3.1 Periodicity Flashcards
periodicity
the repeating trends in physical and chemical properties
what change happens across each period?
elements change from metals to non-metals
how can electron configuration be written in short?
the noble gas before the element is used to abbreviate it
e.g. Li is 1s2 2s1
can also be written as Li is [He] 2s1
1st IE
the energy required to remove 1 electron from each atom in 1 mole of the gaseous element to form 1 mole of gaseous 1+ ions
1st IE og Magnesium equation example
Mg (g) -> Mg+ (g) + e-
what factors affect ionisation energy?
-atomic radius
-nuclear charge
-electron shielding
explain the changes in 1st IE across period 3
1st IE increases across period 3 because there is:
- an increase in nuclear charge
- a decrease in atomic radius
- the same electron shielding
– which means more energy is needed to remove the 1st electron
1st IE dips at Al because its outer electron is in a 3p orbital which is a higher energy level than Mg’s 3s orbital, therefore less energy is needed to remove an electron
1st IE dips at S because one 3p orbital contains 2 electrons -> repulsion between paired electrons -> less energy needed to remove one
why does 1st IE decrease between group 2 to 3?
because in group 3 the outermost electrons are in p orbitals whereas in group 2 they are in s orbital, so the electrons are easier to be removed
why does 1st IE decrease between group 5 to 6?
because in group 5 the outermost electrons in p orbital are single electrons and in group 6 the outermost electrons are spin paired with some repulsion, therefore the electrons are slightly easier to remove
does 1st IE increase or decrease between the end of one period and the start of the next? why?
it decreases because:
-atomic radius increases
-electron shielding increases
does 1st IE increase or decrease down a group? why?
it decreases because:
-shielding increases -> weaker attraction
-atomic radius increases -> distance between outer electrons and nucleus increases -> weaker attraction
-increase in number of protons is outweighed by an increase in atomic radius and shielding
properties of giant metallic lattices
-high melting and boiling point
-good electrical conductors
-malleable
-ductile
describe the structure, forces and bonding in every element across period 2
Li & Be
-giant metallic; strong attraction between positive ions and delocalised electrons; metallic bonding
B & C
-giant covalent; strong forces between atoms; covalent bonding
N2, O2, F2, Ne
-simple molecular; weak IMF between molecules; covalent bonding within molecules and IMF between molecules
describe the structure, forces and bonding in every element across period 3
Na, Mg, Al
-giant metallic; strong attraction between positive ions and delocalised electrons; metallic bonding
Si
-giant covalent; strong forces between atoms; covalent bonding
P4, S8, Cl2, Ar
-simple molecular; weak IMF between molecules; covalent bonding within molecules and IMF between molecules
explain why first IE show a general increase across period 3?
-atomic radius decreases
-nuclear attraction
-electron shielding stays the same
(no. of protons increases so greater nuclear charge)
explain why the general increase in first IE across period 3 is not followed for Mg to Al
-due to orbitals
-an Mg outer electron is removed from 3s orbital but Al outer electron is removed from 3p orbital
-Al (3p) has higher energy level than Mg (3s) so Al electron is easier to lose
explain how first IE decrease down groups
-nuclear attraction decreases
-atomic radius increases
-increased shielding
state and explain the trend in atomic radius from Li to F
trend-nuclear charge increases
explanation
-greater nuclear attraction to outer electrons
-outer electrons experience the same shielding
-atomic radius decreases and so nuclear charge increases
explain the difference in MP of P and Cl
-P has more electrons (high Mr) which means it has greater london forces
-more energy is needed to overcome the intermolecular forces
describe the bonding in Mg and Si
-Mg is held together by metallic bonding - electrostatic attraction between delocalised electrons and positive metal ions
-Si is held together by covalent bonds - electrostatic attraction between shared pair of electrons and nuclei of the bonded atoms
along a period…
1st IE increases due to:
-decreasing atomic radius
-decreasing electron shielding
-and greater forces of attraction
down a group…
1st IE decreases due to:
-increasing atomic radius
-increasing electrons shielding
-which reduces effect of electrostatic attraction
successive ionisation energies…
increase, because atomic radius decreases and there is greater attraction between outer electrons and nucleus
high ionisation energy
means there is a strong attraction between the electron and nucleus so more energy is needed to overcome attraction and remove the electron
G2 elements react with water to produce:
metal hydroxide + hydrogen
G2 metals react with oxygen to produce:
a solid white oxide
G2 metals react with dilute acid to produce:
salt + hydrogen
G2 oxides and hydroxides are:
bases
-and most of are soluble in water so they are alkalis
calcium hydroxide is used to:
neutral is acidic soils in agriculture
magnesium hydroxide is used in milk of magnesia to?
-neutralise excess acid, producing salt + water
calcium carbonate is used in?
building materials like limestone
halogens get less reactive as you go down the group because:
-atomic radius increases
-increased electron shielding
-makes it harder for larger atoms to attract the electron needed to form an ion (despite increased charge on nucleus) so larger atoms are less reactive
-they also become less oxidising
why are halogen atoms oxidising agents?
they react by gaining an electron in their outer shell to form 1- ions
this means they are reduce, and therefore they oxidise another substance - so they are oxidising agents
test for halide ions
- add 2 drops of dilute nitric acid to remove ions that might interfere with the test
- then add silver nitrate solution
ionic eq:
Ag+ (aq) + X- (aq) -> AgX (s)
where X = Cl, Br or I
-colour of precipitate identifies the halide
Cl-white
Br-cream
I-yellow
then to be extra sure we can add ammonia solution
Cl-dissolves in dilute NH3(aq)
Br-dissolves in conc. NH3(aq)
I-insoluble in conc. NH3(aq)
disproportion reaction
when a halogen is simultaneously oxidised and reduced
bleach equation
Cl2(g) + 2NaOH(aq) -> NaClO(aq) + H2O(l)
(sodium chlorate(I) solution=bleach)
Cl goes from 0 to +1 and gets reduced
Cl also goes from 0 to -1 and gets oxidised
adv and dis of using chlorine to sterilise water:
adv-kills bacteria
dis-Cl gas is very toxic
alternatives include ozone or UV light
