Chapter 8. Chemical Bonding: Basic Concepts Chemical bond: is an attractive force that holds 2 atoms together and forms as a result of interactions between electrons found in combining atoms We rarely deal with isolated atoms Chemical bonds are broken and formed in reactions Properties of substances often determined by bonds Bonding lowers the potential energy 1
Chemical Bonds Three basic types of bonds: Ionic Electrostatic attraction between ions Covalent Sharing of electrons Metallic Metal atoms bonded to several other atoms
Ionic Bonds - Transfer of electrons from metal to nonmetal to form ions that come close together in sold ionic compound - Electrostatic attractions of closely packed, oppositely charged ions in a regular 3-D array - Formed when metal (that easily loses electrons) reacts with non-metal (that easily gain electrons). - No. e - lost by metal = No. e - gained by nonmetal
Properties of Ionic Compounds All solids at room temperature are brittle, hard, & rigid Melting points greater than 300 C Liquid state conducts electricity, solid state does not If soluble in water then good electrical conductor Chemical formula is empirical formula - simply giving the ratio of ions based on charge balance (no separate molecules) Ions arranged in a pattern called a crystal lattice maximizes attractions between + and ions
Electrostatic forces and the reason ionic compounds crack
Electrical conductance and ion mobility Solid ionic compound Molten ionic compound Ionic compound dissolved in water
Electron Arrangements and Ion Charge Metals form cations by losing enough e - to form same configuration as previous noble gas Nonmetals form anions by gaining enough e - to from same configuration as next noble gas Completely fill outer s & p orbitals Atom Atoms Electron Config Ion Ions Electron Config Na [Ne]3s 1 Na +1 [Ne] Mg [Ne]3s 2 Mg +2 [Ne] Al [Ne]3s 2 3p 1 Al +3 [Ne] O [He]2s 2 2p 4 O -2 [Ne] F [He]2s 2 2p 5 F -1 [Ne]
Electron Configuration of Ions of the Representative Elements Na 1s 2 2s 2 2p 6 3s 1 = [Ne]3s 1 Na + 1s 2 2s 2 2p 6 3s 1 = [Ne] Cl 1s 2 2s 2 2p 6 3s 2 3p 5 = [Ne] 3s 2 3p 5 Cl _ 1s 2 2s 2 2p 6 3s 2 3p 6 = [Ne] 3s 2 3p 6 = [Ar]
Electron Configuration of Ions of the Representative Elements In polyatomic ions, two or more atoms are bound together by predominantly covalent bonds and the whole acts as a charged species when the ion forms an ionic compound with an ion of opposite charge. E.g. NH 4 + and NH 4 Cl CO 3 2- and Na 2 CO 3
Electron Configuration of Ions of the Representative Elements In forming ions, transition metals lose the valence-shell s electrons first, then as many d electrons as are required to reach the charge of the ion. For example, Fe has the electron configuration [Ar]3d 6 4s 2. To form the Fe 2+ ion, the two 4s electrons are lost giving: [Ar]3d 6 To form the Fe 3+ ion, the two 4s electrons and one d electron are lost giving: [Ar]3d 5 Which element forms a 1+ ion that has the electron configuration [Kr]4d 8?
