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Bonding

Bonding. Year 11 DP Chemistry. What is a bond?. A chemical bond is a force that holds atoms together making a new substance. Ionic Bonds result from electrostatic attraction between oppositely charged ions

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Bonding

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  1. Bonding Year 11 DP Chemistry

  2. What is a bond? A chemical bond is a force that holds atoms together making a new substance • Ionic Bonds result from electrostatic attraction between oppositely charged ions • Covalent Bonds result from electrostatic attractions involving electrons and positively charged nuclei • Metallic Bonds result from electrostatic attractions between delocalised electrons and a lattice of positively charged nuclei

  3. What type of bond? Ionic Bond: As a rule of thumb, we say that the difference between the electronegativity values needs to be high (i.e. greater than 1.7) to be ionic. They form between cations on the left and anions on the right of the Periodic Table. Covalent Bond: If the difference between the electronegativity values of two highly electronegative atoms is low, a covalent bond is formed. They tend to form between non-metals, but sometimes metals are involved (eg Al2Cl6) Metallic Bond: If the difference between the electronegativity values of two highly electropositive atoms is low, a metallic bond is formed. These form between metals of the same or different type of atom

  4. Electronegativity The relative tendency of an atom to attract bonding electrons to itself on the Pauling Scale

  5. Formation of ions Groups 1,2,3 Groups 5,6,7 Gain electrons Group 5 – gains 3 e- to gain a full valence shell Group 6 – gains 2 e- to gain a full outer shell Group 7 - ?? Lose electrons • Group 1 – loses one e- to gain a full valence shell • Group 2 – loses 2 e- to gain a full valence shell • Group 3 – loses ??

  6. Transition ions • Ions of the transition elements can form more than one ion • This is due to s and d orbitals having similar energy levels For example – Fe forms two ions [Ar]4s23d6 Can you deduce which two ions and why? Fe(II) – losing two 4s electrons F(III) – losing two 4s e- and one 3d e- to give a half-filled d

  7. Exercises • Using a Periodic Table, determine which ions are formed from the following elements? Na, Al, O, Ca, F, N • Describe the size of each of the above ions relative to their original atoms.

  8. Ionic Bonding Oppositely charged ions are formed by electron transfer due to a large electronegativity difference(> 1.7 difference) Na has a low electronegativity relative to Cl, so ions are formed by a transfer of an electron to achieve a full valence shell for both atoms. These oppositely charged ions then form a bond.

  9. Sodium Chloride Lattice • Ionic compounds form a repeating crystal lattice of positive and negative ions. The compound is neutral • The electrostatic attraction is very strong so ionic compounds form solids at room temperature This shows a model of a NaCl lattice with alternating positive and negative ions

  10. Ionic formulas • The formulas for ionic compounds are found by balancing the overall charges to zero (neutral) • First determine the charge on each ion • Second determine the number of each ion needed to make the compound neutral • Write the positive ion first indicating the number of each using subscripted numbers Examples: Li+ F- LiF Mg2+Cl-  MgCl2

  11. Polyatomic ions Some ions contain more than one element and the charge on the ion is spread (delocalised) over the entire ion. They have specific names and act as a single unit. Important ones to know: NH4+ (ammonium) NO3- (nitrate) OH- (hydroxide) SO42- (sulfate) CO32- (carbonate) PO43- (phosphate) HCO3- (bicarbonate or hydrogen carbonate) Ionic compounds form in the same way with polyatomic ions. Here we see Na2SO4. Notice the sulfate did not change formula.

  12. Exercises • Write the formulas formed by the following pairs of elements: (Ca, S); (K, N); (Ca, P); (Fe(II), O); (Fe(III), O) • Write the formulas for the following compounds: potassium hydroxide, magnesium nitrate, sodium hydrogen sulfate • Use electronegativity values to determine the type of bonding between Al and O; Al and Br.

