Le tableau périodique Flashcards
What are the valence electronic configurations of each group?
Group - Valence e- config
1 - ns1
2 - ns2
13 - ns2np1
14 - ns2np2
15 - ns2np3
16 - ns2np4
17 - ns2np5
18 - ns2np6
Describe trend of atomic radius across period 3
- atomic radius decrease across period
Across period, - no of proton increase, so nuclear charge increase
- no of e- increase but added to same outermost shell, so shielding effect remain relatively constant
- eff nuclear charge increase, cause stronger e-static attract n btw nucleus & valence e-
- Hence, valence e- closer to nucleus
Why is radius of Ar (grp 18) not compared commonly?
- impractical to measure radius of isolated atom. Half distance btw ctr (nucleus) of pair of atoms (differs for each type of bonding)
- since Ar is monoatomic, neighbouring atoms are not actually bonded, meaning e- cloud overlap is minimal
=> different from other elements compared
Describe trend of ionic radius across period 3
- Anions (P, S, Cl) larger than cations (Na, Mg, Al, Si)
Because
- anions hv 1 more shell of e- than cations
- valence e- in anions experience greater shielding effect AND less strongly attracted to nucleus - Among cations (which r isoelectric), ionic radius decrease w atomic no.
Similarly, among anions (iso electronic oso), ionic radius decrease w atomic no.)
Because:
For cations Na to S,
- isoelectronic (10 e-: 1s² 2s² 2p^6)
- nuclear charge increases as proton no increase
- thus, valence e- more strongly attracted to nucleus
For anions P to Cl,
- isoelectronic (18 e-: 1s² 2s² 2p^6 3s² 3p^6)
- nuclear charge increases as proton no increase
- hence valence e- more strongly attracted to nucleus
Why are cations always smaller than parent atoms?
Both cations & parent atoms hv same nuclear charge (since no of proton same). Cations hv 1 less shell of e- than parent atoms. So, valence e- in cations more strongly attracted to nucleus
Why are anions always larger than parent atoms?
anion hv more valence e- than parent atom. So, inter-electronic repulsion experienced by valence e- of anion is hence greater than that in parent atom. Since both species hv same amt nuclear charge (since no of proton same), resultant e-static attractive force of nucleus on valence e- in anion is smaller than in parent atom, causing greater anionic size
How does 1st ionisation energy change across period 3?
-general increase
across period
- no of protons increase, so nuclear charge increase
- no of e- increase but added to same outermost shell, so shielding effect remain relatively constant
- eff nuclear charge increase, so stronger e-static attract n btw nucleus & valence e-
- So, more energy needed to remove valence e-
Anomalies:
1. Al hv lower than expected 1st IE
- 3p valence e- to be removed from Al hv higher energy than 3s valence e- in Mg
- so less energy required to remove 3p e- in Al than 3s e- in Mg
- S hv lower than expected 1st IE
- inter-electron repulsion exist btw paired e- in 3p orbital of S atom.
- so, less energy needed to remove valence e- from S
How does electronegativity change across period 3?
- increase
Across period,
- no of protons increase so nuclear charge increase
- no of e- increase but added to same outermost shell, so shielding effect remain relatively constant
- eff nuclear charge increase, so stronger e-static attract n btw nucleus & shared e-
- so electro-vity increase
How does melting point change across period 3?
(1) for metals, increase fr Na to Al (relatively high)
bcos
- Na, Mg and Al hv giant metallic structure w strong e-static force of attract n (metallic bonds) btw metal cations & sea of delocalised e-. Large amt energy needed to break strong metallic bonds, so high mp
- increase in mp bcos increase in metallic bond strength
- >can b attributed to
1. increase in no of valence e- contributed per atom
2. increase in charge density due to decrease cationic radius & increasing charge of cation
(2) Si hv highest mp
bcos
- hv giant molecular structure, Si atoms held tgt by numerous strong covalent bond
- large amt energy needed to overcome strong, extensive covalent bonding btw atoms in giant 3-D molecular structure
(3) S8>P4>Cl2>Ar (relatively low)
bcos
- simple molecular structure, molecules (or atoms for Ar) held tgt by weak id-id attract n, needing small amt energy to overcome. So, mp low
- size of e- cloud: Ar<Cl2<P4<S8
- extent of distor n of e- cloud: Ar<Cl2<P4<S8
- extent of weak id-id attract n: Ar<Cl2<P4<S8
- energy required to overcome intermolecular weak id-id attract n: Ar<Cl2<P4<S8
- mp:Ar<Cl2<P4<S8
How does electrical conductivity change across period 3?
