Foundations in chemistry

atoms

made up of protons, neutrons and electrons

atomic mass unit

1/12th of the mass of a single carbon-12 atom

proton

+1 charge, mass - 1

neutrons

0 charge, mass - 1

electrons

-1 charge, mass - 0

nucleus

contains protons and neutrons, positively charged, contains most of the mass of an atom

electron shells

further split into sub-shells, each has different energies, occupying most of the space of the atom

mass number

number of protons + number of neutrons

proton number

the number of protons in its nucleus

isotopes

atoms with the same proton number but a different mass number, ergo having different number of neutrons

they have different physical characteristics to each other

relative atomic/molecular mass (A/M_r)

the average mass of an element/molecule compared to 1/12th of the mass of a single carbon-12 atom

mass spectrometry

• when a sample has passed through the mass spectrometer, a spectrum is produced by the spectrometer

• the spectrum produces lots of peaks, the peak of the mass/charge ratio is the mass/charge value of the molecule

• smaller peaks will cluster around the molecular ion peak, which are the same molecules with different isotopes

• any smaller and lighter peaks in the spectrum are because of fragmentations as it can fragment in the spectrometer

• the relative atomic mass can be calculated as the isotopic masses and relative abundances are identified

groups

the charge on an elemental ion, elements tend to lose or gain elements in order to achieve a full shell of electrons

metallic groups

left-hand side of the periodic table, tends to lose electrons in order to achieve a full shell of electrons and forming positive ions

non-metallic groups

right-hand side of the periodic table, tends to gain electrons in order to achieve a full shell of electrons and forming negative ions

specific ionic compounds

nitrate ion - NO₃^-

hydroxide ion - OH^-

carbonate ion - CO₃^2-

sulfate ion - SO₄^2-

ammonium ion - NH₄⁺

Titration

an experimental technique for finding the concentrations of solutions

mole

6.02 × 10²³ particles, which is the Avogadro’s constant

moles = mass / M_r

concentration

moles per unit volume, moldm^-3

conc = moles / vol

empirical formula

simplest whole number ratio of the atoms in a molecule

molecular formula

the actual number of atoms of each element in a molecule

hydrated salts

produced when compounds precipitate out of solution, to remove the water from the salt, techniques such as heating, placing in a vacuum chamber or drying would be needed

Avogadro’s law

at the same temp and pressure, one mole of two different gases will be the same volume

at RTP, 24 dm³ is the volume taken up by 1 mole of any gas

volume of a gas

moles of the gas x 24, only true if volume is in dm³ and gas is at RTP

Ideal gas equation

pV = nRT

p - pressure/pa

V - volume/m³

n - moles

R - gas constant (8.314)

T - temp/K

percentage yield is usually not 100%

• an incomplete reaction can happen as it can be very slow or reversible, with reactants turning back into products

• some of the chemicals will remain stuck to the glassware and will be wasted

• it might be difficult to fully separate the product from the reaction mixture

• side reactions can happen, meaning the intended product is not created

• if the reaction is carried out in water, extracted product may be ‘wet’, maybe leading to a percentage yield over 100%

