chapter 8 periodicity

order of the periodic table

increasing atomic number

structure of the periodic table

• alkali metals

• alkaline earth metals

• transition metals

• below the staircase metals

• in the staircase metalloids

• above the staircase non metals

• noble gases

the staircase on the periodic table

left = non metals

right= metals

on the stairs metalloids

what are metalloids

elements that have a combination of metallic and non metallic properties

properties of metalloids

• shiny

• conducts electricity

• e.g silicon

history of the periodic table

dimitri mendeleev left gaps in the periodic table for undiscovered elements. newly discovered element fit the gaps he left years ago and he accurately predicted their properties

blocks of the periodic table

s,d,p and f

s orbital block

elements in the s orbital have their highest energy electrons in the s orbital e.g sodium

d block

elements in the d block have their highest energy electrons in the d orbital e.g iron

p orbital block

all elements in the p block have their highest energy electrons in the p orbital e.g carbon

characteristics of transtition metals

form compounds where they have partly filled orbitals

exceptions to the rule of transition eleements and why

scandium and zinc are not transition metals because they do not forma any compounds where their d orbitals are partly filled

origins or the terms s,d,p,f

when elements are heated they give out light energy at certain wavelengths. thi makes line appear in the spectrum of light.

s,d,p,f were words used to describe the lines

s= sharp

d=diffuse

p=principle

f= fine

groups

verticle column of elements

simalar properties

same number of electrons in outer main shells

reactivity

metals get more reactive going down

non metals get less reactive going down

d orbital block elements are unreactive

why helium and hydrogen are atypical

helium and hydrogen

helium;is not a p block element but is placed with noble gases

hydrogen; forms +1 ions like group 1 metals but does not react similarly because its not a metal so its placed on its own

periodicity

the repetition of properties of elements in each period

properties of elements in period 3 in group 1,2,3

Na, Mg, Al form ionic compounds and have gaint stuctures

properties of elements in period 3 group 4

silicon 4 electrons in the outer shell and forms bonds with 4 other silicon atoms. so it has a gaint covalent structure

silicon is a metalloid so it has both metallic and non metallic properties

properties of period 3 elemnts in group 5,6,7

P,S,Cl get reduced to form ionic compounds or share electrons to form covalent compounds

general trends across period 3

• melting poin increases until group 5 where is dcreases

• shielding stays the same

• nuclear charge increases

trends in melting and boiling points of period 3 group 1 to group 4 elements

• increases because they have metallic structures

• high melting and boiling points

• except silicon which is a gaint covalent structure

trends in melting points period 3 group 5 to group7 elements

• low melting points

• simple/molecular structures

why does the melting point increase from sodium to aluminium

• strength of metallic bond increases

• charge on the ion increases

• more delocalised electrons

why do the non metals in period 3 have low boiling points

• molecular structure

• smaller van der waals forces

• more electrons so they cant pack as closely together

why is silicon the exception

gaint structure with a high melting point

so stronger van der waals forces

there are more electrons

and they can pack closely together because 1 Si atom bonds to 4 others

define atomic radii

half the distance between the centres of a pair of atoms used

why is the term atomic radius used

there is no clear point at which the electron cloud density drops to 0

how are the sizes of atoms measured

using their atomic radii

what does the atomic radius depend on?

• the type of bond formed

• e.g covalent, ionic, metallic, and the strength of van der waals

why are noble gases left out of comparisons of atomic sizes

noble gases do not bond covalently with one another because they have full outer main shells

general trend in atomic radius

• periodic property

• atoms get larger down any group

• decreases across each period

• there is a jump when starting the next period

general trend in atomic radius across a period

the size atoms decreases across a period

why does the atomic radius decrease across a period

• there are more protons

• increased nuclear charge

• increased pull towards the nucleus

• same shielding

why does the atomic radius increase down a group

• more shielding

• greater distance between the valence electrons and the nucleus

define first ionisation energy

energy required to convert 1 mol of gaseous atoms of an element into 1 mole of gaseous cations

general trend in first ionisation energy of the first 20 elements across a periods

increases

the discovery of argon

Willian ramsey discovered argon

and the whole group of noble gases

general trend in first ionisation energy of the first 20 elements down a group

decreases

why does first ionisation energy increase across a period

• more protons

• same shielding

• gets harder to lose one electron

why does the first ionisation energy decrease going down

• more shielding going down

• greater distance between the nucleus and valence electron

• easier to remove valence shell electron

why is there a drop in ionisation energy from 1 period to the next

• new main shell

• bigger atomic radius

• weaker attraction of valence shell electron to the nucleus

exception between group 2 and 3 first ionisation energy

first ionisation energy decreases from 2 to 3 because the valence electron(3p) is removed from a higher energy level than in group 2 so magnesium has a higher first ionisation energy than aluminium.

exception between group 5 and 6 first ionisation energy

group 5 has spin pair repulsion group 6 does not

successive ionisation energy

• Reduced electron shielding.

• Increased effective nuclear charge on the remaining electrons.

• Decreased distance of remaining electrons from the nucleus.

• Significant jumps when electrons are removed from a lower energy shell.

why does magnesium have a higher melting point than sodium

• greater nuclear charge

• smaller atoms

• more delocalised electrons

• stronger attraction between ions and delocalised electrons

phosphorus

P4

oxidation state 3 +

covalent molecular structure with weak van der waals