3.2: Periodicity Flashcards
(79 cards)
1
Q
What is periodicity?
A
a regular repeating pattern of atomic physical and chemical properties
2
Q
What is the trend in melting points across period 3?
A
- increases
- due to increase in nuclear charge
- more delocalised eolectrons to increase metallic bond strength
- more energy needed to break these bonds
3
Q
Na, Mg, Al
What are the trends in melting and boiling point across Period 3?
A
1.Na, Mg, Al
- giant metallic latice with positive metal ions surrounded by a sea of delocalised electrons
- strong electrostatic forces, increasing nuclear charge
- Mp- 371,922,933 l Bp- 1156,1380,2740
4
Q
Si
What are the trends in melting and boiling point across Period 3?
A
2.Si
- giant covalent/ macromolecular molecule
- many strong covalent bonds
- Mp- 1683 l BP- 2628
melting and boiling point fluctuate
5
Q
P4, S8, Cl2
What are the trends in melting and boiling point across Period 3?
A
3.P4,S8,Cl2
- simple molecular
- weak intermolecular forces, VDW forces, more molecules, stronger forces as larger molecule
- Mp- 317, 392, 172 l Bp- 553, 718, 238
melting and boiling point fluctuate
6
Q
Ar
What are the trends in melting and boiling point across Period 3?
A
4.Ar
- atomic
- very weak forces between atoms
- Mp- 84 l Bp- 87
melting and boiling point fluctuate
7
Q
What is the trend in ionisation energies across Group 2?
A
- Melting point decreases down the group (Mg has anomalylously low b.p. due to differnt crystal structure)
- Atomic radius increases
- Solubility of salts increases
- Reactivity with water increases
- Ionisation energy decreases
8
Q
What is the trend in melting points across Group 2?
A
- Size of the atoms increase
- Distance between the nucleus and “sea” of delocalised electrons increases
- Strength of metallic bond decreases
- Energy required to overcome the bond decreases
- Melting points decrease down the group
- Magnesium has an unexpectedly low melting point
9
Q
Why does reactivity increases down Group 2?
A
- Increased atomic radius/ size of atoms/ number of electron shells
- Greater distance between nucleus and outer shell electrons
- Weaker attraction between outer electrons and nucleus / more shielding
- Easier to remove outer electrons
10
Q
What happend when you react Group 2 metals with water?
A
By reacting group 2 metal with water, metal hydroxides and hydrogen are formed.
M(s) + 2H2O (l) → M(OH)2 (aq) + H2 (g)
Metal is oxidised
Hydrogen is reduced
Magnesium also reacts with steam at high temperatures to produce magnesium oxide:
Mg(s) + H2O(g) → MgO(s) + H2(g)
11
Q
What is the pattern in solubility in Group 2 hydroxides?
A
Group 2 hydroxides, X(OH)2, are all white solids that become more soluble down the group.
- Magnesium hydroxide- almost insoluble
- Calcium hydroxide- sparingly soluble
- Strontium hydroxide- more soluble
- Barium hydroxide-dissolves to produce strong alkaline solution (because of the OH- ions)
12
Q
What is the pattern in solubility in Group 2 sulfates?
A
By reacting group 2 metal with sulphuric acid, metal sulphates are formed.
Reaction equation:
Mg (s) + H2SO4 (aq) → MgSO4 (aq) + H2 (g)
The group 2 sulphates get less soluble down the group (Mg-Ba).
BaSO4 is insoluble.
13
Q
Acidified BaCl2 solution
How do you test for sulphate ions?
A
- Sulphate ions in solution, (SO4^2-), are detected using acidified barium chloride solution.
- The test solution is acidified using a few drops of dilute hydrochloric acid, and then a few drops of barium chloride solution are added.
- A white precipitate of barium sulfate forms if sulfate ions are present:
BaCl2(aq) + Na2SO4(aq) → 2NaCl(aq) + BaSO4(s) - The hydrochloric acid is added first to remove any carbonate ions that might be present - they would also produce a white precipitate, giving a false positive result.
14
Q
What are different uses of Group 2 in medicine, agricuture and titanium extraction?
A
- Magnesium hydroxide (Mg(OH)2), milk of magnesia, acts as an “antacid”- neutralises excess stomach acid- as a treatment for indigestion.
- Barium sulphate (BaSO4) is used in x-rays to image a patient’s stomach and intestines as it is opaque to x-rays.
- Calcium hydroxide (Ca(OH)2) can be added to solid to raise the pH level to 6-7, the optimum for most plants.
