Chapter 10 Periodicity Flashcards
Properties of the Elements in Period 3: Atomic radius
- Across the period, the atomic radii decrease
- Note that radii is the plural of radius
Properties of the Elements in Period 3: Ionic radius
- Metals produce positively charged ions (cations) whereas nonmetals produce negatively charged ions (anions)
- The cations have lost their valence electrons which causes them to be much smaller than their atoms
- Because there are fewer electrons, this also means that there is less shielding of the outer electrons
- Going across the period from Na+ to Si4+ the ions get smaller due to the increasing nuclear charge attracting the outer electrons in the second principal quantum shell nucleus (which has an increasing atomic number)
- The anions are smaller than their original atoms because each atom has gained one or more electrons in their third principal quantum shell
- This increases the repulsion between electrons while the nuclear charge is still the same
- Going across P3- to Cl– the ionic radii decreases as the nuclear charge increases across the period and less electrons are gained by the atoms (P gains 3 electrons, S 2 electrons and Cl 1 electron)
Atomic radius
is the distance between the nucleus and the outermost electron of an atom
-In metals this is also called the metallic radius and in non-metals, the covalent radius
how is the atomic radius measured
The atomic radius is measured by taking two atoms of the same element, measuring the distance between their nuclei and then halving this distance
Ionic radius
is the distance between the nucleus and the outermost electron of an ion
Melting points(k) of the elements across Period 3 table
84(k)
structure of Period 3 elements table
- Na= Giant metallic
- Mg= Giant metallic
- Al= Giant metallic
- Si=Giant molecular
- P=Simple molecular
- S= Simple molecular
- Cl=Simple molecular
- Ar= Simple molecular
Bonding of Period 3 elements table
- Na= metallic
- Mg=metallic
- Al=metallic
- Si= Covalent
- P=Covalent
- S=Covalent
- Cl=Covalent
- Ar= -
what do Na, Mg and Al have in common
are metallic elements which form positive ions arranged in a giant lattice in which the ions are held together by a ‘sea’ of delocalised electrons around them:
- The electrons are free to move around and are not bound to an atom
- The electrons in the ‘sea’ of delocalised electrons are those from the valence shell of the atoms
- Na will donate one electron into the ‘sea’ of delocalised electrons, Mg will donate two and Al three
- As a result of this, the metallic bonding in Al is stronger than in Na
why do melting points increase from Na - Al
This is because the electrostatic forces between a 3+ ion and the larger number of negatively charged delocalised electrons are much larger compared to a 1+ ion and the smaller number of delocalised electrons in Na
why does Si have the highest melting point in period 3
due to its giant molecular structure in which each Si atom is held to its neighbouring Si atoms by strong covalent bonds
P, S, Cl and Ar are non-metallic elements and exist as
simple molecules
- The covalent bonds within the molecules are strong, however between the molecules there are only weak instantaneous dipole-induced dipole forces
- It doesn’t take much energy to break these intermolecular forces
- Therefore, the melting points decrease going from P to Ar
Electrical conductivity
refers to how well a substance can conduct electricity
Properties of the Elements in Period 3: Electrical conductivity
- Na = 0.218
- Mg= 0.224
- Al=0.382
- Si= 2 x 10^-10
- P=10^-17
- S= 10^-23
- Cl = -
- Ar = -
going from Na to Al, there is an increase in the number of valence electrons that are donated to the ‘sea’ of delocalised electrons
Because of this
here are more electrons available in Al to move around through the structure when it conducts electricity, making Al a better electrical conductor than Na
-Due to the giant molecular structure of Si
there are no delocalised electrons that can freely move around within the structure
- Si is therefore not a good electrical conductor and is classified as a semimetal (metalloid)
- The lack of delocalised electrons is also why P and S cannot conduct electricity
delocalised means
that the electrons are free to move around and are not bound to an atom
metallic bonding in Al is stronger than in Na because of
- The electrons in the ‘sea’ of delocalised electrons are those from the valence shell of the atoms
- Na will donate one electron into the ‘sea’ of delocalised electrons, Mg will donate two and Al three electrons