(1) increase fr Na to Al (high electrical conductivity)
- due to presence of delocalised e- as mobile charge carriers to conduct electricity/heat
- no of delocalised e- per atom in metallic structure increase (fr Na to Al)
(2) Si, a metalloid, hv slight electrical conductivity
- silicon is a semiconductor
(3) non-metals (P to Ar) do not conduct electricity
- due to absence of delocalised e- & free mobile ions to conduct electricity
How does highest oxidation number of element across period 3 change (oxides and chlorides)?
- O.N. in oxides & chlorides r +ve as all less electro-ve than either oxygen or chlorine
- trend: increase in O.N. across period (correspond to no of valence e- available for bond form and max O.N. = no of valence e-)
- for P & S, in oxides & chloride where exhibit max O.S., central atom hv expanded octet (possible due to available vacant, energetically accessible 3d orbitals, can be used for bonding)
What bonds and structure do period 3 elements form with oxides? Why is this so?
(element: bond, structure)
- Na to Al: ionic, giant ionic
- Si: covalent, giant molecular
- P,S: covalent, simple molecular
As electro-ve (EN) diff increase across period, non-metallic elements eg P, S, Cl, O hv small diff in electro-vity -> more energetically feasible to share e-, form covalent bond than transfer e- form ionic.
Due to larger diff in EN values btw metal & non-metal, complete transfer of e- occur, cause form ions and ionic bonding
How does melting point of period 3 element oxides change?
trend:
- mp of Na2O, MgO, Al2O3 high
- mp: Na2O < MgO
- mp: MgO > Al2O3
bcos giant ionic structure, strong e-static attract n btw opp charge ions. Large energy need to overcome attract n
lattice energy formula, ionic bond directly proportional to strength
Mg2+ higher ionic charge, smaller ionic radius than Na+. Ionic bond MgO stronger than Na2O, need more energy to break
Al3+ high charge density, polarise O2-, so Al2O3 ionic bond hv some covalent character, weaken ionic bond strength ->lower mp
- mp of SiO2 high
bcos giant molecular, numerous strong covalent bonds btw Si & O atoms that need Large amt energy to break - mp of P4O10 & SO3 low
- mp: P4O10 > SO3
bcos simple molecular, weak id-id attract n btw molecules needing small amt energy to overcome
P4O10 larger e- cloud size, so stronger id-id, needing more energy to overcome
How do period 3 oxides react w water? Give equation when appropriate
Na2O: dissolve completely, form strongly alkaline sol n,
Na2O (s) + H2O (l) –> 2NaOH (aq),
pH 13
MgO: dissolve partially, form weakly alkaline sol n,
MgO (s) + H2O (l) ⇌ Mg(OH)2 (s) [NOTE: Mg(OH)2 oso partial soluble, result in weakly alkaline sol n)
pH 9
Al2O3: insoluble (due to high lattice energy)
pH 7 (same as pure water)
SiO2: insoluble (due to strong covalent bonds)
pH 7
P4O10 and SO3: dissolves, form strongly acidic sol n
P4O10 (s) + 6H2O (l) –> 4H3PO4 (aq) [phosphoric(V) acid)
SO3 (g) + H2O (l) –> H2SO4 (aq)
pH both 2
How do period 3 oxides react with acid and/or base? Give equation
(basic) Na2O & MgO: react w acid form salt, water
Na2O (s) + 2HCl (aq) –> 2NaCl (aq) + H2O (l)
MgO (s) + 2HCl (aq) –> MgCl2 (aq) + H2O (l)
(amphoteric) Al2O3 (bcos ionic w covalent character) react w both acids & bases
Al2O3 (s) + 6HCl (aq) –> 2AlCl3 (aq) + 3H2O (l)
Al2O3 (s) + 2NaOH (aq) + 3H2O (l) –> 2NaAl(OH)4 (aq) [sodium aluminate]
(acidic) SiO2, P4O10, SO3 react w base form salt, water
SiO2 (s) + 2NaOH (conc.) [hot & conc.] –> Na2SiO3 (aq) [sodium silicate] + H2O
P4O10 (s) + 12NaOH (aq) –> 4Na3PO4 (aq) + 6H2O (l)
SO3 (g) + 2NaOH(aq) –> Na2SO4 (aq) + H2O (l)
How do period 3 METAL hydroxides react with acid and/or base? Give equation
(basic) NaOH & Mg(OH)2: react with acid to form salt, water
NaOH (s) + HCl (aq) –> NaCl (aq) + H2O (l)
Mg(OH)2 (s) + 2HCl (aq) –> MgCl2 (aq) + H2O (l)
(amphoteric) Al(OH)3: react w both acids and bases
Al(OH)3 (s) + 3HCl (aq) –> AlCl3 (aq) + 3H2O (l)
Al(OH)3 (s) + NaOH (aq) [in excess] –> NaAl(OH)4 (aq) [sodium aluminate, colourless sol n]
How does melting point of period 3 chlorides change? Explain
1.
