Notes: Covalent Bonding

Transcription

1 Name Chemistry Pre-AP Notes: Covalent Bonding Period The main focus of this unit is on the covalent bond; however, we will briefly treat the ionic and metallic bond as well. I. Chemical Bonding Overview A. Definition A is a force of attraction that holds atoms or ions together to form a compound. The attractive force is between the nucleus of one atom/ion and the of another atom/ion. B. Purpose of Bonding Chemical bonding is driven by the octet rule: atoms will,, or electrons in order to achieve the electron configuration of the closest noble gas. Unfilled or partially filled valence orbitals are inherently unstable. ( Unstable means possessing high potential energy.) Atoms will bond to reduce their, thereby becoming more stable. C. Types of Chemical Bonds There are 3 types of chemical bonds:,, and. Bond type is determined by electrons of participating atoms. bonds involve the of electrons (lose/gain) between a metal and a nonmetal. bonds involve the of electrons between nonmetals/metalloids. bonds involve a mobile electron sea between metal atoms. II. Ionic Bonding A. Formation of Ions an atom that loses electrons to achieve an octet will form a. Ex: Sodium will lose valence electron to form the ion An atom that gains electrons to achieve an octet will form an. Ex: Chlorine will gain valence electron to form the ion 1

2 B. Formation of Ionic Compounds Involve a metal ion and a nonmetal ion; may also involve polyatomic ions. Anions and cations are attracted to each other by forces, also known as Coulombic forces. These forces of attraction between oppositely-charged ions are called bonds. Ex: Formation of sodium chloride Formation of aluminum bromide C. Properties of Ionic Compounds A is a 3-D system of points showing the positions of the ions that make up the ionic compound: The electrostatic (Coulombic) forces between the ions are very strong. Therefore, ionic compounds: are found in nature as solids. have melting and boiling points. are hard,, and. Ionic compounds will conduct when or (dissolved in water). Electrolyte: substance that conducts electricity when molten or in solution Why does the ionic compound have to be either dissolved in water or molten in order to conduct electricity? (In other words, why won t the dry salt conduct?) 2

3 III. Metallic Bonds A. Metallic Bonds Metals characteristically have few in their highest energy levels. Metals therefore have vacant orbitals (typically and/or block orbitals). The vacant orbitals can be, allowing valence electrons to travel freely throughout the metal. This is described as an electron sea. The electrons are, meaning they do not belong to any one nucleus. bonding is the chemical bonding that results from the attraction between the nuclei of metals atoms and the surrounding sea of. B. Properties of Metals The freedom of to roam in metals accounts for the high and conductivity of metals. Metals exhibit a (shiny) appearance because they can absorb a wide range of light frequencies. Metals are (they can be hammered into thin sheets). Metals are (they can be pulled into a wire). IV. Nature of Covalent Bonds A. Definition/Properties In the covalent bond, electrons are by the bonding atoms to achieve octets for the atoms. Compounds that consist of covalently bonded atoms are compounds. The smallest particle of a molecular compound is called a. The chemical formula for a molecular compound is the formula. elements (and sometimes metalloids) tend to form covalent bonds. 3

4 H 2 O molecule sucrose (C 12 H 22 O 11 ) molecule Relatively weak intermolecular forces are responsible for attractions between molecules in a molecular compound. More on this in the next unit! Therefore, molecular compounds: often exist as or vaporize easily at room temperature have relatively melting points and boiling points are relatively (if solid) B. Types of Covalent Bonds 1. covalent bond -- involves shared pair of electrons Ex: H 2 O 2. covalent bond -- involves shared pairs of electrons Ex: CO 2 3. covalent bond -- involves shared pairs of electrons Ex: N 2 V. Lewis Structures A. Creating Lewis Structures Lewis structures are depictions of molecules that show valence electrons as. Shared pairs of electrons (i.e. ) are drawn between the atoms sharing them. Unshared or pairs of electrons are represented by dots located on one atom only. 4

5 Drawing Lewis structures for molecular compounds 1) Find the needed number of electrons (N) for each atom in the compound and add them up. N will generally be 8 for each atom, with the following common exceptions: H =2, Be = 4, B = 6 2) Find the available number of electrons (A) for each atom in the compound and add them up. A is the number of valence electrons for each atom. 3) If the substance is a polyatomic ion, adjust A by using the ion s charge. For example, in nitrate, you would add 1 electron to A because of the -1 charge. In ammonium, you would subtract 1 electron from A because of the +1 charge. 4) Find the shared electrons (S) by this formula: S = N - A 5) Draw a skeleton of the molecule. In most cases, there is one different atom. Put that one in the middle and surround it by the others. 6) Place the shared (S) electrons in between atoms. 7) The S electrons are part of the A electrons. Figure out how many more electrons you need to add in order to have A total electrons in the structure. Place those as unshared pairs in order to give each atom an octet. 8) Check: is the total # of dots in the structure = to A? Does every atom have an octet? Ex: Draw the Lewis structure for methane. N = A = S = Ex: Draw the Lewis structure for phosphorus trichloride. N = A = S = Ex: Draw the Lewis structure for the carbonate ion. N = A = S = 5