Variation in First Ionisation Energy
-The first ionisation energy (IE1) is the energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions
—-Eg. the first ionisation energy of Na is:
Na(g) → Na+(g) + e–
IE1 values of the Period 3 elements table (IE1(KJ mol^-1)
- Na 494
- Mg=736
- Al=577
- Si=786
- P=1060
- S=1000
- Cl=1260
- Ar=1520
There is a general increase in IE1 across a period. why
- The nuclear charge increases
- The atomic radius decreases
- There are stronger attractive forces between the nucleus and outer electrons
- It therefore gets harder to remove any electrons
- Small ‘dips’ are observed between Mg – Al and P – S
Ceramics
is a rigid material that is made up of an infinite 3D network of sintered metals bonded to carbon, nitrogen or oxygen
Sintering is when a
- a powdered material is heated below its melting point
- This results in the formation of new bonds between the powder grains to form one large mass
Examples of common ceramics are
magnesium oxide, aluminium oxide and silicon dioxide
-Though silicon is a non-polar silicon dioxide it is still considered a ceramic
-These compounds all have giant ionic (magnesium oxide and aluminium oxide) or giant covalent (silicon dioxide) structures
-There are strong bonds and electrostatic forces between the atoms or ions that hold the giant lattice structures together
This affects the physical properties of ceramics
Physical properties of ceramics
are very hard and strong but brittle
—-This is due to the strong bonds and forces that hold the atoms and ions together in the 3D structure
—-However, if the 3D structure is distorted the ceramic will easily break
-The presence of strong covalent or ionic bonds in the structure also means that ceramics have high melting points
—-A lot of energy is required to break these bonds
-Ceramics are good electrical insulators as they do not conduct electricity
—-Covalently bonded ceramics have no delocalised electrons that can move around
Ionic bonded ceramics have no ions that can freely move around
Applications of ceramics: Magnesium oxide
- Furnace linings due to their high melting points
- Heating elements – for example, electrical cookers as ceramics are good electrical insulators
Applications of ceramics: Aluminium oxide
- High temperature and voltage electrical insulators due to its high melting point and electrical insulating properties
- Replacement hip joints as aluminium oxide is highly durable
Applications of ceramics: silicon dioxide
Furnace linings due to the high melting point of the giant covalent structure
Reaction of sodium with water
-Sodium reacts vigorously with cold water:
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
- The sodium melts into a ball and moves across the water surface until it disappears
- Hydrogen gas is given off
- The solution formed is strongly alkaline (pH 14) due to the sodium hydroxide which is formed
Reaction of Magnesium with cold water
-Magnesium reacts extremely slowly with cold water:
Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H2(g)
-The solution formed is weakly alkaline (pH 11) as the formed magnesium hydroxide is only slightly soluble
-When magnesium is heated, it reacts vigorously with steam (water) to make magnesium oxide and hydrogen gas:
Mg(s) + H2O(g) → MgO(s) + H2(g)
Oxidation Number of the Period 3 Oxide
Oxygen is more electronegative than any of the Period 3 elements
-The Period 3 elements therefore have positive oxidation states in their oxides and the oxygen has a negative oxidation state of -2
The Pauling scale shows
the electronegativities of the elements in the periodic table. Oxygen has a higher electronegativity than any of the Period 3 elements which is why the Period 3 elements will have positive oxidation states and the oxygen a negative oxidation state in the oxides of Period 3 elements
Reaction of Period 3 Oxides & Water
look in book
Aluminium oxide is amphoteric
which means that it can act both as a base (and react with an acid such as HCl) and an acid (and react with a base such as NaOH
Acidic & basic nature of the Period 3 oxides
- Na2O=basic
- MgO=basic
- Al2O3=amphoteric
- SiO2=acidic
- P4O10=acidic
- SO2, SO3 = acidic
Reactions of the Period 3 oxides with acid/base table: Na20
Na2O(s) + 2HCL(aq) -> 2NaCl(aq) + H2O(l)
Reactions of the Period 3 oxides with acid/base table: MgO
MgO(s) + 2HCL(aq) -> MgCl2(aq) + H2O(l)
used in indigestion remedies by neutralizing the excess acid in the stomach
Reactions of the Period 3 oxides with acid/base table: Al2O3
