elements of life

the atom

atoms are made up of protons, neutrons and electrons

electrons

have a -1 charge

they are found in shells

shells make up the majority of the atom

nucleus

most of the mass of the atom is concentrated at the nucleus

the diameter of the nucleus is very small compared to the rest of the atom

the nucleus is where protons and neutrons are found

mass= 1/2000

neutron

mass=1

has a charge of 0

proton

mass=1

has a charge of +1

nuclear symbols show numbers of subatomic particles

mass number (top number)= the total of protons and neutrons in the nucleus

atomic number (bottom number)= number of protons in the nucleus - all atoms of the same element have the same number of protons

ions have different numbers of protons and electrons

negative ions

negative charge means theres more electrons than protons

the number next to the - indicates number of electrons added

positive ions

positive charge means theres fewer electrons than protons

the number next to the + indicates number of electrons lost

isotopes are atoms of the same element with different numbers of neutrons

isotopes= atoms of an elements with the same atomic number but different mass numbers

caused by the numbers of neutrons in an atom

can cause a variation in physical properties which do depend on the mass of the atom (density, boiling points)

atomic models

the accepted model of the atom has changed throughout hustory

some ancient Greeks thought matter was made from invisible particles

at the beginning of the 9th century Dalton described atoms as solid spheres and said that different spheres make up different elements

new experiments constantly changed ideas

experimental evidence showed that atoms weren't solid spheres

in 1897 JJ THOMPSON did a series of experiments which disproved Dalton's theory

his measurements of CHARGE and MASS showed that an atom must contain smaller negatively charged particles- electrons

created PLUM PUDDING MODEL- negative electrons embedded into a sphere of positive charge

Rutherford proved plum pudding wrong

In 1909 Rutherford with his students Geiger and Marsden conducted the alpha scattering (gold leaf) experiment

they fired alpha particles at a thin sheet of gold

they expected (due to plum pudding) for the alpha particles to be very slightly deflected by the positive sphere

in fact most of the alpha particles passed straight through and a small number were deflected backwards

so Rutherford created the nuclear model

nuclear model

tiny positively charged nucleus at the centre of atom- where most of mass is concentrated

nucleus is surrounded by a cloud of freely orbiting electrons

most of atom is empty space

the nuclear model was modified

Mosely discovered the charge of the nucleus increased from one element to another by one

Rutherford also discovered that the nucleus contained positively charged protons

Chadwick discovered the neutron

the Bohr model

electrons existed in fixed orbitals (shells)

each shell had fixed energy

when an electron moves between shells electromagnetic radiation is absorbed or emitted

because the energy of the shells are fixed, the radiation will have a fixed frequency

said that a shell can only hold a certain number of electrons and the electrons determined reactivity

relative mass

relative masses are masses of atoms compared to carbon 12

relative atomic mass= the average mass of an atom of an element (on a scale where an atom of carbon 12 is 12

relative isotopic mass= the mass of an atom of an isotope of an element (on a scale where carbon 12 is 12)

relative molecular mass= the average mass of a molecule or formula unit (on a scale where carbon 12 is 12)

relative masses can be measured by a MASS SPECTROMETER

1) VAPORISATION= the sample is turned into gas using an electrical heater

2) IONISATION= the gas particles collide with high energy electrons to ionise them. electrons are knocked off the particles leaving them as positive ions

3) ACCELERATION= the positive ions are accelerated by an electric field

4)DETECTION= time taken for positive ions to reach the detector is measured. (depends on charge and mass o ion)- light, highly charged ions will reach detector first

mass spectrum

the y axis= abundance of ions or the relative isotopic abundance for elements

the x axis= relative mass

the mole and equations

a mole is a very large number of particles

amount of substance is measured in moles

1 mole= 6.02x 1023 (avogadros constant

number of moles= number of particles you have / number of particles in a mole (avogadros constant)

molar mass is the mass of one mole

molar mass= relative molecular/formula mass

number of moles= mass of substance / molar mass

in a solution concentration is measured in mol/dm-3

number of moles= concentration x volume(cm3) / 1000

ionic equations

only reacting particles are included

empirical and molecular formulas

empirical and molecular formulas are ratios

e formula gives the smallest whole number ratio of a compound

m formula gives the actual numbers of atoms in a molecule

percentage yield in never 100%

percentage yield = actual yield/ theoretical yield x100

titrations

a standard solution has a known concentration

1) work out how many moles of solute needed by using moles= concentration x volume/ 1000

