inorganic chemistry Flashcards
what are the different areas of the table in terms of electronic arrangement
s, block, p-block, d - block and the f block
how do we know if an element is in the s-block
an element is in the s block if its highest energy electron is in the s-orbital
e.g. Na 1s2 2s2 2p6 3s1
how do we know if an element is in the p- block
an element is in the p block if its highest energy electron is in the p - orbital
e.g. C
1s2 2s2 2p2
how do we know if an element is in the d -block
an element is in the d - block if the highest energy electron is in the d - orbital
e.g. Fe
1s2 2s2 2p6 (…) 3d6
what are the differences between the transition metals and the d - blocks
the transition metals and d - block elements are not exactly the same
e.g. scandium and zinc are, not transition metals because they do not form any compounds in which they have partly filled d- orbitals, which is the characteristics of transition metals
what are the origins of the terns s,p,d, and f
- when elements are heated they give out light at certain wavelengths as excited electrons fall back from one energy level to a lower
This causes lines to appear in the spectrum of light they give out
The letters s, p, d and f stands for words that were used to describe the lines
S- FOR SHARP
P - FOR PRINCIPLE
D - FOR DIFFUSE
F - FOR FINE
what is the s-group
a group is a vertical column of elements
The elements in the same group form a chemical “family” - they have similar properties
elements in the same group have the same number of electrons in the outer main level
what is the reactivity of the elements in s - block
in the s-block, elements (metals) get more reactive going down a group
what is the general reactivity of the non -metals
non - metals elements tend to get more reactive going up a group
what is the general reactivity of the transition elements
transition elements are a block of rather unreactive metals
this is where the most useful metals are found
what are periods
are horizontal rows of elements the periods are numbered stating from Period 1 which contains only hydrogen and helium
period 2 contains the elements lithium to neon and so on
There are trends as you go across a period
what does the term periodicity/ periodic means
- recurring regularly
what can periodicity be explained by ( the recurring trend in the periodic table)
the electron arrangement of the elements
e.g. the elements in group 1, 2 and 3 (sodium, magnesium and aluminium) are metals
They have GIANT STRUCTURES
they lose their outer electrons to form ionic compounds
Silicon in group 4 has four electrons in its outer shell with it forms four covalent bonds
The elements have METALLIC properties and are classed as a SEMI - METALS
the elements in group 5,6 and 7 are NON - METALS
either accept electrons to form ionic compounds or share their electrons to form COVALENT COMPOUNDS
Argon in Group 0 is a NOBLE GAS - it has a full outer shell and is UNREACTIVE
what is the “clear cut” in terms of the trend in the melting and boiling points
there is a clear break in the middle of the figure between elements that have a high melting point on the left and those with low melting points on the right
what can explain the trends of the melting and boiling points in the periodic table
these trends are due to their structures
- giant structures tend to have high melting and boiling points (on left)
- molecular or atomic structures (on right) tend to have low melting and boiling points
why do the melting points of the metals increase from sodium to aluminium
due to the increased strength of the metallic bonding
as you go from the left to the right the charge on the ion increases so more electrons join the delocalised electron “sea” that holds the giant metallic structure together
what does the melting points of the non -metals depend on
the size of the van Der Waals forces between the molecules
This, in turn, depends on the number of electrons in the molecule and how closely the molecule is packed together
as a result, the melting points of these non-metals are
S4 > P4 > CL2
what do the atomic radii tell us
it tells us the size of atoms
why can’t you measure the radius of an isolated atom
you can’t measure the radius of an isolated atom because there is no clear point at which the electron cloud density around it drops to 0
instead, half the distance between the centres of a pair of atoms is used
why does the atomic radius of an element differ
- atomic radius is a general term
it depends on the type of bond that is forming - covalent, ionic, metallic, van der Waals
the covalent radius is most commonly used as a measure of the size of the atom
why are noble gases often left out of comparisons of atomic radii