-
Magnesium: TiO2 is converted into TiCl4 by heating with carbon and chlorine, TiCl4 is then reduced by Mg
** TiCl4(g) + 2Mg(l) -> Ti(s) + 2MgCl2(l)** - titanium is reduced, magnesium is oxidised
15
Q
What are Group 7 elements?
A
- halogens
- highly reactive non-metals of group 7
- Flourine- pale yellow
- Chlorine- green
- Bromine- red-brown
- Iodine- grey
16
Q
What is the pattern in electronic configuration, melting and boiling points and ionisation energies across Group 7?
A
- boiling points increase down the group due to increase in strength in VDW forces
- electronegativty decreases down the group
- reactivity decreases down the group as it’s harder to gain an electron as there’s more shielding and a larger atomic radius so further from electrosatic forces and opposing charge.
17
Q
When will a halogen displace a halide from solution?
A
if the halide is below it in the periodic table
- add Cl2 (aq) (colourless)- no reaction in KCl (colourless) water l forms orange solution (Br2) in KBr (colourless) water l forms brown solution in KI (colourless water)
- add Br2 (aq) (orange)- no reaction in KCl (colourless) water l no reaction in KBr (colourless) water l forms brown solution in KI (colourless water)
- add l2 (aq) (brown)- no reaction in KCl (colourless) water no reaction in KBr (colourless) water l no reaction in KBr (colourless) water l no reaction in KI (colourless) water
18
Q
What happens when you mix cold, dilute, aqueous sodium hydroxide with chlorine gas?
A
- makes sodium chlorate (I) solution NaClO(aq)
- bleach
19
Q
What is disproportionation and an example?
A
A reaction where the same element is both reduced and oxidised
e.g. Chlorine and water:
Cl2 + H2O 🡪 2H+ + ClO- + Cl-
Chlorate ions, ClO- , kill bacteria so chlorine is added to water to make it safe.
20
Q
Why is chlorine water added to water (benefits and disadvantages)?
A
chlorate (I) ions kill bacteria so adding chlorine to water can make it safe to drink or swim in
advantages:
- kills disease-causing microorganisms
- some chlorine persists in water, preventing reinfection down the supply
- prevent algae growth eliminating bad tastes, smells and discolouration caused by organic compounds
disadvantages:
- irritates the respiratory system
- liquid chlorine causes chemical burns
- any organic compounds in water can form chlorinated
hydrocarbons which are carcinogenic
however, cancer risk small compare to untreated water (e.g. cholera)
21
Q
How do you test for Halides?
A
- Halogens (group 7) can be tested for using silver nitrate solution with nitric acid.
- Ag + (aq) + X-(aq) 🡪 AgX(s)
- Silver Flouride- no precipitate
- Silver Chloride- white precipitate (forms slowest l most soluble in NH3 (dilute))
- Silver Bromide- cream preciptate (2nd most soluble in NH3 (concentrated))
- Silver Iodide- yellow precipitate (forms fastest l least soluble in NH3 (dilute)))
22
Q
What is the reducing power of the halides?
A
- The reducing power of a species is related to the reactivity.
- Reducing power increases across the group
- As reduction is gain of electrons, the reducing power of a species is how easily it can reduce (add electrons to) another species.
- The reducing power of the halides increases down the group as they are more easily able to lose electrons to another species and reduce them.
23
Q
NaF or NaCl
What occurs in the reaction of the halides with sulphuric acid?
A
NaF(s) + H2SO4(aq) 🡪 NaHSO4(s) + HF(g)
NaCl(s) + H2SO4(aq) 🡪 NaHSO4(s) + HCl(g)
- Misty fumes will be seen.
- HF and HCl not strong enough reducing agents so reaction stops there.
24
Q
NaBr
What occurs in the reaction of the halides with sulphuric acid?
A
NaBr(s) + H2SO4(aq) 🡪 NaHSO4(s) + HBr(g)
2HBr(g) + H2SO4(aq) 🡪2H2O(l) + Br2(g) + SO2(g)
- Misty fumes will be seen.
- HBr is a stronger reducing agent so reacts in a redox reaction.
- Choking fumes
- Orange fumes
25
# NaI
What occurs in the reaction of the halides with sulphuric acid?
**NaI(s) + H2SO4(aq) 🡪 NaHSO4(s) + HI(g)**
**2HI(g) + H2SO4(aq) 🡪 2H2O(l) + I2(s) + SO2(g)**
**6HI(g) + SO2(g) 🡪 2H2S(g) + 3I2(s) + 2H2O(l)**
- Misty fumes will be seen.
