The Bohr-Rutherford Model of the Atom

1. The Bohr-Rutherford Model of the Atom • The nucleus is made up of protons (p+) and neutrons (no) and is surrounded by rings of orbiting electrons (e-)

2. Standard atomic notation • Also called isotope notation X = element symbol A = atomic mass = # protons + # neutrons Z = atomic number = # protons

3. Isotopes • What is the difference between these two atoms?

4. Introduction: BONDING • Chemical reactions involve the formation of new products • Bonds between atoms or ions in the reactants must be BROKEN (this requires energy, and is therefore an ENDOTHERMIC process) • Bonds are then FORMED between atoms or ions to make the products of the reaction (this releases energy, and is said to be an ENDOTHERMIC process)

5. Intramolecular vs Intermolecular Forces • The bonds that are present between atoms within a molecule are due to INTRAmolecular forces of attraction – relatively STRONG forces of attraction that hold the molecule together • Forces of attraction between molecules, INTERmolecular forces of attraction, are what keeps molecules together as a solid, or liquid; • INTERmolecular forces of attraction must be overcome if a substance is to change its physical state from solid to liquid to gas: The strength of these intermolecular forces is affected by the type of bonding within the molecule

6. We will begin with a study of the types of bonding found in compounds, and then look at the characteristics of these bonds and how it affects the PHYSICAL PROPERTIES (like solubility, melting point, boiling point, conductivity and hardness) of different compounds

7. Ions • Ions are elements that have gained or lost electrons • Ions are commonly found dissolved in water, such as in the cytoplasm or plasma of the blood • Elements in the same family tend to form the same type of ion (e.g.: Na+, Li+, K+, Rb+) • Some important ions are Ca2+ (used for muscle contraction), Na+ and K+ (nerve and muscle function), Fe2+ and Fe3+(in hemoglobin) and H+ (required for synthesis of ATP)

8. Chemical Bonding • Electrons orbit the nucleus of an atom at a great distance compared to the size of the particles • Analogy: If an apple represented the size of an atom’s nucleus and it was placed at the center of the earth’s core, the valence electrons would be orbiting close to the surface of the earth’s crust • The valence electrons therefore are the part of the atoms that interact in chemical reactions to form compounds

9. Ionic Bonds • Ionic Compounds form when electrons are transferred from one atom to another to form ions with complete outer shells of electrons (ie the same stable electron configuration as the inert gases.) • Ionic BONDS form between a metal and a nonmetal eg NaCl • Metal tends to lose electrons which are transferred to the nonmetal • Metals form a cation (+) and nonmetals form an anion (-)

10. Ionic Bonds Formation of NaCl Formation of MgF2

11. Ionic Compounds • Positive and negative ions are held together in a strong lattice framework through the strong electrostatic attraction of oppositely charged ions • Ionic compounds have high melting points and are solids at room temperature • http://wps.prenhall.com/wps/media/objects

12. Ionic substances • These result in a lattice of ions rather than individual molecules, so we refer to MgF2 and NaCl as formula units, not molecules. • Properties of ionic substances as a result of the strong, rigid lattice that can dissociate when dissolved in water: • Crystalline solids at room temperature • Hard and brittle • High melting and boiling points • Conduct electricity when in liquid form • Most are soluble in water

13. Ionic Compounds are crystalline: Foodsubs.com Docbrown.info

14. Dissolving of NaCl in water: • http://www.youtube.com/watch?v=EBfGcTAJF4o&feature=related • http://www.youtube.com/watch?v=dr4sFNzUVzI&feature=related • NaCl solutions conduct electricity • http://www.youtube.com/watch?v=aELPrWzixeU&feature=related

15. The chemical formula of simple ionic compounds can easily be determined if one knows how many electrons an atom needs to lose or gain to achieve a stable inert gas electron configuration – this is related to its group number = # of valence electrons already present • In an ionic compound, the sum of positive and negative charges = zero to create a neutral compound e.g. MgS, Al2O3 CaBr2

16. Group # indicates charge for families 1,2,3 and 5,6,7

17. Formulas and Names of Ionic Compounds • Simple Ionic Compounds: metal + nonmetal • Eg sodium and bromine

18. Names and Formulas of ionic compounds continued……. • Calcium and chlorine

19. Simple Ionic Compounds • To name, state the name of the metal element first, the nonmetal element second and change the ending of the nonmetal element to “ide”. No prefixes are used. • eg NaCl sodium chloride • Eg MgBr2 magnesium bromide

20. Words  formula • Use the crossover rule to determine the formula, using the valence or ionic charge of each ion. The resulting compound should be neutral (net charge = zero) • Eg aluminum oxide eg lithium bromide

