Periodicity, Group 2 And 7 Flashcards
What’s the trend in atomic radius across period 3
The atomic radius decreases across period 3.
This is bc the increasing nuclear charge across the period, drawing outer electrons (same level/diff sub shells) closer to nucleus
What’s the trend in melting point across period 3 (metals and Si)
The graph showing melting point across the period can be split into 4
- The melting points of metallic elements increase
- from Na to Al number of outer electrons increase
- so more electrons can be delocalised (= greater attraction
Between positive ions and delocalised electrons)
-
- sizes of ions decreases from Na+ to Al3+
- again leading to greater attraction
..
2. Melting point of silicon is very high
- silicon has a giant covalent structure; lots of energy required to break
The strong covalent bonds
What is the trend in first ionisation energy in period 3
The first ionisation energy across period 3 generally increases.
—> this increase is caused by increase in nuclear charge
With no increased shielding
The atomic radius hence is decreasing; outer electron
Is closer to nucleus with greater charge (more energy needed to remove)
..
First ionisation energies (G3 and 6)
- atoms of G3 elements show a decrease in ionisation energy
- due to division of energy level into sub shells
—> like aluminium’s outer 3p1 electronics further from nucleus
(less energy to ionise)
- atoms of G6 elements show a decrease in ionisation energy
- as electrons start pairing in the p subshell
- the additional repulsion between the two electrons lowers energy
Needed to remove(decreasing ionisation energy)
..
First ionisation energies G1 and G0
- G1 atoms have lowest ionisation energy in every period
- as have greatest atomic radius and lowest nuclear charge
- G0 atoms have highest
- as have smallest atomic radius and highest nuclear charge
What are the trends down group 2 (atomic radius)
Atomic radius
The outer electrons is pulled in closer to nucleus
When attraction between increases, so radius = smaller
Also, increasing number of electrons = more repel
-
- as G2 descends, outer electrons is less tightly held
- as shielded from increasing nuclear charge
- which new electrons filling more energy levels
- also provide extra repulsion and outer electron
- moves further from nucleus (radius increases)
What are reactions of G2 metals with water
Be doesn’t react with water
Mg reacts very slowly with water; readily with steam
Mg + 2H2O = Mg(OH)2 + H2 (water)
Mg + H2O = MgO + H2O (steam)
» MgO instead of OH with steam as OH is unstable
at higher temperatures,
thermally decomposes to give O and H2O
..
Ca, Sr and Ba react with cold water to produce metal OH and H2
Reactivity increases as you descend group
Eg. Ca + H2O = Ca(OH)2 + H2
Ca reacts w water releasing bubbles of H2 in exothermic
-> white ppt Ca(OH)2 forms and remaining solution’s alkaline
—> since ppt is slightly soluble in water
Reactions with Sr and Ba are similar
As they get more reactive, amt of ppt decreases as more soluble
Down group.
What’s the solubility if G2 sulfates in water
They become less soluble down the group
Mg2+ no ppt
Ca2+ thin white ppt
Sr2+ white ppt
Ba2+ thick white ppt
..
PRACTICAL use of Ba salts to test for sulfate
Add unknown solution
- add dilute HCl (removes other ions like CO32-
- add aq BaCl2 solution
- if white ppt formed, sulfate ions present
Simplest ionic eq: Ba2+ + SO42- = BaSO4
What are selected uses of G2 elements and their compounds (Mg)
Magnesium
- Mg is used in extraction of titanium
» titanium has applications in aerospace, marine and motor
Vehicle industries due to very high corrosion resistance
» very difficult to extract titanium from its ores
- in theory C can be used to extract from oxide but
titanium forms a brittle carbide
-
- but can be extracted using Mg
- TiO2 is converted to the chloride - subsequently
Reduced to Ti by reaction with Mg
TiCl4 + 2Mg = Ti + 2MgCl2
..
