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Chapter 11 Chemical Bonding

Chapter 11 Chemical Bonding. Forces that hold atoms together. The Nature of Bonding. There are several major types of bonds. Ionic, covalent and metallic bonds are the three most common types of bonds. Covalent bonds – electrons are shared between atoms.

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Chapter 11 Chemical Bonding

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  1. Chapter 11 Chemical Bonding Forces that hold atoms together

  2. The Nature of Bonding • There are several major types of bonds. Ionic, covalent and metallic bonds are the three most common types of bonds. • Covalent bonds – electrons are shared between atoms. • Ionic bonds – electrons are transferred between atoms, creating cations and anions. • Metallic bonds – two or more metals bonded together.

  3. The Nature of Covalent Bonding • There are two different types of covalent bonds, polar covalent and nonpolar covalent. • polar covalent – electrons are not shared equally between the two bonded atoms. The electrons are pulled toward the more electronegative of the elements. • nonpolar covalent – electrons are shared equally between the two bonded atoms.

  4. Electronegativities

  5. The formation of a bond between two hydrogen atoms. Source: Andrey K. Geim/High Field Magnet Laboratory/University of Nijmegen

  6. Probability representations of the electron sharing in HF. (a) What the probability map would look like if the two electrons in the H–F bond were shared equally. (b) The actual situation, where the shared pair spends more time close to the fluorine atom than to the hydrogen atom.

  7. The Nature of Covalent Bonding • Ionic bonds are formed when there is an electronegativity difference (DEN) greater than 2.0. • Polar covalent bonds form when there is a DEN between 0.5 and 1.7. • Nonpolar covalent bonds form when there is a DEN between 0 and 0.49.

  8. The Nature of Covalent Bonding • If the DEN is between 1.7 and 2.0, an ionic bond will form if a metal is one of the elements, and a polar covalent bond will form if only nonmetals or metalloids are present.

  9. The Nature of Covalent Bonding • What type of bond is formed between the following elements? • N and O K and F • Mg and Cl P and F • C and H

  10. • •  H F  Bond Polarity • Covalent bonding between unlike atoms results in unequal sharing of the electrons • One end of the bond has larger electron density than the other • The result is bond polarity • The end with the larger electron density gets a partial negative charge • The end that is electron deficient gets a partial positive charge

  11. The three possible types of bonds: (a) a covalent bond formed between identical atoms; (b) a polar covalent bond, with both ionic and covalent components; and (c) an ionic bond, with no electron sharing.

  12. Dipole Moment • Bond polarity results in an unequal electron distribution, resulting in areas of partial positive and partial negative charge • Any molecule that has a center of positive charge and a center of negative charge in different points is said to have a dipole moment (two different poles of charge).

  13. Dipole Moment • If a molecule has more than one polar covalent bond, the areas of partial negative and positive charge for each bond will partially add to or cancel out each other • The end result will be a molecule with one center of positive charge and one center of negative charge • The dipole moment effects the attractive forces between molecules and therefore the physical properties of the substance

  14. (a) The charge distribution in the water molecule. (b) The water molecule behaves as if it had a positive end and a negative end, as indicated by the arrow.

  15. (a) Polar water molecules are strongly attracted to positive ions by their negative ends. (b) They are also strongly attracted to negative ions by their positive ends.

  16. Polar water molecules are strongly attracted to each other.

  17. Electron Configuration in Ionic Bonding • Metals tend to lose their valence electrons, leaving a complete octet in their next-lowest energy level. • Sodium – (1 valence electron) loses 1 electron and becomes Na+1. • Na ([Ne]3s1)  1e- + Na+1([Ne]) • Calcium – (2 valence electrons) loses 2 electrons and becomes Ca+2. • Ca ([Ar]4s2)  2e- + Ca+2([Ar])

  18. Electron Configuration in Ionic Bonding • Nonmetals tend to gain or share valence electrons to complete an octet in their highest energy level. • Oxygen – (6 valence electrons) gains two electrons to become O-2 . • O ([He]2s22p4) + 2e-  O-2 ([He] 2s22p6) • Phosphorus – (5 valence electrons) gains three electrons to become P-3. • P ([Ne]3s23p3) + 3e-  P-3 ([Ne] 3s23p6)

  19. Formation and Properties of Ionic Compounds • Ionic bonds – forces of attraction that bind cations and anions together. • Ionic compound – consists of electrically neutral group of ions joined by electrostatic forces. • Example: Sodium chloride

  20. Formation and Properties of Ionic Compounds • At room temperature, most ionic compounds are crystalline solids, where ions are arranged in various 3-D patterns. • Because of the large attractive forces of the ions to each other the compounds become very stable and have high melting points.

  21. Sodium Chloride Crystals

  22. The structure of lithium fluoride.

  23. Electron Configuration in Ionic Bonding • Scientists have learned that all of the elements within each group behave similarly because they have the same number of valence electrons. • Valence electrons - # of electrons in the highest occupied energy level of an atom. • The number of valence electrons is related to the group numbers on the periodic table.

  24. Electron Configuration in Ionic Bonding • Group 1 elements = 1 valence electron. • Group 2 elements = 2 valence electrons. • Groups 3-12 elements = 2 valence electrons. • Group 13 elements = 3 valence electrons. • Group 14 elements = 4 valence electrons. • Group 15 elements = 5 valence electrons. • Group 16 elements = 6 valence electrons. • Group 17 elements = 7 valence electrons. • Group 18 elements = 8 valence electrons.

