Periodic table Flashcards
s block
Groups 1&2- outer electrons in the s orbital
P block
groups 3-0
Outer electrons in p orbital
D block
transition metals- outer electrons in d orbital
ionisation energy in the periodic table
General increase across a period. There is an increase in nuclear charge in the same energy level so there is little extra shielding and therefore a greater attraction between a nucleus and outer electrons
ionisation energy in Group 2 and 3
Decrease between Group Two and three. Group-Three elements the electron is in a new subshell of slightly higher energy level and is partially shielded by s-electrons
ionisation energy between Group Five and six
Decrease between five and six. In group six electron is removed from an orbital containing a pair of Electrons. The repulsion between these electrons makes an electron easier to remove. in group five the electron is removed from a singly occupied orbital
ionisation energy down a group
Decreases down a group. Her electron has increased shielding from inner electrons and is further from the nucleus. Outweighs the increase in the nuclear charge
Electronegativity across a period
bonding elecrons
Increases across a period. increase In nuclear charge but the bonding electrons are always shielded by the same inner electrons so there is a greater attraction between nucleus and bonding pair
Electronegativity down a group
bonding elecrons
Decreases down a group. Bonding electrons have increased yielding from the nucleus so the attraction between nucleus and bonding electrons decreases
Metals- melting and boilng temperatures
Have metallic bonding, there is an increase because the metallic bonding gets stronger. Ions have a greater charge and there is an increased number of delocalized electrons
Giant covalent structures eh. Silicon boron
Each atom is bonded covenantly to four other atoms and a large amount of energy is needed to break all these bonds
Simple molecular substances eh. P, S, Cl
Although the Covalent bonds between atoms in the molecules are strong the intermolecular forces holding the molecules together are weak and do not need much energy to break
Argon
Lowest melting boiling temperature because it exists as separate atoms held together by very weakened Induced dipole - induced dipole forces
Electron transfer- half equations
oxidation: Eg. Mg—> Mg2+ + 2e
Reduction: Eg. Cu2+ + 2e —> Cu
Group one metals with cold water
Form a hydroxide and hydrogen EG. Sodium
2Na + 2H2O—> 2NaOH + H2
Lithium’s reaction with water
Floats on the water, gently fizzing
Sodium’s reaction with water
Sodium melts into a bowl that dashes around the surface
Potassium’s reaction with water
Potassium melts into a ball and catches fire
Cesium’s reaction with water
caesium explodes and shatters the glass container
Group two metals reaction with water
React less vigorously and hydroxide and hydrogen are formed
eg. Calcium
Ca + 2H2O —>Ca(OH)2 + H2
Calcium’s reaction with water
Calcium produces a steady stream of bubbles and the liquid goes cloudy as a white participate of calcium hydroxide forms
Barium’s reaction with water
barium produces greater Efveresscence and the solution is clearer since barium hydroxide is more soluble
Magnesiums reaction with water
Magnesium reacts with steam to produce the oxide and hydrogen
Trends in reactivity
g2
Reactivity increases as you go down the group
this is because when the S block metals react they lose electrons to form positive ions.
Since ionisation energy decreases down a group, the energy needed to form positive ions decreases. This leads to lower activation energies and therefore faster reactions
G1 more reactive than G2
Because Group One metals lose only one electron while group 2 will lose two electrons
Reaction with acids
All group two metals react vigorously with hydrochloric acid to produce a colourless solution of metal chloride and bubbles of hydrogen Eg. Mg
Mg + 2HCL—> MgCl2 + H2
Reactivity increases down a group
Only magnesium can react with sulfuric acid because other members have been soluble sulphates
Group 1 metals are too reactive to be added directly to acids
Group 2- Reaction with oxygen
Apart from magnesium, all group two metals tend to burn with a characteristic flame. All G2 metals burned to form solid white oxides
Eg. 2Mg +O2—> 2MgO
groups 1- reaction with oxygen
Group One metals also form white solids and burn with a characteristic flame
Eg. 4Li + O2 —> 2Li2O
Oxides and hydroxides - s block metal oxides
( metal oxides are basic) (non metal oxides are acidic)
All the s block metal oxides are strong bases and they neutralise acids to form a salt and water
Eg. MgO + 2HCl —> MgCl2 +H2O
Other group two hydroxides are not very solid balls so saturated solutions of these hydroxides are only weakly basic because the concentration of hydroxide ion is very low
G1- oxides and hydroxides
G1 oxides And barium oxide react w water to form a soluble hydroxide
Eg. Na2O + H2O —> 2NaOH
HydroxideS are soluble so are alkalis
test for cations
all S-block elements apart from magnesium can be identified by a flame testable to stop a clean metal wire or splint is moistened with hydrochloric acid dipped in the compound and held in a non-luminous Bunsen flame
Flame tests
Li+- Red
Na+- orange yellow
K+- lilac
Mg2+- no colour
Ca2+- brick red
Sr2+- crimson
Ba2+- apple green
Solubility in water
All Group One compounds are soluble but many group two compounds are not
Trends for group two compounds - solubility
- All nitrates are soluble
- All carbonates are insoluble
- The hydroxides become more soluble as you go down the group so magnesium hydroxide is insoluble whilst barium is soluble
Mg2+(aq) + 2OH-(aq) —> Mg(OH)2 (s) - The sulphates become less soluble as you go down the group
Ba2+(Aq) + SO42- (aq)–> BaSO4 (s) - all pretcipirates ate white
Thermal stability of hydroxide and carbonates- Group 2 hydroxides’
All group two hydroxides decompose on heating to the oxide and steam Eg.
