Chemical Bonds

1. Chemical Bonds Force that holds atom together

2. Topics • Stability in Bonding • Valence Electrons • Reactivity • Types of Bonds • Ionic, Metallic, Covalent, Hydrogen bond • Writing Formulas and Naming Compounds • Binary compound • Polyatomic ion • Hydrate

3. Reactivity • Definition: Reactivity describes how likely an element is to form bonds with other elements. • Reactivity of an element is determined by its valence electrons.

4. Valence Electrons • Definition: Valence electrons are electrons in the outermost energy shell (level) that are available to be lost, gained or shared when elements form chemical bonds. • Elements in the same group of the Periodic Table have similar chemical properties because of their valence electrons.

5. Element GroupsAnd Valence Electrons

6. Stability in Bonding • Some elements combine chemically and no longer have the same properties they did before forming a compound. • A chemical formula is composed of symbols and subscripts indicating the number of atoms of an element in a compound. • E.g. NaCl, H2O, CaCO3, CO2

7. Atoms form compounds when the compound is more stable than the separate atoms. • Noble gas are more chemically stable. • Elements that do not have full outer energy shell are more stable in compounds. • Atoms can lose, gain, or share electrons to get a stable outer energy level • A chemical bond is the force that holds atoms together in a compound.

8. Types of Bonds

9. Intramolecular Force • Definition: force and its property within the structure of a single molecule. • Ionic, Covalent, and Metallic bond

10. Intermolecular Force • Definition: force that act between stable molecules or between functional groups of macromolecules. • Example: dipole-dipole interaction, hydrogen bonds, di-sulphide bonds Cystein-Cystein

11. Chemical bond • Definition: The force that holds atoms together in a compound. • Ions are charged particles because it has more or fewer electrons than protons. • When an atom loses an electron, it becomes a positively charged ion; a superscript indicates the charge, e.g., H+, Ca2+, Li+ • When an atom gains an electron, it becomes a negatively charged ion, e.g., HCO3-, OH-

12. Ionic Bond • Definition: Force of attraction between opposite charges of the ions, often forming (crystal) lattice. • An ionic compound is held together by the ionic bond, often joining a metal to a nonmetal • The result of this bond is neutral compound. • The sum of the charges on the ions is zero.

13. Covalent Bond • Definition: the force of attraction between atoms sharing electrons. • Molecules are neutral particles formed as a result of sharing electrons. • Atoms can form double or triple bonds depending on whether they share two or three pairs of electrons. • Covalent bonds are usually involved when two nonmetals form a compound or when a nonmetal bond with a metalloid.

14. Electrons shared in a molecule are held more closely to the atoms with the larger nucleus. • A polar molecule has one end that is slightly negative and one end that is slightly positive although the overall molecule is neutral. • In a nonpolar molecule electrons are shared equally.

15. Main Difference • Ionic bonds form when atoms lose or gain electrons; covalent bonds form when atoms share electrons.

16. Metallic Bond • electromagnetic interaction between delocalized electrons, called conduction electrons and gathered in an "electron sea", and the metallic nuclei within metals. • Form metallic bond joins metal to metal. • This type of bonding is collective in nature. There is no single metallic bond.

17. Electronegativityaka Electron Affinity • Electronegativity measures the ability of an atom to attract electrons. Fluorine has the highest electronegativity.

18. Ionization Energy • Definition: The amount of energy needed to remove an electron from a neutral atom is called ionization energy. • Moving from left to right, ionization energy increases across a period. • Moving from top to bottom in a group, it generally decreases.

19. Drawing Chemical Bond • Bohr’s Model • Valence shells are drawn in circles for s suborbital (Group 1, 2) and p suborbital (Group 13 thru 18) • Protons and Neutrons are indicated in nucleus. Electrons in shells travel in pairs. • Lewis (Electron) Dot Diagram • Easier way to illustrate electrons on outermost energy shell and their reactivity. • Charged atoms and molecules are placed in brackets with their charges indicated at top right corner.

20. Octet Rule • Octet rule is a simple chemical rule of thumb that states atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. • This rule is applicable to main group elements.

21. Valence Shells

22. Bohr’s Model

24. Bond Using Bohr’s Model

27. Electron Dot (Lewis) Diagram

28. Chemical Nomenclature • Nomenclature: The system of naming in arts or science • Writing chemical formulas and naming compounds • Binary Metal-Nonmetal Compounds • Binary Nonmetal-Nonmetal Compounds • Polyatomic ions • Hydrate

30. Oxidation Number • Oxidation Number – # of e- gained, lost, or shared to become stable • Determine the oxidation number from the periodic table.

31. Balancing Oxidation Number • Ba + F

32. Binary Metal-Nonmetal Compounds • Composed of two elements • When writing binary compounds: • Use the least common multiples of oxidation number. • Compound is neutral. • Charges on the ions must be balanced. • Use first element, the root name of the second element, and the suffix -ide.

33. If the metal has more than one possible charge, indicate the ion with charge in roman numerals • FeCl2 - Iron (II) chloride. • Alternatively, common names may be used. Add –ous for the lower charge and –ic for the higher charge: • FeCl2 – ferrous chloride • FeCl3 – ferric chloride

34. More examples

35. Binary Nonmetal-Nonmetal Compounds • Add –ide to the second element • Use Greek prefixes for number of atoms: • Examples: • CO – carbon monoxide • CO2 – carbon dioxide • N2O5 – dinitrogen pentoxide

36. Common names: –ous and –ic (-ic has greater charge, OR has fewer atoms)

37. Polyatomic Ion • The compound contains three or more elements. • When writing names: • Positive ions first, followed by negative ion • Use the least common multiple of oxidation number, and put parenthesis around the polyatomic ion before adding a subscript

38. Balancing Oxidation Number • K + OH Ca + NO3

39. Polyatomic Compounds • Ammonium ion NH41+ • -ide ions: CN1- cyanide, OH1- hydroxide • Oxyanions • -ate are more oxygen • NO21- nitrite • NO31- nitrate • Some oxyanions have extra hydrogen • SO42- sulfate • HSO41- hydrogen sulfate (or bisulfate)

40. Oxyanions (continued) • If more than two possibilities

41. Naming compounds with polyatomic ions • Positive charge species on left • Negative charge species on right • Use parenthesis as needed

42. Hydrate • Compound with water chemically attached to its ions • Two types: • Hydro Acids: Hydro + Halogen name + ic • HCl – hydrochloric acid • HF – hydrofluoric acid • Oxoacids: polyatomic ion + acid • HNO3 – nitric acid, from nitrate • HNO2 – nitrous acid, from nitrite