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UNIT 7: BONDING. Why Bond?  Elements bond in order to get a full valence shell  Elements want to have the same number of electrons as the noble gases.

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Presentation on theme: "UNIT 7: BONDING. Why Bond?  Elements bond in order to get a full valence shell  Elements want to have the same number of electrons as the noble gases."— Presentation transcript:


2 Why Bond?  Elements bond in order to get a full valence shell  Elements want to have the same number of electrons as the noble gases  The Octet Rule: in forming compounds, atoms tend to achieve the electron configuration of a noble gas **** only VALENCE electrons are involved in bonding

3 Types of Bonds  Ionic  Covalent  Metallic

4 Review of Metals  Metals lose electrons to form positive ions (cations)  Ionic radius is smaller than atomic radius  Group 1: Alkali metals  Lose 1 electron  Group 2: Alkaline Earth  Lose 2 electrons

5 Review of Nonmetals  Gain electrons to form negative ions (anions)  Ionic radius is larger than atomic radius  Group 17: Halogens  Gain 1 electron  Group 16: Chalcogens  Gain 2 electrons

6 Ionic Bonds  The force of attraction that holds ions of opposite charge together  This attraction is between a metal and a nonmetal  Forms from the TRANSFER of electrons  Electrons are transferred from METAL  NONMETAL  Large difference in electronegativity  Overall charge of compound is neutral (+ and – cancel out)  Ex) NaCl

7 Animation  s/chang_7e_esp/bom1s2_11.swf s/chang_7e_esp/bom1s2_11.swf

8 Writing Lewis Dot Structures for Ionic Compounds  Lewis Dot diagrams help us visualize what is happening to valence electrons when a bond forms  Remember- Follow the octet rule!! Ex) Li and F Steps: 1) Draw all elements and their individual Lewis Dot diagrams. 2) Draw an arrow indicating the transfer of electrons 3) Redraw your bonded compound with appropriate charges

9 More Practice… Draw the Lewis Dot Diagrams for the following ionic compounds: BaO K 2 S CaCl 2

10 What if you are not given the formula? Draw the Lewis Dot diagram for the ionic bonding between sodium and sulfur.

11 Ionic Compounds  An ionic compound exists as a collection of positively and negatively charged ions arranged in repeating patterns  A formula unit is the lowest whole-number ratio of ions in an ionic compound NaCl (1:1 ratio), MgCl 2 (1:2 ratio)

12 Properties of Ionic Compounds (salts)  Ionic bonds are the strongest bonds, so all are solids  Hard  High melting points and boiling points  Soluble in water  Cannot conduct electricity as a solid  Can conduct electricity as a liquid or in an aqueous solution  They are electrolytes- substances that conduct electricity when dissolved in water

13 Covalent Bonds  A bond between two nonmetals that involves a sharing of electrons (“tug of war”)  Can have EQUAL or UNEQUAL sharing  Small or no difference in electronegativity  Overall charge of a covalent compound is neutral  Covalent compounds are called molecules  Ex) O 2, CO, CCl 4, H 2

14 Types of Covalent Bonds  There are 3 types of covalent bonds Covalent Bonds Non-Polar Covalent Polar covalent Coordinate Covalent

15 Non-Polar Covalent Bond  Electrons are shared equally  Same atoms- same electronegativity  All diatomics have non-polar bonds! Remember: HOFBrINCl  Two of the same non-metal  Ex) Br 2

16 Polar Covalent Bond  Electrons are shared unequally  Different atoms with different electronegativity values  Ex) HF  The more electronegative atom attracts electrons more strongly and gains a slightly (-) charge  The less electronegative atom has a slightly (+) charge

17 Bond Polarity Bond Polarity: refers to a separation of charge in a bond Ex) HF : partial (+) charge on H and partial (-) charge on F This separation of charge is often called a dipole ***The greater the difference in electronegativity, the greater the polarity.