even metals can form covalent molecules such as Na2 in the gas phrase
since noble gases do not bond covalently with one another they do not have covalent radii and so they are often left out of comparisons in atomic radii
why do the radii of atoms decrease across a period
- we can explain this trend by looking at the electronic structures of elements in a period
e. g. sodium to chlorine in period 3
as you move from Na to Cl you are adding protons to the nucleus and electrons to the outer main level, the third shell ( the same shell), the charge increases from + 11 to +17
this increased charge pulls the electrons in closer to the nucleus
There are no additional electrons shells to provide more shielding so the size of the atom decreases as you go across the period
why do the radii of atoms increase down a group
going down a group in the Periodic table, the atoms of each element have one extra complete main level of electrons compared with the one before
-so going down the group, the outer electron main level is further from the nucleus and the atomic radii increases
what is the first ionisation energy
the first ionisation energy is the energy required to convert a mole of isolated gaseous atoms into a mole of singly positively charge gaseous ions, that is, to remove one electron from each atom
what is the general trend for ionisation energies across a period
the ionisation energy increases going down any period
this is because the number of protons increases but the shielding and the number of shells stays the same
this increased nuclear charge means that it gets increasingly difficult to remove an electron
why does the first ionisation energy decreases going down the group
the number of oof filled inner levels increases down the group
This results in an increase in shielding
Also, the electron to be removed is at an increasing distance from the nucleus and is therefore held less strongly
Thus the outer electrons get easier to remove going down a group because they are further away from the nucleus
why is there a drop in ionisation energy from one period to the next
moving from neon in Period 0 with electron arrangement 2,8 to sodium, 2,8,1 there is a sharp drop in the first ionisation energy
This is because at sodium a new main level starts and so there is an increase in atomic radius, the outer electron is further from the nucleus, less strongly attracted and easier to remove
a detailed look into the trends in ionisation energy in period 3
the first ionisation energy drops between GROUP 2 and GROUP 3 so that aluminium has lower ionisation energy than magnesium
the ionisation energy drops again slightly between GROUP 5 ( phosphorus) and GROUP 6 (sulfur)
similar patterns occur in other periods
how do we explain the drop in the first ionisation energy across GROUP 2 and GROUP 3
magnesium loses an electron in 3s while magnesium loses an electron in 3p which is a higher energy level than magnesium so the electron can be lost easier
explain the drop in the first ionisation energy between GROUP 5 and GROUP 6
an electron in a pair will be easier to remove than one in an orbital on its own orbital because to os being repelled by the other electron
phosphorus has no paired electrons in a p -orbital because each p -electron s in its own orbital
sulfur has two of its p – electrons paired in a p - orbital so one of these will be easier to remove than an unpaired one due to the repulsion of the other electron in the same orbital
what are group 2 elements sometimes called
elements in group 2 are called the alkaline earth metals
this is because their oxides and hydroxides are alkaline
which block is group 2 in
like group, they are s- block elements
they are similar to group 1 but are less reactive
what are the physical properties of group 2
electron arrangement
2 electrons in outer s- orbital
This s- orbital becomes further from the nucleus going down the group
the sizes of the atoms
the atoms get bigger going down the group. The atomic (metallic) radii increases because each element has an extra filled main level of electrons compared with the one above
explaining the trend in melting points in group 2
group 2 elements are metals with high melting points typical of a giant metallic structure
this is because going fon the group, the electrons in the “sea” of delocalised electrons are further away from the positive nuclei
As a result, the strength of the metallic bonds decreases going down the group
For this reason, the melting points of group 2 elements decrease slightly going down the group staring with calcium
explain the ionisation energies in group 2
