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periodic table things Flashcards

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1
Q

How does atomic radius vary across period 3?

A
  • As the number of protons increase across the period, the nuclear charge increases
  • Successive e- are added to the same outermost electron shell and thus shielding effect remains approximately the same
  • Thus, the effective nuclear charge increases and atomic radius decreases across the period from Na to Cl

(Argon not included in trrend because it cannot be properly measured)

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2
Q

When should you use ‘outermost electron further/closer to nucleus thus electrostatic attraction between nucleus and outermost electron decrease/increase’ and ‘effective nuclear charge’ to explain trends?

A
  • When there is a difference in number of electron shells, compare distance between outermost electron and nucleus to comment on strength of electrostatic forces of attraction/ -> energy required to remove valence e-
  • ***BUT you still need to mention how the nuclear charge and shielding effect changes
  • Effective nuclear charge is mostly for when there are the same number of electron shells.
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3
Q

Are cations bigger/smaller than their respective atoms? Explain Why.

A
  • Cations are smaller than their respective atoms
  • Cations have one less electron shell than the neutral atoms
  • The shielding effect is weaker and nuclear charge remains the same
  • The outermost e- are closer to nucleus
  • Electrostatic attraction between valence electron and nucleus increases
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4
Q

Are anions bigger/smaller than their respective atoms? Explain Why. (Grp 17 and period 3)

A
  • Anions are bigger than their respective atoms due to an *INCREASE IN REPULSION DUE TO ADDED e-
  • Since nuclear charge remains the same (no mention of shielding effect*)
  • The electrostatic attraction between the nucleus and outermost e- decreases
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5
Q

How do you compare the sizes of anions and cations? (Period 3)

A
  • Anions are bigger than cations because they contain one more electron shell.
  • There is a significant increase in shielding effect and nuclear charge increases too due to more protions in anions than in cations.
  • *Outermost e- are further away from nucleus in anions than in cations
  • Hence the outermost e- are less attracted to nucleus in anions than cations
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6
Q

What is the trend of ionic radii across period 3?

A
  • From Na+ to Si4+ and P3- to Cl-, the nuclear charge increases as there is an increase in number of protons
  • The shielding effect remains approximately the same as the number of e- in each ion is the same
  • Hence the effective nuclear charge increases across period 3 and the ionic radii decreases (within their own block)
  • There is a sharp increase from cation to anion because anions contain one more electron shell than cations
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7
Q

How does first ionisation energy vary across period 3?

A
  • As number of protons increases, the nuclear charge increases
  • Successive e- are added to the same outermost shell thus shielding effect remain approximately the same
  • Effective nuclear charge increases
  • First ionisation energy generally increases across period 3

Exceptions:

  • Al has smaller 1st i.e. than Mg due to increased shielding provided by filled 3s subshell (less energy required to remove 3p electron)
  • S has smaller 1st i.e than P because less energy is required to removed paired 3px electron from S due to inter electron repulsion (shown all the 3p orbitals)
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8
Q

What is the electronegativity trend across period 3 like?

A

Electronegativity is the tendency to attract a bonding pair of electrons

  • As number of protons increase in the atoms across period 3, the nuclear charge increases
  • As successive e- are added to the same outermost shell, shielding effect reamins approximately the same
  • Effective nuclear charge increases across period
  • Electronegativity increases across period 3 (due to stronger attraction)
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9
Q

How does the melting point of metal elements (Na, Mg, Al) vary across period 3?

A

Melting point is due to the forces of attraction between the elements
- The elements possess giant metallic structure where strog metallic bonds exists between cations (Na+, Mg2+, Al3+) and sea of delocalised e-
- Melting points of metals are generally higher than other elements
- No. of valence e- used in metallic bonding increases from 1 to 3 (strength of bonding increases with the number of valence e- involved)
- Smaller radius -> stronger attraction and bond
∴ Na

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10
Q

Explain the melting point of Si. How does it compare to that of the other elements in period 3?

A
  • Si has giant molecular structure where strong covalent bonds exist between Si atoms in a 3D network
  • Large amounts of energy needed to overcome strong covalent bonds
  • ∴ Has the highest melting point out of all the elements
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11
Q

How does the melting point of non metal elements (P4, S6, Cl2, Ar) vary across period 3?

