an electron

• covalent bonding

• eg. NH₄⁺

why are noble gases unreactive?

• high ionization energies

• stable electron shells (full)

lattice enthalpy

strength of ionic bond

the … the ionic charge, the … the attraction between ions, the … the lattice enthalpy

higher … higher … higher

the … the ionic radius, the … the attraction between ions, the … the lattice enthalpy

higher … lower … lower

ionic compounds

• high MP

• high BP

• low volatility

• solid at room temperature

• soluble in water, insoluble in non-polar

• conduct electricity only when molten

• generally brittle

% ionic character

difference in electronegativity/3.2

covalent bond

• sharing of electrons

• electrostatic attraction between shared pair of electrons and positively charged nuclei

• forms a molecule

… and … are exceptions to the octet rule

• Boron and Beryllium

• (to Be or not to B)

the … the atomic radius, the … the bond length

larger … longer

the … the bond length, the … the bonds

shorter … stronger

coordination bond

• covalent bond in which both shared electrons originate from the same atom

• eg. CO

VESPR model

valense shell electron pair repulsion

electron domains

• electron pair

• double/triple bonds are 1 ED

EDG vs MG

• electron domain geometry = positions of all EDs

• molecular geometry = positions of bonded EDs

2 EDs

• 180˚

• linear shape

3 EDs

• EDG

  • 120˚

  • triangular planar

• MG

  • 2 bonded, 1 lone = 117˚, bent/v-shaped

4 EDs

• EDG

  • 109.5˚

  • tetrahedral

• MG

  • 3 bonded, 1 lone = 107˚, trigonal pyramidal

  • 2 bonded, 2 lone = 104.5˚, bent/v-shaped

electronegativity

measure of the ability of an atom to attract electrons in a covalent bond

the … the electronegativity, the … the pull on electrons

higher … higher = polar

bond dipole

• bond with 2 partially separated electric charges

• more electronegative = delta -

• less electronegative = delta +

the … the electronegativity, the … the bond polarity

higher … higher

net dipole

dipoles do not cancel each other out

carbon allotropes

• diamond; C bonded to 4 other Cs

• graphite; C bonded to 3 other Cs, layers held by weak london forces

• graphene; single layer of graphite

• C₆₀; fullerene, sphere of 60 Cs

SiO₂

silica or SAND!!!!!!!!!

london dispersion forces

• opposite ends of temporary dipoles

• weakest forces of attraction

• increased strength with increased molecular size

• only forces between non-polar

dipole-dipole

• opposite ends of permanent dipoles

• stronger than london

dipole-induced dipole

permanent dipole creates a temporary dipole

hydrogen bonding

• strongest type of bonding

• only when H bonds to F, N, or O

• large electronegativity difference

R_f

distance moved by component / distance moved by solvent

metallic character

loss of control over outer shell electrons

metallic bond

electrostatic attraction between a lattice of cations and delocalized electrons

metal properties

• good electrical conductivity

• good thermal conductivity

• malleable

• ductile

• high MP

• shiny, lustrous appearance

alloys

• homogenous mixtures containing at least one metal, and held together by metallic bonding

• presence of atoms of different sizes disrupts regular structure + prevents atoms from slipping across each other

atom economy

molar mass of desired product/molar mass of all reactants x 100%

effective nuclear charge

attraction of nucleus to outer electrons

the … the effective nuclear charge, the … the attraction between outer electrons + nucleus

higher … higher

ionic radii … down a group as the number of electron energy levels …

increase … increases

first ionization energy

• measure of attraction between nucleus and outer electrons

• energy required to remove 1 mole of gaseous atoms in the ground state

electron affinity

energy change when 1 mole of electrons is added to 1 mole of gaseous atoms)

chemical properties are determined by …

the number of valence electrons in their outer energy level

different types of oxides

• oxides of metals = ionic + basic

• oxides of non-metals = covalent + acidic (ACID RAIN!)

• amphoteric oxides = both acidic + basic

SO₂ acid rain

NO acid rain

ocean acidification

oxidation state

the charge the atom would have if the compound was composed of ions

functional groups

atoms/groups of atoms present in organic compounds responsible for a compounds physical properties and chemical reactivity

same functional group = same …

class

alkane

-ane

alkene

alkenyl

-ene

alkyne

alkynyl

-yne

alcohol

-OH

hydroxyl

-anol

ether

alkoxy

-oxyalkane

aldehyde

carbonyl (aldehyde)

-anal

ketone

carbonyl (ketone)
-anone

carboxylic acid

carboxyl (acid)

-anoic acid