Presentation is loading. Please wait.

Presentation is loading. Please wait.

Unit 7: Bonding.

Similar presentations

Presentation on theme: "Unit 7: Bonding."— Presentation transcript:

1 Unit 7: Bonding

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

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 K2S CaCl2

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) , MgCl2 (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) O2, CO, CCl4, H2

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) Br2

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) F2

20 What if we have more than two non-metals?
Ex) H2O Ex) NH3

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) CH4 Ex) CCl4

22 Practice Problem Draw the Lewis Structure for CH3Br

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 CO2

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) NH4+

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? F2
Cl2 Br2 I2

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. H2O, NH3 ***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! (C12H22O11)

40 Network Solids Properties: A special case of covalent bonding
Atoms held together in a very strong covalent network Ex. Carbon (Diamond) and SiO2 Properties: Hard High m.p. and b.p. Poor conductors

41 Metallic Bonding 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”

42 Bond Energy Ex) F(g) + F(g)  F2 (g) + ENERGY
Bonds do not break and form spontaneously- an energy change is required When a bond is broken, energy is ABSORBED (required) Ex) F2 (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)  F2 (g) + ENERGY

Download ppt "Unit 7: Bonding."

Similar presentations

Ads by Google