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Covalent Bonds and Molecular Forces Chapter 6. Sharing electrons Sodium atom reacts with chlorine gas to form the ionic compound sodium chloride, NaCl,

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Presentation on theme: "Covalent Bonds and Molecular Forces Chapter 6. Sharing electrons Sodium atom reacts with chlorine gas to form the ionic compound sodium chloride, NaCl,"— Presentation transcript:

1 Covalent Bonds and Molecular Forces Chapter 6

2 Sharing electrons Sodium atom reacts with chlorine gas to form the ionic compound sodium chloride, NaCl, is an example of this type of reaction. The reaction of hydrogen and oxygen to form water is another kind of rearrangement where electrons are shared.

3 Molecular and Atomic orbitals The simplest example of sharing electrons occurs mainly in diatomic molecules such as H₂, and O₂. When two hydrogen atoms approach each other, the positive nucleus of each atom attracts its own electron and the electrons of the other atom. At the same time the positive nuclei of the two atoms repel each other. Likewise the electron cloud of the atoms repel. Since they are both of the same atom neither has enough attraction to take an electron from the other. Instead of forming ions, the 2 hydrogen atoms share electrons. The shared electrons moving about in space surrounding the the two nuclei are in molecular orbital.

4 Molecular orbital is a region where an electron pair is most likely to exist as it travels in the three dimensional space around the nuclei. A bond formed when two or more valence electrons are attracted by the positively charged nuclei of two atoms and are thus shared between both atoms.

5 Potential energy curve for H₂ As a covalent bond forms between two atoms, they reach a distance from each other at which the attractive and repulsive forces are balanced and the energy is at the minimum.

6 As the two hydrogen atoms come nearer the potential energy of the combination becomes lower and lower until it reaches the minimum value of -436kJ/mol at a distance of 75pm.At the lowest energy, the H-H combination is most stable because lower energy means greater stability. At the distance of 75pm, the repulsion between the like charges equals the attraction of the opposite charges.This is the bond length.

7 Diatomic Molecules Bond length- The distance between two bonded atoms at their minimum potential energy; the average distance between two bonded atoms. The energy required to break a bond between two atoms is the bond energy.

8 Electronegativity and bonding The tendency of an atom to attract bonding electrons to itself when it bonds with another atom is electronegativity. To help explain why some combinations of atoms form ionic bonds and some form covalent bonds, this concept was developed by Linus Pauling. In general electronegativity decreases down a group and increases across a period.


10 Polar and non polar covalent bonds In a molecule such as H₂, the atoms are identical, so they pull on the bonding electrons with the same force. The electrons are shared equally. Such a covalent bond, in which the bonding electrons are shared equally, is called a nonpolar covalent bond. In other words the electronegativities of two atoms are equal. If the electronegativities are greatly different, an ionic bond is formed.

11 There is a bond between these two extremes in which electrons are shared but not equally. These bonds are called as polar covalent bonds.

12 In the previous example of polar covalent bond oxygen attracts electrons more strongly than the other atoms. Polar molecules have both positive and negative charges. Example hydrogen fluoride. The electronegativity of fluorine is much higher than the electronegativity of hydrogen. The fluorine atoms attracts electron much more than hydrogen atoms.

13 The hydrogen having its electron pulled away has a partial positive charge and the fluorine has a partial negative charge. This is not an ionic bond. A molecule that has a partial positive charge on one end and partial negative charge on the other end is called a dipole.

14 Homework Page 201 Q.5 and Q.6

15 Electron Dot Structures Valence electrons are electrons in the outermost energy level of an atom, where it can participate in bonding. Lewis Structure is a structure in which atomic symbols represent nuclei and inner shell electrons, and dots are used to represent valence electrons. Consider a chlorine atom, which has the electronic configuration 1s²2s²2p⁶3s²3p⁶.

16 Only the electrons in the outermost energy level are involved in bonding., so in the Lewis structure only seven valence electrons are represented by dots.

17 Rules for Drawing Lewis Structures with many Atoms 1. Hydrogen or halogen atoms often bind to only one other atom and are usually on the outside or the other end of the molecule. 2. The atom with the lowest electronegativity is often the central atom. These atoms often have fewer than seven electrons and may form more than one bond. 3. When placing valence electron s around an atom, place one electron on each side before pairing any electrons.

18 Class Practice Draw Lewis structure for iodine monochloride, ICl and hydrogen bromide HBr. Draw Lewis Structure for formaldehyde CH 3 OH.

19 Resonance Structures A possible Lewis dot structure of a molecule for which more than one Lewis structure can be written.

20 Class Practice Page #211 Q.4 all

21 Naming Covalent compounds The most common naming system uses prefixes, roots and suffixes. Example Carbon dioxide and carbon monoxide. Prefixes and suffixes are usually attached to root words. For binary compounds the root word is the name of the element. The first element named is usually the one first written in the formula which is the least electronegative element.

22 If the molecule contains only one atom of the first element given in the formula, the prefix mono is omitted in the name of the compound. For example, to distinguish between the two oxides of carbon, the prefixes mono and di are used.