- HI is a strong reducing agent so reacts with the sulphuric acid, and then also reduces the SO2.
- H2S- toxic smells of bad eggs
26
What is the reaction of sodium with water?
**2Na(s) + H2O(l) -> 2NaOH(aq) + H2(g)**
- reacts in water in rapid and violent reaction
- sodium fizzes (effervescence) and floats on surface of water
- exothermic reaction
- resulting solution __NaOH__ is strongly alkaline, solution of __pH 13-14__
27
What is the reaction of magnesium with water?
__Mg(s) + 2H2O(l) -> Mg(OH)2(aq) + H2(g)__
- reacts less rapidly and violetly than sodium
- very slow reaction at room temp
- results in weak alkaline as Mg less soluble in water than Na
- **pH 10**
28
What is the reaction of chlorine with water?
**Cl2 (aq) + H2O (l) -> HCl (aq) + HClO (aq)**
- redox reaction/ disproportionation (Cl both oxidised and reduced)
- gas produces, aqueous solution
- dissolves, producing very pale green solution
29
What is the reaction of sodium and oxygen?
| Na
**2Na(s) + 0.5O2(g) -> Na2O(s)**
- exothermic, Na reacts directly with oxygen
- burns bright yellow flame, white smoke (Na peroxide)
30
What is the reaction of magnesium and oxygen?
| Mg
**2Mg(s) + O2(g) -> 2MgO(s)**
- Mg ribbon burns readily in O2, producing MgO
- gives bright white flame, leaves white powder, exothermic
31
What is the reaction of aluminium and oxygen?
| Al
**4Al(s) + 3O2(g) -> 2Al2O3(s)**
- exothermic, reacts slowly w/ O2 unless in powder form
- burns bright, produces white powder Al2O3 (white flame)
32
What is the reaction of silicon and oxygen?
| Si
**Si(s) + O2(g) -> SiO2(s)**
- exothermic, reacys w/ O2 slowly (red/orange flame)
- must be heated strongly in oxygen
33
What is the reaction of phosphorus and oxygen?
| P4
**P4(g) + 5O2(g) -> P4O10(s)**
- exothermic, red phosphorus must be heated before reacting
- white P4 reacts spontaneously w/ O2 producing white smoke (**phosphorus sun**- white light)
34
What is the reaction of sulfur and oxygen?
| S8
**S(g) + O2(g) -> SO2(g)**
- exothermic, blue flame
- smoke produced, not fully oxidised
35
What is the reaction of sulfur and oxygen at high temp. and with a catalyst?
| S8
**S(g) + 3/2O2(g) -> SO3(l)**
36
What is the reaction of Na2O and water?
| giant ionic
**Na2O(s) + H2O(l) -> Na+(aq) + 2OH-(aq)**
- base, neutralising acids forming alkaline solutions
- strong alkaline soluion **pH 14**
37
What is the reaction of MgO and water?
| giant ionic
**MgO(s) + H2O(l) -> Mg(OH)2(s) ⇌ Mg2+(aq) + 2OH-(aq)**
- base, neutralising acids, form alkaline solutions
- sparingly soluble in water, **pH 9-10**
38
What is the reaction of Al2O3 and water?
| giant ionic (covalent character)
- **insoluble in water**, so no reaction in water
- Al2O3 insoluble due to high lattice enthalpy
39
What is the reaction of SiO2 and water?
| giant covalent
- **insoluble in water**, so no reaction in water
- SiO2 insoluble as giant covalent molecule with millions of strong covalent bonds
40
What is the reaction of P4O10 and water?
| molecular
**P4O10(s) + 6H2O(l) -> 4H3PO4(aq)**
**H3PO4(aq) ⇌ H2PO4-(aq) + H+ (aq)**
- forms acidic solutions with non-metal oxides
- reats violently with water
- produces acidic phosphorus acid solution
- H2O molecules attatch to delta positive P atoms, leading to release of H+ ions from H2O (**pH 0-1**)
41
What is the reaction of SO2 and water?
| molecular
**SO2(g) + H2O(l) -> H2SO3(aq) ⇌ HO3-(aq) + H+(aq)**
- SO2 fairly soluble, produce acidic solutions **(pH 2-3)**
42
What is the reaction of SO3 and water?
| molecular
**SO3(g) + H2O(l) -> H2SO4(aq) -> HSO4-(aq) + H+(aq)**
- SO3 reacts violently w/ water, producing sulfuric acid (**pH 0**)
43
How does Na2O and MgO react with acids/bases?