Magnesium Hydroxide
- sparingly soluble; sold as a suspension in water
- (in this form, known as milk of magnesia)
- taken to alleviate constipation and (as antacid)
To neutralise excess acid in the gut
- solutions are slightly alkaline due to low solubility
Not irritating the oesophagus
What are selected uses of G2 elements and their compounds (Ca)
Calcium hydroxide
- in solid form is known as slaked lime , neutralising acidic soil
- pH of soil affects its physical , chemical, and biological properties
> affecting crop yield
-
- as soil pH decreases, soluble Al3+ and Mn ion lvls increase
Becoming toxic (restrict root growth, curling and black spot on leaf)
- low pH restrict bacterial growth - wch help leguminous plants fix
N2 from air
—>reducing soil acidity prevents all these effects and increases availability
Of other important nutrients (P4)
—> BUT soil texture improves, allowing plants with small seedlings
An opportunity to break through soil surface
..
Calcium Oxide
- neutralises sulfur dioxide - may be produced as a byproduct
Of fossil fuel combustion (prevents acid rain)
—> CaCo3 can also be used
What are selected uses of G2 elements and their compounds (BaSO4)
Barium sulfate
- can be eaten as part of a ‘barium meal’
- barium is good at absorbing x rays, so when BaSO4
Reaches gut, the outline of gut can be located with X-rays
- although barium ions are toxic, this is harmless as BaSO4
Is insoluble and not absorbed into blood
Physical halogen properties - appearance?
- fluorine
Is a poisonous yellow gas - chlorine
Is a poisonous, dense, yellow green gas - bromine
A caustic and toxic red-brown volatile liquid (red brown vapour) - iodine
Shiny grey-black solid
Sublimes to form violet vapour on gentle heating
-Astatine
No stable isotopes (longest lived isotope has 8.1hr half life)
So hasn’t properly been seen.
—> A mass large enough to be seen wd immediately vaporise
By heat generated by its own radioactivity
It’s dark, with a higher melting point than iodine
Halogen properties - diatomicity and bond enthalpy
Halogens are diatomic molecules
Due to formation of single covalent bond between atoms
—> bonding between atoms is important
As bind enthalpy bust be overcome to allow halogens to react
- bond between halogen atoms is due to attraction of electrons
By the two nuclei - as atomic number increases, atomic radius increases, so
Distance between electrons and nucleus increases
» true for chlorine, bromine and iodine
- - but fluorine rather than having highest bond enthalpy has
A value close to iodines. - as fluorine atoms are so close together that one electron pairs
Round each atom repel each other so bond weakens
What are trends in G7 (colour, atomic radius, reactivity)
All the following trends occur as atomic number increases in G7
1. Colour of elements gets darker
2. Atomic radius increases
- on moving down group, there’s more energy lvls of electrons
- outer energy lvls of electrons are further from nucleus
- increases size of atoms down group
- Reactivity decreases
- outer electrons are further from nucleus as energy lvl added
Every time u descend the group - there’s increase in shielding
- atoms gain electrons less easily as there is a weaker attraction between
Nucleus and incoming electron; smaller halogens = more reactive
- outer electrons are further from nucleus as energy lvl added
What are redox reactions of halogens (Cl2 and Cold dilute aq NaOH)
> > Sodium chlorate and Sodium chloride formed
This reaction may be applied to other halogens.
The chlorate ion (ClO-) is a halate ion (pattern in formulae)
- chlorate (ClO-)
- bromate (BrO-) etc..
- sodium iodate (NaIO)
Chlorine reacts with cold dilute aq NaOH to produce ClO- (chlorate ion)
Ionic eq: 2OH- + Cl2 = Cl- + ClO- + H2O
Symbol eq: 2NaOH + Cl2 = NaCl + NaClO + H2O
..
Observations..
