Chemistry Key Revisionary Items

ionic formula of sulfate

SO4 2-

ionic formula of nitrate

NO3 -

ionic formula of phosphate

PO4 3-

ionic formula of carbonate

CO3 2-

ionic formula of hydroxide

OH -

ionic formula of hydrocarbonate

HCO3 -

ionic formula of ethanoate

CH3COO -

ionic formula of ammonia

NH4 +

oxidation state of fluorine

always -1

oxidation state of oxygen

usually -2 unless F2O or peroxide

oxidation state of chlorine

usually -1 unless with oxygen or fluorine

oxidation state of hydrogen

usually +1 unless with -1 in metal hydrides

oxidation states of group 1 metals

\+1

oxidation states of group 2 metals

\+2

relative mass and charge of protons

1, +1

relative mass and charge of neutrons

1, 0

relative mass and charge of electrons

negligible, -1

how to determine mass number?

total number of protons and neutrons

how to determine atomic number?

number of electrons or protons

define isotope

elements with the same atomic number but different mass numbers, or the same number of protons but different number of neutrons

define radiation

when isotopes are unstable and their nuclei decompose into new elements, releasing radiation

nature of alpha particles

4/2 He 2+

nature of beta particles

0/-1 e/B

nature of gamma rays

0/0 high energy electromagnetic radiation

penetration power of alpha

stopped by paper

penetration power of beta

stopped by 0.5cm aluminium

penetration power of gamma

stopped by more than 2cm lead

describe alpha emission

helium nucleus produced, atomic number decreases by 2, mass decreases by 4

describe B- decay

neutron becomes proton, negative e- is emitted, increase in atomic number, no change in mass, neutrino released

describe B+ decay

proton becomes neutron, positive e- is emitted, decrease in atomic number, no change in mass, neutrino released

describe electron capture

electron from closest energy level falls into nucleus, causing a proton to become a neutron, emitting a neutrino, so the atomic number decreases, no mass change

describe gamma emission

high energy released, energy of the nucleus decreases, no new products are formed

define half life

the time taken for the mass of an element to decrease by half

define orbital

pathway of an electron around the nucleus

define principle quantum number

distance of orbital from nucleus

define spin

when 2 electrons occupy the same subshell, the electrons pair up with the opposite spin and repel each other

define first ionisation energy

the energy required to remove one mole of electrons from one mole of its gaseous atoms

define second ionisation energy

the energy required to remove one mole of electrons from one mole of its unipositive ions

define electromagnetic radiation

energy associated with electric and magnetic fields travelling as a wave

equation for frequency

wavelength / speed of light

equation for energy

plancks constant x frequency

define balmer series of the hydrogen spectra

4 coloured lines that converge towards the purple end of the spectrum, the lines correspond to electronic transitions to n=2, the colour of the lines correspond to its frequency which is determined by the size of the energy gap

define the lyman series of the hydrogen spectra

appears in the UV region of the electromagnetic spectrum and corresponds to the transitions n=1, as the energy increases, the lines get closer together until they converge to a limit, when the electron is lost and the atom is ionised

define relative atomic mass (Ar)

the average mass of one atom of the element relative to 1/12th the mass of one atom of carbon-12

define relative isotopic mass

mass of a specific isotope of an element relative to 1/12th the mass of one atom of carbon-12

define relative molecular mass (Mr)

the sum of all the relative atomic masses of all the atoms present in its formula

define empirical formula

the simplest formula of a compound showing the ratio of all the different atoms

define molecular formula

shows the actual number of atoms in a compound

define a mole

the amount of a substance that contains the same number of particles as there are carbon atoms in 12g of carbon-12

define avogadro’s number

the number of particles in one mole of a substance

define moalr mass

mass of one mole of a substance

define solution

a homogeneous mixture of 2 or more substances in which the proportions are identical throughout

state boyle’s law

at a constant temperature the volume of a fixed mass of gas is inversley proportional to the pressure PV = constant

state charle’s law

at a constant pressure the volume of a fixed mass of gas is directly proportional to the temperature in kelvins V/T = constant

state avogadro’s gas law

equal volumes of gas at the same temperature and pressure contain the same number of molecules V is directly proportional to n

combining boyle, charles and avogadros laws makes

constant = PV/T

equation for density of gas

(m/v) = MrP/RT

positive ions are more covalent if …

large, highly charged, as they have higher density

negative ions are more covalent if….

