Achievement Standard 2.4. COLVALENT BONDING: LEWIS STRUCTURES AND SHAPES OF MOLECULES. LEWIS STRUCTURES. Developed from observation that noble gases have a stable electron configuration. Atoms that share electrons form an electron-pair bond that helps to attain a stable noble-gas structure.
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Achievement Standard 2.4 COLVALENT BONDING: LEWIS STRUCTURES AND SHAPES OF MOLECULES
LEWIS STRUCTURES • Developed from observation that noble gases have a stable electron configuration. • Atoms that share electrons form an electron-pair bond that helps to attain a stable noble-gas structure. • Octet Rule: The principle that bonded atoms (except H) tend to have a share in eight valance electrons.
LEWIS STRUCTURES • In writing Lewis structures, only the valance electrons are used. • There are two kinds of electron pairs: • Shared electrons form covalent bonds (indicated by a line) • Unshared pairs of electrons (indicated by two dots)
WRITING LEWIS STRUCTURES • Count the total number of valance electrons available (A) by first adding the group numbers of atoms present. (For group numbers with two digits, use last digit only!!) • For polyatomic anion: also add a value equal to charge to number of electrons. • For polyatomic cation: also subtract a value equal to charge from number of electrons.
WRITING LEWIS STRUCTURES • Calculate the total number of electrons needed (N) for each atom to have its own noble gas structure (two for H and eight for all atoms C and beyond). • Subtract the number in step 1 from the number in step 2. This represents the number of shared, or bonding electrons (S) S = N - A
WRITING LEWIS STRUCTURES • To predict arrangement of atoms in molecules and ions, use the following as a guide: • Many molecules and ions will have a central atom and two or more terminal atoms bonded to it. • Hydrogen is almost always a terminal atom; oxygen and the halogens are often terminal atoms. • The central atom is often the first atom presented in a formula (e.g., S in SO2 or C in CH4) Note exceptions
WRITING LEWIS STRUCTURES • From step 3, assign two bonding electrons (shared pairs) to each connection between atoms in the molecule or ion. The remaining electrons represent a pool of electrons you use for any lone pairs and/or to make multiple bonds, if necessary.
WRITING LEWIS STRUCTURES • Place lone pairs about each terminal atom (except H) to satisfy the octet rule. • If the central atom is not yet surrounded by four electron pairs, convert one or more terminal atom lone pairs to make multiple bonds between central and terminal atoms. In general, C, N, O and S have a tendency to form multiple bonds (double or triple).
EXCEPTIONS TO OCTET RULE • Electron-deficient molecules • Odd electron species (NO or NO2) • Central atom with only two or three bonds (BeF2 or BF3) • Expanded octets: where central atom is surrounded by more than 4 pairs of valance electrons (PCl5 or SF6 or XeF4) Note: this only occurs in Level 3 examples
SHAPES OF MOLECULES • Molecular geometry refers to the three dimensional shape of molecules. • Shapes can be predicted based on electron repulsion. • VSEPR or valence shell electron pair repulsion theory helps to predict molecular geometry • In essence, VSEPR says that the electron pairs surrounding an atom repel one another. Therefore, the electron pairs are oriented to be as far apart as possible.
VSEPR THEORY Consider a central atom (A) bonded with terminal atoms (X), with no unshared electron pairs. The common species with 2 to 6 electron pairs and their geometries would be: AX2 180o linear AX3 120o triangular planer AX4 109.5o tetrahedron And for Level 3 shapes: AX5 90o, 120o, 180o triangular bipyramid AX6 90o, 180o octahedron
VSEPR THEORY Now consider what happens when the central atom (A) has both terminal atoms (X) and unshared pairs of electrons (E): • AX2 180o linear • AX3 120o triangular planar AX2E 120o bent (angular) • AX4 109.5o tetrahedral AX2E2 109.5o bent And the Level 3 example AX3E 109.5o triangular pyramid
VESPR THEORY VSEPR theory also applies to: • Species with double and triple bonds; multiple bonds are treated like single bonds in terms of shape. (e.g. carbon dioxide, dinitrogen oxide) • Species with no single central atom. (e.g. Acetylene, ethylene)
POLARITY: BONDS • Nonpolar: A symmetrical distribution of electrons between atoms. Nonpolar bonds form whenever two atoms joined are identical. (e.g. H2) • Polar: An unsymmetrical distribution of electrons between atoms. Polar bonds form whenever two atoms joined are different. (e.g. HCl)
POLARITY: MOLECULES • A polar molecule is one that contains positive and negative poles (a dipole). • A dipole for a simple molecule such as HF can be illustrated as: H F • A nonpolar molecule is one that contains no positive and negative poles.
POLARITY: PREDICTION • For simple molecules with two atoms, determining polarity is easy: • All diatomic molecules are nonpolar. • Molecules with two different atoms are polar.
POLARITY: PREDICTION • To determine polarity for molecules with more than two atoms you need to know bond polarity and molecular shape. (BeF2, H2O and CCl4). • If the polar A—X bonds in a molecule are arranged symmetrically in 3D (has 2 or more planes of symmetry) around the central atom A, the molecule is nonpolar. • Polarity can also be determined by drawing the molecule in the shape of a + , using vectors for bonds. If the vectors add to zero, then the molecule is nonpolar.