Topic 3 Periodicity - IB CHEMISTRY
Terms in this set (116)
what is periodicity?
repeating pattern of physical and chemical properties
what is one of the fundamental properties of an element?
the group in the Periodic Table tells what about an element?
tells the number of electrons in the valence outer energy level
the period in the Periodic table tells what about an element?
the number of energy levels of the element
which trends explain chemical properties of elements?
atomic + ionic radii
how are trend in both physical and chemical properties explained?
using the concept of effective nuclear charge
1. energy levels
2. distance from nucleus
4. nuclear charge
5. effective nuclear charge
effective nuclear charge trend across a period
Increase across a period (due to increasing nuclear charge with no accompanying increase in shielding effect).
effective nuclear charge trend down a group
Decrease down a group (although nuclear charge increases down a group, shielding effect more than counters its effect).
provide an example of how to structure previous answer properly (in general terms)
the nuclear charge of an atom is given by the atomic number and so increases by successive elements in the table as a proton is added to the nucleus. The outer electrons which determine which determine many of the physical and chemical properties do not experience the full attractive of this charge as they are shielded from the nucleus and are repelled by inner electrons. The presence of increasing number of inner electrons reduces the attraction of nucleus for the outer electron. Therefore, the effective nuclear charge experienced by the outer electrons is less than the full nuclear charge.
as period 3 is crossed from left to right, what happens to the protons and electrons?
1 of each is added
what happens to the effective nuclear charge?
it increases with the nuclear charge as there is no change in the number of inner electrons as the valence outer energy sub level can hold 10 electrons
describe effective nuclear charge across a period and down a group.
the effective nuclear charge experienced by an atom's outer electrons increases with the group number of the element. It increases across a period but remains approximately the same down a group.
The atomic radii is measured how?
as half the distance between neighbouring nuclei
State the trend of atomic radii down a group? explain
increases as the number of occupied energy levels increases. The attraction between the nucleus and the outer electrons become weaker increasing the atomic radii. Moreover, more electrons are being shielded from the nucleus making a weaker attraction which creates a bigger atomic radii. Nuclear charge increases but effective nuclear charge decreases because there's enough shielding effect to counteract effective nuclear charge's effect.
State the trend of atomic radii across a period? explain
decreases because even though the elements have the same occupied energy levels, the attraction between the nucleus and the outer electrons increases as the nuclear charge increases. So as the attraction gets stronger, the outer energy level becomes closer to the nucleus.
compare sizes of positive ions and their parent atom and explain the answer.
positive ions are smaller than their parent atoms because the creation of the positive ions means the loss of the outer energy level.
how is the radii of the ion increased in negative ions compared to the parent atom?
more electrons are added which increases electron repulsion between the electrons in the outer energy level and causing them to move further apart increasing the radii of the outer energy level.
explain the trend of ionic radii down a group
increases as the number of electron increases increasing the number of occupied energy levels. The attraction of the nucleus and the outer electron therefore becomes weaker and electron shielding is increased which all contributes to increase of ionic radii. Nuclear charge increases down the group but effective nuclear charge decreases because shielding more than counteracts effective nuclear charge's effect.
explain the trend of ionic radii across a period
decreases from group 1 to 14 for positive ions. Positive ions from different groups such as Na+, Mg2+, Al3+ and Si4+ all have the same electron configuration. the decrease in ionic radius is due to the increase in nuclear charge with atomic number across the period. The increased attraction between the nucleus and the electrons pulls the outer energy level closer to the nucleus.
increasing from group 14 to 17 for negative ions. negative ions such as Si4-, P3-, S2- and Cl- have the same electron configuration. The decrease in ionic radii is due to the increase in number of electrons across the period. This increases the shielding by inner electrons and electron-electron repulsion which would decrease the force of attraction between the nucleus and the outer electrons, making the radii bigger.
what are isoelectronic ions?
ions which have the same number of electrons
Define first ionisation energy and write the equation
is the energy required to remove one mol of electrons from 1 mol of gaseous atoms in their ground state.
explain the trend of ionisation energy across a period
increases because across a period there's an increase in no. protons which causes a stronger attraction between the nucleus and the outer electrons. Shielding stays roughly the same because the outer electrons are being removed from the same energy level. There's an increase in effective nuclear charge which causes an increase in the attraction between outer electrons and the nucleus and makes the electrons more difficult to remove.
explain the trend of ionisation energy down a group
decreases. Although the nuclear charge increases as the number of protons increase, the effective nuclear charge is the same because the shielding by the inner electrons counteracts it. because the electron removed is from the energy level furthest from the nucleus, the increased distance between the electron and the nucleus reduces the attraction between them.
define first electron affinity
the energy change when 1 mol of electrons is added to 1 mol of gaseous atoms to form 1 mol of gaseous ions
X(g) + e- ---> X-(g)
this process is generally exothermic
define second electron affinity
energy absorbed when 1 mol of electrons is gained by 1 mol of singly charged negative ions in the gas state
why is the second electron affinity for oxygen endothermic?