larger

negative ions are more ionic if …

small, singly charged

compounds are more ionic if …

large differences in electronegativity, or further apart on periodic table

ionic substance + water →

doesnt react, dissolves to form aqueous ions

covalent substances + water →

reacts, doesnt dissolve

define electronegativity

the tendency of an atom to attract a shared electron pair

define poalr bond

has a slightly positive end, and a slightly negative end, the more electronegative atoms pull electron density towards them, casuing an uneven distribution of electron density across the molecule

difference between intermolecular forces and intramolecular forces

inter are weak between molecules, intra are strong between atoms in a molecule

define temporary dipoles

very weak intermolecular forces that result from instantaeneous uneven distributions of electron desnity within the atoms of a molecule

describe the forces in temporary dipoles

attractive forces

define dipole-dipole forces

unsymmetrical molecules have atoms with different electronegativities bonded together, so are polar, and have permanent dipoles, where the more electronegative atom is delta negative and the less electronegative atom is delta positive

describe the forces in dipole-dipoles

attractive force

define hydrogen bonding

form between the hydrogen atom of one molecule and the lone pairs of small, highly electronegative atoms

how do H-bonds form?

the hydrogen atoms are small and become highly delta positive when bonded to more electronegative atoms, this attracts a lone pair of small highly electronegative atom, so the hydrogen is sandwiched between 2 highly electronegative atoms, covalently bonded to one, and h-bonded to the other

VSEPR theory states

electron paris around a central atom repel each other, with the electron pairs arranging themselves to be as far apart as possible, determining the shape of a molecule

2BP shape, bond angle and e.g.

linear, 180 degrees, BeCl

3BP shape, bond angle and e.g.

trigonal planar, 120 degrees, BF3

4BP shape, bond angle, and e.g.

tetrahedral, 109.5 degrees, CH4

5BP shape, bond angle, and e.g.

trigonal bipyradmidal, 90/120 degrees, PCl5

6BP shape, bond angle, and e.g.

octahedral, 90 degrees, SF6

3BP 1LP shape, bond angle and e.g.

pyramidal, 108 degrees, NH3

2BP 2LP shape, bond angle and e.g.

bent, 104 degrees, H2O

structure of NaCl

ionic strucutre, each Na+ is surrounded by 6 Cl-, so coordination number 6:6

strucure of CsCl

ionic structure, each Cs+ is surrounded by 8 Cl-, so coordination number 8:8

describe structure of diamond

covalent structure, each carbon is bonded to 4 other carbons by strong covalent bonds, all 4 valence electrons are bonded so no delocalised electrons

describe structure of graphite

covalent structure, each carbon is bonded to 3 other carbons by strong covalent bonds, the 4th valence electrons are delocalised

describe the structure of iodine

covalent structure, crystalline structure, when 12 molecules arrange themselves in a regular pattern, with strong covalent bonds and weak intermolecular temporary dipole forces

describe the structure of ice

covalent structure, hydrogen bonded, can form many different structures

describe the structure of metals

when metal atoms are close together in solid state, each atom loses control over one or more of its valence electrons, they become delocalised. there is an attractive force between the delocalised electrons and the positive metal ions, acting like glue holding together the ions in a symetrical lattice

ionisation across a period

increases as nuclear charge increases as protons increases

ionisation down a group

decreases, as valence electrons are further and shielded from the nucleus, so there is less force of attraction

trends in MP/BP across a period

increases as the 3 metals increase the number of electrons contributed to delocalised sea of electrons, so larger attractive force, and group 4 is giant molecule so high MP due to strong covalent bonds between all atoms,

followed by a sharp fall, then a gradual decline, as the remaining elements have molecular structures so MP is determiend by size of van der waals

MP/BP down a group

1/2 = decreases as larger , so positive nuclear charge is further from delocalised electrons, so bond is weaker and easier to break

7 = increase as atoms are larger, so there is a greater surface area, so larger van der waals

electronegativity across a period

increases as number of protons increases so bonding electrons are more strongly attracted

electronegativity down a group

decreases, as bonding electrons are further from nucleus and there is increased shielding

flame test for lithium

red

flame test for sodium

golden yellow

flame test for potassium

lilac

flame test for calcium

brick red

flame test for strontium

crimson

flame test for barium

apple green

flame test for magnesium

colourless