Group 1 elements- Alkali metals:
Lithium(Li) Sodium(Na) Potassium(K) Rubidium(Rb) Caesium(Cs) Francium(Fr)
~All alkali metals are silver-colored, soft, low density metals, high electropositivity, good conductors of heat and electricity.
Atomic Size -increases gradually
Reason: The number of shells occupied with electron increases.
Density -increases gradually
----------Float on wate---------I--------Sink in water------------
Reason: The increase in atomic mass is bigger than the increase in volume.
Boiling and melting points -decreases gradually
Reason: As the atomic size increases, the metallic bonding between the atoms of alkali metals becomes weaker. Hence, less heat energy is required to overcome the weaker metallic bonding during melting or boiling.
Hardness -softer
Reactivity -increases
Reason: As the atomic size increases, the single valence electron becomes futher away from the nucleus. This causes the single valence electron can to be more weakly pulled by the nucleus. Thus, the single valence electron can be release more easily to achieve a stable electron arrangement.
Electropositivity -increases
Reactions of alkali metals with water:
Alkali metal + water → Alkali metal hydroxide + hydrogen gas
Example:
2K (s) + 2H2O (l) → 2KOH (aq) + H2 (g)
Reactions of alkali metals with oxygen:
Alkali metal + oxygen → Alkali metal oxide
Example:
4K (s) + O2 (g) → 2KO2H (s)
Reactions of alkali metals with chlorine and bromine:
Alkali metal + chlorine → Alkali metal chloride
Example:
2K (s) +Cl2 (g) → 2KCl (s)
Alkali metal + bromine → Alkali metal bromide
Example:
2K (s) +Br2 (g) → 2KBr (s)

Group 2 elements- Alkali earth metals:
Beryllium (Be) Magnesium (Mg) Calcium (Ca) Strontium (Sr) Barium (Ba) Radium (Ra)
The alkaline earth metals are silvery colored, soft, low density metals, which react readily with halogens to form ionic salts, and with water, though not as rapidly as the alkali metals, to form strongly alkaline (basic) hydroxides. For example, where sodium and potassium react with water at room temperature, magnesium reacts only with steam and calcium with hot water:

Mg + 2H2O → Mg(OH) 2 + H2

Beryllium is an exception: It does not react with water or steam, and its halides are covalent.

All the alkaline earth metals have two electrons in their outermost shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions.

Group 3 to group 12 elements- Transition metals:
There are several common characteristic properties of transition elements:
They often form colored compounds.
Example:
Compound of transition elements Colour
Cobalt chloride crystal Pink
Copper(II) sulphate crystal Blue
Iron(II) sulphate crystal Pale yellow
Iron(III) sulphate crystal Brown
They can have a variety of different oxidation states.
Example:
Transition metals Oxidation number in compounds
Iron +2, +3
Nickel +2, +3
Copper +1, +2
Manganese +2, +3, +4 , +6, +7
Chromium +2, +3,
At least one of their compounds has an incomplete d-electron subshell.
They are often good catalysts.
Example:
Nickel acts as catalyst in the hydrogenation of alkene to form the corresponding alkane.
Cn H2n+H2--Ni→ Cn H2n+2
They are silvery-blue at room temperature (except copper and gold).
They are solids at room temperature (except mercury).
They form complex ions (aqua ones included).
Example:
Tetraamminecopper(II) ion,[Cu(NH3)4]2+
They are often paramagnetic.

Poor metals
The trivial name poor metals (or post-transition metals) is sometimes applied to the metallic elements in the p-block of the periodic table. Their melting and boiling points are generally lower than those of the transition metals and their electronegativity higher, and they are also softer. They are distinguished from the metalloids, however, by their significantly-greater boiling points in the same row.
"Poor metals" is not a rigorous IUPAC-approved nomenclature, but the grouping is generally taken to include aluminium, gallium, indium, tin, thallium, lead and bismuth; germanium, antimony and polonium are occasionally included, although these are usually considered to be metalloids or "semi-metals". Elements 113 to 116, which are currently allocated the systematic names ununtrium, ununquadium, ununpentium and ununhexium, would likely exhibit properties characteristic of poor metals; however as yet insufficient quantities of them have been synthesized to examine their chemical properties.

