topic 3- periodicity Flashcards
what does a periodic table consist of?
groups (vertical columns) and periods (horizontal rows)
how is the periodic table organised?
by order of increasing atomic number
elements in the same group…
have the same number of valence electrons
elements in the same period…
have the same number of electron shells
location of metals
right side
location of non-metals
left side
location of metalloids
border between metals and non metals
what are metalloids (give egs)
chemical elements whose physical and chemical properties fall in between the metal and non-metal categories (boron, germanium, silicon)
alkali metals
group 1
halogens
group 17
noble gases
group 18
groups 3-12
transition metals
top row of ‘island’ at bottom
lanthanides
bottom row of ‘island’ at bottom
actinides
first 2 blocks
outer electron is in s shell
middle block
outer electron is in d shell
far right block
outer electron in p shell
bottom block
outer electron in f block
period number (n)=
outer energy level occupied by electrons
as you go down a group, atomic radius…
increases:
- more shells (increased shielding/distance)
- increase in nuclear charge outweighed
as you go across a period, atomic radius…
decreases:
- shielding stays the same; nuclear charge increases
- outer electrons experience stronger attraction to the nucleus
define ionic radius
the distance between the nucleus of an ion and the outermost shell of the ion.
as you go down a group, ionic radius…
increases:
- there is an extra complete inner shell of electrons
as you go across a period, ionic radius…
decreases:
- nuclear charge increases
- electrons feel a stronger attraction to the nucleus
define ionisation energy
The minimum energy required to remove one electron from each atom in a mole of gaseous atoms.
as you go down a group, ionisation energy…
decreases:
- shielding and distance increase
- this outweighs the increase in nuclear charge
as you go across a period, ionisation energy…
increases:
- all electrons are in the same shell
- nuclear charge increases
define electronegativity
a measure of the tendency of an atom to attract a bonding pair of electrons
as you go down a group, electronegativity…
decreases:
- shielding and distance increase
- electrons experience less attraction to the nucleus
as you go across a period, electronegativity…
increases:
- increased nuclear charge
- shielding is almost the same
define electron affinity
the enthalpy change when one mole of gaseous atoms gain an electron each, forming negative ions
as you go down a group, electron affinity…
decreases:
- nuclear charge decreases
- as a result, it becomes more difficult to add an electron - resulting in a decrease in the amount of energy released.
as you go across a period, electron affinity
increases:
- atomic radius gets smaller
- nuclear charge gets larger.
what are the similarities in the electron configurations of the alkali metals?
they all have 1 electron in their valence shell
what are the differences in the electron configurations of the alkali metals?
as you go down the group, there are more electron shells.
describe the chemical and physical properties of alkali metals
- they get more reactive as you go down the group
- they lose an electron to become a positive ion
Lithium- observations and relative reactivity
- effervescence
- metal disappears
- least reactive
Lithium- ease of ignition/flame colour
- least easy
- red
Lithium- products formed
lithium hydroxide + hydrogen
Sodium- observations and relative reactivity
- effervescence
- metal floats and moves on water
- metal disappears
Sodium- ease of ignition/flame colour
- middle easy
- yellow/orange
Sodium- products formed
sodium hydroxide + hydrogen
Potassium- observations and relative reactivity
- effervescence
- floats and moves quickly on water
- lilac flame
- most reactive
Potassium- ease of ignition/flame colour
- lilac flame
- easiest
Potassium- products formed
potassium hydroxide + hydrogen
why are group 1 metals called alkali metals?
their reactions with water result in the formation of an alkaline solution
write balanced equations for the reactions of the alkali metals with water
2Li (s) + 2H2O (l) -> 2LiOH (aq) + H2 (g)
2Na (s) + 2H2O (l) -> 2NaOH (aq) + H2 (g)
2K (s) + 2H2O (l) -> 2KOH (aq) + H2
describe and explain the trend in reactivity of group 1 metals
more reactive as you go down the group:
- more electron shells so increased shielding
- easier for outer electron to be lost
halogens are ——– oxidising agents
very good (they gain electrons easily)
what is the oxidising power of halogens?
the extent to which a halogen atom is able to attract an electron to form a halide ion
the halogens become —– reactive down the group
less
oxidising power of halogens —— down the group
decreases
explain why chlorine is less reactive than fluorine
chlorine has a greater atomic radius than fluorine (more shielding);
electrons experience less attraction to nucleus;
fluorine gains electron more easily;
less oxidising power
why do halogens have a low melting and boiling point?