- bases
- neutralise acids to form salt and water
**Na2O(s) + H2SO4(aq) -> Na2SO4(aq) + H2O(l)**
**MgO(s) + 2HCl(aq) -> MgCl2(aq) + H2O(l)**
44
How does Al2O3 react with acids/bases?
**amphoteric**- acts as an acid and a base
**Al2O3(s) + 6HCl(aq) ->2AlCl3(aq) +3H2O(l)**
**Al2O3(s) + 2NaOH(aq) +H2O(l) -> 2NaAl(OH)4(aq)**
45
How does SiO2 react with acids/bases?
- reacts as a weak acid w/ a strong base
**SiO2(g) + 2NaOH(aq) -> NaSiO3(aq) + H2O(l)**
46
How does P4O10 react with acids/bases?
- reacts w/ alkali, reacts in 3 stages due to 3 OH- groups
- each H turns into OH- ion replaced by sodium ion
**P4O10(s) + 6H2O(l) -> 4H3PO4(aq)**
**H3PO4(aq) + 3NaOH(aq) -> Na3PO4(aq) + 3H2O(l)**
47
How does SO2 react with acids/bases?
- first reacts w/ NaOH then sodium sulphate(IV)
**SO2(aq) + NaOH(aq) -> NaHSO3(aq)**
**NaHSO3(aq) + NaOH(aq) -> Na2SO3(aq) + H2O(l)**
48
What are structures of different anions formed when period 3 reacts with acids and bases?
- PO43-: phosphate (V)- basic
- SO32-: sulfate (IV)- acid
- SO42-: sulfate (VI)- acid
49
What is a transition metal?
a metal that contains an incomplete d-subshell in atoms or one of its common ons (can form 1 or more stable ions w. incomplete d-subshells)
**Ti -> Cu**
50
Why are some elements not transition metals?
- **Sc-** forms only one ion (Sc3+) which has an empty d-subshell
- **Zn-** forms only one ion (Zn2+) which has a full d-subshell
51
What are the physical properties of transition metals?
- conductive/ unreactive (high density)
- high melting and boiling point
- strong, hard shiny, same ionic radii
52
Why do transition metals have these physical properties?
due to an incomplete d-subshell
53
What are the chemical properties of transition metals?
- **variable oxidation states-** trasition metals have more than one oxidation state in their compounds,ca take part in many redox reactions
- **catalytic action**
- **form complex ions**
- **coloured ions-** distinctive colours caused by compound absorbing energy corresponding to light in the visible region of the spectrum
54
What is a complex ion?
| ligand,complex,co-ordinate no.
- formed when a transition metal ion is surrounded by its ions, or other molecules, collectively called ligands (w/ a coordinate bond)
- **ligand-** particle w/ lone pair of electrons bonded to metals by co-ordinate bond
- **complex-** metal ion w/ co-ordinately bonded ligands
- **coordination no-** no. of co-ordinate bonds formed from ligands to the metal ion
55
What is a lewis base and acid?
base- lone pair donor
acid- lone pair acceptor
56
What are the different types of ligands?
- **monodentate-** only form one coordinate bond (e.g. Cl-, OH-, CN-, H2O, NH3)
- **bidentate-** ligands forming 2 coordinate bonds to the metal ion, lower co-ord no.(no. of bonds) (e.g. 1,2-diaminoethane (NH2CH2CH2NH2), ethanedioate ion (C2O42-)
- **multidentate-** form many co-ordinate bonds to metal ion (e.g. EDTA4-, porphysin (in haemoglobin))
57
What is haemoglobin and how is it a complex ion?
- globular protein w/ 4 Fe2+ centres, multidentate ligand
- **each iron has porphyrin ligand 4 out of 6 co-ordinate sites**
- **other 2 co-ordinate sites bonded to rest of haemoglobin and 6th oxygen** ( binds as ligands)
- 5th nitrogen acts as ligand as part of complex protein (globin)
- oxygen not a good ligand, so given up by cells easily
- **if CN/ CO inhaled, substitutes O2 as CO better ligand, can bond irreversibly to iron (toxic, prevents O2 transfer)**
58
How is optical isomerism shown in complex ions?
- when 3 bidentate ligands co-ordinately bond to central metal atom
- exist as non-super imposable mirror images
59
How does cis-trans isomerism shown in complex ions?
- octahedral complex w/ 4 monodentate ligands of one type, 2 monodentate ligands of other type
- sqaure planar, two ligands, one type different to other 2
60
What is ligand substitution?
