Chemical Bonding

1. Chemical Bonding Comparison of Properties Ionic Compounds Covalent Compounds Metals

2. Macroscopic properties of matter vary greatly due to the type of bonding Properties of Matter

3. Types of Bonding Ionic • Metallic • Covalent • ionic Metallic Covalent

4. What is a chemical bond? • An attractive force that holds two atoms together • Can form by • The attraction of positive ion to a negative ion or • The attraction of the positive nucleus of one atom and the negative electrons of another atom

5. Bond • the interaction between two or more atoms that allows them to form a substance different from the independent atoms. • involves the outer (valence) electrons of the atoms. • These electrons are transferred from one atom to another or shared between them.

6. Can be between atoms of different elements to make a compound, like the two hydrogen atoms and one oxygen atom in a water molecule.

7. Sulfur • can also be between atoms of a single element. • Sulfur is an example of an element that has its most stable form as a small molecule, in this case a ring of eight sulfur atoms.

8. Chemical Bond Energy Considerations • A chemical bond forms when it is energetically favorable • when the energy of the bonded atoms is less than the energies of the separated atoms.

9. Bonding • Chemical compounds are formed by the joining of two or more atoms. • A stable compound occurs when the total energy of the combination has lower energy than the separated atoms. • The bound state implies a net attractive force between the atoms ... a chemical bond.

10. Energy Changes in Bonding • When bonds are formed, energy is released. • Demonstrations: • Formation of an Ionic Compound: Mg + O2 • Formation of a Molecular Compound: S +O2

11. Breaking Bonds • In order to break bonds energy must be added, usually in the form of heat, light, or electricity. • Demonstration: Electrolysis of water http://www.youtube.com/watch?v=HQ9Fhd7P_HA • Demo: Decomposition of Nitrogen Triiodidehttp://www.youtube.com/watch?v=z5vsQ8sPgX4

12. In chemical bonds, atoms can either transfer or share their valence electrons.

13. Ionic Bonds • When one or more atoms lose electrons and other atoms gain them in order to produce a noble gas electron configuration, the bond is called an ionic bond.

14. Bonding Animation: • http://www.youtube.com/watch?v=QqjcCvzWwww • Intro to bonding • http://www.youtube.com/watch?v=31bec8Hl7wo&NR=1&feature=endscreen

16. Ionic Solids • Ionic solids are solids composed of ionic particles (ions). • These ions are held together in a regular array by ionic bonding. • Ionic bonding results from attractive interactions from oppositely charged ions. • In a typical ionic solid, positively charged ions are surrounded by negatively charged ions and vice-versa. • The close distance between these oppositely charged particles results in very strong attractive forces. • The alternating pattern of positive and negative ions continues in three dimensions. • The regular repeating pattern is analogous to the tiles on a floor or bricks on a wall. • called the crystal lattice.

17. Ionic Compounds • Crystalline solids (made of ions) • High melting and boiling points • Conduct electricity when melted or dissolved in water • Demo: Electrolytes • Many are soluble in water but not in non-polar liquid

18. Comparison of Conductivity

19. Common Ionic Compounds • NaCl - sodium chloride - table salt • KCl - potassium chloride - present in "light" salt (mixed with NaCl) • CaCl2 - calcium chloride - driveway salt • NaOH - sodium hydroxide - found in some surface cleaners as well as oven and drain cleaners • CaCO3 - calcium carbonate - found in calcium supplements • NH4NO3 - ammonium nitrate - found in some fertilizers

20. Ionic compounds in the solid state are held together by electrostatic attractions between opposite charges. • Sodium chloride (table salt), silver sulfide (silver tarnish), and hydrated iron (III) oxide (rust) are examples of ionic compounds.

21. Covalent (Molecular) Compounds • Gases, liquids, or solids (made of molecules) • Low melting and boiling points • Poor electrical conductors in all phases • Many soluble in non-polar liquids but not in water

22. Molecular (Covalent) Substances

23. Covalent Network Solids • Covalent because combinations of nonmetals • Interconnected • very hard and brittle • Insoluble • Extreme melting and boiling points Diamond

24. Covalent Bonds • involve the sharing of a pair of valence electrons by two atoms • Such bonds lead to stable molecules if they share electrons in such a way as to create a noble gas configuration for each atom

25. Covalent bonding can be visualized with the aid of a Lewis Structure

26. Polar Covalent Bonds • Covalent Bonds in which the sharing of the electron pair is unequal • the electrons spend more time around the more nonmetallic atom • In such a bond there is a charge separation with one atom being slightly more positive and the other more negative……. will produce a dipole moment.