O-(g) + e- ----> O2-
as the added electron is repelled by the negatively charged oxide
how is electronegativity and electron affinity different?
electronegativity refers to bonding electrons (covalent molecules) and electron affinity relates to individual atoms (ionic ions)
explain the trend for electron affinity down the group
The greater the distance between the nucleus and the incoming electron, the less the attraction and so the less energy is released as electron affinity.
explain the trend for electron affinity across a period
values become more exothermic across a period with some exceptions. we expect that the nuclear charge increases, increasing attraction between the nucleus and the incoming electron.
describe the electron affinity between F and Cl
there's a decrease in electron affinity which is unusual.
explain the anomaly above
The incoming electron is going to be closer to the nucleus in fluorine than in any other of these elements, so you would expect a high value of electron affinity.
However, because fluorine is such a small atom, you are putting the new electron into a region of space already crowded with electrons and there is a significant amount of repulsion. This repulsion lessens the attraction the incoming electron feels and so lessens the electron affinity.
A similar reversal of the expected trend happens between oxygen and sulphur in Group 6. The first electron affinity of oxygen (-142 kJ mol-1) is smaller than that of sulphur (-200 kJ mol-1) for exactly the same reason that fluorine's is smaller than chlorine's.
the electron affinity reach a maximum for which elements?
the electron affinity reach a maximum for group 2 and group 5 elements
Why is group 6's electron affinity less than group 7's?
It's simply that the Group 6 element has 1 less proton in the nucleus than its next door neighbour in Group 7. The amount of screening is the same in both.
That means that the net pull from the nucleus is less in Group 6 than in Group 7, and so the electron affinities are less.
the ability of an atom to attract electrons in a covalent bond
an element with high electronegativity has what electron pulling power?
strong and vice versa for low => weak
explain the trend across a period
increases because across the period elements gain protons which increases their nuclear charge which results in an increased attraction between the nucleus and the bond electrons.
explain the trend down a group
As we descend the group there is an increasing distance between the nucleus and shared pairs of electrons, therefore the attraction for these electrons decreases due to the increase in number of electron. The increase in nuclear charge down a group is counteracted by the increased shielding caused by the extra occupied main energy levels within the atom.
Group 1 and 2 are referred to as what because of thier low electronegativity values
Group 5, 6 and 7 are referred to as what because of thier low electronegativity values
metals have ............... ionisation energies and electronegativity than non-metals
what is the melting point?
the temperature at which a pure solid is in equilibrium with its pure liquid at atmospheric pressure
explain melting point trend down group 1
decreases because the elements get larger so the attractions between the nucleus and the outer electrons decreases making it easier to break them.
explain melting point trend down group 17
increases because the elements get larger so they get more intermolecular forces.
explain melting point trend across a period
generally increase until group 14 and then fall till group 18 because the bondings change.
how are chemical properties of an element determined?
by the number of valence electrons in their outer energy level
state the visuals of group 18
why are they unreactive?
unable to lose or gain electrons because they have a complete valence energy levels with 8 electrons
they don't form positive ions because they have highest ionisation energy
they dont't form negative ions because extra electrons would have to be added to an empty outer energy level where they would experience a negligible effective nuclear force.
describe physical properties of group 1
good conductors of electricity and heat
grey shiny surfaces when cut
describe the chemical properties of group 1
form ionic compounds with non-metals
what does the low density suggest about metallic bonding?
big ions with +1 charge so the attraction between the ion and the delocalised electron is weak and therefore has weak metallic bonding
state the trend for density
increases down the group but the mass increases faster than the volume with increasing atomic number.
why are group 1 metals reactive?
because their 1st ionisation energies are relatively low so their valence electron is relatively easy to lose to form an ion with a singly positive charge.
describe the reactivity of group 1
increases down the group as the elements get bigger. The outer electron is more shielded from the nucleus because it's in a higher energy level. This more than compensates for the increase in nuclear charge. The valence electron feels a weaker attraction to the nucleus (weaker effective nuclear charge) so is more easily lost.
reactions of group 1 with oxygen
form metal oxides
Alkali metals + water ------>
hydrogen + metal hydroxide
how do the metals react with water?
how does Li react in water?
floats and reacts slowly
releases hydrogen but keeps in shape
how does Na react in water?
reacts vigorously with a release of hydrogen
the heat produced is sufficient enough to melt the unreacted metal which forms a small ball that moves around on the water surface
how does K react in water?
reacts even more vigioursly to produce enough heat to ignite the hydrogen produced.
