Chemistry Unit 3 Periodicity
2/21/2012 1:35:00 AM
Periodicity is the repeating of similar physical and chemical properties of
an element. means that it varies in a repetitive way as you move through
the Periodic Table
Chemical and Physical properties of elements are periodic functions of
their number of protons.
Group no. is equal to number of valence (outer shell) electrons
Elements in same group share same no. of valence electrons
Elements in same group share physical and chemical properties
Period no. is equal to no. of main energy shells (orbits) filled by
electrons in neutral atom of element
Elements in same period have different no. of valence electron but are all
contained in same shell as other elements in period
Atomic Radii: Size of an atom. Radii increases down a group, and
decreases across a period (ignoring noble gases). Difficult to measure as
precise boundary is unknown. Measured by determining distance between
nuclei of molecules and dividing it by amount of atoms in molecule
(generally two as they should be molecule of the same element i.e. H2)
It can be considered as the distance from the nucleus to the
outermost electrons of the Bohr atom.
Ionic Radius: Radius of an ion. Generally smaller than atomic radius
when it is a cation (positive ion).
Isoelectronic: Sharing electronic configuration but not proton number. i.e.
Sodium cation Na+ and Neon Ne are configured as 2,8
More positive charge equals smaller radius.
Anions (negative ions) are generally larger than corresponding atom.
Greater negative charge equals greater radius.
Ionic radius increases down a group.
Electronegativity is power of an atom to attract electrons to it during
bonding. Top right hand of periodic table is high electronegativity, bottom
left is low electronegativity.
Moving Across the Period, we see an increase in effective nuclear charge,
resulting in smaller atomic radii and increased Ionization Energies.
Moving down a group, we see a decrease in first ionization energy due to
greater atomic radii.
Melting point of an element is a reflection of how strongly it bonds to
other atoms and how strong its attractive force is. High melting
temperatures entail stronger bonds than lower temperatures
Melting point decrease down a group as bonds get weaker and size
increases.
First Ionization Energy is the energy needed to remove the most loosely
held (outermost) electron from each of one mole of gaseous atoms:
X (G)
X+(G) + eIonization energy is a measure of the energy needed to pull a particular
electron away from the attraction of the nucleus. A high value of
ionization energy shows a high attraction between the electron and the
nucleus.
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.
Attraction is Governed by:
Charge of Nucleus:
More protons = greater positive charge = greater electrostatic effect
The Distance of the electron from the Nucleus:
Attraction falls off very rapidly with distance. Electron closer to the
nucleus is much more strongly attracted than one further away.
General Trend for Ionization Energies:
Ionization energies increase across a period.
Removal of Successive Electrons:
X
(G)
X2+(G) + e-
X
(G)
X3+(G) + e-
Energy necessary for loss of successive electrons increases progressively.
The first four ionisation energies of aluminium, for example, are given by
1st I.E. = 577 kJ mol-1
2nd I.E. = 1820 kJ mol-1
3rd I.E. = 2740 kJ mol-1
4th I.E. = 11600 kJ mol-1
A huge jump in energy indicates breaking into a new shell and can be
used to identify group number of species. Here the jump occurs after the
3rd ionization indicating its belonging to Group 3.
In order to form an Al3+(g) ion from Al(g) you would have to supply:
577 + 1820 + 2740 = 5137 kJ mol-1
Metallic Bonding occurs between Groups 1,2,3
Electrostatic attraction between Metal Cation and its delocalized negative
electron holds metal together.
High melting point suggest strong bonds between metals
Valence electrons become delocalized
Ionic Bonding occurs between Metals and Non-Metals (Electrostatic
Attraction between Positive and Negative Ions)
Covalent Bonding occurs between Two Non-Metals (Involves sharing of
Electrons)
The two chlorine atoms are
said to be joined by a covalent
bond. The reason that the two
chlorine atoms stick together
is that the shared pair of
electrons is attracted to the
nucleus of both chlorine
atoms.
Most Covalent bonds exist as Di-Atomic Molecules like O2, H2
An example of an exception is:
The hydrogen has a helium structure, and the
chlorine an argon structure.
CH4
NH3
H2O
Trends in Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a
bonding pair of electrons. Measured on Pauling scale.
The most
electronegative element is
fluorine. If you remember
that fact, everything
becomes easy, because
electronegativity must
always increase towards
fluorine in the Periodic
Table.
As you go across a period electronegativity increases.
Going down a group decreases whilst going up a group increases
electronegativity.
Metallic bonds are stronger. Atomic radii increases down a group for
metals. Electronegativity for metals is weaker and therefore melting point
is lower as you descend a group
Covalent bonds are weaker than metallic bonds. Dispersion force
increases as you go down a group, increasing melting point.