=Al2O3(s) + 3H2SO4(aq) -> Al2(SO4)3(aq) + 3H2O(l)
(reacts with acid to form a salt and water)
=Al2O3(s) + 2NaOH(aq) -> 3H2O(l) + 2NaAl(OH)4(aq)
(reacts with hot, concentrated alkali to form a salt)
Reactions of the Period 3 oxides with acid/base table: SiO2
=SiO2(s) + 2NaOH(aq) -> Na2SiO3(aq) + H2O(l)
reacts with hot, concentrated alkali to form a salt and water
Reactions of the Period 3 oxides with acid/base table: P4O10
P4O10(s) + 12NaOH(aq) -> 4Na3PO4(aq) + 6H2O(aq)
Reactions of the Period 3 oxides with acid/base table: SO2,SO3
SO2(g) + 2NaOH(aq) -> Na2SO3(aq) + H2O(l)
SO3(g) + 2NaOH(aq) -> Na2SO4(aq) + H2O(l)
Period 3 hydroxide: Al(OH)3
is amphoteric and can acts both as an acid and base
Al(OH)3(s) + 3HCl(aq) → AlCl3(s) + 3H2O(l)
Al(OH)3(s) + NaOH(aq) → NaAl(OH)4(aq)
Period 3 hydroxide: NaOH
is a strong base and will react with acids to form a salt and water:
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
Period 3 hydroxide: Mg(OH)2
is also a basic compound which is often used in indigestion remedies by neutralising the excess acid in the stomach to relieve pain:
Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + 2H2O(l)
The oxides of P and S which show purely covalent bonding produce
acidic solutions with water because when these oxides react with water, they form an acid which donates H+ ions to water
—-Eg. SO3 reacts with water as follows:
SO3(g) + H2O(l) → H2SO4(aq)
—-The H2SO4 is an acid which will donate a H+ to water:
H2SO4(aq) + H2O(l) → H3O+ (aq) + HSO4–(aq)
what behaviour is very typical of a covalently bonded oxide
are insoluble and when they react with hot, concentrated alkaline solution they act as a base and form a salt for example Ai and Si
what behaviour is very typical of an ionic bonded metal oxide
ionic bonded metal oxide can also react with acidic solutions to form a salt and water example:Al
-This behaviour of Al proves that the chemical bonding in aluminium oxide is not purely ionic nor covalent: it is amphoteric
Reaction of Period 3 Chlorides & Water: observations and pH of NaCl and MgCl2
white solids dissolve to form colourless solutions with pH of 7 for NaCl and 6.5 for MgCl2
Reaction of Period 3 Chlorides & Water: observations and pH of Al2Cl6, SiCl4, PCl5 and SCl2
Chlorides react with water giving off white fumes of hydrogen chloride gas ph= -Al2Cl6=3 -SiCl4=2 -PCl5=2 -SCl2=2
Sodium & magnesium chloride
NaCl and MgCl2 do not react with water as the polar water molecules are attracted to the ions dissolving the chlorides and breaking down the giant ionic structures: the metal and chloride ions become hydrated ions
Aluminium chloride exists in two forms:
- AlCl3 as a giant lattice and with ionic bonds
- Al2Cl6 as a dimer with covalent bonds
When water is added to aluminium chloride
the dimers are broken down and Al3+ and Cl– ions enter the solution
- The highly charged Al3+ ion becomes hydrated and causes a water molecule that is bonded to the Al3+ to lose an H+ ion which turns the solution acidic
- The H+ and the Cl– form hydrogen chloride gas which is given off as white fumes
Silicon chloride is hydrolysed in water
releasing white fumes of hydrogen chloride gas in a rapid reaction
SiCl4(l) + 2H2O(l) → SiO2(s) + 4HCl(g)
-The SiO2 is seen as a white precipitate and some of the hydrogen chloride gas produced dissolves in water to form an acidic solution
Phosphorus(V) chloride is hydrolysed in water
-PCl5 also gets hydrolysed in water
PCl5(s) + 4H2O(l) → H3PO4(aq) + 5HCl(g)
-Both H3PO4 and dissolved HCl are highly acidic
Electronegativity
is the power of an element to draw the electrons towards itself in a covalent bond
-Going across the period, the electronegativity of the elements increases
Electronegativity across Period 3 table
- Na=0.9
- Mg=1.2
- Al-1.5
- Si=1.8
- P=2.1
- S=2.5
- Cl=3.0
- Ar= -
Electronegativity across Period 3 according to atomic number
As the atomic number increases going across the period, there is an increase in nuclear charge
Across the period, there is an increase in the number of valence shells however the shielding is still the same as each extra electrons enters the same shell
As a result of this, electrons will be more strongly attracted to the nucleus causing an increase in electronegativity across the period
If the chemical and physical properties of an element are known
the position of that element in the Periodic Table can be predicted
Similarly, predictions can be made about the physical and chemical properties of elements if the position of the element in the Periodic Table is known
