Periodic Trends

These elements bond by metallic bonds (intramolecular forces) to form giant metallic structures.

The size of the atoms increases down the group

Physical properties of Group II metals

Ionization
The process of removing an electron from an isolated atom (or an ion) to form a cation. First ionization:M(g) M+(g) + 1e-
Second ionization:M+ (g) M2+(g) + 1e-

Ease of ionization
All the Group II elements have two electrons in their outer shell. They can lose these electrons to form positive ions. e.g. Ca Ca2+ + 2e-

In the neutral atoms the electrons are held to the nucleus by the electrostatic attraction of the protons in the nucleus. The further the electrons are away from the nucleus, the more weakly they are held. First and second ionization energies are the energies required to remove the first and second electron respectively.

The larger the Group II atom, the smaller the ionization energies. The larger
the atom, the easier it is to ionize it.

Reactivity
The reactivity of Group II elements depends on their ease of ionization. Reactivity of Group II INCREASE down the group.

Reactivity with oxygen

Magnesium metal + oxygen gas magnesium oxide
Mg (s) + O2 (g) MgO (s)unbalanced
2Mg (s) + O2 (g) 2MgO (s)balanced

Calcium metal + oxygen gas calcium oxide
Ca (s) + O2 (g) CaO (s)unbalanced
2Ca (s) + O2 (g) 2CaO (s)balanced

Barium metal + oxygen gas barium oxide
Ba (s) + O2 (g) BaO (s)unbalanced
2Ba (s) + O2 (g) 2BaO (s)balanced

Read:Chemistry for CSEC by Tania Chung-Harris and Mike Taylor

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Reactivity with water

Magnesium metal + cold water magnesium hydroxide + hydrogen gas Mg (s) + H2O (l) Mg(OH)2 (aq) + H2 (g)unbalanced Mg (s) + 2H2O (l) Mg(OH)2 (aq) + H2 (g) (very slow reaction) balanced

Magnesium metal + steam magnesium oxide + hydrogen gas Mg (s) + H2O (g) MgO (s) + H2 (g) balanced

Calcium metal + cold water calcium hydroxide + hydrogen gas
Ca (s) + H2O (l) Ca(OH)2 (s) + H2 (g) unbalanced Ca (s) + 2H2O (l) Ca(OH)2 (s) + H2 (g) balanced

Barium metal + cold water barium hydroxide + hydrogen gas Ba (s) + H2O (l) Ba(OH)2 (aq) + H2 (g) unbalanced Ba (s) + 2H2O (l) Ba(OH)2 (aq) + H2 (g) balanced

Reactivity with dilute hydrochloric acid
Magnesium metal + hydrochloric acid magnesium chloride + hydrogen gas Mg (s) + HCl (aq) MgCl2 (aq) + H2 (g) unbalanced Mg (s) + 2HCl (aq) MgCl2 (aq) + H2 (g) balanced

Calcium metal + hydrochloric acid calcium chloride + hydrogen gas Ca (s) + HCl (aq) CaCl2 (aq) + H2 (g) unbalanced Ca (s) + 2HCl (aq) CaCl2 (aq) + H2 (g) balanced

Barium metal + hydrochloric acid barium chloride + hydrogen gas Ba (s) + HCl (aq) BaCl2 (aq) + H2 (g) unbalanced Ba (s) + 2HCl (aq) BaCl2 (aq) + H2 (g) (Explosive. do not try this)

Periodic Trends in Group VII

Elements in Group VII
The elements in Group VII are called halogens
The atoms are:F, Cl, Br, I, At
The molecules are:F2, Cl2, Br2, I2, At2
Fluorine, Chlorine, Bromine, Iodine, Astatine
Astatine is radioactive

Physical state at room temperature and pressure (r.t.p.)
The halogens all form molecules with two atoms in each (diatomic molecules) F2, Cl2, Br2, I2

The size of the atoms increases down the group

The larger the atoms, the higher the temperature needed to melt or boil the element. The melting points and boiling points of the halogens increase steadily as we go from one element to the next down Group VII.

Fluorine is a pale yellow gas
Chlorine is a green gas
Bromine is a red-brown liquid (vaporizes easily to red-brown vapour) Iodine is a grey-black solid (sublimes easily to purple vapour)

Ease of ionization
All the Group VII elements have seven electrons in their outer shell. They can lose these electrons to form positive ions. NOT FEASIBLE

However – the atoms Group VII elements usually GAIN one electron each electrons to form negative ions

Electron Affinity / Electron Gain / Electron Capture
The process of accepting an electron on an atom in the gaseous state to form an anion. First electron affinity:X(g) + 1e- X-(g) Second electron affinity:X-(g) + 1e- X2- (g)

Ease of electron gain
All the Group VII elements have seven electrons in their outer shell. Each atom can gain ONE electron to form a negative ion. e.g. Cl(g) + e- Cl-(g)

Since chlorine exists naturally as diatomic molecules the equation is best written as:

Cl2(g) + 2e- 2Cl-(g)

In the neutral atoms the electrons are attracted to the nucleus by the electrostatic attraction of the protons in the nucleus. The smaller the atom, the more strongly electrons are attracted.

The smaller the Group VII atom, the easier it is to gain an electron. .

Reactivity
The reactivity of Group VII elements depends on their ease of electron gain. Reactivity of Group VII INCREASE up the group.

Strength of oxidizing power of group VII elements
To be done next term

Displacement reaction for group VII elements
To be done next term

Periodic Trends in Period 3

Elements in Period 3
Na, Mg, Al, Si, P, S, Cl, Ar.
Sodium, Magnesium, Aluminium, Silicon, Phosphorus, Sulphur, Chlorine, Argon

Note the position of Magnesium from Group II and Chlorine from Group VII

Gradation from metal to non-metal
Sodium, magnesium and aluminium are metals. The atoms bond by metallic bonds to form giant metallic structures (lattices).

Silicon is a metalloid. The atoms bond by covalent bonds to form giant atomic (covalent) structure.

Phosphorus, sulphur, chlorine and argon are non-metals

Phosphorus atoms bond by covalent bonds (intramolecular forces) in groups of four (P4) to form simple covalent (molecular) structures. In solid phosphorus the P4 particles are held together by weak INTERMOLECULAR FORCES.

Sulphur atoms bond by covalent bonds (intramolecular forces) in groups of eight (S8) to form simple covalent (molecular) structures. In solid sulphur the S8 particles are held together by weak INTERMOLECULAR FORCES.

Chlorine atoms bond by covalent bonds (intramolecular forces) in groups of two (Cl2) to form simple covalent (molecular) structures. The DIATOMIC molecules have very weak INTERMOLECULAR FORCES of attraction between them.

Argon does not bond so this element exists as FREE ATOMS.

INTERMOLECULAR FORCES
The atoms of a simple molecule are bonded covalently, however the separate molecules in the crystal lattice are held together by weak intermolecular forces (van der Waals forces of attraction, hydrogen bonds etc.). These weak intermolecular forces of attraction allow the molecules to separate from each other easily when heated. For example, if a small amount of energy is supplied to solid iodine, the weak van der Waals forces of attraction are disrupted and the solid sublimes. The shiny black iodine crystals are converted into purple vapour. The crystals are soft with low melting points and boiling points.

Please note:
The energy required to break the weak intermolecular forces (van der Waals forces) is not sufficient to break the strong intramolecular forces (covalent bonds).

Examples of simple molecular crystals (simple covalent)
Solid carbon dioxide (dry ice), ice (solid water), naphthalene, solid iodine, solid sulphur, solid phosphorus.

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