6 Practice 1. Draw Lewis structures for a) ammonia b) hydrogen cyanide c) boron trichloride B. Resonance Structures Resonance structures are structures that occur when it is possible to write 2 or more valid Lewis structures for the same molecule or ion. Ex: ozone, O 3 Experimental data indicate that the 2 bonds in ozone are the same, BUT double bonds are shorter than single bonds! So the explanation for this is that the actual bonds are of those in the 2 resonance structures. The extra electron pair in ozone is delocalized over the two bonding regions. So each bond spends about half the time being single and half being double. Practice 1. Draw the resonance structures for the nitrate ion. C. Structural Formulas 1. Structural formulas -- shows the pairs of shared electrons with a straight line to represent each shared pair of electrons. Lone pairs are still depicted as dots. Ex: CH 4 O 2 H 2 O CO 3 2-6

7 VI. VSEPR Theory and Molecular Geometry (shape) The alence hell lectron air epulsion Theory explains the threedimensional shape of molecules. The VSEPR theory states that because electron pairs each other, the molecular shape results from valence electron pairs positioning themselves as far apart as possible. Ex. methane, CH 4 : bond angles o, not 90 o Note: lone pairs (unshared pairs) of electrons repel the bonding pairs of electrons more than other bonding pairs do, causing the bonding angles to be smaller than expected. Ex. H 2 O has o bond angle NH 3 has 107 o bond angles Use the Lewis structure for the molecule and the chart on the next page to predict molecular geometry (shape): Practice Use VSEPR Theory to predict the molecular geometry (shape) of - - a) CO 2 b) ClO 3 c) NO 3 d) SCl 2 e) CCl 4 7

8 # of atoms bonded to central atom # of lone pairs on central atom Molecular Geometry (shape) Sketch (note: bonds may not all be single) General Form and Example 2 0 Linear AX 2 CO Bent 120 o AX 2 E 1 SO Bent o AX 2 E 2 H 2 O 3 0 Trigonal planar AX 3 BF Trigonal pyramidal AX 3 E 1 NH Tetrahedral AX 4 CH 4 8

9 VII. Polarity of Bonds A. Determination of Bond Type To determine type of bond: (not applicable for metallic bonding*) Remember that electronegativity is the measure of an atom s ability to attract bonding electrons. See chart of values below. 1) Determine the difference in electronegativity values between the 2 atoms (abs value) 2) If the difference is: > 1.70 bond (electrons transferred) bond (electrons shared unequally) < 0.36 bond (electrons shared equally) *Bonds between 2 metals are always classified as metallic. Summary A large electronegativity difference leads to an ionic bond. A moderate electronegativity difference leads to a polar covalent bond. Little to no electronegativity difference between two atoms leads to a nonpolar covalent bond. Even though we use electronegativity difference to classify the type of bond, the real truth is that there is no clear-cut division between types of bonds. It s really more of a spectrum than a black-and-white distinction. In a nonpolar covalent bond, the electrons are held on average exactly half way between the atoms. In a polar covalent bond, the electrons have been dragged slightly towards one end. How far does this dragging have to go before the bond is considered ionic? There is no real answer to that. Sodium chloride is a typical ionic solid, but even here the sodium hasn't completely lost control of its electron. Because of the properties of sodium chloride, however, we tend to classify it as if it were purely ionic. 9

10 Ex: Determine the type of bond between cesium and sulfur. Ex: Determine the type of bond between carbon and oxygen. Ex. Determine the type of bond between carbon and hydrogen. Ex: Determine the type of bond between fluorine and fluorine. Ex: Determine the type of bond between zinc and copper (in the alloy brass). Practice 1. Determine the type of bond (ionic, nonpolar covalent, polar covalent, or metallic): a) C-H b) O-H c) Na-Cl d) H-H e) Fe-Fe B. Polar vs. Nonpolar Covalent Bonds The bonding pairs of electrons in covalent bonds are located between the of the atoms sharing the electrons. When the pull of each nucleus for the electrons is equally strong, the electrons are shared equally and are located on average halfway between the two nuclei. This type of bond is a covalent bond. When the pull of one nucleus for the electrons is stronger than the other, the electrons spend more time closer to the more electronegative nucleus. This type of bond is a covalent bond. The more electronegative atom acquires a partial charge. The less electronegative atom therefore acquires a partial charge. These partial charges are indicated by the following symbols: δ for the more electronegative element. δ + for the less electronegative element. Ex: H 2 O The water molecule contains two dipole moments, sometimes just called dipoles. The dipole moments can be drawn in with arrow notation: 10

11 VIII. Polarity of Molecules Note that just because a molecule contains polar bonds, it may or may not be classifed as a polar molecule. The polarity of a molecule will depend on: (1) existence of polar bonds (if none of the bonds are polar, the molecule is nonpolar*) (2) shape of the molecule (3) orientation of the polar bonds *there are exceptions to this (such as O 3 ), but you will not be tested over them To help determine if the molecule will be polar, look for lines of symmetry. If the molecule is symmetrical in its 3-D shape, then the dipoles will cancel and the molecule is nonpolar. If the molecule is asymmetrical and therefore the dipoles do not cancel, the molecule is polar. Note that water is polar because its dipoles do not cancel. Helpful hint: The following molecular geometries are symmetrical and have a high probability of leading to nonpolar molecules, assuming equivalent dipoles around the central atom: Linear (triatomic) Trigonal planar Tetrahedral For the following examples, (1) determine the molecular geometry; (2) determine the bond polarities and draw in any dipoles that you find; (3) classify the molecule as overall polar or nonpolar. Ex: HCl CO 2 CF 4 NH 3 Practice Determine if the following molecules will be polar or nonpolar. a) H 2 S b) CBr 4 c) PCl 3 d) BF 3 e) CS 2 11