1. elements of life

electrons

negatively charged particles which move around the nucleus in shells

nucleus

central part of the atom with most the mass made up of protons and neutrons

mass number

larger number which is the total of protons and neutrons

atomic number

smaller number which is the number of protons or electrons

negative ions

have more electrons than protons

postive ions

have fewer electrons than protons

isotope

elements with the same atomic number but different mass number due to different amounts a neutrons, with the same chemical properties but different physical properties (density, diffusion)

ancient greeks

though that all mater was made from indivisible particles

John dalton

at the start of the 19th century it was suggested atoms were solid spheres

JJ thomson

in 1897 experiments concluded that atoms must have smaller negatively charged particles, electrons, he created the plum pudding model

Ernest Rutherford

Geiger-Marsden experiment where alpha particles were fired at gold sheet and where deflected and some went through, nuclear model with a positive centre and cloud of electrons

Henry Moseley

charge of nucleus increased by one form element to element, lead to Rutherford discovering the proton

James Chadwick

conclude there are neutrons with a mass but no charge

Bohr model (Niels Bohr)

proposed that elections are in fixed orbits with a fixed energy, when electrons move shells they emit or absorb electromagnetic radiation

quantum model

model which uses quantum mechanics to predict where electrons will be

relative mass

the molecular mass is relative to carbon 12

mass spectrometer

1. sample is vaporised
2. ionisation to positive ions by bombardment of high energy electrons
3. the ions are accelerated by an electric shield
4. time for the ions to reach detector is measured with lighter ones been quicker

calculate relative atomic mass

(abundance x mass) ÷ 100

mole

amount of a substance

avogadros constant

6\.02 x10 ^23

number of moles

= number of particles you have ÷ avogadros constant

number of moles

= mass ÷ Mr

concentration

= moles ÷ volume

empirical formula

smallest whole number ratio of a compound, worked out from moles and then dividing by smallest

water of crystallisation

compounds with incorporated water

hydrated

compounds containing water of crystallization

anhydrous

compounds without incorporated water

calculate water of crystallization

find moles of water lost from mass

find moles of compound from mass

divide by smallest

percentage yield

actual yield ÷ theoretical yield x 100

standard solution

a solution with a known concentration

6, 3, 4, 9, 8, 2, 7, 1, 5

order statements making standard solution


1. place stopper in bottle and invert
2. rinse the beaker and stirring rod with distilled water and add it to flask
3. work out how many grams of solute needed using mass = moles x mr
4. measure out the mass of the solute first measuring beaker
5. check meniscus again and add water if needed
6. work out moles of solute using using moles = conc x volume ÷ 1000
7. top up flask with distilled water till at meniscus using pipette
8. top up solution into volumetric flask using funnel
9. add distilled water to beaker to dissolve solute

diluting solution

making a standard solution from more conc solution

volume to use

final conc ÷ initial conc x volume required

1, 4, 7, 2, 3, 5, 6,

1. add acid to burette to the 0cm3 line
2. carry out rough titration to find the endpoint
3. work out amount of acid used
4. add alkali to flask using 25cm3 pipette
5. carry out accurate titration till you get concordant results
6. calculate mean
7. add indicator to flask

methyl orange

turns yellow to red/peach when adding acid to alkali

phenolphthalein

turns red to colourless when adding acid to alkali

s

sub shell with 2 electrons

p

sub shell with 8 electrons

d

sub shell with 10 electrons

f

sub shell with 14 electrons

spin pairing

electrons spin in opposite directions

s orbital

spherical orbitals

p orbital

dumbbell shaped orbital, made up of 3

s block

group 1 and 2

d block

transition metals

p block

group 3-8

ionic bonding

electrons are transferred making 2 ions and then hold each other together by electrostatic attractions, non metal and metal

dot and cross

diagram showing electron distribution in shells, use square brackets and charge for ionic compounds, shells overlap for covalent

giant ionic lattice

compounds which form crystals with regular repeating structure

ionic properties

conduct when molten or dissolved

high melting points

soluble in water

covalent bonds

share electrons been 2 non metals, help by electrostatic attraction between nuclei and electrons as well as repulsion between nuclei