Electronegativity is stronger. Atomic radii increases down a group
Increased reactivity means decreased electronegativity and vice versa
Na + 2H2O = 2NaOH + H2
Na + O2 = Na2O
Halogens Are Salt Formers
Na + Cl2 = NaCl Metal + Halogen = Metal Halide
Metal + Flourine = Metal Flouride Na +F2 = NaF
Metal + Bromine = Metal Bromide Li + Br2 = LiBr
Metal + Iodine = Metal Iodide K + I2 = KI
2Na + 2HCl = 2NaCl + H2
Li + H2SO4 = Li2SO4 + H2
When bonded to an atom with higher electronegativity, an atoms
oxidization number can change
Strength of Acid is determined by rate of disassociation.
Hydrogen + Halogen = Hydrogen Halide H2 + F2 = HF Weak Acid does not
disassociate in water
H2 + Cl2 = HCl Strong Acid
H2 + Br2 = HBr
H2 + I2 = HI
H + Cl (in solution) = H+ + Cl- disassociate/ionize
No Precipate indicates FAgCl is white precipate. Indicates Cl.
AgBr is a creamy precipate. Indicates Br.
AgI is a yellow precipate. Indicates I.
Silver Nitrate is Ag+NO3- (Aq)
Displacement Reactions
Hydrogen
Halide
Cl2
ClBr2
BrI2
ICompounds conduct electricity if they have delocalized electrons or are
have an electric charge.
Alkali Earth Metals do not exist in a pure form
MgO (electrolysis) = Mg + O2
Chemistry Unit 3 Periodicity
2/21/2012 1:35:00 AM
3.1.1 Describe the arrangement of elements in the periodic table
in order of increasing atomic number
Elements are arranged in the periodic table in order of their increasing
atomic number (Z) or number of protons, which is a fundamental
property any element.
Groups 1, 2 and 3 are Metals. (Easier for them to lose electrons)
Group 5, 6, 7 and 8 are Non-Metals. (Easier for them to gain electrons)
Metalloids react differently according to what they are reacting with.
3.1.2 Distinguish between the terms group and period
Groups are columns of the periodic table whilst periods are rows of the
periodic table.
Groups are numbers represented by roman numerals at the top of the
periodic table. This number refers to the number of valence electrons in
the atom which determines the chemical properties of the element.
Period numbers represent the number of energy levels or shells of the
atom.
3.1.3 Apply the relationship between the electron arrangement of
elements in their position in the periodic table up to Z = 20
Electron arrangement indicates in which period an element sits. If it has
three occupied energy levels like sodium, it will be in period 3. This is true
for all of the first twenty elements. Therefore, if an unknown element of
z=18 needed to be located, one could a certain its position as Group 8 or
0, and Period 3. As it would have the configuration 2, 8, 8, this indicates
three energy levels being occupied, hence it will be in period 3, as well as
8 valence electrons (or 0 as it is a noble gas), indicating group 8 or 0.
3.1.4 Apply the relationship between the number of electrons in
the highest occupied energy level for an element and its position
in the Periodic Table
The number of electrons in the highest occupied energy level (valence
electrons) indicates in which group the element is positioned. Therefore is
an element has 2 electrons in its highest occupied energy level, or two
valence electrons, it will be in Group 2.
Chemistry Unit 3 Periodicity
2/21/2012 1:35:00 AM
3.2.1 Define the terms first ionization energy and
electronegativity.
First Ionization Energy is the energy needed to remove the most loosely
held (outermost) electron from each of one mole of gaseous atoms or
from an isolated gaseous atom:
X (G)
X+(G) + eIt is a measure of the attraction between the nucleus and the outer
electrons.
Electronegativity is power of an atom to attract electrons in a covalent
bond. It is related to ionization energy as it is also a measure of the
attraction between the nucleus and its outer electrons, which in this case
are bonding electrons. Top right hand of periodic table is high
electronegativity, bottom left is low electronegativity.
3.2.2 Describe and explain the trends in atomic radii, ionic radii,
first ionization energies, electronegativities and melting points for
the alkali metals (Li -> Cs) and the halogens (F -> l)
Trends in physical properties are governed by the effective nuclear
charge, which increases left to right across a period due to increased
nuclear charge and no change in inner electron configuration, whilst
remaining approximately the same down a group as the increase charge
is offset by additional electrons.
The Atomic radii for alkali metals and halogens increases down a group as
a result of the addition of occupied energy levels which results in no
change in effective nuclear charge, therefore highlighting weaker
attraction between nucleus and outer electrons, allowing them to move
further away, enlarging the atom.
Atomic Radii decreases across a period due to increased effective nuclear
charge, which strengthens the attraction between the nucleus and
outermost electrons, causing them to be pulled closer to the nucleus and
shrinking the atom.