in all their reactions, atoms in Group 2 lose their two outer electrons to form ions with two positive charges
SO, an amount of energy equal to the sum of the first and second ionisation energies is needed for complete ionisation
Both the first ionisation energy and the second ionisation decreases going down the group - it takes less energy to remove the electrons as they become further away from the positive nucleus
The nucleus is more effectively shielded by more inner shells of electrons
what is oxidation
oxidation is the loss of electrons so in all reactions the group 2 metals are oxidised
what is the trend of group2 metals reactions with water
with water you see a trend in reactivity:
- metals you get more reactive going down:
These are also redox reactions
The basic reaction (where M is any metal in Group 2_)
M(s) +2H2O - M(OH2)aq +H2
why would magnesium hydroxide important
magnesium hydroxide is a milk of magnesia and is used in indigestion remedies to neutralise excess stomach acid
how does magnesium react with cold water
magnesium reacts very slowly with cold water but rapidly with steam to form an alkaline oxide and hydrogen
Mg(s) + H2O - MgO(s) + H2
how does calcium react with water
calcium react in the same way but more vigorous, even with cold water
how do strontium and barium react with water
strontium and barium reacts more vigorous
what is calcium hydroxide sometimes called
calcium hydroxide is sometimes slaked lime and is used to treat acidic soil
Most plants have an optimum level of acidity or alkalinity in which they thrive
e.g. grass grows best at a pH of 6
what are the hydroxides
the hydroxides are all white solids
why do all the metals in group 2 form (OH)2
because it is in group 2 and needs 2 hydroxides
what are the trends in solubilities in group 2 hydroxides
Mg(OH)2 (magnesium hydroxide) (milk of magnesia) is almost insoluble
Ca(OH)2 calcium hydroxide, is sparingly soluble and a solution is used as lime water (slightly soluble)
Strontium hydroxide Sr(OH)2 is more soluble
Barium hydroxide Ba(OH)2 dissolves to produce alkaline solution:
Ba(OH)2(s) + aq - Ba2+(aq) + 2OH-(aq) (soluble)
why are sulfates important
due to the trend in solubility in sulfates, barium sulfatesis virtually insoluble
This means that it can be taken by mouth as a barium meal to outline the gut in medical x - rays ( the heavy barium atom is very good at absorbing x - rays)
the test is safe, despite the fact that is barium compounds are highly toxic, because barium sulfate is so insoluble
how do we test for barium sulfates
the insolubility of barium sulfate is used in a simple test for sulfate ions in solution
The solution is first acidified with nitric or hydrochloric acid
The barium chloride solution is added to the solution under test and if sulfate is present a white precipitate of barium sulfate is formed
what are the solubilities of the sulfates
MgSO4 - soluble
CaSO4 - slightly soluble
SrSO4 - insoluble
BaSO4 - insoluble
what are group 7 elements
group 7 on the right - hand side of the Periodic Table
- it is made up of non -metals
- as elements they exist as diatomic molecules F2, Cl2, Br2 and I2 called the halogens
what are the physical properties of group 7
gaseous halogens vary in appearance
e.g.
Fluorine appears clear/ pale yellow gas
Chlorine - a greenish gas
Bromine - a red-brown liquid
iodine - a black solid
they get darker and denser going down the group
they all have a characteristic “swimming - bath” smell
why are a number of fluorines properties untypical
- many of these untypical properties stem from the fact that the F - F bond is unexpectedly weak, compared with the trend fo the rest of the halogen
this is because the small size of the fluorine atom leads to repulsion between non - bonding electrons because they are so close
what is the reactivity of the halogens
F - very reactive, toxic
Cl - very reactive, toxic
Br - very reactive, toxic
I grey crystals, reactive, toxic
what are the clear trends in the halogens
the size of the atom increases as it goes down
the electronegativity of an atom decreases as it goes down
the melting points increases as it goes down
the boiling points increases as it goes down
why does the atom get larger as it goes down a group
atoms get bigger going down the group because each element has one extra filled main level of electrons compared with one above it
what is electronegativity
electronegativity is a measure of the ability of an atom to attract electrons, or electron density, towards itself within a covalent bond
what does electronegativity depend on
electronengativity depends on the attraction between the nucleus and bonding electrons in the outer shell
