Understanding Points
Structure 3.1.1—The periodic table consists of periods, groups and blocks.
Structure 3.1.2—The period number shows the outer energy level that is occupied by electrons.
Elements in a group have a common number of valence electrons.
Structure 3.1.3—Periodicity refers to trends in properties of elements across a period and down a
group.
Structure 3.1.4—Trends in properties of elements down a group include the increasing metallic
character of group 1 elements and decreasing non-metallic character of group 17 elements.
Structure 3.1.5—Metallic and non-metallic properties show a continuum. This includes the trend
from basic metal oxides through amphoteric to acidic non-metal oxides.
Structure 3.1.6—The oxidation state is a number assigned to an atom to show the number of
electrons transferred in forming a bond. It is the charge that atom would have if the compound
were composed of ions.
The periodic table
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Elements on the periodic table go from left to right in increasing atomic number by 1
•
Z and # of e- increase by 1 as you move from left → right
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Period number = outer energy level occupied by electrons
•
Group number = number of valence electrons
The periodic table consists of:
•
7 Periods (rows)
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n of the outermost energy level occupied by e-
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Elements in the same period have varying chemical/physical properties
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18 groups (columns)
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# of valence electrons in outermost energy level
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Elements in the same group have similar chemical/physical properties
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4 blocks (s, p, d, f)
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Elements which have occupied subshells orbitals
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Metal, non-metal, and metalloid regions
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From left → right, metallic character ↓
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Groups 1, 2, and d, f, block contain metals, right-most part of p block contain non-metals, and the in between contain metalloids
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Element Groups
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Group 1 Alkali metals
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Group 17 Halogens
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Group 18 Noble gases
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D-block Transition metals
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F-block 1st row Lanthanoids
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F-block 2nd row Actinoids
Periodicity: repeating pattern of physical and chemical properties across a period /down a group
Trends in Physical Properties | Comparison of Physical and Chemical Properties |
Group 1 and 17
Down Group | Group 1 Alkali Metals | Group 17 Halogens |
Melting Point | ↓ ˙.˙ delocalised e- is further away from nucleus
Metallic bonding ↓
M.P. ↓
“Electrostatic attraction b/w cations and delocalised e-s” | ↑ ˙.˙ molecular size ↑
London dispersion force ↑
M.P. ↑
“Electrostatic attraction temporary b/w non-polar molecules with temporary induced dipoles (due to random shift in e- position)” |
Reactivity | Ionisation energy ↓ ˙.˙ valence e- being ionised is further away from nucleus
More easily lost and ionised to form ionic compound | electronegativity
↓ ˙.˙ e- being accepted is further to nucleus
Less easily accepted to ionise to form ionic compound |
Reactions of Group 1 with water
Reaction with Water | Reactivity | Product | Observation |
Li | Slow | Alkaline solution (LiOH, NaOH, KOH) + H2 gas | Floats on water |
Na | Vigorous | Melts into a small ball (large heat) | |
K | V. Vigorous | Hydrogen gas ignites into lilac flame |
https://www.youtube.com/watch?v=uixxJtJPVXk (Youtube illustration)
Group 1 with halogens | Group 17 - Displacement reaction |
“Salt production” | 2NaBr(aq) + Cl2(g) -> NaCl(aq) + Br2(g)
Cl2 is more reactive than Br2 and displaces Br from Na
*Observation - Solution turns colourless to orange/brown |
Period 3 Oxides
Na2O(s) | MgO(s) | Al2O3(s) | SiO2(s) | P4O10(s) | SO3(g) | |
Acid-base character | Basic
Na2O(s) + H2O(l) → 2NaOH(aq)
MgO(s) + H2O(l) → Mg(OH)2(aq) | Amphoteric
(Insoluble) | Insoluble | Acidic
P4O10(s) + 6H2O(l) → 4H3PO4(aq)
SO3(g) + H2O(l) → H2SO4(aq)
NO2(g) + H2O(l) → HNO3(aq) |
Oxidation state
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Represents the charge on an atom in a compound if it were composed of ions
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Describes the number of electrons shared or transferred when forming a bond
Oxidation state rules:
1.
All elements in their natural form have O.S.= 0
2.
For a given ion with only one atom, O.S.=charge of the ion
a.
e.g. Na+ → +1, Cu2+ → +2, O2- → -2
3.
For a polyatomic ion, the sum of the O.S. of the atoms in the ion = charge of the ion
4.
For a neutral compound, the sum of O.S. of the atoms = 0
5.
Metal O.S. in a compound
a.
Group 1 metals have O.S= +1 in a compound
b.
Group 2 metals have O.S.= +2 in a compound
c.
Transition metals can have varying O.S.
6.
O.S. to know for common atoms in a compound:
a.
Hydrogen O.S. = +1
i.
except when in metal hydrides where O.S.= -1 (e.g. KH, O.S. of H = -1)
b.
Fluorine O.S.= -1
c.
Oxygen O.S. = -2
i.
Except when in OF2 (O.S. of O= +1) and in peroxides (e.g. H2O2, O.S. of O = -1)
d.
Halogens O.S. = -1
i.
Except when in compounds with O, F where O.S. >0
Transition metals have variable O.S. → expressed as roman numerals in parentheses
Compound | Oxidation state | Name (oxidation numbered) |
FeO | Fe (+2) | iron(II) oxide |
Fe2O3 | Fe (+3) | iron(III) oxide |
Cu2O | Cu (+1) | copper(I) oxide |
CuO | Cu (+2) | copper(II) oxide |
MnO2 | Mn (+4) | manganese(IV) oxide |
MnO4– * | Mn (+7) | manganate(VII) ion |
K2Cr2O7* | Cr (+6) | potassium dichromate(VI) |
Cr2O3 | Cr (+3) | chromium(III) oxide |
NO | N (+2) | nitrogen(II) oxide |
NO2 | N (+4) | nitrogen(IV) oxide |
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no. of valence e-s with assumption that shared electron pair in a covalent bond belongs entirely to a more electronegative element, e.g. in S2O32-, 0 & +4 or -2 & +6 O.S. for average value of of +2