S3.1 The periodic table: classification of elements Flashcards
what does the periodic table consist of
periods, groups, blocks
how are elements arranged
- arranged in increasing atomic number
- atoms with the same number of shells are together
- atoms with similar outer shell electron configurations are together
helium and hydrogen are unusual
so are put into a group based on similarities in physical and chemical properties
He = 0
H = group of its own
s block
only s electron in outer shell
p block
minimum 1 p electron in outer shell
d block
minimum 1 d electron and 1 s electron in outer shell, no p or f electron in outer shell
f block
minimum 1 f electron and 1 s electron in outer shell, no d or p electron in outer shell
atomic radius trend down a group
- General increase
- More shells
- Electrons in inner shells repel electrons in outer shells
- Shielding from positive nuclear charge
- Less attraction between outer shell electron and nucleus → larger atoms
atomic radius trend across a period
- General decrease
- Increased atomic number, increased nuclear charge, more electrons added to same principal quantum number shell
- Larger nuclear charge, greater nuclear attraction between outer shell electron and nucleus → smaller atoms
when does the ionic radius increase
with increasing negative charge (more repulsion, more electrons while constant nuclear charge)
when does the ionic radius decrease
with increasing positive charge (greater electrostatic force of attraction, less electrons while constant nuclear charge)
electronegativity
ability of atom to attract a pair of electrons towards itself in a covalent bond
electronegativity trends across a period
- Increase
- Greater nuclear charge, same shielding
- Nucleus has increasingly strong attraction to electrons → smaller atomic radii
electronegativity trends down a group
- Decrease
- Nuclear charge increases, increased shielding, larger atomic radii
- Decrease in attraction between nucleus and outer electron, decreased effective nuclear charge
electron affinity trends
Don’t include noble gases as they don’t form negative ions
Most exothermic electron aff = group 17
Strongest pull on electron is correlated with greater energy released when -ve ions form
electron affinity trend down a group
- general decrease
- As atoms become larger, the attraction for an additional electron is less, the effective nuclear charge reduces due to shielding
- Becomes less exothermic (fluorine is an exception)
- fluorine has a very small electron affinity, it is such a small atom, additional electron in 2p subshell experiences repulsion with other valence electrons
group 1: alkali metals
- Form high pH alkaline solutions when reacted with water
- End in ns1 electron configurations
alkali metals physical properties
- Soft, easy to cut, softer denser down group
- Shiny silver surface when cut
- Conduct heat and electricity
- Low MP, low densities
- MP decreases down group, increased atomic radius, weaker metallic bonding
alkali metals chemical properties
- Readily react with oxygen and water vapour in air (kept in oil to stop)
- Similar reactions with water- vigorous, produce alkaline metal hydroxide and hydrogen gas
lithium + water
2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g)
- Lithium floats, slowly reacts
- Hydrogen gas released
- Lithium keeps shape
sodium + water
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
- Vigorous hydrogen gas release
- Heat produced melts unreacted metal, forms small ball that moves on surface
- Highly alkaline solution due to NaOH
potassium + water
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
- Bubbles of hydrogen gas
- Melts into shiny ball, dashes around surface
- Hot enough to ignite hydrogen gas
- K burns with lilac flame
alkali metals with halides
- Vigorous reactions
- Product is an alkali halide salt
2Na (s) + Cl2 (g)–> 2NaCl (s) - Increasingly vigorous down group: larger atom, ns1 electron further from nucleus, is shielded, easier to lose electron
what do metallic oxides form when reacting with water
hydroxides
what do non metallic oxides form when reacting with water
oxoacids
what does the acid-base character of oxides give evidence of
chemical trends
how do oxides change across a period
basic → amphoteric → acid
amphoteric
can act as acid and base
oxides: trend across a period
Oxides become less ionic
Due to increased electronegativity
oxides: trend down a group
Oxides become more ionic
Due to decreased electronegativity
oxidation state
- a number assigned to an atom to show the number of electrons transferred in forming the bond
- it is the change the atom would have if the compound were composed of ions
what will always have a negative value
the more electronegative species
sum of oxidation numbers in a compound
= 0
sum of oxidation numbers in an ion
= charge of the ion
group 1 oxidation state
+1
group 2 oxidation state
+2
fluorine oxidation state
-1
hydrogen oxidation state
+1
hydrogen in metal hydrides oxidation state
-1
oxygen oxidation state
-2
oxygen in peroxides oxidation state
-1
oxygen in F2O oxidation state
+2
uncombined element oxidation state
0
naming transition metal compounds
- Stock notation (roman numerals) if variable oxidation number
E.g. iron (+2, +3), iron (II) oxide - generic names persist and are acceptable
generic names: nitrate
NO3-
generic names: nitrite
NO2-
generic names: sulfate
SO4 2-
generic names: sulfite
SO3 2-