dative covalent bonding

covalent bond where both bonding electrons are from one element

giant covalent structures

structures like carbon and silicon dioxide

properties of giant covalent structures

very high melting point

very hard

good thermal conductors

don’t dissolve

can’t conduct

metallic bonding

bonding between metals where they lose electrons making them delocalised leaving a positive ion, attraction between ion and electrons hold them together

metallic bonding properties

high melting points due to sea of delocalised electrons

ions able to slide over each other

good thermal and electrical conductors due to sea of delocalised electrons

insoluble due to high bond strength

shape

depends on the number pf pairs of electrons in outer shell

electron effect

pairs of electrons are negative and will repel each other to get as fair apart from each other as possible

lone pairs

pairs of electrons that repel more than bonding pairs

wedges

shows bonds pointing out page

dashes

shows bonds pointed into page

linear molecules

2 bonding pairs bond angle of 180 degrees

trigonal planar

3 bonding pairs, bond angle of 120˚

non linear

2 bonding pairs with double bonds cancel out effect of 1 lone pair, bond angle of 120˚

tetrahedral

4 bonding pairs, bond angle of 109.5˚

trigonal pyramidal

3 bonding pairs, 1 lone pair, bond angle of 106˚

bent

2 bonding pairs, 2 lone pairs, bonding angle of 104.5˚

trigonal bipyramidal

5 bonding pairs, bonding angle of 120˚ and 90˚

octahedral

6 bonding pairs, bonding angle of 90˚d

determine shapes

draw a dot and cross diagram, allows you to see numbers of bonding and lone pairs

calculate shape

1. find group number of central atom (only one of)
2. add number of atoms around atom to its group number
3. divide by 2 to get electron pairs

if there is the same number of electron pairs as surrounding atoms then no lone pairs.

if ion then subtract or add its charge from its group

group

column where all elements have the same number of electrons in outer shell

period

row where all elements have the same number of shells

melting points

a trend which increases across a period peaking at group 4, related to number of delocalised electrons, giant covalent structures and intermolecular forces

first ionisation enthalpy

energy needed to remove 1 electron from each atom in one mole of gaseous atoms to become 1 mole of gaseous 1+ ions

lower enthalpy

easier it is to remove on outer electron and form an ion

factors effecting first ionisation enthalpies

atomic radius- dist electrons from nucleus

nuclear charge- more positive nucleus attracts electrons more

electron shielding- inner electrons shield outer electrons.

group 1 and 2 first ionisation enthalpy

trend which decreases down a group due to increasing shells causing shielding

across a period

direction where first ionisation enthalpy increases due to increasing protons

s block metals

block with low first ionisation enthalpies due to few protons a limited outer shells

hydroxide

product of water and group 2 metal, also produces H2

oxides

product of when oxygen combines with a group 2 metal

strong alkaline solutions

when group 2 oxides react with water forming metal hydroxides which dissolve, expect magnesium oxide, increases down group

neutralise acids

both are bases so react with acids to form salts and water

group 2 hydroxide

solubility increases down group (single negative charge)

group 2 oxide

solubility decreases down group (double negative charge)

thermal decomposition

when heat is added to group 2 carbonates they form the oxide and CO2

thermal stability increases

change in thermal stability down group due to smaller cations (group 2) distorting the large carbonate more due to charge density

salts

neutral ionic compounds with a cations and anions

acids

substances with a pH less than 7

bases

substances with a pH more than 7

HN4+

ammonium

NO3-

nitrate

HCO3-

hydrogencarbonate

(SO4)2-

sulfate

(CO3)2-

carbonate

cation first

order when naming salts

most sulfates

soluble except barium, calcium, lead which form white ppt

lithium, sodium, potassium, ammonium and nitrates

all soluble salts

most chlorides, bromides, iodides

soluble salts except silver halides, copper iodide (white ppt), lead chloride and bromide (white ppt) and lead iodide (yellow ppt)