Covalent Bonds Atoms bond by sharing pairs of electrons Commonly found between nonmetal atoms e.g. Cl 2, O 2, N 2, H 2 etc Example H 2 : overlap of 1s orbitals 2H(g) H 2 (g) H = -432 kj Most atoms form covalent bonds by sharing enough electrons to satisfy octet rule 11
Covalent Bonding In these bonds, atoms share electrons. There are several electrostatic interactions in these bonds: Attractions between electrons and nuclei Repulsions between electrons Repulsions between nuclei
Bond Polarity Nonpolar covalent bond the electrons are shared equally between the two atoms, e.g. in Cl 2 and N 2 Polar covalent bond one of the atoms exerts a greater attraction for the bonding electrons than the other, e.g. in H 2 O
Bond Polarity Covalent bonding between unlike atoms results in unequal sharing of the electrons One end has larger electron density than other The result is bond polarity End with larger e - density gets partial - charge End that is e - deficient gets partial + charge H F Dipole Moment (μ) 14
Bond Polarity and Electronegativity The ability of an atom in a molecule to attract electrons to itself - Electronegativity On the periodic chart, electronegativity increases as you go: from left to right across a row( period) from the bottom to the top of a column
Electronegativity Relative ability of bonded atom in a molecule to attract shared electrons Larger electronegativity: atom attracts more strongly Values 0.7 to 4.0 EN: across period (left to right) on Periodic Table EN: down group (top to bottom) on Periodic Table Larger difference = more polar bond negative end toward more electronegative atom 16
Electronegativity values for selected elements 17
Electronegativity and Bond Type Difference in electronegativity ( EN) provides a good estimate of bond type: 0 to 0.4 non-polar covalent 0.4 1.7 polar covalent > 1.7 ionic 18
Electronegativity values for selected elements EN NaCl = 3.0 0.9 = 2.1 19
Dipole Moments When two atoms share electrons unequally, a bond dipole results. The dipole moment,, produced by two equal but opposite charges separated by a distance, r, is calculated: = Qr It is measured in debyes (D).
Polar Covalent Bonds Although atoms often form compounds by sharing electrons, the electrons are not always shared equally. Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.
Polar Covalent Bonds The greater the difference in electronegativity, the more polar is the bond.
Covalent Bonds Single Bond: atoms share 2 e - (1 pair) eg., C-C bond order = 1 Double Bond: atoms share 4 e - (2 pair) eg., C=C bond order = 2 Triple Bond: atoms share 6 e - (3 pair) eg., C C bond order = 3 23
Covalent Bonds Bond Strength: Triple > Double > Single For bonds between same atoms: C N > C=N > C N Bond Length: Single > Double > Triple For bonds between same atoms: C N > C=N > C N 24
Comparison of Ionic vs Covalent Property NaCl CCl 4 State solid liquid Melting point ( C) 801-23 Molar heat of fusion (kj/mol) Molar heat of vaporization (kj/mol) 30.2 2.5 600 30 Electrical conductivity Good Poor 25
Lewis Symbols The Lewis symbol of an element consists of the chemical symbol for the element plus a dot for each valence electron. Sulfur, for example, has the electron configuration [Ne]3s 2 3p 4. Therefore it has the following Lewis symbol: S which clearly depicts the six valence electrons.
What is the electron configuration & Lewis Symbol for the following Elements? B Ar Si O
Element Electron Configuration Lewis Symbol B [He] 2s 2 2p 1 :B Ar [Ne]3s 2 3p 6 :Ar: : : Si [Ne]3s 2 3p 2 :Si: O [He]2s 2 2p 4 :O:
Octet Rule Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons. An octet of electrons consists of full s and p subshells in an atom. Note: There are many exceptions to the octet rule, but it remains useful for introducing many important concepts of bonding.
Lewis Structures Lewis structures are representations of molecules showing all electrons - bonding and nonbonding.
Lewis Structures: Multiple Bonds When two electron pairs are shared, two lines are drawn, i.e. double bond. O :: C :: O or O = C = O When three electron pairs are shared, three lines are drawn, i.e. triple bond. :N ::: N: or :N N: N N N = N N N 147 pm 124 pm 110 pm
Lewis Symbols Al [Ne] 3s 2 3p 1 How many valence e - s? 3 Al F [He] 2s 2 2p 5 F 32
Drawing Lewis Structures PCl 3 1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. 5 + (3 x 7) = 26 If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge.
Drawing Lewis Structures 2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond, a dash (representing two electrons).
Drawing Lewis Structures 3. Complete the octets around all the atoms bonded to the central atom.
Drawing Lewis Structures 4. Place any leftover electrons on the central atom. Remember we have 26 electrons
Drawing Lewis Structures 5. If there are not enough electrons to give the central atom an octet, try multiple bonds.
