periodicity, group 2 and 7 Flashcards
random periodic table info
Horizontal- period ( no of shells)
vertical- group (no of es in outer shell)
group 1- alkali metals
2- alkaline earth metals
middle- transition metals
7- halogens
8-noble gases
1+2 = S block
transition- d block
3+4+5+6+7+8= p block except He
all others= F block
Arrangement of elements
arranged with increasing proton number
elements are in periods showing gradual trend in physical+ chemical properties across period
trends repeat across each period- periodicity
elements in groups have similar chem and physical properties- same no of electrons in outer shell
metals to the left and non metals to the right with a zig zag starting above aluminium to seperate them
classifying the elements
metal to the left
these are good conductors
non metals to the right
these are non conductors
metalliods/semi metals have properties of both- poor conductors
elements in the same group also have the same no of outer shell electrons so react in a similar way as this is what is involved in chem reactions
u can also see which orbitals contain there highest energy outer shell electron by seeing which block they reside in
1+2 are s block
transition- d block
3+4+5+6+7+8- are P block except He
rest are F block
each period and group helps as well so if it is in period 3 then it is in 3 whatever so 3s or 3p. Then the groups show the no of electrons in that outer shell so group 5 period 2 would have have 5 electron in 2p
trends in atomic radii (periods)
across period)
protons added to nucleus so nuclear charge increases
electrons added to same shell
nuclear attraction on outer shell electrons increases across period
electron shells are drawn inwards by nucleus making atoms across period smaller
so trend- atomic radii decrease left to right across period
trend in atomic radii (down group)
no of shells increase
outer electrons added to new shell which is further from nucleus
shielding effect by inner shell electrons increase down group
increase in distance between outer e and nucleus + shielding effect outweigh increase in nuclear charge
nucleus attraction decreases down group
trend- atomic radius increases down group
for questions about trends in atomic radius
Now- nuclear charge
Do- distance/ atomic radii
something- shielding effect
amazing- nuclear attraction
trends in melting and Boiling points
depends on structure of substance and strength of bonds
A substance has high Mp or BP if the bonds broken are strong. The structure is giant
low Mp and BP are opposite weak bonds .simple structure
trend between period 2 and 3 is high boiling point then low then high then low
there is a sharp decrease between group 4( high) and 5 (low) this is because structure changes here from giant to simple
happens in both 2 and 3
group 1+2 elements have giant metallic structure containg strong bonds that take lots of energy to break
group 4- giant covalent (strongest bonds so highest BP e.g silicon)
5-8 - simple Covalent/ simple molecular weaker bond and less energy to break
if they are both the same structure one may have higher BP due to stronger London forces because it has more electrons all together in the molecule e.g. P4 has more than Cl2 so has higher BP
ionisation energy
process of ionisation produces positive ions by removing electrom from outer shell, energy is required for this to overcome electrostatic attraction between nucleus + and electron. This is ionisation energy
First ionisation energy
First ionisation energy is the energy requires to remove 1 electron from each atom in one mole of GASEOUS atoms to form one mole of GASEOUS 1+ ions
state symbols for equation will always be GASEOUS
e.g.
O (g) ——–> O+ (g) + e-
factors effecting ionisation energy
outer shell electrons are always removed first as they experience the least nuclear attraction
thus nuclear attraction depends on
NUCLEAR CHARGE
DISTANCE BETWEEN O E- AND NUCLEUS
ELECTRON SHEILDING
Nuclear charge- ionisation energy
the more protons in nucleus the greater nuclear charge
greater nuclear charge stronger nuclear attraction on Outermost electron
higher nuclear charge more energy needed to overcome attraction between nucleus and outer electrons
distance between O e- and nucleus (ionisation energy)
As the distance between them increases the attraction between them decreases
the weaker the nuclear attraction the less energy needed to remove the outer electron
electeon shielding (ionisation energy)
electron shielding is the repulsion between electrons of diff inner shells.