27. Types of Covalent bonds • Pure Covalent (also called non-polar covalent) bonds are ones in which both atoms share the electrons evenly • By evenly, we mean that the electrons have an equal probability of being at a certain radius from the nuclei of either atom. • Polar covalent bonds are ones in which the electrons have a higher probability of being in the proximity of one of the atoms • Determined by Electronegativity Difference

28. Electronegativity • the periodic property that indicates the strength of the attraction an atom has for the electrons it shares in a bond. • Atoms with high electronegativities tend to hold tightly to their electrons or to form negative ions. • These elements are found to the upper right on the periodic table. • Atoms with low electronegativities tend to have a lower attraction for their electrons and may form positive ions. • These elements are found to the lower left on the periodic table.

29. Pure covalent or Non-polar covalent bond • Electronegativity difference of 0.3 or less in between the two atoms. • A pure covalent bond can form between two atoms of the same element (such as in diatomic oxygen molecule) • or atoms of different elements that have similar electronegativies (such as in the carbon and hydrogen atom in methane).

30. Polar Covalent Bond • A is a pair of electrons shared between two atoms with significantly different electronegativities (from 0.3 to 1.7 difference). • These bonds tend to form between highly electronegative non-metals and other non-metals, such as the bond between hydrogen and oxygen in water.

31. Ionic Bonds • In compounds that have elements with very different electronegativities (greater than 1.7 difference), the electrons can be considered to have been transferred to form ions.

32. Many of the properties of a compound, such as solubility and boiling point, depend, in part, on the degree of the polarity of its bonds.

33. Examples to Determine Bond Character • Using electronegativity in the prediction of the polarity of a chemical bond. • sodium bonded to chlorine • Difference between the electronegativities of Na(0.9) and Cl(3.0) are so great that they form an ionic bond. • The hydrogen molecule (2 H atoms bonded to each other) • zero electronegativity difference, form a non-polar covalent bond.

34. Bond Character • Nonpolar-Covalent bonds (H2) • Electrons are equally shared • Electronegativity difference of 0 to 0.3 • Polar-Covalent bonds (HCl) • Electrons are unequally shared • Electronegativity difference between .3 and 1.7 • Ionic Bonds (NaCl) • Electrons are transferred • Electronegativity difference of more than 1.7

35. Diatomic Molecules • hydrogen gas H2 • the halogens: • chlorine Cl2 • fluorine F2 • bromine Br2 • iodine I2 • Nitrogen N2 • Oxygen O2 Pneumonic Device to remember the diatomic molecules: Professor BrINClHOF

36. Metals and Metallic Bonding • Typical Properties of Metals • Malleable • Ductile • Good Conductors of Heat and Electricity • Generally high melting and boiling points

37. Metallic Bonds • The properties of metals suggest that their atoms possess strong bonds • yet the ease of conduction of heat and electricity suggest that electrons can move freely in all directions in a metal • The general observations give rise to a picture of "positive ions in a sea of electrons" to describe metallic bonding.

38. Metal Properties • Malleable and Ductile • Strong and Durable • Good conductors of heat and electricity. • Their strength indicates that the atoms are difficult to separate… strong bonds • but malleability and ductility suggest that the atoms are relatively easy to move in various directions. • The electrical conductivity suggests that it is easy to move electrons in any direction in these materials. • The thermal conductivity also involves the motion of electrons. All of these properties suggest the nature of the metallic bonds between atoms.

39. Metallic Bonding the Electron Sea Model • Explained by the Electron Sea Model • the atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic cations. • delocalized electrons are not held by any specific atom and can move easily throughout the solid. • A metallic bond is the attraction between these electrons and the metallic cation.

40. A mixture of elements that has metallic properties is called an alloy. • Two types of alloys • An interstitial alloy is one in which the small holes in a metallic crystal are filled by other smaller atoms. • A substitutional alloy is one in which atoms of the original metal are replaced by other atoms of similar size.