A
  • Melting point is due to intermolecular forces of attraction between MOLECULES
  • Elements possess simple molecular structure consisting of P4, S8 and Cl2 molecules or Ar
    atoms* and has weak id-id forces between molecules
  • Small amount of energy required to overcome these forces ∴ low MP
  • the more the number of e- in the molecule the larger the electron cloud = more easily polarised ∴ stronger id-id
  • Hence according to number of e- the MP varies accordingly: S6>P4>Cl2>Ar
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12
Q

What is the electrical conductivity trend of the elements in period 3?

A

Na, Mg, Al:

  • GIant metalic structure where cations (Na+, Mg2+, Al3+) are surrounded by delocalised, mobile elctrons which are able to act as charge carriers.
  • No. of valence e- used for bonding increases from 1 to 3 ∴ condcutivity increases

Si:

  • Metalloid semi conductor which posses properties of metals and non metals
  • Low electrical conductivity usually but increases with higher temp. and when mixed with other elements

P4, S6, Cl2, Ar:

  • Simple molecular structure where valence e- are localised within covalent bond ∴ no charge carriers
  • Complete insulators
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13
Q

How does volatility of elements in Group 17 differ down the group? due to id-id interactions

A

Volatitility of elements is how easily it becomes vapour has to do with MP/BP ∴ discuss the intermolecular forces of attraction!!

  • Halogens have simple molecular structure with weak id-id forces of attraction between the molecules
  • Down the group, the no. of electrons per molecule increase -> e- cloud becomes more polarisable due to larger size
  • ∴ the id-id becomes stronger and more energy is required to overcome them -> MP/BP increases and volatility decreases!
  • Physical state changes from gas->liquid-> solid (@rtp Cl2 gas, Br2 liquid, I2 solid)
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14
Q

How does atomic/ionic radius vary down group 17?

A
  • Atomic radius increases down the group
  • Nuclear charge increases as more protons are added to the nucleus, shielding effect increases significanty as number of e- shells increase
  • Outermost electron is further away from nucleus and thus electrostatic attraction between them decreases
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15
Q

How does first ionisation energy change down group 17?

A
  • First ionisation energy decreases down the group
  • The nuclear charge increases as protons are added to the nucleus and shielding effect increases significantly as the number of e- shells increase
  • Effective nuclear charge decreases
  • The distance between outermost e- and nucleus is large ∴ the electrostatic attraction between them decreases
  • Less energy is required to remove the outermost e-
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16
Q

What is the trend of electronegativity in group 17?

A

Electronegativity is the tendency to attract a bonding pair of electrons

  • It decreases down the group
  • Down the group, nuclear charge increases as number of protons added to nucleus increases
  • The shielding effect increases significantly as the number of e- shells increase
  • Effective nuclear charge decreases, less attraction between outermost e- and nucleus, less easy to attract e-
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17
Q

What is the reactivity of elements down Group 1 like?

A

Group 1 elements are reducing agents (donate e-) and tend to be oxidised themselves

  • Down the group, the nuclear charge increases as there are more protons. Shielding effect increases significantly as there are more e- shells
  • The outermost e- is further away from the nucleus and has a weaker electrostatic attracton.
  • Less energy is need to remove the outermost e- as IONISATION ENERGY decreases down the group
  • ∴ INCREASE EASE OF LOSING VALENCE E- > Stronger reducing power -> reactivity of the elements increase down the group
18
Q

What is the reactivity of elements down Group 17 like?

A

Group 17 elements are oxidising agents (accept e-) and tend to be reduced themselves

  • Down the group, the nuclear charge increases as there are more protons. Shielding effect increases significantly as there are more e- shells.
  • ADDITIONAL e- are added to outermost electron shell which is increasingly further from nucleus down the group
  • Less attraction between outermost e- and nucleus ∴ ease of gaining e- decreases
  • Hence oxidising power decreases down the group -> less reactive
19
Q

Explain the thermal stability of Group 17 hydrides

A

Halides are HX compounds. Upon heating they decompose into H2 gases and X2 gases.