- diatomic molecules
- London forces
- easily overcome
How do we test for halides?
add an aqueous solution of silver nitrate, AgNO3 (aq) then try to dissolve in ammonia, NH3
give the symbol equations for the reaction of silver nitrate with halides
AgNO3 (aq) + NaX (aq) -> NaNO3 (aq) + AgX
Ag+ (aq) + X- (aq) -> AgX (s)
testing for chlorides result
- white precipitate formed
- fully dissolves in dilute NH3
testing for bromides result
- cream precipitate
- partially dissolve in concentrated NH3
testing for iodides
- yellow precipitate
- does not dissolve in concentrated NH3
chlorine solution
very pale green
bromine solution
orange
iodine solution
brown
why use cyclohexane when determining the reactivity of halogens?
the halogens are very soluble in the cyclohexane layer which will then change colour depending on the dissolved halogen
halogens in cyclohexane
chlorine and bromine= same
iodine= purple
symbol equations for chlorine displacing bromine
Cl2 (aq) + 2NaBr (aq) -> 2NaCl (aq) + Br2 (aq)
Cl2 (aq) + 2Br- (aq) -> 2Cl- (aq) + Br2 (aq)
why are halogen displacement reactions redox?
more reactive halogen oxidises a less reactive halide ion
why is chlorine the most reactive in terms of electrons?
it has the strongest affinity for electrons and will remove electrons from bromide ions and iodide ions
give the general equation for halogens reacting with alkali metals
2M (s) + X2(g) -> 2MX (s)
what is the product of halogens + alkali metals?
- colour
- solubility
salts
- white/colourless
- soluble in water and form colourless, neutral solutions
define disproportionation
a reaction in which the same element is both reduced and oxidised.
give 2 examples of disproportionation
- chlorine as a disinfectant in water
- chlorine in bleach
Cl2 (aq) + H2O (l) -> HCl (aq) + HClO (aq)
Cl2 (aq) + 2NaOH (aq) -> NaCl (aq) + NaClO (aq) + H2O (l)
describe the trend in acid- base nature of the oxide of an element across a period
Acidity increases from left to right, ranging from strongly basic oxides on the left to strongly acidic ones on the right, with an amphoteric oxide (aluminum oxide) in the middle
metals are ——, non metals are ——-
basic, acidic
balanced symbol equation for the reaction of Na2O with water
Na2O + H2O → 2NaOH
balanced symbol equation for the reaction of MgO with water
MgO(s) + H2O(l) → Mg(OH)2(aq)
balanced symbol equation for the reaction of P4O10 with water
P4O10(s) + 6H2O (l) -> 4H3PO4 (s)
NaOH pH and solubility
13/14 (strong base), solubility
Mg(OH)2 pH and solubility
9, not very soluble
H3PO4 pH
1/2 (H+ ions)
balanced symbol equation for the reactions of nitrogen and sulphur oxides with water
2NO2 (g) + H2O (l) → HNO2 (aq) + HNO3 (aq)
SO2 (g) + H2O (l) → H2SO3 (l)
use the products of the reactions of nitrogen and sulfur oxides with water to explain the pH of the resulting solution
- pH 2/3
- pH 1
acid rain- formation of sulphur based acids
Fossil fuels are often contaminated with small amounts of sulfur impurities:
- When these contaminated fossil fuels are combusted, the sulfur in the fuels get oxidised to sulfur dioxide
- S (s) + O2 (g) → SO2 (g)
- Sulfur dioxide may be further oxidised to sulfur trioxide
- 2SO2 (g) + O2 (g) ⇌ 2SO3 (g)
The sulfur dioxide and sulfur trioxide then dissolve in rainwater droplets to form sulfurous acid and sulfuric acid
- SO2(g) + H2O (l) → H2SO3 (aq)
- SO3 (g) + H2O (l) → H2SO4 (aq)
acid rain- formation from nitrogen oxides
The temperature in an internal combustion engine can reach over 2000 °C
Here, nitrogen and oxygen, which at normal temperatures don’t react, combine to form nitrogen monoxide:
N2 (g)+ O2 (g) ⇌ 2NO (g)
Nitrogen monoxide reacts further forming nitrogen dioxide:
2NO (g) + O2 (g) ⇌ 2NO2 (g)
Nitrogen dioxide gas reacts with rain water to form a mixture of nitrous and nitric acids, which contribute to acid rain:
2NO2 (g) + H2O (l) → HNO2 (aq) + HNO3 (aq)
Lightning strikes can also trigger the formation of nitrogen monoxide and nitrogen dioxides in air
Nitrogen dioxide gas reacts with rain water and more oxygen to form nitric acid
4NO2 (g) + 2H2O (l) + O2 (g)→ 4HNO3 (aq)
- When the clouds rise, the temperature decreases, and the droplets get larger
- When the droplet containing these acids are heavy enough, they will fall down as acid rain
incomplete and complete combustion of methane
CH4 + 2O2 → CO2 + 2H2O
2CH4 + 3O2→ 2CO + 4H2O
CH4 + O2→ C + 2H2O
catalytic converters
- used in car exhaust boxes to reduce air pollution
- the transition metal catalysts facilitate the conversion of these pollutants into less harmful products
steps of a catalytic converter
- Molecules of carbon monoxide and nitrogen monoxide are absorbed onto the surface
- The bonds in both molecules are weakened causing them to react together to form carbon dioxide and nitrogen
- The products are then desorbed from the surface of the catalyst
2NO (g) + 2CO (g) → N2 (g) + 2CO2 (g)
CH3CH2CH3 (g) + 5O2 (g) → 3CO2 (g) + 4H2O (g)
effects of acid rain
Acid deposition can react with metals and rocks (such as limestone) causing buildings and statues to get damaged
Apart from acid deposition directly falling on leaves and killing plants, acid particulates can block stomata ( plant pores) and prevent gaseous exchange
When acid rain falls on rivers and lakes the pH can fall to levels that are unable to support life