Ca(OH)2(s) —> CaO (s) + H2O (g)
Thermal stability increases down a group so Hydroxides have to be heated more strongly before they will decompose
Himal stability of hydroxide and carbonates - group two carbonates
all group two carbonates decompose on heating to the oxide and carbon dioxide Eg.
MgCO3(s) —-> MgO(s) + CO2(g)
Fairmont stability increases as you go down the group
Presence of CO2 turns limewater cloudy
Chemistry of group seven halogens and halides
at room emp
Halogens produce salt called halides
At room temperature chlorine is a green gas, bromine a red brown liquid and iodine a grey solid
Volatility down a group
Ask the number of Electrons increases with atomic number, there is an increase in the induced dipole - induced dipole intermolecular forces holding the diatomic molecule together. Therefore the melting and boiling temperatures increase as you go down the group.
A substance that forms vapours easily are called volatile. Since with the low boiling temperature has high volatility and volatility decreases down the group
Trends in reactivity with group seven
The halogens react by gaining electrons to form negative halide ions. They gain electrons during reactions or halogens are reduced and they oxidise with other substances
As you go down the group’s outer electrons are shielded more and further from the nucleus so it’s harder to attract electrons and both reactivity and oxidising power decreases down the group
Reaction with a metal and halogen to form a white metal halide
eg.
Sodium + cchlorine gas= white sodium chloride
2Na +Cl2–> 2NaCl
Reaction with iron wool
Ironwall also bends directly in chlorine or bromine vapour to give iron (III) halide
2Fe +3Br2—> 2FeBr3
When it burns in iodine vapour it only produces iron (II) iodide since iodine is less reactive and it’s a weaker oxidising agent
Fe+ I2 —> FeI2
Displacement reactions
A halogen in a higher position in the group will oxidise a halide ion from a lower in the group since oxidising powers decrease down a group
Eg. Cl2(aq) + 2Br- (aq) —> 2Cl- (aq) + Br2(aq)
0 -1 -1 0
Task for hslide ions silver nitrate test
If you start from a solid it must be first dissolved in water
A few drops of nitric acid is added first to make sure that any other anions especially carbonates are removed, as they would also form precipitates
Silver nitrate test observation
Cl- forms white precipitate
Br- forms cream
I- forms pale yellow
Eg. Ag+(aq) + Cl-(aq) —-> AgCl(s)
Aqueous ammonia to the precipitate
Sometimes it can be difficult to tell the difference between the precipitates so to distinguish between them aqueous ammonia is added to the precipitate
Aqueous ammonia observation
AgCl- Precipitate dissolves in dilute ammonia
AgBr- Precipitate does not dissolve much in dilutes ammonia but dissolves in concentrated ammonia
AgI- Participate is insoluble in dilutes and concentrated ammonia
Chlorine added to water
Chlorine is commonly added to water as a gaseous element and the equilibrium is established
Cl2 + H2O –> HCl + HOCl
User of the chlorate ion
ClO- Kills bacteria and other microbes so adding chlorine to water makes it safe to drink or swim. Chlorination is used to prevent the outbreak of serious diseases like typhoid and cholera
- Highly toxic and can react with naturally occurring organic compounds found in water supply to form chlorinate hydrocarbons which can cause liver and kidney cancer
- “mass medication”
Fluoride added to water
To reduce tooth decay by preventing cavities. Water fluoridation reduces cavities in children’s but its effectiveness in adults is less clear
- Although fluoridation can cause dental fluorosis leading to tooth discoloration there is no clear evidence of other adverse effects
- “ mass medication”