- yellow green gas forms colourless solution
This is a redox as Chlorine is oxidised from
0 in Cl2 to 1 in NaClO and to -1 in NaCl
- the solution with NaClO is used as bleach. ClO- ions are
Responsible or the bleaching action of solution - NaClO kills bacteria and other microorganisms
Redox reactions of halogens. (Chlorine and water)
Chlorine is slightly soluble in water resulting in pale green solution
- some chlorine reacts with water to form mixture of
*HCl
*HOCl (cloric acid) - responsible for strong germicidal
and bleaching power of wet chlorine
*HClO (hypochlorus acid)
Cl2 + H2O = HCl + HClO
Explains use of chlorine to sterilise pools and drinking water
- HClO kills microorganisms; safe using as very little Cl2 used
- as HClO kills microorganisms, position of equilibrium moves
From left to right so little chlorine remains once HClO is used
-
- also the benefits of using Cl2 in drinking water and pools outweigh
- potential health risks to humans if not present. In high conc= toxic
Society assesses pros and cons of adding chemicals to water supplies
Sodium fluoride/other fluoride compounds may be added to drink water
As fluoride ions prevent tooth decay ,strengthening enamel
—> some may object as there may be other health risks with fluoride
As it can be viewed as mass medication
what happens when halogens react with other halide ions in solution
the reactions be explained by examining oxidising ability of halogens, and reducing ability of halide ions.
- Cl and Br react with water so are partially soluble in water
-
- chlorine in water is a solution called chlorine water; same with bromine water.
- chlorine water is pale green depending on conc (colourless at low conc)
- bromine water is orange
-
- iodine is slightly soluble in water (pale yellow solution)
- dissolves in a iodide ion solution to form brown solution
what are the oxidising power/ability of halogens
reduces as atomic number increases
> oxidants/ing agents causes oxidations
» so causes losses of electrons, but gains electrons/ reduces itself
-
-halogens gain electrons to fill shell
- electron gained by F atom completes an energy level closer to nucleus than Cl
–> electron F gains is has
stronger attraction to nucleus
as it’s closer than in Cl
–> electron gained by F is
subjected to less shielding (fewer electrons between electron and nucleus)
..
oxidising ability of halogens decreases down group
what would happen in displacement reactions?
- F2 displaces all other halides from solution of a halide compound
- Cl2 displaces bromide and iodide from solutions of these halide ions
- Br2 displaces iodide in solution
whats the reaction bromine reacts with aqueous iodide ions
colour change in solution isnt as clear since it simply darkens from oarnge to brown
- Br2+ 2I- = 2Br- + I2
half equations
- Br2 + 2e- = 2Br-
- 2I- = I2 + 2e-
whats the reducing ability of halides
icreases as atomic number increases
- reducants/ing agents lose electrons so another can gain electrons/be reduced
- they readily donate electrons, the ease wch a reductant loses electrons
- dtermines how effective it is
halide ions lose from outer shell.
- electron lost by iodide ion comes from shell further from nucleus than bromide
–> electron iodide loses has weaker attraction to nucleus (further compared to other halide ions)
–> electron lost is subjected to more shielding than others, as more electrons between.
..
fluoride and chloride ions have little reducing ability as not as able to donate electrons
reaction of solid halides (sodium fluoride) and conc sulfuric acid
solid halide salts react with conc sulfuric acid
- depending on reducing ability of halides
- NaF + H2SO4 = NaHSO4 + HF
- ionic
- F- + H2SO4 = HSO4 - + HF
this isnt a redox.
sodium hydrogen sulfate and hydrogen fluoride made
–> misty white fumes of HF
reaction of solid halides (sodium chloride) and conc sulfuric acid
2 NaCl + H2SO4 = NaHSO4 + HCl
- ionic
- Cl- + H2SO4 = HSO4 - + HCl
not redox.
–> misty white fumes form
reaction of solid halides (potassium bromide) and conc sulfuric acid
- KBr + H2SO4 = KHSO4 + HBr
2HBr + H2SO4 = Br2 + SO2 +
2H2O
first not redox; other is redox.
–> misty white vapour, red brown vapour
redox as Br oxidised=-1 HBr;0 Br
S reduced = +6 to +4
..