real yield

the mass of a desired product obtained from a reaction

theoretical yield

the maximum mass of a product that could possibly be created from a reaction

atom economy

M_r of the desired products / total M_r of all reactants

Acids

substances that form hydrogen ions when they dissolve in water/aqueous solutions

they’re also proton donors

bases

substances that form OH^- ions when they dissolve in water/aqueous solutions

they can neutralise acids to form salt and water and are proton acceptors

alkali

a base that dissolves in water

pH scale

the measurement tool for the acidity and alkalinity of a substance

water dissociation

H₂O ⇌ H⁺ + OH^-

strong base

fully ionises in aqueous solution

weak base

doesn’t fully ionise in aqueous solution

strong acid

fully ionises in aqueous solutions

weak acid

doesn’t fully ionise in aqueous solutions

neutralisation reactions

acid + base ⇌ salt + water

acid + metal oxide ⇌ salt + water

acid + metal carbonate ⇌ salt + water + carbon dioxide

concordant results

values that within 0.1 cm³ of each other

indicator

changes colour when all of the unknown reactant is used up

OIL RIG

Oxidation is loss of electrons; reduction is gain of electrons

oxidising agents

gains electrons by taking electrons from another compound, so its reduced

reducing agents

loses electrons by giving electrons to another compound, so its oxidised

oxidation state

shows how many electrons an atom has gained or lost

oxidation state rules

• fluorine is always -1

• oxygen is always -2, unless its a compound of oxygen and flourine, which the fluorine rule takes priority, and in a peroxide

• hydrogen is always +1, except in metal hydrides, where it is -1

• a pure element is always 0

electron shells

electrons are arranged in shells, as per the Bohr model

these shells are defined by principal quantum number, ‘n’, the higher the n of an electron, the further from the nucleus it orbits

sub-shells

shells are split into sub-shells that have slightly different energies, a shell with a given n will have n sub-shells

orbitals

a region of space where there is a 95% chance an electron is located

sub-shells are composed of orbitals that have the same energy

Types of orbitals

s sub-shells have 1 s orbital

p sub-shells have 3 p orbitals

d sub-shells have 5 d orbitals

f sub-shells have 7 f orbitals

orbitals of exactly the same energy are called degenerate

orbital rules

orbitals are filled lower energy first

each orbital can hold a maximum of 2 electrons

orbitals with the same energy are filled singly before in pairs

ionic bonds

electrostatic attraction between ions of opposite charges

noble gas configurations

ions form to make a noble gas configurations, a full outer shell

giant ionic lattice structure

• made of repeating units of identical structure

• they dissolve in water as it’s polar

• they can conduct electricity when molten/dissolved, as ions can move and carry charge when molten

• they have high melting points as ionic bonds are very strong

covalent bonds

shared pair of electrons between atoms

dative bonds (coordinate bonds)

covalent bond in which both electrons in the bond come from one atom

repulsions

electrons will try to stay as far apart as possible as they repel each other; determining the geometry of a molecule

repulsion strengths

lone pairs are held closer to the nucleus of an atom, repelling each other more as they’re physically closer

lone pair - lone pair

lone pair - bonding pair

bonding pair - bonding pair

geometry of a molecule

linear - 2 electron pairs, 180 degrees

trigonal planar - 3 electron pairs, 120 degrees

tetrahedral - 4 electron pairs, 109.5 degrees

trigonal pyramidal - 3 electron pairs + 1 lone pair, 107 degrees

octahedral - 6 electron pairs, 90 degrees

(lone pair causes electrons to repel stronger, hence decreasing bond angles by 2.5 degrees)

electronegativity

The ability of an atom to attract the electron pair is called electronegativity

in a bond between 2 unlike atoms, 1 atom will have a stronger attraction electron pair than the other, meaning the electron pair won’t be at the centre of the bond

This causes a buildup of a partial charge on one atom,δ

Pauling scale

a scale to measure the values of electronegativity

permanent dipole

the partial charge difference between two atoms due to values of electronegativity

polar molecules

for molecules to be polar, it needs to have polar bonds and if there is a charge separation between one side of the molecule and the other, there will be a permanent dipole

intermolecular forces

• induced dipole-dipole interactions - weakest

• permanent dipole-dipole interactions

• hydrogen bonds - strongest (weaker than covalent bonds)

induced dipole-dipole interactions (London forces)

formed from temporary dipoles, exerting forces on neraby molecules, pushing away/attracting towards the electrons, creating a induced dipole interaction

strength depends on the number of electrons in a molecule as it’ll have larger fluctuations in electron density

permanent dipole-dipole interactions

exists between two permanently polar molecules, the partial charge of one molecule will attract the partial charge of another molecule

hydrogen bonds

hydrogen can form a strong dipole-dipole interaction to either oxygen, nitrogen or fluorine when bonded to those atoms

this develops a strong partial charge and a high charge density to form a strong bond with any strong partial charged atom