Produces a lilac coloured flame and moves fast on water
reactions of the metals with halogens
form ionic halide precipitates
what is a redox reaction?
one which involves oxidation and reduction
loss of electrons
the alkali metals are what kind of agents?
reducing agents because they can readily donate an electron
the halogens are what kind of agents?
oxidising agents because they can readily accept an electron
state the physical properties of group 17
change from gas to liquid to solid states
explain their trend of melting and boiling point
going down the group, melting and boiling points increase which result in a change in physical state at room temperature from gas to liquid to solid. The forces between molecules are weak intermolecular forces called London dispersion forces. As you go down the group, the number of electrons in the molecules increases and this results in an increase in the strength of the dispersion forces. This means that more energy is needed to separate molecules apart which produce an increase in melting and boiling points.
the halogens are more soluble in which liquids
non-polar organic solvents such as hexane than water. Hexane is less dense than water so it floats above water.
state the chemical properties of group 17
v reactive non-metals shown by the readiness to accept electrons as illustrated by their very exothermic electron affinities
form ionic compounds with metals and covalent compounds with other non-metals
explain their trend in reactivity
Reactivity decreases down the group as the atomic radius increases and the attraction for outer electrons decreases from the nucleus as the outer energy level is further from the nucleus and more shielded from the nucleus by inner electrons. These factors more than counteracts the increase in nuclear charge hence an electron is less easily gained.
why do their oxidising ability decrease down the group?
atomic radius increases down the group. small atoms accept electrons less easily than large ones for the reasons already discussed.
Reaction of chlorine with iron
heated iron wool reacts with chlorine gas to produce a reddish-brown solid, iron (III) chloride.
The iron glows very brightly as it is a very exothermic reaction.
reactions of bromine with iron
Iron is heated constantly this time as bromine vapour is passed over it (from the liquid in the test tube).
the iron glows but not as brightly as when it reacts with chlorine
redish-brown slid produced
Reaction of iodine with iron
heated iron wool reacts with iodine vapour if the iron is heated constantly as the iodine vapor passes over it.
reddish-brown solid formed
2KBr(aq) + Cl2(aq) ----->
and explain the results
2KCl(aq) + Br2(aq)
write the ionic equation and 2 half-equations
this is because chlorine nucleus has a stronger attraction for an electron than a bromine nucleus because of its smaller atomic radius and so takes the electron from bromide ion. The chlorine has gained an electron from the bromide ion. The chlorine has gained an electron and so forms the chloride ion, Cl-. The bromide ion loses an electron to form bromine.
from colourless to orange
2KI(aq) + Cl2(aq) ------>
2KCl-(aq) + I2(aq)
from colourless to dark orange/brown
the colour is seen clearly if hexane is added as the iodine then moves to the upper organic layer and turns into a pink solution
2KI(aq) + Br2(aq) ------>
2KBr(aq) + I2(aq)
from colourless to red-brown
if hexane is added, the iodine moves to the upper organic layer and is seen as pink solution
2KCl(aq) + I2(aq) ----->
2KI(aq) + Cl2(aq)
from colourless to red-brown, this indicates that there was no chemical reaction as the colour seen is simply that of the aqueous iodine added. if hexane is added, the iodine moves to the upper organic layer and is seen as pink solution
state and explain the halogen test on litmus paper
chlorine bleaches moist litmus paper rapidly. It turns red first and then is bleached because acid is made (HCl and HClO). HCl is stronger. HClO is also a bleach so it bleaches the paper.
Bromine slowly bleaches moist litmus paper but iodine has no effect on litmus paper.
test for Iodine
use starch solution
turns from orange-red to blue-black
how to tell the difference in the colours of Br and I solutions?
add a hydrocarbon solvent
iodine forms a purple solution and bromide a dark orange
how do halogens react with silver?
state the equation
forming insoluble salts
Ag+(aq) + X-(aq) ------> AgX(s)
test for halide ions
adding silver nitrate and then ammonia
AgCl - white ppt soluble in dilute (aq) ammonia
AgBr - cream ppt partially soluble in dilute (aq) ammonia, soluble in concentrated ammonia
AgI - pale yellow ppt insoluble in concentrated ammonia
If the ppt are allowed to stand under sunlight (not reacting with ammonia), the ppt gradually darken. AgCl darkens quickly than AgBr and AgBr darkens more quickly AgI. This is due to a decomposition reaction:
2AgCl(s) -------> 2Ag(s) + Cl2(g)
appearance of period 3 chemicals at rtmp
Na, Mg and Al - silvery, shiny solids
Si - grey solid
P - red / white solid
S - pale yellow, brittle solid
Cl - yellow - green gas
Ar - colourless gas
reactions of heating Mg, Al and S in air
Mg - burns with a brilliant white flame to leave a white powder
Al - foil shrivels, white solid coating on surface
S - melts easily to yellow liquid, red liquid, becoming more viscous. Burns with a blue flame to leave a colourless gas.