Building Lewis structures of molecules HCN as an example... Step 1. Count the total number of valence electrons H has 1 C has 4 N has 5 Total of 10 Step 2. Place one e - pair between each BONDED atom H C N We have 6 e - left All atoms must have an octet or duet Step 3. Add electrons to terminal atoms first The H OK it has its duet... Next... 38
Step 3. Building Lewis structures of molecules Add remaining electrons to terminal atoms first Add 6 electrons in pairs to give the N an octet. H C N Step 4. Add any electrons left over to central atom We have none left! Step 5. Check for an acceptable Lewis Structure Do all atoms have an octet??? IN THIS CASE 39
Building Lewis structures of molecules H C N No! Both C and N need an octet.. the C and N have to share more than one pair of e - bring e - pairs from outer N atom to form shared pairs to give C its octet!!! 40
Building Lewis structures of molecules Step 5. Check for an acceptable Lewis Structure bring electron pairs from outer N atom to form shared pairs to give C its octet!!! H C N Still no octet on C Do it again!!!! H C N H C N three electron pairs between the C and N 41
Building Lewis structures of molecules three electron pairs between the C and N H C N Lewis (electron dot) structure of HCN There is a triple bond.. Also written H C N Another possible structure is. 42
Another structure H N C Lets do the Lewis structure... Step 1. Count the total number of valence electrons C has 4 N has 5 H has 1 Total of 10 Step 2. Place one e - pair between each atom H N C Step 3. Place remaining electrons on terminal atoms until their octet complete We get. 43
Step 3. Another structure Place remaining electrons on terminal atoms until their octet complete H N C Step 4. No electrons left. Step 5. Check for acceptable Lewis structure. The N does not have an octet... We bring electron pairs from outer C atom to form shared pairs to give N its octet!!! Again we need a triple bond. 44
Lewis structure of HNC H N C three electron pairs between the C and N this is called a triple bond.. Also written H N C How can we choose? H C N The octet rule is obeyed!!. FORMAL CHARGE.. 45
Formal Charge Guidelines Generally choose the Lewis structure in which the atoms bear formal charges closest to zero Generally chose the Lewis structure in which any negative charges reside on the more electronegative atoms
Drawing Lewis Structures - Formal Charge Assign formal charges: For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. Subtract that from the number of valence electrons for that atom: The difference is its formal charge.
CN NCS O 3
[:C N:] Valence Electrons 4 5 Electrons assigned to atom 5 5 Formal Charge -1 0
: : Drawing Lewis Structures The best Lewis structure is the one: with the fewest charges that puts a negative charge on the most electronegative atom. S S S -2 0 +1-1 0 0 0 0-1
Drawing Lewis Structures The best Lewis structure is the one: with the fewest charges that puts a negative charge on the most electronegative atom. -2 0 +1-1 0 0 0 0-1
Resonance Structures More than one Lewis structure that differs only in position of e - s Lone pairs & multiple bonds in different positions Actual molecule is combination (or blending) of all resonance forms (delocalized) Actual structure is average of resonance structures O S O.... O S O 52
RESONANCE We use a double headed arrow between the structures. O O O N N N O O O O O O The electrons involved are said to be DELOCALIZED over the structure. The blended structure is a RESONANCE HYBRID 53
QUESTION Which of the following molecules exhibit resonance? 1 CO 2 2 ClO 3-3 O 3 4 Cl 2 CO 5 F 2 O 54
Resonance Structures This is the Lewis structure we would draw for ozone, O 3.
Resonance Structures This is at odds with the true, observed structure of ozone, in which: both O O bonds are the same length, and both outer oxygens have a charge of ½.
Resonance Structures One Lewis structure cannot accurately depict a molecule such as ozone. We use multiple structures called resonance structures, to describe the molecule.
Resonance Structures Just as green is a synthesis of blue and yellow, ozone is a synthesis of these two resonance structures.
Resonance Structures In truth, the electrons that form the second C O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. They are not localised, but rather are delocalised.