2) work out how many grams of solute is needed using mass= moles x molar mass

3) weigh the empty beaker and then the beaker with the correct mass of solute

4) add a small amount of distilled water and stir until dissolved

5) tip solution into a volumetric flask

6) rinse the beaker with distilled water then add to flask

7) top up flask to 250cm3 with distilled water

8) stopper the bottle and turn upside down 5 times

making a standard solution from a more concentrated solution

vol to use= final conc / initial conc x vol required

use this equation to get desired concentration

then transfer this into volumetric flask

titrations must be accurate

allow you to find out exactly how much acid is needed to neutralize and alkali

you measure out some alkali using a pipette and put in a flask with indicator (phenolphthalein)

first do a rough titration to get an idea where point of neutralization is

take a reading of initial volume of acid and slowly add acid to the alkali swirling solution and stop when there is a colour change

work out amount of acid needed to neutralize the alkali (the titre)

repeat and work out mean

indicators

methyl orange- turns yellow to red

phenolphthalein - turns red to colourless

titration calculations

number of moles = concentration x volume / 1000

number of moles= mass / molar mass(mr)

electronic structure

electron shells are made up of sub shells and orbitals

electrons move around nucleus in shells which are given numbers known as PRINCIPAL QUANTUM NUMBERS

Shells further away from nucleus have a greater energy level than shells closer to the nucleus

s sub shell

1 orbital

maximum number of electrons= 2

p sub shell

3 orbitals

maximum number of electrons = 6

d sub shell

5 orbitals

maximum number of electrons= 10

f sub shell

7 orbitals

maximum number of electrons = 14

orbitals have shapes

electrons in each orbital have to spin in opposite electrons (spin pairing)

s orbitals = spherical

p orbitals = dumbbell shapes

l1423d

using the periodic table to workout electron configurations

s block= groups 1 & 2

p block= boron to neon and below

d block= transition metals

for an ion

write electronic structure of atom

add or remove electrons from the highest occupied sub shell

4S SUB SHELL FILLS BEFORE 3D SUB SHELLS BUT EMPTIES BEFORE IT

ionic bonding

ionic bonding= when ions are joined by electrostatic attractions

ions are formed when electrons are transferred from one atom to another

generally the charge on the ion is equal to its group number

electrostatic attractions= forces that hold positive and negative ions together (very strong)

giant ionic lattices

a lattice is just a regular structure

the structure is giant because it is the same basic unit repeated over and over again

in the lattice ions with different charges attract eachother and the ions with the same charge repel

ions arrange themselves in order to maximize attractions and minimise repulsions

ionic structure influences properties

ionic compounds conduct electricity when molten or dissolved (not solid)

the ions in a liquid/ solution are free to move and carry charge

ionic compounds have high melting points

giant ionic lattices are held together by strong electrostatic forces

lots of energy is needed to overcome these forces

ionic compounds are often soluble

water molecules are POLAR

water molecules pull ions away from lattice and cause it to dissolve

covalent bonding

covalent bonds hold atoms in molecules together

two atoms share electrons o they both have full outer shells

both positive nuclei are attracted electrostatically to the shared electrons

repulsion between positive nuclei - maintains a balance in forces

molecular substances have some typical properties

low melting points

intermolecular forces between molecules are weak

don't conduct electricty (no electrons free to move)

insoluble in water

the polar water molecules are more attracted to eachother than the molecular substance

