periodic table things Flashcards
How does atomic radius vary across period 3?
- 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)
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?
- 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.
Are cations bigger/smaller than their respective atoms? Explain Why.
- 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
Are anions bigger/smaller than their respective atoms? Explain Why. (Grp 17 and period 3)
- 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
How do you compare the sizes of anions and cations? (Period 3)
- 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
What is the trend of ionic radii across period 3?
- 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
How does first ionisation energy vary across period 3?
- 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)
What is the electronegativity trend across period 3 like?
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)
How does the melting point of metal elements (Na, Mg, Al) vary across period 3?
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
Explain the melting point of Si. How does it compare to that of the other elements in period 3?
- 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
How does the melting point of non metal elements (P4, S6, Cl2, Ar) vary across period 3?
- 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
What is the electrical conductivity trend of the elements in period 3?
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
How does volatility of elements in Group 17 differ down the group? due to id-id interactions
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)
How does atomic/ionic radius vary down group 17?
- 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
How does first ionisation energy change down group 17?
- 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-
What is the trend of electronegativity in group 17?
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-
What is the reactivity of elements down Group 1 like?
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
What is the reactivity of elements down Group 17 like?
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
Explain the thermal stability of Group 17 hydrides
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
State and explain the variation of oxidation number of the oxides and chlorides of period 3 elements.
- 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
State and explain the variation in bonding in Period 3 element oxides and chlorides in terms of electronegativity
- 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
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
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
Explain the acid/base behaviour of the oxides
- 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 !
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
- 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
Na2O reaction with water
# 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
Describe the acid/base behaviour of the period 3 metal hydroxides
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-
MgO reaction with water
# 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-
Al2O3 reaction with water
# 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
SiO2 reaction with water
# 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
P4O10 reaction with water
# 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
SO3 reaction with water
# 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
Al2O3 acid/base behaviour
# 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
SiO2 acid behaviour
Only reacts with CONC. NaOH with heating # SiO2 + 2NaOH -> NaSiO3 + H2O
P4O10 acid behaviour
P4O10 + 12NaOH -> H2O + 4Na3PO4 + 6H2O
** PO4 3-
SO3 acid behaviour
SO3 + 2NaOH -> Na2SO4 + H2O
H3PO4 acid behaviour
a hydroxide product of P4O10 rxn with water # H3PO4 + 3NaOH -> Na3PO4 + 3H2O
H2SO4 acid behavior
a hydroxide product of SO3 rxn with water
NaCl reaction with water
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)
MgCl reaction with water
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
AlCl3/Al2Cl6 reaction with water
(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
SiCl4 reaction with water
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
PCl5 reaction with water
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