23 Home work page 213 Q.7. b and d Q.9. all Q.11. a and c

24 Molecular shapes The shape of a molecule can be predicted by the Lewis Structure. In a molecule of only two atoms, such as HF, or H₂, only as linear shape is possible. Molecules of more than two atoms, molecular shapes will vary. Example CO₂ and SO₂ their formulas are similar then why carbon dioxide is linear, while sulfur dioxide is bent?

25 Different possible shapes

26 Molecular geometry based on electron pairs CO₂ SO₂

27 There is a simple model that can be used to determine the three dimensional arrangement of the atoms in a molecule. This model is based on the valence shell electron pair repulsion (VSEPR) theory. According to this theory you can predict the shape of a molecule by knowing the electron pairs around a central atom.

28 Steps in determining the geometry of a molecule or polyatomic ion. 1.For a molecule,count the number of electron pairs surrounding the central atom. Each single or multiple bond counts as one electron group. Each nonbonding electron pair counts.

29 There are two double bonds around the central carbon atom. Therefore there are two electron groups around this atom. Electron groups have negative charge, and like charges would repel each other and will remain as far apart as possible. If a central atom has 2 electron groups they will be linear.

30 In SO₂ one of the electron pairs is a lone pair. The nonbonding electrons repel the bonding electrons, causing the three electron pairs to orient in a trigonal planar geometry.

31 When a molecule consists of a central atom bonded to three other atoms its shape will be trigonal planar as long as there are only three electron groups determining the geometry.

32 SO₃ can have resonance structures. Because of the resonance structure the three electron clouds surrounding the sulfur atom in SO₃ are identical in order to be as far as possible the groups arrange like three spokes of a wheel, extending out from the sulfur atom. This geometry is called as trigonal planar.The angle between them will be 120⁰.

33 When there are 4 molecules surrounding a central atom the electron pairs are farthest away when they orient themselves towards the corner of a tetrahedron.

34 If the pairs are all equivalent (if there are four identical bonds on the central atom) all four of the angles between the bonds are 109.5⁰.

35 Class Practice Page 217 Q.1 & Q.2.

36 Shape and property The shape of a molecule affects the chemical properties. Many of these properties depend on the polarity of the molecule. For molecules containing more than two atoms, molecular polarities depends on both the polarity of each atom and its orientation. The bond polarities in a water molecule add together, causing a molecular dipole. In carbon dioxide, bond polarities extend in opposite direction, cancelling each other.

37 H 2 O, CO 2, and CH 4 molecules

38 The double bonds between C and CO 2 are polar because oxygen attracts electrons more strongly than does carbon. However, the linear shape of the molecule causes the two bond dipoles to act in opposite directions, canceling each other and causing the molecular polarity to be zero. In the water molecule, the polar H-O bonds are oriented at a 105⁰ angle to each other, which creates a dipole for the molecule.

39 Properties Of Compounds Covalent compounds melt at a lower temperature than ionic compounds. Ionic compounds consists of ions each of which is attracted to all ions of opposite charges. These attractions hold the ions tightly in a crystal lattice that can be disrupted only by heating to very high temperatures.

40 Intermolecular forces The attraction that exists between molecules are called as intermolecular forces. If there are no intermolecular forces between molecules then the substance exists as gases.

41 Intramolecular force An intra molecular force is any force that holds together the atoms making up a molecule. Intra molecular forces of attraction (covalent) are stronger than the intermolecular forces of attractions. The stronger the intermolecular forces, the higher the melting and boiling point of the substance.

42 Dipole Forces Dipole forces affect the melting and boiling points. In a polar molecule we have one end of the molecule having partial positive charge and the other end having a partial negative charge.The positive end of a molecule can attract the negative end of another molecule holding the two molecules together. This force that exists between the two positive and negative ends is called as dipole force.

43 The dipole forces help the molecules to exist as a solid or a liquid. Oxygen and methane are non polar molecules but molecules of water and ethyl acetate are polar. Because of the dipole forces these have higher melting and boiling points.

44 Hydrogen bonds are stronger dipole forces. Hydrogen atom bonded to an atom that is more electronegative. HF has a very high boiling point and HCl has the lowest. There is a strong dipole force between HF molecules due to large electronegativity difference between H and F.

45 Hydrogen also has just one electron, and when that electron is pulled away then there are no electrons to protect the nucleus so the proton in the nucleus is attracted to the electron rich fluorine end of another HF molecule. Hydrogen bonds are usually formed with small atoms with a high electronegativity, like oxygen, fluorine and nitrogen. The HCl, HI, and HBr molecules are polar but they are much larger than HF, the distance between the molecules is greater and so the hydrogen bonds are weaker.

46 Water’s unique properties. Water has a high boiling point as the molecules of water are held together by hydrogen bonds. Water (H₂O) has a higher b.p than hydrogen sulfide (H₂S) as the electro negativity difference between H and S is 0.4 and that of H and O is 1.2. As a result, the hydrogen bonds in H₂O is stronger than H₂S.

47 London forces/ Van der Waal’s forces In case of noble gases the boiling points increases in this way. Higher boiling point indicate the addition of electrons in the atoms and hence strong bonds. This was explained by Fritz London.

48 London forces are an attraction between atoms and molecules caused by the formation of instantaneous dipoles in the atoms and molecules because of the unequal distribution of electrons around the nucleus or nuclei.

49 Homework Page 227 Term Review all Page 228 14, 16 and 17.

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