Periodic Table Flashcards
Explain the trend of atomic radius across the period.
Across the period, nuclear charge increases as number of protons increase. Successive electrons are added to the same shell hence increase in shielding effect is negligible. Thus effective nuclear charge increases across the period and valence electrons are more strongly bonded to nucleus. Hence atomic radius decreases across the period.
Explain the trend of atomic radius down the group.
Down the group, number of protons increase hence nuclear charge increases. The number of electrons increases and number of electron shells increase, leading to an increase in shielding effect. Thus effective nuclear charge decreases and valence electrons are less strongly bonded to the nucleus. Hence atomic radius increases down the group.
Compare the cation in radius to its corresponding parent atom.
Cation and atom has same number of protons hence similar nuclear charge. However the cation has one lesser electron shell than the atom and experiences weaker shielding effect. Thus valence electrons of cation are more strongly attracted to nucleus and has a smaller cationic radius
Compare the radius of anion and corresponding parent atom
Both anion and atom has similar nuclear charge as they have the same number of protons. However anion has more electrons than atom and has stronger shielding effect. Thus effective nuclear charge in anion is weaker and valence electrons are less tightly attracted to the nucleus. Hence anion has larger anionic radius than parent
Explain the trend of ionic radius across the period
Across the period, number of protons increases hence nuclear charge increases. They have similar shielding effect as ions are isoelectronic. Effective nuclear charge increases and valence electrons are more strongly attracted to the nucleus thus ionic radius decreases across period
Explain the trend of first ionisation energy across the period.
Across the period, number of protons increase hence nuclear charge increases. Increase in shielding effect is negligible as successive electrons are added to the same electron shell. Thus effective nuclear charge increases, and valence electrons are more strongly attracted to the nucleus. More energy is required to remove the first valence electrons hence first ionisation energy increases across the period.
Irregularities in trends for first ionisation energy
Mg>Al and P>S in period 3
Be>B and N>O in period 2
Explain trend of first ionisation energy down the group
Down the group, number of protons increases hence nuclear charge increases. Number of electrons and electron shells increase hence shielding effect increases. Thus overall effective nuclear charge decreases and valence electrons are less strongly attracted to the nucleus. Hence energy required to remove valence electrons decreases and first ionisation energy decreases
Trend of electronegativity across the period
Same explanation for why atomic radius decreases across the period and hence ability to attract bonding electrons to itself increases.
Trend of electronegativity down the group
Same explanation for why atomic radius down the group. Hence ability to attract bonding electrons to itself decreases
Definite electronegativity
Electronegativity is the ability of an atom in a molecule to attract bonding electrons to itself
Draw the graph of melting points across period 3 atoms and explain
Na<Mg<Al: Na, Mg, Al have giant metallic structure with strong metallic bonds between positively charged cations and a sea of delocalised negatively charged electrons. From Na-Mg-Al, cationic radius decreases and charge increases. Hence larger sea of delocalised electrons, and metallic bond gets stronger. Hence increase in mp.
Si: Si has a giant molecular structure consisting of atoms held by strong covalent bonds. Large energy is required to…
P4<S8<Cl2: Are simple molecules held by weak idid attractions between molecules. S8 has most number of electrons hence most polarisable electron cloud, strongest idid, highest mp
Draw the graph of electrical conductivity for period 3 elements and explain
Na<Mg<Al: Metals have delocalised electrons. Charge increases, sea of delocalised electrons increases.
Si: Semi conductor: low electrical conductivity
Others: no delocalised electrons cannot conduct electricity
Across period 3 oxides, the oxides become less ionic and more covalent. Why
Across the period, electronegativity increases hence the electronegativity diff between element and O decreases. Thus bond is less polar and have more covalent character
Draw the trend for MP of period 3 oxides and explain
Na2O<MgO: Sodium oxide and magnesium oxide have giant ionic lattice structure with ionic bonds between oppositely charged ions () and (). Compare ionic bond strength using LE.
MgO>Al2O3: High charge density of Al3+ polarises the electron cloud of oxygen hence Al2O3 has covalent characteristics, which weakens the ionic bonds present between oppositely charged ions. Thus lesser energy required…
SiO2>P4O10 and SO3: SiO2 giant molecular structure, strongest covalent bonds, most energy required.
P4O10>SO3: More electrons, more polarisable electron cloud, stronger idid
Explain the reactions of period 3 oxides with water with equations and state the pH.
Na2O: Reacts vigorously with strongly alkaline solution
MgO: Slow reaction. MgO insoluble in water, forms weakly alkaline solution
Al2O3: Does not react due to highly exothermic lattice energy
SiO2: Does not react due to strong covalent bonds between Si and O atoms in giant molecular structure
P4O10: Reacts vigorously to form colourless, acidic solution
SO3: Reacts vigorously and exothermically to form strongly acidic solution
Explain the nature of oxides across period 3 with equations
Na2O,MgO: basic
Al2O3: Amphoteric
SiO2,P4O10,SO3: acidic
Draw the trend for MP of period 3 chlorides and explain
NaCl>MgCl2: Mg has larger ionic radius, compare LE
AlCl3: ionic+covalent
SiCl4<PCl5: PCl5 got more electrons, more polarisable electron cloud
Explain period 3 chlorides reaction with water using equations and state the pH
NaCl: 7, dissolves in water to form colourless, neutral solution. Na+ and Cl- forms ion-dipole interaction with water molecules. Na+ and Cl- have low charge density and does not undergo hydrolysis.
MgCl2: 6.5, dissolves in water to form colourless, slightly acidic solution. Mg2+ has higher charge density than Na+ and undergoes partial hydrolysis. The high charge density of Mg2+ polarises the electron cloud of water, weakening the O-H bond, releasing small amount of H+ ions, causing the solution to be acidic
AlCl3: 3, dissolves in water to form colourless, acidic solution. Al3+ has higher charge density than Mg2+ and undergoes a larger extent of hydrolysis, thus releasing more H+ ions, causing solution to be more acidic
SiCl4: 2, reacts vigorously with water. Undergoes full hydrolysis in water to give a white solid and a colourless, strongly acidic solution. Complete hydrolysis can take place due to presence of empty and energetically accessible 3d orbitals in Si which can be attacked by lone pair of electrons on oxygen atom of water molecules
PCl5: 2, reacts vigorously with water. Undergoes full hydrolysis in water to give colourless, strongly acidic solution. Complete hydrolysis can take place due to presence of empty and energetically accessible 3d orbitals in P that can be attacked by lone pair of electrons from O atom in water
