Chemical Bonds: Formation of Compounds

Transcription

1 hapter 11 s 1 st shell 2 nd shell d hemical Bonds: Formation of ompounds dsubshell p Periodic System consists of periods (n= 1, 2, ) called shells, Each shell contains subshells(s, p, d, f ). Elements are classified in groupshaving similar chemical characteristics. Noble gases have filled sand psubshell, effectively having 8 e-in the valence shell. f Transition metals populate 4s before 3d. Before we continue with filling 4p, 10 e-must go into 3d. We put them in a loop since we are. stalled filling 4 th shell. detour f subshell can be placed in another loop. This Periodic table has 8 groups.

2 Ionization energy is the energy required to remove an e-from the atom. It increasesleft to right across a period Atomic properties and increases bottom to top, up a group. The behavior exactly oppositeto that of atomic radii! Ionization Energy vs. Atomic Number Atomic radii increase right to left across the period, and top to bottom down the group.

3 hemical Bonding A moleculeis a collection of atoms bound together. It is considered as an elementif all atoms are of the same type (e.g. 2 ), or a compoundif it is made of different atoms. A bond between two atoms of hydrogen will occur spontaneously if the atoms are within a close proximity, i.e. when attractive forces between the nucleus of one atom and the e - of the other overcome repulsive forces among their e -, thus pulling the atoms together. In a covalent bond between two hydrogen atoms, each atom apparently has 2 electrons in its shell. That gives each atom the electronic configuration of the closest noble gas, e. Electrons are shared between two atoms. ovalent bond can be presented as any of the following ways: : Remember that the line stands for a pair of e -! 74 pm 1 pm = m The bond that is formed is called covalent bond. (o-means partner, valentrefers to valence electrons). It always releases energy. Filled-shell e - (or core e - ) are almost never involved in the bond as they are too close to their own nucleus. 2 e atoms will never form a bond because the energy of e 2 > 2 e.

4 Molecules and the Octet Rule Elements want to have the electronic configuration identical to that of noble gas (8e-). When form a molecule, atoms achieve the octet(8e - ) by sharing e - with other atoms. ydrogen is an exception, as it only needs one more electron to fill its 1s orbital. ow many electrons an element needs to satisfy the octet rule could be found in two ways: 1) From the position of the element in periodic table with respect to noble gas; 2) From short electron configuration: 1) Ois two places left of Ne needs 2 e- 2)O: [e] 2s 2 2p 4 has 6 e-, needs two. : [e] 2s 2 2p 2 Needs 4 e- a: [Ar] 4s 2 Needs 6 e- Gaining six electrons is difficult. For a to obtain octet, it is easier to loose two e-. All atoms in the molecule have the electronic configuration identical to that of the closest noble gas. Valence electrons Nonmetals usually gain electrons, metals loose them. Four atoms form covalent bonds with the atom Electron from atom ore electrons Electron from atom : 8e - : 2e - O: 8e - : 2e -

5 Step 1: Find the total number of valence e - for each atom by adding e-in the short electron configuration. For ions, adjust e-count accordingly (subtract e - for cation, add for anion). Step 2: Assume that the first nonhydrogen atomin the formula of the group is the central (less electronegative)atom. onnect peripheral atoms with the central atom with single bonds. The central atom must form multiple bonds, hence can never be the central atom. Step 3: Subtract 2e - for each bond from total #e - to get#e - in lone pairs. Step 4: Put in the remaining electrons, two at a time, as lone pairs. Satisfy octet to the terminal atoms first; ifthere are any e - pairs left, put them on the central atom. Step 5: heckthat each atom has octetsatisfied (doublet for ). If not, move e - pair(s) from the adjacent atom to form multiple bonds. Practice with O 2-3, SO 3, etc. Drawing Dot Diagrams O O 2 [e] 2s 2 2p 2 [e] 2s 2 2p 2 [e] 2s 2 p 4 [e] 2s 2 2p 4 4e x 6e - 4 e e - Total: 16e - Total: 10e - O Technically, no central atom here, but the rule applies as is less electronegative atom. 10e 2e= 8e - in lone pairs : O : O O 16e 4e= 12e - in lone pairs :O.. O : :O.. O : :O.. O : : O :. O. = = O.. : O Ξ O : : Ξ O : : O Ξ O : Equivalent structures, or resonance forms. or or

6 O 3 2-4e+ (3 x 6e)+ 2e= 24e - O O O 24e (3 x 2e) = 18e - : O O: :. O. : Octet rule not satisfied for in lone pairs : O O: :. O. : total Which is the central atom? In symmetrical molecules, BOT! 2x4e+4x1e=12e - total Two more resonance structures for O 3 2- ion are possible. 12e (5x2e) = 2e - in lone pairs : Octet rule not satisfied for has no e - to give in for double bonds : l : 2 2 l 2 2x4e+ 2x1e+ 2x7e= 24e - total cannot pull electrons from l. : l : : l : Two more structures can be made: : l: : l: : l : These are NOT resonance structures! SO 3 is identical to this; try it yourself.