mp NaCl, MgCl2 high
mp NaCl > MgCl2
bcos
giant ionic structure, strong e-static force attract n btw opp charge ions that need Large amt energy to overcome
Mg2+ hv higher charge density, polarises Cl- anion e- cloud to some extent, so some covalent character in ionic bond. Weaken e-static force attract n btw Mg2+ & Cl-
- mp AlCl3 low (compared to NaCl, MgCl2, sublimes)
bcos
Al3+ hv high charge density, polarises large Cl- anion e- cloud to vv large extent. So, Al-Cl bond is covalent, hv simple molecular structure w weak id-id that need small amt energy to overcome
3.
mp SiCl4, PCl5 low
mp SiCl4 < PCl5
bcos
simple molecular structure w weak id-id attract n that need small amt energy to overcome
size e- cloud: SiCl4 < PCl5
extent of distort n of e- cloud: SiCl4 < PCl5
extent of weak intermolecular id-id attract n: SiCl4 < PCl5
energy needed to overcome weak intermol id-id attract n: SiCl4 < PCl5
How do period 3 chlorides react with water?
NaCl dissolve readily, form neutral sol n (no rxn)
NaCl (s) –> Na+ (aq) + Cl- (aq)
pH 7
MgCl2 dissolve readily form slightly acidic sol n (due to slight hydrolysis of aq cation)
MgCl2 (s) –> Mg2+ (aq) + 2Cl- (aq)
[Mg(H2O)6]2+(aq) ⇌ [Mg(H2O)5(OH)]+ (aq) + H+(aq)
pH 6.5
AlCl3 dissolve readily form strongly acidic sol n (due to appreciable hydrolysis of aq cation)
AlCl3 (s) –> Al3+ (aq) + 3Cl- (aq)
[Al(H2O)6]3+ (aq) ⇌ [Al(H2O)5(OH)]2+ (aq) + H+ (aq) {= H3O+}
*NOTE: w limited water, white solid formed, white fumes of HCl observed (Al(OH)3 (s) oso form)
pH 3
SiCl4 & PCl5 react w water (readily hydrolyse) form white fumes (HCl) & strongly acidic sol n
SiCl4 (l) + 2H2O (l) –> SiO2 (s) + 4HCl (g)
PCl5 (s) + H2O (l) [cold/limited amt] –> POCl3 (l) + 2HCl (g)
OR
PCl5 (s) + 4H2O (l) [hot/excess] –> H3PO4 (aq) + 5HCl (g)
both pH 2
Why is NaCl neutral but MgCl2 & AlCl3 give acidic sol n in water?
- Al3+ (aq) sol n acidic bcos high charge density, further polarise, weaken O-H bond, so appreciable hydrolysis of Al3+ (aq) release H+ ions ==> [Al(H2O)6]3+ (aq) ⇌ [Al(H2O)5(OH)]2+ (aq) + H+ (aq)
- Mg2+ (aq) sol n oso slightly acidic as hydrolysis of Mg2+ (aq) occurs to a much smaller extent due to lower charge density of Mg2+ ion
- Na+ (aq) ions do ont undergo hydrolysis at all due its low charge density
Why chlorides like SiCl4, PCl5 undergo hydrolysis in water? Why not CCl4 ?
- hydrolysis essentially lewis acid-base rxn. Si & P in chlorides hv vacant, energetically accessible 3d orbitals, accept lp of e- fr H2O molecules to undergo hydrolysis
- Carbon, being in period 2, no hv vacant energetically accessible 3d orbitals to accept lp of e- fr H2O to undergo hydrolysis
How does atomic/ionic radius change down a group?
trend: increase
down grp,
- no of proton increase, so nuclear charge increase
- but, no of valence e- shell & shielding effect oso increase
- valence e- experience weaker e-static force attract n to nucleus
- so, valence e- further fr nucleus
How does 1st IE change down group?
trend: decrease
down grp
- no of proton increase, so nuclear charge increase
- no of e- shell & shielding effect oso increase
- so valence e- experience weaker e-static force attract n to nucleus
- Greater ease in lose e-/less energy needed to remove valence e-
How does electronegativity change down group?
trend: decrease
down grp
- no of proton increase, so nuclear charge increase
- no of e- shell & shielding effect oso increase
- eff nuclear charge decrease, so weaker e-static attract n btw nucleus & shared e-
- lesser ease in gaining e-
How does melting point and boiling point change down group 2 and 17 respectively?