this shielding effect reduces net nuclear attraction between nucleus and o electrons
the more inner shells greater the shielding effect and the weaker the nuclear attraction
so less energy needed to remove electrons
trends in first ionisation energy ( down group) + explanations
First ionisation energy decreases down a group
as there are more shells- so more shielding effect
the atomic radius increases
the increased shielding effect and distance from nucleus far outweigh increase in nuclear charge
therefore nuclear attraction on outer shell electrons decreases
less energy need to remove outer electrons
Trends in first ionisation energy (across period) + explanations
First ionisation energy shows a gener increase across periods
as the outer electrons fill the same shell so no change in shielding
the no of protons increases so nuclear charge increases
the atomic radius decreases
therefore there is a greater nuclear attraction on outer electrons
more energy is needed to remove the outer electron
how to answer a trends in ionisation energy question
Now- nuclear charge
Do- distance/ atomic radii
Something- Shielding
Amazing- nuclear attraction
Everyday- energy needed
Anomaly in first ionisation energy trend
Be and B there should be an increase in ionisation energy here but instead there is a decrease
this is because the 2p subshell in B has higher energy than the 2s subshell in Be
the 2p1 electeon in B needs less energy to be removed giving B a lower first ionisation energy (same for Mg and Al)
also happend with N and O there is a decrease where there should be an increase
This is due to the electeon pairing In P orbital of O nitrogen p orbital contains 1 electron and Os contains 2 electeons paired.
the 2 paired electrons repel eachother meaning it is easier to remove 1 of then so less energy is needed to remove the 2p electeon from O rather than N ( same in S and P)
Group 2
Alkaline earth metals
called this as they all form hydroxide that are alkaline
Electron configuration- They are S² block elements as there highest energy orbital occupies an S orbital
group 2- physical properties
giant metallic structure- high mp and BP as lots of energy required to break strong metallic bonds and electrostatic forces of attraction between ions and delocalised electrons
good electrical conductors as they have mobile ions and delocalised electrons
group 2- reactions/ oxidation states
2 electrons in outer shell
so they lose 2 electrons to form a di-positive ion
reactivity increases down group
loses electrons so is oxidised
react with non metals to form ionic bonds e.g. CaCl2
group 2 trends
Atomic radii-
increases down group
more shells
more shielding effect
increased distance and shielding outweigh increase in nuclear charge
so nuclear attraction on outer electrons decreases down group causing atomic radius to increase
ionisation energy-decreases down
atomic radii increases
more shells
more shielding effect by inner e-
increase distance and shielding outweigh increase in nuclear charge
so nuclear attraction decreases on outer e-
less energy nessecary to remove outer e- so 1st ionisation energy decreases- second ionisation energy similar reason
reactivity- increases down group
same as ionisation
but because ionisation energy decreases less energy us nessecary to remove 2 electrons from outer shell and so the reactivity increases down group
group 2- reaction with oxygen
group 2 metals react vigorously with oxygen to produce ionically bonded oxide
they can be burned in air to produce oxide
metal + oxygen——-> metal oxide
giant metallic+simple covalent—>giant ionic
when they burn the colour of the flame is linked to the metal
group 2- reaction with water
react with room temp water to form metal hydroxide
these are weak bases with ph 10-12 usually
reduce the hydrogen in water to form H gas
Metal + water —–>metal hydroxide + H2
solid dissolves and effervescence
Mg reacts slowly with water and as u move down group they react more vigorously as reactivity increases down group
Mg does react vigorously with steam
Mg + steam —-> MgO + H2
reactions of g2 elements and compounds
group 2 oxides and hydroxides are bases (proton acceptors)
metal + acid ——> salt + H2
redox reaction
H+ is oxidisng agent
metal is reducing agent
group 2 oxide+ water—> alkaline solution of the metal hydroxide
MgO + H20 —> Mg(OH)2 Aq (not redox)
Group 2 hydroxide + water —>dissolve to form alkaline solution