This in turn, depends on a balance between the number of protons in the nucleus (nuclear charge) and the distance between the nucleus and the bonding electrons plus the shielding effect of inner shells off electrons
what do the shared electrons in hydrogen halides do
the shared electrons in the H-X bond get further away from the nucleus as the atoms get larger going down the group
This makes the shared electrons further from the halogens nucleus and increases the shielding by more inner shells of electrons
why do the melting and boiling points increases going down group 7
increases as you go down the group
This is because larger atoms have more electrons and this makes the van der Waals forces between the molecules stronger
the lower the boiling point …
the lower the boiling point, the more volatile the element
So chlorine which is a gas at room temperature is more volatile than iodine which is a solid
what are the trend in oxidising ability in the halogens
usually react by gaining electrons to become negative ions with a charge of -1
reactions are redox reactions - halogens are oxidising agents known
what are displacement reactions
halogens will react with metal halides in solution in such away that the halide in the compound will be displaced by a more reactive halogen but not by a less reactive
this is called a displacement reaction
give an example of a displacement reaction
e.g. chlorine will displace bromide ions, but iodine will not
in this redox reaction the chlorine is acting as an oxidising agent, by removing electrons from Br- and oxidising 2Br- to Br2 (the oxidation number of the bromine increases from -1 to 0)
why can you not investigate fluorine in an aqueous solution
because it reacts with water
how can you extract bromine from seawater
the oxidation of a halide by a halogen is the basis of a method for extracting bromine from seawater
seawater contains small amounts of bromide ions which can be oxidised by chlorine to produce bromine:
Cl2(aq) + 2Br-(aq) - Br2(aq) +2Cl-
ho do we extract iodine from kelp
iodine was discovered in 1811
It was extracted from kelp, which is obtained by burning seaweed
Some iodine is still produced in this way
Salts such as sodium chloride, potassium chloride, and potassium chloride and potassium sulfate are removed from the kelp by washing wit water
The residue is then heated with manganese dioxide and concentrated sulfuric and iodine is liberated
2I- +MnO2 +4H+ - Mn2+ +2H2O +I2
what are halide ions
halide ions can act as reducing agents
these reactions the halide ions (give away) electrons and become halogens molecules
what is the trend in reducing ability of the halides
it is linked to the size of the ions
the bigger the ion, the more easily it loses an electron
This is because the electron is lost from the outer shell which is further from the nucleus as the ion gest larger so the attraction to the outer electron is less
This trend can be seen in the reactions of solid
Sodium halides with concentrated sulfuric acid
how does sodium halides react with concentrated sulfuric acid
the products are all different and reflect the reducing powers of the halides ions shown above
how does sodium chloride react with concentrated sulfuric acid
in this reaction, drops of concentrated sulfuric acid are added to solid sodium chloride
steamy fumes of hydrogen chloride are seen
the solid product is sodium hydrogensulfate
the reaction is:
NaCl(s) + H2SO4 - NaHSO4(s) + HCl(g)
This is not a redox reaction because no oxidation state has changed
The chloride ion is too weak a reducing agent to reduce the sulfur (oxidation state = +6) in the sulfuric acid
this reaction can be used to prepare hydrogen chloride gas which, because of this reaction was once called a salt gas
similar reaction with sodium fluoride - the fluoride ion is an even weaker reducing agent than the chloride ion
how does sodium bromide react with concentrated sulfuric acid
you will see steamy fumes of hydrogen bromide and brown fumes of bromine
- colourless sulfur dioxide is also formed
Two reactions occur
first sodium hydrogen sulfate and hydrogen bromide are produced (in a similar acid- base reaction to sodium chloride)
NaBr(s) +H2SO4(l) - NaHSO4(s) +HBr(g)
However, bromide ions are strong enough reducig agents to reduce the sulfur acid to sulfur dioxide
The oxidation sate of the sulfur is reduced from +6 to +4 and that of bromine increases from - 1 to 0
This is a REDOX reaction
the reactions are exothermic and some of the bromine vaporiese
how does sodium iodine react with concentrated sulfuric acid
you see steamy fumes of hydrogen iodine, the black solid of iodine, and the bad egg smell of hydrogen sulfide
Yellow solid sulfur may also be seen