Nonmetals
The nonmetals are generally to:
Hydrogen (H)
In Group 14: Carbon (C)
In Group 15 (the pnictogens): Nitrogen (N), Phosphorus (P)
Several elements in Group 16, the chalcogens: Oxygen (O), Sulfur (S), Selenium (Se)
All elements in Group 17 - the halogens
All elements in Group 18 - the noble gases
Common properties considered characteristic of a nonmetal include:
poor conductors of heat and electricity when compared to metals
they form acidic oxides (whereas metals generally form basic oxides)
in solid form, they are dull and brittle, rather than metals which are lustrous, ductile or malleable
usually have lower densities than metals
they have significantly lower melting points and boiling points than metals
non-metals have high electronegativity

Semi-metals- Metalloids
There is no rigorous definition of the term, however the following properties are usually considered characteristic of metalloids:
metalloids often form amphoteric oxides.
metalloids often behave as semiconductors (B,Si,Ge) to semimetals (eg. Sb).
The concepts of metalloid and semiconductor should not be confused. Metalloid refers to the properties of certain elements in relation to the periodic table. Semiconductor refers to the physical properties of materials (including alloys, compounds) and there is only partial overlap between the two.
Some allotropes of elements exhibit more pronounced metal, metalloid or non-metal behavior than others. For example, for the element carbon, its diamond allotrope is clearly non-metallic, however the graphite allotrope displays limited electric conductivity more characteristic of a metalloid. Phosphorus, tin, selenium and bismuth also have allotropes which display borderline behavior.

In the standard layout of the periodic table, metalloids occur along the diagonal line through the p block from boron to astatine. Elements to the upper right of this line display increasing nonmetallic behaviour; elements to the lower left display increasing metallic behaviour. This line is called the "stair-step" or "staircase." The poor metals are to the left and down and the nonmetals are to the right and up. In addition, the halogens are found at the bottom right.

Lanthanides and actinides
The chemical properties of the lanthanides (elements 57-71) and the actinides (elements 89-103) are even more similar to each other than the transition metals, and separating a mixture of these can be very difficult. This is important in the chemical purification of uranium concerning nuclear power.

Group 17 elements- Halogens
Fluorine(F) Chlorine(Cl) Bromine(Br) Iodine(I) Astatine(At)
~Halogens exist as diatomic molecules, low density, high electronegativity, good oxidising agents, weak conductors of heat and cannot conduct electricity.
Atomic Size -increases gradually
Reason: The number of shells occupied with electron increases.
Density -increases gradually
Reason: The increase in atomic mass is bigger than the increase in volume.
Boiling and melting points -increases gradually
Reason: As the molecular size increases, the force of attraction(van der Waals forces) become stronger. Consequently, more heat energy is required to overcome the stronger of attraction during melting or boiling.
Colour -Pale yellow-----Greenish-yellow-----Reddish-brown------purplish-black
Reactivity -decreases
Reason: As the atomic size increases, the single valence electron becomes futher away from the nucleus. Therefore, the strength of the nucleus of a halogen atom to attract one more electron into the outermost shells to achieve an octet electron arrangement decreases.
Electronegativity -decreases
Reason: The number of shells occupied with electrons in the atom with halogens increases. This cause the outermost occupied shell to become further away from nucleus. Hence, the strength of the nucleus to attract outermost shell electrons becomes weaker.
Reactions of halogens with water:
Example:
Cl2 (g) + H2O (l) → HCl(aq) + HOCl (aq)
Chlorine Water Hydrochloric acid Hypochlorous acid
Reactions of halogens with iron:
Example:
2Fe(s) + 3Cl(g) → 2FeCl3 (s)
Iron Chlorine Iron(III) chloride
Reactions of chlorine with sodium hydroxide solution:
Example:
Cl2(g) + 2NaOH(aq) → NaCl(aq) + NaCl(aq) + H2O
Chlorine Sodium hydroxide Sodium chloride Sodium chlorate Water

Group 18 elements- Noble gases:
Helium(He) Neon(Ne) Argon(Ar) Kryton(Kr) Xenon(Xe) Radon(Rn)
~All noble gases exist as single atom(monoatomic) which with a duplet and octet electron arrangement, colourless gases, insoluble in water, low density, weak conductors of heat and cannot conduct electricity.

Atomic Size -increases gradually
Reason: The number of shells occupied with electron increases.
Density -increases gradually
Reason: The increase in atomic mass is bigger than the increase in volume.
Boiling and melting points -increases gradually
Reason: As the molecular size increases, the force of attraction(van der Waals forces) become stronger. Consequently, more heat energy is required to overcome the stronger of attraction during melting or boiling.