  • The stronger the H-X covalent bond, the higher the thermal stability (more difficult to decompose)
  • Quote H-X bond energy to explain the the decreasing thermal stability
  • A longer H-X bond is weaker, hence as the atomic radius of the halogen atoms increases, the H-X bonds becomes weaker
  • Thermal stability decreases down the group as halogen atom increases in size
20
Q

State and explain the variation of oxidation number of the oxides and chlorides of period 3 elements.

A
  • Oxidation state depends on the number of valence e- used for bonding.
  • Oxidation number of the element is a positive value because O and Cl are both more electronegative than the element.
  • Oxidation number increases from +1 to +6 down period 3
  • Maximum oxidation number corresponds with the number of outermost electrons = Group number (syllabus only focuses on max possible)
  • P and S show multiple O.N because they can expand octet by excitation of outer e- into vacant and energetically accessible 3d orbitals
21
Q

State and explain the variation in bonding in Period 3 element oxides and chlorides in terms of electronegativity

A
  • Electronegativity of elements increase (Increase to the right)
  • Electronegativity difference between elements and O decrease across period 3
  • Large electronegativity difference -> ionic bond bc electrons are completely transferred (no sharing)
  • Therefore bonding changes from ionic to covalent and the oxide structures change from giant ionic to giant molecular and simple molecular structure
22
Q

Explain the variation in melting point of the period 3 oxides / Suggest the types of structure and bonding present in period 3 oxides based on their chemical and physical properties

A

Na2O, MgO, Al2O3: (more exothermic = higher lattice energy)

  • Giant ionic crystal lattice structure where strong ionic bonds exist between oppositely charged (Na+,Mg2+,Al3+) Cations and (O2-) anions
  • Large amounts of energy is needed to overcome the strong ionic bonds thus high MP
  • MgO has higher MP than Na2O since lattice energy of MgO is more exo. as Mg has greater charge and smaller ionic radius than Na=
  • Al2O3 has lower MP than MgO bc its lattice energy i less exothermic since it has some covalent character due to high charge density of Al3+ ion which polarises O2- ion easily

SiO2:

  • Giant molecular structure where strong covalent bonds exist between Si and O atoms in a rigid 3D network
  • Large amounts of energy required to overcome strong covalent bonds

P4O10 and SO3:

  • Simple molecular structure where weak id-id exists between molecules
  • Small amount of energy required to overcome weak bonds
23
Q

Explain the acid/base behaviour of the oxides

A
  • Nature of oxide depends on electronegativity of element in period 3
  • Metals with low electronegativity form ionic oxides (large electronegativity difference) -> basic
  • Non metals form covalent oxides (small electronegativity difference) -> acidic
  • Al3+ has high charge density and polarised electron cloud of O2- ion thus forming ionic oxide with covalent character -> amphoteric !
24
Q

Explain the variation in melting point of the period 3 chlorides / Suggest the types of structure and bonding present in period 3 chlorides based on their chemical and physical properties

A
  • As bonding changes from giant ionic to covalent bonding, the structure changes from ionic to simple molecular structure-> MP/BP decreases
  • AlCl3 is a covalent molecule as Al3+ ion has high charge density and polarises Cl- ion
    (solid at room temp, dimerises when sublime)
  • PCl5 is ionic in solid state
25
Q

Na2O reaction with water

A 
# Na2O + H2O -> 2NaOH : dissolves readily to form pH 13 solution
- Hydration energy evolved from ion dipole interactions between Na+ ions with water and O2- ions with water is able to overcome ionic bonds between Na+ and O2- ions
26
Q

Describe the acid/base behaviour of the period 3 metal hydroxides

A

NaOH: Basic

Mg(OH)2: Basic

Al(OH)3: Amphoteric

  • OH reacts as base, proton acceptor : Al(OH)3 + 3H+ -> Al3+ + 3H2O
  • reacts as acid, Al(OH)3 + OH- -> Al(OH)4-
27
Q

MgO reaction with water

A 
# MgO + H2O -> Mg(OH)2 : dissolves slightly to form alkaline solution of pH 8
- Hydration energy evolved from the interactions between Mg2+ ions with water and O2- ions with water is sufficient to overcome ionic bonds between ionic bonds between Mg2+ and O2-
28
Q