dative covalent bonding= where both electrons come from one atom

giant covalent and metallic structures

some covalently bonded substances have giant structures

have a network of covalently bonded atoms- electrostatic attractions holding the atoms together are stronger than in simple covalent molecules

e.g. diamond, graphite, silicon dioxide

properties of giant covalent structures

very high melting points

needs a lot of energy to break bonds before substance melts

very hard

strong bonds throughout lattice

good thermal conductors

vibrations travel easily through stiff lattices

insoluble

don't contain ions

cant conduct electricty

no delocalised electrons to carry charge throughout the structure (all are bonded)

metals have giant structures

metal elements exist as giant metallic lattice structures

the electrons in the outer shell are delocalised (leaves a positive metal ion)

positive metal ions are attracted to the delocalised negative electrons- form a lattice of packed positive ions in a sea of electrons

properties

high melting points

strong metallic bonding (the more electrons there are the stronger the bonding)

the size of the metal ion and the structure of the lattice effect it

ductile

no bonds holding specific ions together

the metal ions slide over eachother when the structure is pulled

good thermal conductors

delocalised electrons can pass kinetic energy to eachother

good electrical conductors

delocalised electrons are free to move and carry charge

impurities can reduce this (electrons transfer to impurities are form anions)

insoluble

strength of metallic bonds

shapes of molecules

molecular shape depends on electron pairs around central atom

electron pairs repel eachother

electrons are negatively charged so repel as much as possible

lone pairs repel more than bonding pairs

the greatest angles are between lone pairs of electrons

bonding pair angles are often reduced because they are pushed together by lone pair repulsion

the shape of the molecule depends on the type of electron pairs as well as the number

electron pair repulsion theory is used to predict shapes

you can use electron pairs to predict shapes of molecules

draw dot and cross diagram first

then work out shape of molecule

types of molecules

2 electron pairs on central atom

LINEAR

treat double bond same as single

3 electron pairs on central atom

TRIGONAL PLANAR

no lone pairs

BENT/ non linear

one lone pair

4 pairs of electrons on central atom

180 degrees angle

117.5 degrees

120 degrees

TETRAHEDRAL

no lone pairs

109.5 degrees

TRIGONAL PYRAMIDAL

one lone pair

107 degrees

BENT/ non linear

2 lone pairs

104.5 degrees

5 electron pairs on central atom

TRIGONAL BIPYRAMIDAL

90 and 120 degrees

6 electron pairs on central atom

OCTAHEDRAL

90 degrees

periodic trends

the periodic table arranges elements by atomic number

all elements within a period have the same number of electron shells

all the elements within a group have the same number of electrons in their outer shells

periodic trends are patterns (periods 2 &3)

metals

melting points increase across period (metallic bonding gets stronger)

bonds get stronger due to increasing number of delocalised electrons and a decreasing ionic radius

creates a higher charge density- attracts ions more strongly

elements w/ giant covalent structures

a lot of energy is needed to break these strong bonds

carbon and silicon have the highest melting points in their periods

simple molecular structures

their melting points depend on intermolecular forces

intermolecular forces are weak and easily broken

have low melting points

noble gases have the lowest melting points as they exist as individual atoms

more atoms in a molecule means stronger intermolecular forces- higher melting point

ionisation is the removal of one or more electrons

first ionisation enthalpy= the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions

X(g) ----> X+ + e-

the lower the ionisation enthalpy the easier it is to remove the outer electron

main things that affect size of ie's

atomic radius

the further the outer shell electrons from the positive nucleus the less they will be attracted

it will be lower

nuclear charge

this the positive charge on the nucleus caused by protons

the more protons the more they will attract electrons

it will be higher

electron shielding

the inner electron shells shield the outer shell electrons from the more attractive force of the nucleus