reaction of dilute acids of Mg, Al, Si and S
Mg - effervesces. gas pops with a lighted splint. Tube gets hot, metal disappears to leave colourless solution
Al - effervesces. gas pops with a lighted splint. metal disappears to leave colourless solution
Si and S - no reaction
reactions of dilute acids and Mg, Al, Si and S
Mg - no reaction
Al - effervesces. gas pops with a lighted splint. metal disappears to leave colourless solution
Si - effervesces. gas pops with a lighted splint. metal disappears to leave colourless solution
S - no reaction
write ionic equations of previous reactions
2Al + 2NaOH + 6H2O ------> 2NaAl(OH)4 + 3H2
2Al + 2OH- + 6H2O ------> 2[Al(OH)4]- + 3H2
Si + 2NaOH + H2O -----> Na2SiO3 + 2H2
Si + 2OH- + H2O ------> SiO32- + 2H2
explain the melting points of Na, Mg and Al oxides
the giant ionic oxides have high melting and boiling points hence are solid at room temperature because the electrostatic forces of attraction between oppositely charged ions in the giant attic structure are strong hence a lot of energy is needed to overcome them.
explain Na, Mg and Al oxides' conductivity of electricity abilities
Conduct electricity when molten as the ions are then free to move and act as mobile charge carriers. The molten oxides are electrolytes and are decomposed to their elements during electrolysis.
explain the melting points of SiO2
it has a high melting point hence is solid at room temperature as a lot of energy is needed to break the strong covalent bonds between silicon and oxygen atoms in the giant structure.
explain SiO2 conductivity of electricity abilities
it doesn't conduct electricity as its bonding is covalent and there are no mobile charge carriers.
explain the melting points of
P, S and Cl oxides
they have low melting and boiling points as they have simple molecular covalent structures. The covalent bonds within the molecules are strong but there are weak dispersion forces between molecules which require little energy to overcome them.
explain P, S and Cl oxides conductivity of electricity abilities
they don't conduct electricity because the are covalent so do not have mobile charge carriers
state an equation for silver chloride and ammonia
AgCl(s) + 2NH3(aq) ------> [Ag(NH3)2]+ + Cl-(aq)
ionic compounds have which structures?
covalent compounds have which structures?
The ionic character of a compound depends on what?
the difference in electronegativity between its elements and oxygen
state the trend of ionic characters of oxides across a period
decreases because electronegativity increases
the oxides of Si, P, S and Cl are covalent because
of the relatively small different in electronegativity values between the elements and oxygen
what structure is silicon dioxide
what structures are P S Cl oxides
simple molecular covalent
state the trend of ionic characters of oxides down a group
increases because electronegativity decreases
oxides of metals are
ionic and basic
oxides of non-metals are
covalent an acidic
state an example of an amphoteric oxide
what are alkalis
bases which are soluble in water
Na2O(s) + H2O(l) ------>
A basic oxide reacts with acid to form
salt and water
Li2O(s) + 2HCl(aq) -----> 2LiCl(aq) + H2O(l)
non-metallic oxides react readily with water to form
P4O10(s) + 6H2O(l) -----> 4H3PO4(aq) - phosphoric (V) acid
P4O6(s) + 6H2O(l) -----> 4H3PO3(aq) - Phosphoric (III) acid
sulphur oxide and water ===>
SO3(l) + H2O(l) -----> H2SO4(aq)
SO2(g) + H2O ------> H2SO3(aq)
what is acid rain
Nitrogen oxides and sulfur oxides mixed with water in clouds
how are nitrogen oxides formed?
nitrogen monoxide NO is formed at high temperatures in internal combustion engines:
N2(g) + O2(g) ---> 2NO(g)
NO is further oxidised:
2NO(g) + O2(g) ---> 2NO2(g)
This then reacts with water to form nitric (V) acid and nitric (III) acid which leads to acid rain:
2NO2(g) + H2O(l) ----> HNO3(aq) = HNO2(aq)
how are sulphur oxides formed?
fossil fuels such as coal or oil contain super compounds, and when the product of these is burned, SO2 is produced. This leads to acid rain.
S(s) + O2(g) -----> SO2(g)
Nitrogen dioxide catalyses this reaction:
2SO2(g) + O2(g) -----> 2SO3(g)
this then reacts with water:
SO3(l) + H2O(l) ------> H2SO4(aq)
what are the consequences of acid deposition
- corrosion of limestone buildings and ironwork
- acidification of lakes and rivers => death of aquatic life
- A fall in pH dissolves aluminium ions from the soil which is toxic to fish
- damages trees