periodicity, group 2 and 7 Flashcards by Rosie Pritchard

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

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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

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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

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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

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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

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for questions about trends in atomic radius

Now- nuclear charge
Do- distance/ atomic radii
something- shielding effect
amazing- nuclear attraction

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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

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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

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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-

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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

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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

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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

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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

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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

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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

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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