- overall equation
2KBr + 3H2SO4 =
2KHSO4 + Br2 + SO2 + 2H2O
–> red brown vapour is only observation
-
- ionic equations include
2Br- + 3H2SO2 =
2HSO4 - + Br2 + SO2 + 2H2O
how is conc sulfuric acid used to test for halide ions
when conc sulfuric acid is added to a solid bromide compound,
misty white fumes of (HBr) are released with some brown Br2 gas
when conc sulfuric acid is added to a solid iodide compound,
misty white fumes (HI),
purple fumes (I2) , and eggy smell (H2S)
not normaly used as sulfuric acid is corrosive
and toxic gases formed.
whats the trend in melting/boiling point across period 3 (non metals apart from Si)
- The melting point of P4, S8, Cl2, and Ar are low
- Phosphorus, sulfur and chlorine are simple covalent molecules
And little energy is required to overcome the weak vdw forces
- Ar exists as atoms and very little energy required to
Overcome weak vdws between atoms
The melting points increase from P to S then decrease
- the attractions between molecules are vdws (non polar molecules)
- if a molecule has more electrons the induced dipole interactions
Will be greater, causing greater vdw between molecules.
» increasing energy required to melt
- Argon’ melting point is very low as has very weak vdws
What are the trends down group 2 (ionisation energy)
Ionisation energy
- First ionisation energy decreases down group
- energy required to remove outer depends in attraction
Between nucleus and outer electron
- G2’s atomic radius increases
- outer electron is tightly held due to shielding
from increasing nuclear charge
- so easier to remove; ionisation energy decreases
What are the trends down group 2 (melting point)
Melting point
Generally melting points of the G2 metals decrease as go down
- G2 elements exist as giant metallic strucs
- a regular arrangement of positive ions
Surrounded by a sea of delocalised electrons
- metallic bond is attraction between the layers
Greater strength of metallic bond, higher the melting point
-
- number of delocalised electrons, metal ion charge and ion size
- are proportional to metallic bond strength
»_space; G2 has the same 2+ charge and number of delo. Per atom
- as group descends, radius increases
- attraction between ion and delo electrons decreases
So strength of metallic bond decreases and melting point decreases
What are trends in G7 (boiling point, electronegativity)
- Boiling points of elements increase
- size of diatomic molecules increase down group
- larger molecules have more electrons, leading to greater
Induced dipole dipole forces - these mean greater vdws forces between molecules
- more energy need to overcome greater vdws between mols
To change from liquid to gas
- Electronegativity decreases
- for smaller atoms, bonding electrons in a covalent bond are
Closer to positive nucleus - so less shielding, meaning positive nucleus attracts bonding pair
of electrons strongly (so greater electronegativity)
- for smaller atoms, bonding electrons in a covalent bond are
halogen redox reactions (chlorine and water in sunlight)
BUT in bright light, Cl2 and water react to produce a solution
With Cl- and O2
2Cl2 + 2H2O = 4 HCl + O2
—> ionic = 2Cl2 + 2H2O = 4H+ + 4Cl- + O2
- The UV radiation from sun breaks down HClO into HCl and O2
»_space; 2 HClO = 2HCl + O2
reactions of solid halides (potassium iodide) and conc sulfuric acid
KI + H2SO4 = KHSO4 + HI
2HI + H2SO4 = I2 + SO2 + 2H20
6HI + H2SO4 = 3I2 + S + 4H2O
8HI + H2SO4 = 4I2 + H2S + 4H2O
–> misty white fumes (HI)
–> purple vapour (I2)
–> yellow solid (S)
–> rotten egg smell (H2S)
–> black solid (I2)
all are redox apart from first
iodine is the oxidation product
and sulfur dioxide, sulfur and hydrogen sulfide are reduction products
reactions of solid halides (potassium iodide) and conc sulfuric acid - what are the ionic equations?
