Module 3 Flashcards
How are elements in periodic table arranged
By increasing atom/proton number
What is periodicity
Repeating trends in physical and chemical properties
Why do elements of dame group have similar chemical properties
Same number of electrons in outer shell/same electron config
What is first ionisation energy
The energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions
Equation for sodiums first ionisation energy
Na (g) —> Na+ (g) + e-
Finish sentence
Ionisation energy is affected by how strongly……. are attracted to ….
Stronger the attraction …. energy required to remove electron so ionisation energy is ….
Electrons
Nucleus
More
High
3 factors affecting attraction between nucleus and electrons
Brief explain why
Nuclear charge: more protons means stronger pull
Atomic radius: distance between outer shell electron and nucleus affects level of attraction
Electron shielding: nucleus and electron attraction partially blocked by inner shell electrons
Why are there large increases in ionisation energy as you jump to different shell
(In same element subsequent further IE)
Big reduction in atomic radius and shielding
Why does first ionisation energy decrease down the group
Increase in nuclear charge is more than cancelled out by increase in atomic radii and shielding which decreases nuclear attraction between nucleus and outer shell electrons so less energy required to remove electron
General trend in 1st ionisation energy across a period
Nuclear charge increases which causes atomic radius to decrease as stronger pull and nuclear attraction between nucleus and outer shell electron increased and require more energy to remove electron
Shielding is same as same shell across period
Explain slight decrease in 1st ionisation energy across period (between some elements)
Ionisation that results in full or half full sets of orbitals require slightly less energy
Explain metallic bonding
Giant lattice structure of close packed metal cations in a sea of delocalised electrons with electrostatic forces
Why are metals insoluble in water
Why do metals conduct electricity in both states
More likely to simply react rather than break apart by solvents
Delocalised electrons act as mobile charge carriers
Explain giant covalent structure
Giant lattice of atoms held by covalent bonds which are strong electrostatic attraction between nuclei and shared/bonded electrons
Why are giant covalent structures insoluble in water
Conduct electricity??
Why does it have high mp/bp
Bonds too strong to pull apart
Electrons localised and no charged particles so cant carry charge
Strong bonds require more energy to break
Explain simple molecular structures
Small groups of covalent bonded atoms with weak intermolecular forces so easy to melt
What are the giant covalent elements in period 2 and 3
B
C
Si
Why does mp drop from group 4 to 5
Why does mo increase from grp 1 to 3
Change from giant to simple structures
Nuclear charge increases so stronger electrostatic attraction in metallic bonding
What is a reducing agent and example
Example of oxidising agent
Reduces other element but itself is oxidised
Eg metals
Eg halogens
Why do trends in reactivity increase down grp 2
1st and 2nd ionisation energy decreases down grp so more reactive
Product of grp 2 oxides and water
Unsaturated: metal and hydroxide ions in solution
Saturated: metal hydroxide solid
Why does solubility of grp 2 increase down the group
Charge is same but mass increases so charge/mass ratio decreases so less electrostatic attraction and easier to pull apart
Approx pH comparison of Mg(OH)2 and Ba(OH)2
Mg is 10
Ba is 13
Ba is more alkaline due to more OH- released in solution
Grp 2 metal and acid product
Salt and hydrogen