it will be lower

decreases down the group

increase across a period

group 2

group 2 elements react with water and oxygen

they react with water to produce hydroxides

gives a metal hydroxide + hydrogen

e.g. M + 2H20 -----> M(OH)2 + H2

they burn in oxygen to form oxides

produces white solids

e.g. 2M + O2 ----> 2MO

group 2 oxides and hydroxides are bases

they form alkaline solutions in water

the oxides of group 2 react with water to form metal hydroxides which dissolve

the hydroxide ions (OH-) make the solution strongly alkaline

More alkaline down the group

they neutralize acids

form solutions of corresponding salts

DEACREASE IN SOLUBILITY DOWN THE GROUP

group 2 carbonates decompose to form CO2 and metal oxides

thermal decomposition= when a substance breaks down when heated

e.g. MCO3 -----> MO + CO2

Thermal stability of carbonates changes down group

the more thermally stable, the more heat needed to break it down

carbonate ions are large anions and can be made unstable by the presence of a cation (polarizes it) and distorts it (the greater the distortion the less stable)

large cations cause less distortion than small ones (have a lower charge density)

so the further down the group the larger the cations, the less distortion caused and the more stable the carbonate

making salts

salts are ionic compounds

acid + base ---> salt + water (neutralization)

salts are formed from cations and anions so product is neutral

cations:
-NH4+
-Cu2+
-Zn2+
-Pb2+
-Fe2+
-Fe3+

anions:
-NO3-
-OH-
-HCO3-
-SO42-
-CO32-

Salts can soluble or insoluble in water

soluble= nitrates, most chlorides and bromides, most sulfate, lithium sodium potassium ammonium salts

insoluble= carbonates, hydroxides

tests for ions

flame tests

1) dip nichrome loop into HCl

2) dip wire loop into sample and place in flame

3) observe colour

colours:
-Li+= crimson
-Na+= orange
-K+= lilac
-Ca2+= brick red
-Ba2+= green
-Cu2+=blue/ green

add sodium hydroxide and look for coloured precip

colours:
-Ag+= brown
-Ca2+= white
-Cu2+= blue
-Pb2+= white
-Fe2+=green
-Fe3+= red brown
-Zn2+= begins white then colourless
-Al3+=begins white then colourless

hcl can detect carbonates

test for limewater

test for sulfates

add dilute HCl followed by BARIUM CHLORIDE SOLUTION

if white precip forms it contains sulfate

litmus paper going blue= ammonia or hydroxide is present

test for halides with silver nitrate

Cl= white precip

Br= cream precip

I= yellow precip

nitrates

heat solution with sodium hydroxide and aluminium foil

if present ammonia is produced

used litmus paper to test for ammonia

atomic spectra and nuclear radiation

electromagnetic spectrum

radiation increases in frequency and decreases in wave length from left to right

radio waves, micro waves, infra-red, visible light, ultra violet, x rays, gamma rays

electrons absorb or release energy

atoms in their ground state have electrons in the lowest possible shells

if an atom takes in energy they can move to higher energy levels further from nucleus

electrons can also release energy by dropping from a higher energy level to a lower one

the energy levels all have certain fixed values- electrons can jump by absorbing or releasing a fixed amount of energy

transition between energy levels produces line spectra

absorption spectra

energy is related to frequency

when em radiation is passed through a gaseous element the electrons only absorb certain frequencies corresponding to different energy levels

the missing frequencies show up as dark lines

emission spectra

when electrons drop to lower energy levels they give out certain amounts of energy

each element has a different electron arrangement so the frequencies of radiation absorbed and emitted are different

has coloured lines

atomic spectra and nuclear radiation

both types of spectra are made up of sets of lines

absorption and emission spectra have lines in the same position for an element

lines represent electrons moving from or to a different energy level

the lines on both get closer as frequency increases

energy is related to frequency

E= h x v

speed and frequency related to wavelength

c= v x y

nuclear fusion releases lots of energy

nuclear fusion=when two small nuclei combine under high temp and pressure to make one larger nucleus

in stars hydrogen nuclei combine to make helium nuclei

can only be done in stars