- I- + H2SO4 = HSO4 - + HI
- 2I- + SO4 2- + 4H+
= I2 + SO2 + 2H2O - 6I- + 8H+ + SO4 2-
= 3I2 + S + 4H2O - 8I- + 10H+ + SO4 2-
= 4I2+ H2S + H2O
The redox reactions happen due to relative reducing powers
of halides/hydrogen halides
- Cl- not as powerful as Br-, so
HCl cant reduce H2SO4 (only displaced by the acid)
-
- Br- in HBr reduce better so some of the H2SO4 is reduced to SO2 and H2O
products are mixed = Br2 and HBr
-
- I- are best reductants, reducing S in in H2SO4 to SO2
and even further to H2S and S
product is mainly I2
how are halide ions identified
can be identified by using Ag+1 ions in a solution
1. spatula of halide dissolves and is acidified first with dilute HNO3
- removes other ions that could react with Ag+1 ions. could form ppt that would interfere
- few cm3 silver nitrate solution added
-(if F- is in solution, no ppt
»_space; AgF is soluble in water, so cant be used to test) - the ppt formed from each other halide ion is insoluble
Ag+ + Cl- = AgCl (white ppt)
Ag+ + Br- = AgBr (cream ppt)
Ag+ + I- = AgI (yellow)
whats a further test with ammonia solution to confirm identity of halide ion (after ppt test)
- silver chloride will dissolve in dilute and conc ammonia solution
> forming colourless solution - silver bromide doesnt dissolve in dilute, but does in conc ammonia solution
> forming colourless solution - silver iodide doesnt dissolve in either
for ammonia solution test to confirm identity of halide ion, WHY does this happen?
redissolution of AgCl in dilute ammonia
- bc of complex formation
- ammonia molecules have a lone electron pair
- form a coordinate bond with lone pair on Ag+ ion
» complex ion is soluble so AgCl dissolves
AgI doesnt dissolve in either
- mostly covalent compound bases on electronegativity values of Ag and I
> so more covalent a compound the less likely it is to form a complex ion with ammonia
..
dissolution of silver chloride..
AgCl + 2NH3 = [Ag(NH3)2]+ + Cl-
silver halides darken when exposed to sunlight , forming black deposit of silver
» photographic plates darken when exposed to light
what is the test for sulfate ions
(may also identify barium ions in a mixture containing SO4 2- ions, giving same result)
- dissolve spatula of sample in dilute HNO3 or HCl
- add a few cm3 of barium chloride solution
- white ppt if sulfate ions present (white ppt is BaSO4)
carbonate ions test
(HCO3 ions also react with the acid producing CO2)
1. place few cm3 of dilute HNO3 in test tube
2. add spatula of sample
3. if effervescence and solid disappears, gas produced when bubbled through limewater makes milky, it’s present
- carbon dioxide is the gas
whats the carbonate test (rules out hydrogen carbonate ions)
- dissolve spatula of sample in deionised water
- add a few cm3 of Magnesium nitrate OR chloride solution
- white ppt forms if present (MgCO3)
testing for Mg2+ or Ca2+ cations
(or can test presence of OH- add Mg(NO3)2 solution to sample and white ppt forms)
- dissolve spatula of sample in deionised water
- add few drops NaOH solution
- then 5cm3 of NaOH , adding until excess
- present if white ppt doesnt dissolve in excess, Mg2+ or Ca2+
- white ppt is Mg(OH)2 or Ca(OH)2
test for ammonium ion (NH4 +)
(also tests for OH- ; add ammonium compound to sample and warm. ammonia gas if present)
..
- few cm3 of NaOH solution in test tube
- add spatula of sample; warm gently
- test any gas given off with red litmus/UI
» or using glass rod dipped in conc HCl - pungent gas produced turns litmus/UI blue
OR white smoke with rod
—> ammonia released from action of alkali (NaOH) on ammonium compound
> ammonia gas is alkaline
> white smoke is ammonium chloride from reaction of NH3 with HCl