Ca(OH)2 + Aq—–> CA 2+(Aq) + 2OH- (Aq)
don’t conduct conducts
solubility and alkalinity of metal hydroxide in water increases down group
solubility of metal sulphates in water decreases down the group barium sulphate is insoluble and uses as dye in hospitals
Commercial uses of G2 hydroxides and sulfates
Mg(OH)2- uses in some indigestion tablets as antacid as it neutralizes excess stomach acid safely( mild alkali)
Ca(OH)2- Used in agriculture to neutralise acidic soil does this by reacting with acid substances in soil
BaSO4- used in medicine for visualizing X rays, it is insoluble in solution so doesn’t release toxic Barium ions into patients bloodstream
thermal decomposition of group 2 carbonates
thermal decomposition is the breaking down of chem substance by heat into atleast 2 chem substances
G2 carbonates are decomposed by heat to form solid metal oxide and CO2
CaCO3 —-> CaO + CO2
Ease of thermal decomposition decreases down the group, thermal stability increases down group
BeCo3 is so unstable that it doesn’t exist at room temp
Group 7 general info
electron config-
P block
p5
7 electrons In outer shell
bonding and structure-
covalent and exist as diatomic molecules
simple covalent
physical properties- group 7
low Mp and Bps as little heat energy is needed to break weak london forces between molecules
increase down group
more shells so more electrons
more/stronger London forces between molecules
more energy needed to break stronger london forces between molecules
how do halogens react
powerful Oxidising agents
remove electrons from other species to get reduced
Ox power is a measure of strength with which on atom can attract and capture an electron to form a halide ion
1/2 X + e- ——> x-
half equation to become anions- reduction
ox number decreases
Trend in reactivity- group 7
reactivity decreases down the group
F is strongest ox agent- reacts with almost everything
ox power decreases down group
more shells so increased atomic radius
greater shielding effect
weaker nuclear attraction on outer e-
nucleus is less able to attract and capture another electron to outer shell
displacement reactions- group 7
displacement reactions are redox
a displacement reaction is a reaction in which a more reactive element displaces a less reactive element from an AQ solution of its halide ions
when writing half equation remove spectator ions also always mention colour change
displaced halogen has charactistic colour
to test which halogen is present- it is added to a AQ solution of diff halides and shake mixture
observations- colour change
water organic solvent
Cl2 pale green pale green
Br2 orange orange
I2 brown purple
example- aqueous(water)
redox-Br2 + 2NaI —–> 2NaBr + I2
ionic- Br2 + 2I- —-> 2Br- + I2
orange—-> brown as I2 is displaced
Equations u have to know- group 7
Cl2 (aq) + H20 (aq) ——> HClO (aq) +HCl (aq) chloric acid
used to kill bacteria in water treatment as small amounts of chlorine kills bacteria
but chlorine is also toxic and forms carcinogenic chlorated hydrocarbons
Famous disproportion equation
Cl2 + 2NaOH —–> NaCl + NaClO + H2O
sodium chlorate(1)
all AQ except H20
only happens in cold dilute aqueous NaOH
famous example of disproportion
makes household bleach
Sodium chlorate (V) or NaClO3
is formed when hot concentrated sodium hydroxide is reached with chlorine instead
Halide test method is also relevant to group 7 but is in the quantitative analysis deck
successive ionisation energy
successive ionisation energies are a measure of the energy required to remove each electron in turn
second is the energy requires to remove 1 electron from each 1+ ion in one mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions
e.g of the equations
2nd- Na+ (g)—–> Na 2+ (g) + e-
evidence for shells+ reason ionisation energy increases
graphs show successive ionisation energy providing evidence of shells
graphs show:
largest increase is between for example 5 and 6 ionisation energies
so 6th electron must be removed from new shell which is closer to nucleus with less shielding
so element must have 5 electrons in outer shell
so must be in group 5
also shows period based on no of shells shown
reason ionisation energy increase with ionisation number:
once electron has been removed there are same no of protons but less electrons
proton: electron ration increases
remaining electrons are more strongly attracted to nucleus
more energy needed to remove each electeon in turn