- where one ligand replaced by another ligand
- new complex formed more stable than original
- ligands in original complex can be partially substituted by others
- the complex ion can change charge or remain the same, depending on ligand
61
What happens if a ligand of the same size is substituted?
- no change in co-ordination no.
- uncharged, no change in geometry complex
- e.g. H2O/ NH3 (co-ordination no. 6)
62
What happens if a bigger or smaller size ligand in substituted?
- if ligands of different sizes are substituted, co-ordination number may change
- charge can also change
- e.g. Cl- can form 4 co-ordinate bonds to most transition metals, larger ligand, change co-ordination no. and charge
63
What is the chelate effect?
- when **bidentate/ multidentate** ligands replace **monodentate** ligands to form a complex (chelation)
- △H is negligible as same no. of similar bonds are broken and formed (substitution)
- when ligands are replaced to form more co-ordinate bonds, increase in **entropy(△S)** as more moles produced creates more disorder
64
How does chelation links to gibbs free energy change?
- feasibility depends on **△G = △H - T△S**
- if △G is negligible and △S is very positive, △G is very negative the reaction is very feasible
- therefore, ligands replaced by those forming fewer co-ordinate bonds aren't feasible
65
What are chelating agents?
- ligands that are very good at bonding to metal ion and difficult to remove
- ligands forming more than one co-ordinate bond
**EDTA4-:** uses lone pairs on 4O2 and both H atoms
- complex ions w/ polydentate ligands are called chelates
- chelates demove d-block metal ions from a solution
66
What are the different shapes of complex ions?
- linear (180°)
- square planar (90°)
- tetrahedral (109.5°)
- octahedral (90°)
67
What is an isomer?
- compounds w/ the same molecular formula, but different arrangement of atoms in space
68
How do geometrical (EZ) isomers arrise in complex ions?
- square planar and octahedral complexes
- if opposite to each other (trans)
- if next to each other/ ion same side (cis)
69
How do optical isomers arrise in complex ions?
- stereoisomers (same structural formula) form 2 non super-imposable mirror images
- 3 bidentate ligans co-ordinately bonded to central metal ion
- octahedral complex
70
Why are there coloured ions in complex ions?
- gap in energy between d orbitals
- when electrons move from excited to ground state, absorb UV energy/ light
- wavelength light not absorbed is reflected and shows the colour
71
What is the equation for energy gap between d orbitals?
**△E= hf = (hc)/ Λ**
energy gap (J)= plancks constant x frequency of light
or
energy gap (J)= (plancks constant x velocity of light (ms-1))/ wavelength(m)
72
What factors affect colours of ions?
- **higher oxidation state-** stronger interaction w/ ligand, higher energy spectrum absorbed
- **co-ordination number-** lower co-ord number, increase strength of metal-ion ligand, alter spilitting of orbitals
- **different ligands-** split d orbital by different amounts of energy, frequency of light absorbed changes
Zn2+, Cu+, Sc3+ aren't coloured as all have a full/ empty d shell
| different metals of same colour can have different colours
73
What if UV/ visible spectrosopy used for?
measure frequency where complex absorbs UV light
- UV light passed through complex, detects drequence of UV light passing
- light that doesn't pass isn't absorbed
- more conc. solution, more light absorbed, used to measure conc.
74
What is colorimetry used for?
- uses light source and detector to measure amount of light of particular wavelength passing through the coloured solution
- determines conc. of coloured ions in solution, uses a blank sample
- **more conc.= more light absorbed**
75
What occurs in colorimetry if colour isn't that prominent?
- **ligand substitution** used to intensify the colour
- **coloured filter** used to maximise light absorbance (of complementary colour)
76
How does colorimetry work?
1. get colour standard (water)
2. place coloured filter, of complementary colour to sample
3. get known concentrations of metal ion solution and measure light absorbed/ transmitted
4. plot calibration curce (conc. x light transmitted/ absorbed)
5. find % light transmitted/ absorbed for unknown sample
6. use calibration curve to find unknown concentration
77
What are uses of transition metal ions variable oxidation states?
- testing for aldehydes
- testing for primary and secondary alcohols
- redox titrations
- catalysts
78
How are variable oxidation states used to test for aldehydes?
- **tollen's reagent [Ag (NH3)2]+ :** silver mirror formed when aldehyde present (converted to carboxylic acid) | **Ag(1+) reduced to Ag(0)**
- **fehlings solution (Cu2+):** brick red (C2O) ppt formed when aldehyde present (converted to carboxylic acid) | **Cu(2+) reduced to Cu(1+)
79