Ionic radii for alkali metals and halogens increase down a group as
additional energy levels occupied by electrons weakens electrostatic
attraction between nucleus and the outermost electrons, allowing them to
move further away, enlarging the atom.
Ionic radii decrease across a period from group 1 to group 4, as the loss
of electrons to form positive ions increases the effective nuclear charge on
the outermost electrons, pulling them closer in to the nucleus, shrinking
the atom. Ionic radii increase across a period from group 4 to 7, as the
gain of electrons to form negative ions results in decreased effective
nuclear charge, allowing the electrons to move further away from the
nucleus, enlarging the atom.
First Ionization Energies increase across a period as the result of
increased effective nuclear charge, which strengthens electro static
attraction and requires more energy to remove an electron from the
atom.
First Ionization Energies decrease down a group as the result of a lack of
any significant or relative change in effective nuclear charge makes it
progressively easier for electrons that are further away from the nucleus
to be removed from the atom.
Electronegativity increases left to right across a period as the result of
increased nuclear charge, which effectively increases the atoms ability to
attract electrons.
Electronegativity decreases down a group as a result of the lack of change
in effective nuclear charge, which makes it progressively easier for
electrons to repel the electrostatic forces from the nucleus.
Melting Points decrease down Group 1, as the metallic bonds grow weaker
as the increase in size of the atoms means nuclei are further away from
the delocalized electrons and become progressively more easier to break.
Melting points increase down Group 7, as the covalent bonds of the
diatomic molecules become progressively stronger, requiring more energy
to break them down.
Melting points increase across a period from Group 1 to 3, as the number
of delocalized electrons in the metallic bonds increase, strengthening the
bonds and requiring more energy to break them down. It continues to
increase across to Group 4, as the giant covalent bonds of the metalloids
are extremely difficult to break down.
Melting points then decrease across a period from Group 5 – 8/0, as the
non-metallic structure of these atoms as the simple molecules have
progressively weaker bonds until we reach Group 8 which does not bond
at all, indicating that little energy is required to melt them.
3.2.3 Describe and explain the trends in atomic radii, ionic radii,
first ionization energies and electronegativities for elements
across Period 3.
Ignoring Argon which is a noble gas, Atomic radii tends to decrease
across Period 3 elements as a result of increased effective nuclear charge,
which strengthens the attraction between the nucleus and outermost
electrons, causing them to be pulled closer to the nucleus and shrinking
the atom.
Ionic radii decrease across Period 3 from group 1 to group 4, as the loss
of electrons to form positive ions increases the effective nuclear charge on
the outermost electrons, pulling them closer in to the nucleus and
shrinking the atom. Ionic radii increase across Period 3 from group 4 to
7, as the gain of electrons to form negative ions results in decreased
effective nuclear charge, allowing the electrons to move further away
from the nucleus, enlarging the atom.
First Ionization Energies increase across Period 3 as the result of
increased effective nuclear charge, which strengthens electro static
attraction and requires more energy to remove an electron from the
atom. A fall is experienced at Aluminium as the 3p1 orbital, which is
occupied by only one electron, is relatively more easier to remove than
the 3s2 orbital before it. Another fall is experienced at Sulphur, where the
repulsion between the two pairs of electrons in the 3p orbital means less
energy is required to remove the electron from the atom.
Electronegativities increase across Period 3, ignoring the noble gas Argon,
as the result of increased effective nuclear charge due to increased
nuclear charge and no offsetting change in the configuration of the atoms
inner electrons.
3.2.4 Compare the relative electronegativity values of two or more
elements based on their positions in the Periodic Table.
The trend for electronegativity is that is increases going to the right and
up. Therefore the element that is closer to the top right hand side of the
periodic table will have a higher relative electronegativity than an atom
that is closer to the bottom left hand side of the periodic table.
2/21/2012 1:35:00 AM
3.3.1 Discuss Similarities and differences in the chemical
properties of elements in the same group.
Alkali metals are very reactive as they can lose their single valence
electron easily – they always form M+ ions.
The reactivity increases down groups 1 as the atomic radius
increases and the ionization energy decreases so it becomes easier to lose
the valence electron.
They react vigorously (exothermic) with water to form a metal
hydroxide solution and hydrogen gas, the resulting solution is alkaline:
2M(s) + 2H2O(l) → 2M+(aq) + 2OH-(aq) + H2(g)
Alkali metals react easily with the halogens to form an ionic
compound:
Ex. 2Na(s) + Cl2(g) -> 2NaCl(s)
The halogens are very reactive as they only require one electron to
complete their valence shell and form a Hal- ion.
The halogens all exist as diatomic molecules joined by a covalent
bond.
Reactivity decreases down the group as the valence shell is further
from the nucleus so it becomes harder to attract an electron.