- Group 2
trend: mp decrease
bcos
- charge density of cations decrease as their size increase
- strength of metallic bond decrease
- less energy needed to overcome e-static attract n btw cations & sea of delocalised e- - Group 17
trend: bp increase down grp (volatility decrease)
- Cl2, Br2, I2 hv simple molecular structures
- size e- cloud: I2>Br2>Cl2
- extent distort n e- cloud: I2>Br2>Cl2
- extent weak intermolecular id-id attract n: I2>Br2>Cl2
- energy required to overcome weak intermol id-id: I2>Br2>Cl2
How do reactivity of group 2 elements as reducing agents change down the group?
trend: become more reactive (more easily oxidised, btr reducing agent, stronger reducing pwr)
down grp
- no of proton increase, so nuclear charge increase
- no of e- shell & shielding effect oso increase
- so, valence e- experience weaker e-static attract n to nucleus
- group 2 atoms lose valence e- form M2+ cations more easily/readily, more reactive as reducing agent
How does thermal stability of carbonates (or nitrates, hydroxides) change down group 2?
observed trend: increase, decomposit n temp increase, ease of thermal decomposit n decrease
Down grp 2
- charge density & polarising pwr of cation decrease as ionic charge remain same while cation (NOT ATOM) size increase
- less distort n of anion e- cloud
- this weaken covalent bond within anion to smaller extent
- more energy needed to decompose CO3 2-
=> more thermally stable down grp
*similar expl n for stability of grp 2 nitrates & hydroxides
Generally, are grp 2 carbonates more thermally stable or less than grp 1 carbonates in same period? Why?
less stable
bcos grp 2 cations hv higher charge density, polarising pwr than grp 1’s
- grp 2 cations distort anion e- cloud more & weaken covalent bond within carbonate anion to a greater extent.
- so, less energy needed to decompose grp 2 carbonates
What are general equations for decomposition of group 2 carbonates, nitrates and hydroxides?
carbonate: MCO3 (s) –> MO (s) + CO2 (g)
nitrates: M(NO3)2 (s) –> MO (s) + 2NO2 (g) + 0.5O2 (g)
hydroxides: M(OH)2 (s) –> MO (s) + H2O (l)
How do reactivity of group 17 elements as oxidising agents change down group?
trend: bcome less reactive (E values bcome less +ve, less easily reduced, weaker oxidising agent, oxidising pwr)
down grp,
- no of protons increase, so nuclear charge increase
- no of e- shell & shielding effect oso increase
- so, additional e- (in valence shell) experience weaker e-static attract n to nucleus
- tendency to accept e- decrease
What are some group 17 element reactions with sodium thiosulfate? What does it show?
- Cl, Br oxidise thiosulfate (S2O3 2-) to sulfate (SO4 2-)
eg 4Cl2 + S2O3 2- + 5H2O –> 2SO4 2- + 10H+ + 8Cl- (OS of S increase +2 to +6) - however, iodine oxidise thiosulfate to tetrathionate (S4O6 2-)
eg I2 + 2S2O3 2- –> S4O6 2- + 2I- (OS of S increase +2 to oni +2.5, *2.5 is avg btw 2 OS of S present)
This shows that, I2 increase OS of S to smaller extent than Cl2 & Br2, so I2 is weaker oxidising agent
Also, from relationship X2 + 2e- ⇌ 2X-, since oxidise pwr is Cl2>Br2>I2, reduce pwr is Cl-<Br-<I- (Opposite)
this trend can be seen in following observation:
- HCl no reduce conc H2SO4
- HBr reduce S fr OS +6 in conc H2SO4 to OS +4 in SO2
2HBr + H2SO4 –> Br2 + SO2 + 2H2O
- HI reduce S fr OS +6 in conc H2SO4 to OS -2 in H2S
8HI + H2SO4 –> 4I2 + H2S + 4H2O
How do ease of thermal decomposition of group 17 hydrides change down group?
trend: increase (thermal stability: HF>HCl>HBr>HI)
down grp
- H-X bond strength decrease
- less energy needed to break H-X bond to decompose HX
How does successive ionisation energy vary generally (for same element)?
Increase
Bcos
- increase amt energy needed remove successive e- fr increasingly +ve ion
- due to increasing strength of e-static attract n btw nucleus & valence e-