  25. Determining Valence Electrons for an Ion or a Compound • 1. Multiply the number of valence electrons by the number of moles of each element. • 2. Add up all the electrons for each of the elements. • 3. If there is a charge and it is negative, add that number of electrons to the total. • 4. If there is a charge and it is positive, subtract that number of electrons from the total. • Total # of electrons should always be an even number!

  26. Determining Valence Electrons Examples • Determine the number of valence electrons in each of the following compounds and ions: • NH4+1 • CH2ClBr • PO4-3

  27. Electron Configuration in Ionic Bonding • Valence electrons are the only electrons involved in bonding, and are the only ones written when drawing electron dot structures. • In forming compounds, atoms tend to achieve the electron configuration of a noble gas, having 8 valence electrons which as known as having a stable octet (octet for 8 valence electrons).

  28. •• •• • •• •• Li• Be• •B• •C• •N• •O::F: :Ne: • • • • • • •• •• •• Li• Li+1:F: [:F:]-1 • •• Lewis Symbols of Atoms and Ions • Also known as electron dot symbols • Use symbol of element to represent nucleus and inner electrons • Use dots around the symbol to represent valence electrons • put one electron on each side first, then pair • Elements in the same group have the same Lewis symbol • Because they have the same number of valence electrons • Cations have Lewis symbols without valence electrons • Anions have Lewis symbols with 8 valence electrons

  29. The Nature of Covalent Bonding • Structural formula – chemical formulas that show the arrangement of atoms in molecules and polyatomic ions. • Octet rule – atoms gain or lose electrons to acquire the stable electron configuration of a noble gas, usually having 8 valence electrons.

  30. . . : + H H H H Lewis Structures • You can represent the formation of the covalent bond in H2 as follows: • This uses the Lewis dot symbols for the hydrogen atom and represents the covalent bond by a pair of dots.

  31. : H H Lewis Structures • The shared electrons in H2 spend part of the time in the region around each atom. • In this sense, each atom in H2 has a helium configuration.

  32. : : . . : : : + H Cl H Cl : : Lewis Structures • The formation of a bond between H and Cl to give an HCl molecule can be represented in a similar way. • Thus, hydrogen has two valence electrons about it (as in He) and Cl has eight valence electrons about it (as in Ar).

  33. bonding pair : : : H Cl : lone pair • An electron pair is either a bonding pair (shared between two atoms) or a lone pair (an electron pair that is not shared). Lewis Structures • Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures.

  34. The Nature of Covalent Bonding • Exceptions to the octet rule: • H needs 2 electrons to be stable • Be needs 4 electrons to be stable • B needs 6 electrons to be stable

  35. The Nature of Covalent Bonding • Steps for Drawing Lewis-dot structures • Determine the number of valence electrons in the molecule. - When drawing determining valence electrons for an ion, add electrons if it an anion, and subtract electrons if it is a cation. • The first element in the compound will be the central atom. Exception: hydrogen will never be the central atom.

  36. The Nature of Covalent Bonding Steps for Drawing Lewis-dot Structures 3. Use one pair of electrons to bond each outer or terminal atom to the central atom. 4. Make all outer or terminal atoms stable using the valence electrons. 5. Put any remaining electrons around the central atom as lone pairs.

  37. The Nature of Covalent Bonding • Draw the Lewis structure for: • NH3 • PO43- • CHFClBr • PF5-2

  38. The Nature of Covalent Bonding • Single covalent bond – a bond in which two atoms share a pair of electrons. • Double covalent bond – a bond in which two atoms share two pairs of electrons. • Triple covalent bond – a bond in which two atoms share three pairs of electrons.

  39. The Nature of Covalent Bonding • If you have used up all of the valence electrons and you still need two more electrons to make the central atom stable, you must have one double bond. • If you still need four more electrons to make the central atom stable, you must have either one triple bond or two double bonds. • Double and triple bonds exist most commonly between C, N, O, and S atoms.

  40. The Nature of Covalent Bonding • Draw Lewis structures for: • NOCl • CO2 • N2 • SiO3-2

  41. The Nature of Covalent Bonding • Resonance structures – molecules or ions that can have two or more different Lewis structures. They must contain a double bond to have any resonance structures. • Resonance structures don’t truly have a single bonds or a double bond, but a hybrid mixture of bonds where the extra bond is spread equally among the other single bonds.

  42. The Nature of Covalent Bonding • Draw Lewis structures for: • NOCl • CO2 • N2 • SiO3-2

  43. The Nature of Covalent Bonding • Single bonds are longer (length between the atoms) than double and triple bonds. • Double bonds are longer than triple bonds. • Single bonds are not as strong as double bonds, and can be broken much easier than double bonds. • Triple bonds are stronger than double bonds.

  44. Bonding Theory • The valence-shell electron pair repulsion (VSEPR) model predicts the shapes of molecules and ions by assuming that the valence shell electron pairs are arranged as far from one another as possible. • To predict the relative positions of atoms around a given atom using the VSEPR model, you first note the arrangement of the electron pairs around that central atom.

  45. Predicting Molecular Geometry • The following rules and figures will help discern electron pair arrangements. • Draw the Lewis structure • Determine how many bonding pairs are around the central atom. Count a multiple bond as one pair. • Determine how many lone pairs, if any, are around the central atom. • All diatomic molecules have a linear shape.

  46. 3 pairs Trigonal planar 4 pairs Tetrahedral 5 pairs Trigonal bipyramidal 6 pairs Octahedral Arrangement of Electron Pairs About an Atom 2 pairs Linear

  47. Molecular Geometry Examples • NH3 • PO43- • CHFClBr • PF5 • SeF6 • NOCl • CO2 • SF2 • N2 • SiO3-2

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