colourless sulfur dioxide is also evolved
SEVERAL REACTIONS OCCUR:
- hydrogen iodine is produced in an acid - base reaction as before
- iodine ions are better reducing agents than bromide ions so they reduce the sulfur in sulfuric acid even further (from +6 to 0 and -2) so that sulfur dioxide, sulfur and hydrogen sulfide gas are produced
during the reduction from +6 to -2, the sulfur passes through oxidation state 0 and some yellow, solid sulfur may be seen
how do we identify metal halides with silver ions
all metal halides (except fluorides) react with silver ions in an aqueous solution e.g. in silver nitrate
to form a precipitate of the insoluble silver halide
- silver fluoride does not form a precipitate because it is soluble in water
- Dilute nitric acid HNO3 or (H+ aq) + NO-3(aq) is first added to the halide solution to get rid of any soluble carbonate CO2-(aq) or hydroxide OH- impurities
- these would interfere with the test by forming insoluble silver carbonate - then add a few drops of silver nitrate solution and the halide precipitate forms
why is the test useful for finding halide ions
the reaction can be used as a test for halides because you can tell from the colour of the precipitate which halide has formed
The colours of silver bromide and silver iodide are similar but if you add a few drops of concentrated ammonia solution, silver bromide dissolves but silver iodide does not
what are the halide precipitates
silver fluoride - no precipitate
SILVER CHLORIDE - white precipitate
dissolves in dilute ammonia
SILVER BROMIDE - cream ppt
dissolves in concentrated ammonia
SILVER IODIDE - pale yellow ppt
insoluble i concentrated ammonia
why do we extract titanium
titanium is a very useful metal because it is:
- abundant
- has a low density
- is corrosion resistant
It is useful for making strong, light alloys for use in aircraft for example
how is titanium extracted
titanium is extracted by reaction with a more reactive metal (e.g. Mg)
steps in extracting titanium:
- TiO2 (Solid) is converted to TiCl4 (liquid) at 900C
- The TiCl4 is purified by fractional distillation in an argon atmosphere
- The Ti is extracted by Mg in an argon atmosphere at 500C
why is titanium expensive
the expensive cost of the Mg
This is a batch process which makes it expensive because the process is slower (having to fill up and empty reactors take time) and requires more labour and the energy is lost when the reactor is cooled down after stopping
- the process is also expensive due to the argon, and the need to remove moisture (because TiCl4 is susceptible to hydrolysis)
- high temperature required in both steps
displacement of halides
potassium chloride & chloride - very pale green solution, no reaction
potassium bromide & chlorine - yellow solution, Cl has displaced Br
potassium iodide &chlorie - brown solution CL has displaced I
what is the colour of the solution after a displacement reaction occurred
the colour of the solution in the test tube shows which free halogen is presented in solution
Chlorine - very pale green solution (often colourless)
Bromine = yellow solution
Iodine = brown solution (sometimes a black solid is present)
how do you make bleach
if you mix bleach with cold, dilute sodium hydroxide solution at room temperature, you get sodium chlorate(I) solution (NaClO)
this is a common household bleach that kills bacteria
what is the reaction between cold dilute sodium hydroxide and chlorine gas called
it is called disproportionation
this is because in the reaction chlorine is both oxidised and reduced
write the reaction between dilute sodium hydroxide and Chlorine gas
2NaOH + Cl2 → NaClO + NaCl + H2O
ClO- is the chlorate (I) ion
Chlorine’s oxidation state is +1 in this ion
In NaCl the chlorine oxidation state is -1
what happens when we mix chlorine with water
when you mix chlorine with water, it undergoes disproportionation
You end up with a mixture of chloride ions and chlorate (I) ions
write the equation that occurs between water and chlorine
Cl2 + H2O ⇔ 2H+ + Cl- + ClO-
What happens when chlorine reacts with water in sunlight
Cl2 + H2O ⇔ 2H+ + Cl- + ½O2
why is it good to add chlorine to drinking water
chlorate ions kill bacteria so adding chlorine or a compound that contains chlorate ions to water can make it safe to drink or swim in it
it kills disease-causing microorganisms
it prevents the growth of algae, eliminating bad tastes and smells, and removes discolouration caused by organic compounds
however there are risks to using chlorine to treat water - chlorine gas is very harmful when breathed in - it irritates the respiratory system
why do we still use chlorine to treat water
the benefits outweigh the negatives/ the risks