The electronegativity and oxidizing power of the halogens decreases
down the group as the atomic radius increases and they are not so able to
attract electrons. This means that a higher halogen will displace a lower
halogen from its salts,
Ex. Cl2 will oxidize iodide ions to form I2 and Cl- ions:
Cl2 + 2I- → I2 + 2Cl- the reverse reaction would not occur.
There are often color changes accompanying halogen reactions: Br2
is orange-brown but Br- is colorless; I2 is violet in a non-polar solvent
and yellow-brown in water but I- is colorless.
Chemical properties of an element are largely determined by its electron
configuration and the number of electrons in the elements outer shell.
Elements of the same group have similar chemical properties as they
have the same number of electrons in their outer shells.
Elements of the same group however, have differing levels of reactivity to
the same substances. Examples can be seen with the Group 1 Alkali
metals. Lithium reacts slowly with water, releasing hydrogen and keeping
its shape. Sodium reacts more violently, producing heat that further melts
the unreacted metal. Potassium is more violent than both of these,
releasing enough heat to ignite the subsequent hydrogen gas and
producing a lilac coloured flame. The result of all these reactions between
Alkali metals and water is the production of an alkaline solution.
Further examples occur when Alkali metals react with Halogens in Group
7. Sodium can react with chlorine to produce sodium chloride, which is
table salt. Potassium reacts with Iodine to produce potassium Iodide.
Lithium reacts with Bromine to produce Lithium Bromide. These reaction
all produce salt.
Similarly, Halogens can react with Halide ions from the same group in
Displacement reactions. The more electronegative of the two elements
will displace the less electronegative element and gain the electron,
forming an ion of a Halogen known as a Halide. The more reactive
halogen displaces the ions of the less reactive halogen.
These reactions highlight that elements in the same group will react
similarly, but in progressively more violent or reactive natures as we
progress through the group (Down group 1, Up group 7), due to differing
electronegativities..
3.3.2 Discuss the changes in nature from ionic to covalent and
from basic to acidic, of the oxides across Period 3.
Ionic nature decreases across the period: Na2O and MgO are ionic,
Al2O3 has ionic bonding with some covalent character (could be
considered giant covalent), SiO2 has giant covalent structure and P2O5,
SO3 and Cl2O7 are covalently bonded in molecules (simple covalent).
Trend is from basic to acidic across the period: Na2O and MgO are
basic, Al2O3 is amphoteric, and the rest are acidic and become more
strongly acidic on the right of the period.
Na2O(s) + H2O(l) → 2Na+(aq) + 2OH-(aq)
MgO(s) + H2O(l) → Mg2+(aq) + 2OH-(aq)
P4O10(s) + 6H2O(l) → 4H+(aq) + 4H2PO4-(aq) ←→ 4H3PO4(aq)
SO3(s) + H2O(l) → H+(aq) + HSO4-(aq) ←→ H2SO4(aq)
Oxides of metals are ionic and basic. Oxides of the non-metals are
covalent and acidic. Oxides of some elements in the middle of the Periodic
Table are amphoteric. Amphoteric oxides show both acidic and basic
properties. Alkalis are bases which are soluble in water. They form
hydroxide ions in aqueous solutions. High melting points are associated
with ionic or covalent giant structures whilst low melting points are
associated with molecular covalent structures.
All elements across Period 3 react with oxygen to form various types of
compounds. Metals will form ionic compounds whilst non-metals will form
covalent molecules. The ionic nature of the compounds decreases from
left to right across the period. The first two elements, Sodium and
Magnesium, will form Ionic Compounds and produce strong Alkaline
Solutions.
Aluminium is amphoteric, indicating its ability to react as a base and/or
an acid. It will react with an Acid to produce a salt solution, much like
Sodium and Magnesium, However it will also react with a base to produce
an Aluminate.
These reactions are due to the elements ionic nature as metals.
Silicon will not react with acids or water, but will form a silicate solution
when it reacts with a base. This is due to its covalent nature.
SiO2(s) + 2OH-(aq) => SiO32-(aq) + H2O(l)
Phosphorus(III) oxide will react with water to produce phosphorous or
phosphonic acid. (Phosphoric III Acid)
Phosphorous (V) oxide will produce Phosphoric (V) Acid
Sulphur Dioxide will react with water to produce Sulphurous Acid
(Sulphuric IV Acid)
Sulphur Trioxide will react with water to produce Sulphuric Acid (Sulphuric
VI Acid)
Dichlorine Heptoxide or Chlorine VII Oxide will react with water to
produce Chloric (VII) Acid also known as Perchloric acid.
Dichlorine Monoxide reacts with water to produce Chloric (I) Acid
Cl2O(l) + H2O(l) => 2HClO(aq)