Al2O3 reaction with water

A 
# Insoluble thus pH 7 solution
- Hydration energy evolved from the interactions between Al3+ ions with water and O2- ions with water is insufficient to overcome strong ionic bonds between Al3+ and O2- ions
29
Q

SiO2 reaction with water

A 
# Insoluble (Giant molecular structure)
- Energy evolved from from forming weak id-id interactions between water and SiO2 molecules is not enough to break strong covalent bonds between Si and O atoms
30
Q

P4O10 reaction with water

A 
# P4O10 + 6H2O -> 4H3PO4 : dissolves to form pH 2 acidic solution
- Interactions w water molecules evolve enough energy to break weak id-id interactions between P4O10 molecules which have simple molecular structure
31
Q

SO3 reaction with water

A 
# SO3 + H2O -> HSO4; dissolves to form pH 1 acidic solution!
- Interactions with water molecules evolve enough energy to break weak id-id interactions between SO3 molecules which have simple molecular structure
32
Q

Al2O3 acid/base behaviour

A 
# as a base: Al2O3 + 6HCl -> 2AlCl3 + 3H2O
# as an acid: Al2O3 + 2NaOH + H2O -> 2NaAl(OH)4 Sodium aluminate *q similar to base rxn of Al(OH)3
33
Q

SiO2 acid behaviour

A 
Only reacts with CONC. NaOH with heating 
# SiO2 + 2NaOH -> NaSiO3 + H2O
34
Q

P4O10 acid behaviour

A

P4O10 + 12NaOH -> H2O + 4Na3PO4 + 6H2O

** PO4 3-

35
Q

SO3 acid behaviour

A

SO3 + 2NaOH -> Na2SO4 + H2O

36
Q

H3PO4 acid behaviour

A 
a hydroxide product of P4O10 rxn with water
# H3PO4 + 3NaOH -> Na3PO4 + 3H2O
37
Q

H2SO4 acid behavior

A

a hydroxide product of SO3 rxn with water

38
Q

NaCl reaction with water

A 
Readily dissolves form hydrated ions to form neutral pH 7 solution 
# NaCl + H2O ->  Na+ + Cl-
- Na+ has low charge density hence it does not undergo hydrolysis (form new bonds with water)
39
Q

MgCl reaction with water

A 
Readily dissolves to form hydrated Mg2+ ions which undergo slight hydrolysis to form solution of pH 6.5
# MgCl2 + 6H2O -> [Mg(H2O)6]2+ + 2Cl- 
# to get the acid part: [Mg(H2O)6]2+ + H2O <=> [Mg(H2O)5(OH)]+ H3O+
- Mg2+ has a relatively high charge density hence it allows for slight hydrolysis
40
Q

AlCl3/Al2Cl6 reaction with water

A 
(i) Readily dissolves to form hydrated Al3+ ion which undergoes hydrolysis in largeeeee amounts of water to get solution of pH3
# AlCl3 + 6H2O -> [Al(H2O)6}3+ + 3Cl-
# [Al(H2O)6}3+ + H2O -> [Al(H2O)5(OH)]2+ + H3O+ 
- Al3+ has high charge density and is able to polarise water molecules resulting in breakage of O-H bonds to form acidic H3O+ ion
(ii) White solid and white fumes of HCl formed*** when in limited amount of water 
# AlCl3 + H2O -> Al(OH)3 + 3HCl
41
Q

SiCl4 reaction with water

A 
Undergoes complete hydrolysis to give strong acidic solution of pH 2 with white fumes of HCl
# SiCl4 + 2H2O -> 4HCl +  SiO2
- When it dissolves, the slightly positive Si atom attracts the slightly negative O atom from water and breaks O-H bond, releasing H+ to make acidic solution
42
Q

PCl5 reaction with water

A 
Undergoes complete hydrolysis to give strong acidic solution of pH 1-2 with white fumes of HCl
# PCl5 + 4H2O -> H3PO4 + 5HCl
# Limited water/ cold water: PCl5 + H2O -> POCl3 + 2HCl
- When it dissolves, the slightly positive P atom attracts the slightly negative O atom from water and breaks O-H bond and releasing H+ to make acidic solution

Chem (8 decks)

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