Matter & Bonding Lesson # 3

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Presentation transcript:

Matter & Bonding Lesson # 3 Molecular Compounds

Ionic vs. Molecular Bonds Recall that in ionic bonding, a positive and negative ion come together and are called a salt, or ionic compound. This involves the transfer of electrons from a metal to a non-metal. When working with only non-metals, a different type of bond forms, called a molecular, or covalent bond. A molecular or covalent bond is a shared pair of electrons held between two non-metal atoms that hold the atoms together in a stable molecule. This bond is also an example of an intramolecular force.

Diatomic Molecules When two of the same non-metal are covalently bonded together, it is called a molecular element, or diatomic molecule (di = two). Chlorine is an example of a molecular element – in nature it exists as a pair: Cl2. Other molecules that can form diatomic bonds are: H, N, O, F, Cl, Br, I – the "hockey stick and puck" on the periodic table.

Properties of Molecular Compounds When two or more different non-metals are covalently bonded together, it is called a molecular compound. Molecular compounds tend to be soft, they do not conduct electricity, they have low melting points (which is why sugar so easily melts and turns black when on a stove), and can exist in all states of matter. Some are soluble in water, others are not.

Naming Molecular Compounds It is much easier to determine names and formulas for molecular compounds, because there is no “criss cross rule” involved, as there is no transfer of electrons to form charges. You simply “name them like you see them” This means, that if given the formula, you can determine the name from the number of each element present, and vice versa. Before this, you need the prefixes associated with each number.

Prefixes 1 – mono 2 – di 3 – tri 4 – tetra 5 – penta 6 – hexa 7 – hepta 8 – octa 9 – nona 10 – deca

Examples N2O Carbon Monoxide Si4Cl6 Nitrogen Trihydride C3H8 Diphosphorus Pentoxide

Drawing Molecular Compounds Similar to ionic compounds, you must draw the Lewis structure for each atom, and determine how many of each atom you need to make the compound stable. The only difference is that the electrons are shared (circled) versus transferred (arrow). You also need to determine the central atom – this is always the non-metal that has the most bonding (lone) electrons.

Examples Carbon + Hydrogen Nitrogen + Fluorine

Examples (continued) Silicon + Oxygen Sulfur + Chlorine