7 The bond between metals and non-metals is usually ionic. Metals give away their e-and become positively charged (cations). Nonmetals accept them and become anions. The ionic bond is formed as a result of attraction between oppositely charged ions. The compound is called ionic compound, and the three-dimensional ordered network of the ions is called ionic lattice. Electronegativity and the Polar ovalent Bond The difference in EN defines the bond. EN = 0, covalent; EN = 1.0, polar covalent (23% ionic); EN = 1.9, polar covalent (60% ionic); EN > 1.9, ionic. Electrons are rarely shared equally between atoms. Electronegativity (EN) is numerical rating of an atoms ability to attract to itself the shared electrons in a covalent bond. Generally, electronegativity of metals is low, and that of nonmetals is high. The least electronegative atom (except!)is the central atom in dot structures. Polar covalent bondis a covalent bond in which e - are shared unequally (large EN). A partial negative charge (δ-) occurs on the more EN atom. A partial positive charge (δ+) occurs on the less EN atom. Ionic and covalent are two extremes at the ends of a continuum bonding types.

8 The Shape of the Molecules Valence Shell Electron Pair Repulsion (VSEPR) theory is the model mostly used to predict molecular shape. Electron pairs on the central atom repel one another. The two dimensional dot structure of methane, 4. gives the angles between electron pairs of 90 o. But the dot structure angles are arbitrary. Electron pairs move further away in three dimensions. The molecule of methane is tetrahedralwith bond angles of o between electron pairs. Four electron pairs around an atom assume tetrahedral arrangement. When there are not enough electrons for single bonds the molecule forms multiple bonds and the structure differs. VSEPR theory treats each multiple bond as a single electron group, because it occupies roughly the same region of space. The number of electron groups around an atom is called the atom s steric number (SN). Dot structures of formaldehyde and acetylene are arbitrarily shown with angles of 90 o. Their true geometry has bond angles of 120 o and 180 o, respectively. :O : formaldehyde Ξ acetylene Dot structures : O : Ξ True geometry

9 VSEPR arrangements of electron groups around an atom having no lone pair electrons Molecular arrangement may differ from that of the electrons. We only see stationary atoms, not the fast moving electrons. Lone pair(s) of electrons are thereforeignored. The molecule of N 3 has a pyramidal shape. O = = O Lone pairs on the central atom are considered in the prediction of the electron arrangement. The dot diagram of ammonia presents the atom in a plane. There are 4 electron pairs around nitrogen in N 3 (3 bonding pairs and a lone pair). They assume tetrahedral arrangement. Electrons in the lone pairoccupy more space than the bonding pairs. They squeeze-n- bond angle to approx. 107 o... N Each lone pair of e-on a period-2 atom compresses the remaining bond angles around that atom by ~2 o.

10 Using VSEPR 1. Draw a dot diagram. 2. ount the number of e- pairs around the central atom, including lone pairs (i.e. the steric number, SN). A multiple bond counts as a single e- group. 3. Find the best arrangement of the electrons using SN. 4. Pretending the lone e- pairs are invisible, describe the resulting shape of the molecule. Practice on 2 O and O 3. -O- bond angle ~105 o. bent O-O-O bond angle 120 o. bent O : O: O: SN=4 SN=3 O.... Molecular polarity depends on the position of the centers of negative and positive charge.

11 W, hapter 11 (p.247): 3, 9, 19, 37, Which one in each pair has the larger radius? Explain. a atom or a 2+ ion; l atom or l - ion; Mg 2+ or Al 3+ ; Na atom or Si atom. K + ion or Br - ion. 9. Using the table of electronegativity values (Table 11. 5) indicate which element is more positive and which is more negative in these compounds: 2 O; Rbl; N 3 ; PbS; PF 3 ; ow many electrons must be gained or lost for the following to achieve a noble gas electron configuration? K atom; Al ion; Br atom; Se atom. 37.Draw Lewis structures for the following: Nl 3 ; 2 O 3 ; 2 6 ; NaNO Use VSEPR theory to predict the shape of these molecules: Si 4 ; P 3 ; SeF 2.