  13. Covalent bonding Electrons are shared between two atoms. These atoms are most commonly non-metals. • The shared electrons are held in place by the positively charged nuclei that are sharing them. • Single bond = 1 shared pair • Double bond = 2 shared pairs • Triple bond = 3 shared pairs

  14. How many bonds? Multiple bonding: Single bonding:

  15. Multiple bonding The more shared pairs the stronger the bond and the shorter the bond. Note the trend in C-C bonds. Note: bond energy is the amount of energy required to break a bond

  16. Coordinate (dative) bonds In this type of covalent bond, the difference is that one of the atoms in the pair donates both of the electrons in the bond. Examples: CO, NH4+, H3O+

  17. Electron Dot Diagrams (Lewis) The dots represent the valence electrons for each element

  18. The Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Fluorine molecule

  19. Hydrogen Chloride Here, the octet rule is satisfied for Cl, but is irrelevant for H which can only hold 2 e-.

  20. Double Bonds - ethene Two pairs of shared electrons Draw the Lewis structure for Acetylene (ethyne) – C2H2

  21. Formation of water Notice the double bond in O2. No other configuration will satisfy the octet rule. Why is H4O not formed?

  22. Resonance • When more than one possible Lewis structure can be drawn, it indicates resonance • The actual structure is an average of the resonance structures Notice that the bond lengths would be different in benzene unless there is a resonance structure like the one represented above

  23. The Octet Rule • 2nd row elements C, N, O, F observe the octet rule. • 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. • 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. • When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

  24. Lewis Diagrams • Determine the number of valence e- for each element • Determine the number of bonds for each element (usually the number of e- needed to fill the valence shell – exceptions AHL) • Central atom is the one with the most bonds • Join the atoms so that all have the correct number of bonds • Each bond has one e- from each atom (unless dative bond), so a pair makes a bond • Once all bonds are made, place the remainder of e- around the appropriate atoms as non-bonding pairs • Check that the octet rule has been satisfied for each element

  25. Completing a Lewis Structure -CH3Cl • Valence e- (C = 4, H = (3)(1), Cl = 7 Total = 14 ) • Number of bonds (C = 4, H = 1, Cl = 1) • Carbon is the central atom (most bonds) • Bonds are arranged and extra e- are added around Cl as non-bonding pairs. H .. .. All octets (ignoring H) have been satisfied. .. H .. C .. Cl .. .. H

  26. Exercises N2 has a triple bond. Using dot diagrams, show why a single or double bond is incorrect. Illustrate how CO and H3O+ contain dative bonds. Draw a Lewis diagram for HCN Draw the two resonance structure of ozone O3 Use a drawing to show how many resonance structures are possible for the nitrate ion (NO3-) Compare the bond lengths and strengths of the two C,O bonds in the carboxyl group below.

  27. Octet rule exceptions There are 3 ways the octet rule breaks down: • Molecules with odd number of electrons • Molecules where an atom has < octet • Molecules where an atom has > octet

  28. Odd number of electrons NO (nitrous oxide) • Total e- = 6+5=11 • Octet can be achieved around the O with a single bond between • Remaining e- around the N • What about a double or triple bond? Try it.

  29. Less than an octet This mostly occurs with H, B and Be BF3 (boron trifluoride) • Total e- = (3x7)+3 = 24 Octets to outer atoms Extra e- (24-24=0) to B Therefore no extra electrons to add. B has only 6 e- (< octet). What about double bonds??? see next…

  30. Less than a octet (cont’d) This would give 3 resonance structures. What would they look like? Add a double bond to BF3 for a possible octet… The above structure would lead to a δ+ on F and a δ- on B. Is this likely considering the electronegativites? NO Because B has only 6 valence electrons, BF3 reacts strongly with compounds that have unshared pairs of electrons

  31. Greater than an octet PCl5 Starting in period 3, expanded valence shells are possible. This is the most likely exception to the octet rule. The octet rule is based upon valence orbitals containing an s and p orbital. This gives 2 + 6 = 8 e- (an octet). In the third shell (n=3) d orbitals become available. P is below. A 3s can be exited to the 3d, which allows for 5 valence shell bonding electrons. Promote one e-

  32. Greater than an octet (cont’d) Sulfur (also in group III) can expand it’s octet to have more than 8 electrons as well. Sulfur can form SF2, SF4, SF6 Go to this website to see an animation on expanded octets in sulfur: http://www.saskschools.ca/curr_content/chem20/covmolec/exceptns.html