18 Challenge Question Which of the following covalent compounds has the greatest degree of polarity? Choice 1) CO Choice 2) HCl Choice 3) NO Choice 4) HBr

19 Drawing Lewis Structures for Covalent Compounds Remember! After drawing your diagram, all atoms MUST have 8 valence electrons (except for hydrogen, which should have 2 electrons) Ex) F 2

20 What if we have more than two non-metals? Ex) H 2 O Ex) NH 3

21 Lewis Structures Continued…  When elements from Group 14 are involved in a covalent bond, they spread their e- out  Carbon tends to form 4 covalent bonds Ex) CH 4 Ex) CCl 4

22 Practice Problem  Draw the Lewis Structure for CH 3 Br

23 Can you have more than one bond?  Yes! So far all we have seen are single covalent bonds in which one pair of electrons is shared  Double covalent bond: a bond between two atoms where 2 pairs of e- are shared Ex) O and O *** Oxygen tends to form 2 bonds!  Triple covalent bond: a bond between two atoms where 3 pairs of e- are shared Ex) N and N *** Nitrogen tends to form 3 bonds!

24 More Practice with Lewis Diagrams.. Draw the Lewis diagrams for CO 2

25 Challenge Problem Draw the Lewis Structure for HCN

26 Let’s Summarize C, H, O and N form how many bonds??

27 Coordinate Covalent Bond  A bond in which both electrons of the shared pair come from the same atom Ex) NH 4 +

28 VSEPR Geometry  Lewis Dot diagrams fail to show the 3-dimensional shapes of molecules  VSEPR Theory (Valence-Shell Electron Pair Repulsion Theory)  Repulsion between e- pairs causes molecular shapes to adjust so that valence e- pairs stay as far apart as possible shapes

29 Molecular Polarity  Molecular polarity is different from bond polarity!!  In a non-polar molecule, electron distribution is even (symmetrical)  In a polar molecule, electron distribution is uneven (asymmetrical)

30 Intermolecular Forces (IMFs)  Only between covalent molecules, never ionic compounds  Weak forces that act between molecules and hold molecules to each other ****IMFs are not bonds!!!  IMFs occur BETWEEN molecules, bonding occurs WITHIN molecules

31 IMFs vs. Bonds

32 Intermolecular Forces of Attraction Van der Waals Forces Dipole interactions Hydrogen Bonds

33 London Dispersion Forces  London Dispersion forces: weakest of all molecular interactions; caused by temporary shifts in charge - Between nonpolar molecules london-forces.shtml - The bigger the atom or molecule  the greater the strength of dispersion forces  the higher the BP

34 Think about it…  Which has the strongest London dispersion forces? F 2 Cl 2 Br 2 I 2

35 Dipole-Dipole Interactions  Dipole interactions: attraction between polar molecules  The positive and negative charges of different molecules attract each other  Ex. HCl

36 Hydrogen bonds: A special case of dipole-dipole interactions  Hydrogen bonds: intermolecular force between the H of one molecule and a highly electronegative atom of another molecule (must be N, O, or F)  Ex. H 2 O, NH 3  ***The high b.p. of water is due to hydrogen bonding

37 Molecule-Ion Attraction  Partial charges on a molecule are attracted to ions  This is what happens when NaCl dissolves in water  Picture:  The hydrogens of water align themselves towards the anion and the oxygen align themselves towards the cation  ssolveslike.htm

38 Molecule-Ion Attraction

39 Properties of Covalent Compounds  Soft  Low melting points and boiling points  Cannot conduct electricity in any phase  Generally insoluble in water  Except sugars! (C 12 H 22 O 11 )

40 Network Solids  A special case of covalent bonding  Atoms held together in a very strong covalent network  Ex. Carbon (Diamond) and SiO 2  Properties:  Hard  High m.p. and b.p.  Poor conductors

41  Holds metals together  Electrons are mobile and move from one atom to another, creating (+) charged metal ions  Charged metal ions are immersed in a “sea of mobile electrons” Metallic Bonding

42 Bond Energy  Bonds do not break and form spontaneously- an energy change is required  When a bond is broken, energy is ABSORBED (required)  Ex) F 2 (g) + ENERGY  F (g) + F (g)  The greater the # of bonds between atoms, the more energy you need to break them  When a bond is formed, energy is RELEASED (given off)  Ex) F(g) + F(g)  F 2 (g) + ENERGY

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