  33. Greater than an octet (cont’d) Other notable expanded octets… PF6- (12 e-) XeF4(12 e-)

  34. Polarity • Polar covalent bonds • Electrons are unequally shared by atoms • Electronegativity difference is 0.5>__<1.7 • Nonpolar covalent bonds • Electrons are equally shared • Electronegativity difference is <0.5

  35. Polar covalent bonds Some bonds are not purely ionic, but they still have a significant difference in electronegativity that leads to one atom pulling the electrons more strongly than the other. Polarity This electronegativity difference leads to a partially positive end δ+ of a molecule and a partially negative end δ- (note: δ is the Greek letter delta and means partial) F is more electronegative, so the electrons spend more time around the F nucleus

  36. Non-polar covalent Small or no difference in electronegativity values leads to non-polar substances. Cl is a highly electronegative element, but there is no difference when one Cl atom bonds to another Cl atom C and H have very little difference in electronegativity, so methane is non-polar

  37. Non-polar compounds Some compounds contain polar bonds, but the polarity is cancelled out due to the structure. O is more electronegative than C meaning each bond is polar towards the O atom, but due to its linear shape, these polarities cancel each other resulting in a non-polar molecule.

  38. Exercise • Show how S can form 3 compounds with F • Predict if these substances contain polar bonds. H2, CCl4, H20 • Use arrows to show the overall polarity of the compounds below?

  39. Molecular geometry VSEPR To determine the shape of covalent molecules, we use the Valence Shell Electron Pair Repulsion Theory (VSEPR) which states: “The geometric arrangement of atoms around a central atom is determined by the repulsion between electron pairs in the valence shell of the central atom.” Stay away! I am repulsed by you Ditto!! e- e- e- e- In other words…

  40. What shape? So, VSEPR theory says that molecular geometry is determined by the shape that keeps e- pairs as far apart as possible Consider CO2 We know C forms four bonds and O forms 2 bonds What arrangement will allow the valence e- around the central atom (C) to be as far apart as possible? The Lewis structure looks like this: Linear

  41. Linear (1800) Linear molecules have two areas of high electron density around the central atom. Other examples of linear molecules : Ethene (C2H4) Ethyne (acetylene) C2H2 Molecular chlorine (Cl2)

  42. 3 pairs around the central atom 3 e- pairs around the central atom leads to a trigonal planar shape as in BF3 Trigonal planar - Angles are1200 If one of those pairs is a non-bonding or lone pair of electrons, the shape is described as bent or v-shaped as in SO2 Bent – angles are less than 1200 due to lone pairs taking up more space than bonding pairs. These angles are 1170

  43. 4 pairs However, in 3-D space, it is possible to allow the electrons to be further apart using a tetrahedral shape with bond angles of 109.50 If we look at the Lewis structure for CCl4, we might assume a flat structure with 900 bond angles

  44. 4 pairs (cont’d)

  45. 5 pairs (AHL) Trigonal bipyramidal (0 lone pairs) See saw (1 lone pair) Angles = 900 & 1200

  46. 5 pairs (cont’d) T-shaped 2 lone pairs Linear 3 lone pairs

  47. 6 pairs (AHL) Octahedral (0 lone pairs) Square pyramid (1 lone pair) Square planar (2 lone pairs) Justify the shape of XeF4. Why are the lone pairs at 1800?What other two shapes are possible with 6 pairs?

  48. Summary of geometries

  49. Allotropes of Carbon Each carbon atom in graphite is bonded to 3 other carbon atoms forming flat sheets of carbon rings. These layers are loosely bonded to each other making graphite soft In diamond, each carbon is bonded to 4 other carbons in a giant repeating lattice. This lattice is non-polar and very strong, making diamond the hardest mineral on Earth. It’s m.p. is over 35000C! The fullerene contains 60 carbons arranged like a soccer ball with alternating 5 and 6-member rings.

  50. Graphite Carbon bonded to three other carbon atoms leaves one valence electron per carbon atom. These electrons are delocalised allowing graphite to conduct electricity The individual layers contain strong covalent bonds, but are only loosely bonded to other layers. This allows them to easily slide over one another making graphite useful in pencils and as a solid lubricant

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