21. For ionic compounds, always reduce the ratio of metal : nonmetal to lowest terms • Eg Magnesium and oxygen

22. Alternatively, use an accounting method instead of the crossover: • Calcium and oxygen or calcium and chlorine

23. Multivalent Metal Ions

24. Polyatomic ions • Some ions are composed of a number of nonmetal elements bonded together covalently, with an overall net charge due to a surplus or deficit of electrons • Eg NH4+ the ammonium ion • Eg SO42- the sulphate/sulfate ion

25. Common Polyatomic ions

26. Memory Aid for Polyatomic Ions • “Nick the Camel ate a Clam for Supper in Phoenix” • First letter: the non metal element • # consonants = # of O’s in formula • # vowels = # of negative charges on the ion • OR • # of hydrogens in the related oxyacid

27. Example: Nick • “N” is for nitrogen N • 3 consonants = 3 O’s NO3 • 1 vowel = 1 negative charge NO3– • or 1 H in the oxyacid: HNO3(aq)

28. Now you try……. • Camel tells us:

30. Nick the Camel gives us the –ate polyatomic ions, aka the “MOTHER IONS” (by Mrs. Wheelihan) • When 1 oxygen is removed from the oxyion, the ending changes from -ate to –ite, BUT THE CHARGE STAYS THE SAME • EG NO3- nitrate and NO2- nitrite • SO42- sulfate and SO32- sulfite

31. This idea can be used to name a number of variations from the “mother ions”

32. Sometimes a hydrogen ion, H+, stays attached to a polyatomic ion • This creates a new polyatomic ion with the prefix “hydrogen” or “dihydrogen” • Eg hydrogen carbonate ion

33. Other examples: • PO43- phosphate • HPO42- hydrogen phosphate • H2PO41- dihydrogen phosphate • NaH2PO4 sodium dihydrogen phosphate

34. Peroxide Ion, O22- • PEROXIDES contain the peroxide ion, O22- instead of the oxide ion, O2- • example, hydrogen peroxide H2O2 • Peroxides are named by placing the prefix “per” before the word oxide in the name of the compound. When writing the formula of a peroxide, simply add an additional oxygen atom to the formula of the normal oxide, representing the peroxide ion, O22- • Eg lithium peroxide Li2O2

35. Covalent Bonds • Form between two nonmetals • Electrons are shared rather than transferred • Macromolecules and organic molecules are covalent molecules held together with covalent bonds. • Examples: lipids, carbohydrates, proteins and nucleic acids.

36. What holds a covalent bond together? • Nuclei repel each other, but are both attracted by the pair of negative electrons being shared.

37. Covalent Bonds Formation of H2 Formation of NH3

38. Other examples, with their Lewis diagrams:

39. Covalent Bonds Formation of O2 – a double bond

40. “Lone pairs” and “Shared pairs” • Non-bonded pairs are called “LONE PAIRS” • Pairs of electrons that are shared between atoms are called: ”BONDING PAIRS” • BOTH LONE PAIRS AND BONDING PAIRS MUST BE SHOWN ON YOUR LEWIS DIAGRAM • ie ALL VALENCE ELECTRONS must be shown

41. Bonding pairs and Lone Pairs Alevelchem.com

42. Lewis Structures of Covalent Molecules • Covalent molecules can be represented in a number of different ways. In Lewis structures, ONLY THE VALENCE ELECTRONS ARE SHOWN. • In Lewis structures, the outside electrons are shown with dots and covalent bonds are usually shown by bars, but there are several ways of drawing Lewis structures (see Figures 5 to 9, Neuss (2007) pages 64, 65)

43. There are a number of different ways to represent Lewis structures of covalent molecules: These can also be drawn with lines or crosses to represent lone pairs and bonding pairs of electrons

44. . Covalent Bonds usually involve the sharing of an electron pair consisting of 1 electron from each molecule involved in the bond Chemteam.info

45. Coordinate or Dative Covalent Bonds • Sometimes the “bonding pair” of electrons both come from the same element. This type of bond is called a “coordinate or dative covalent bond”. • Eg ammonia Chewtychem.wiki.edu

46. Other examples of dative covalent bonds (also called coordinate bonds) are found on page 65 of Neuss (2007) and include carbon monoxide, CO(g); the hydronium ion, H3O+; and aluminum chloride dimer

47. Single Covalent Bond • When 1 pair of electrons is shared between the same two atoms, a SINGLE COVALENT BOND is formed

48. Double Covalent Bonds • When 2 pairs of electrons are shared between the same two atoms, a DOUBLE COVALENT BOND is formed as found in CO2 (g), carbon dioxide or carbon (IV) oxide • Carbon dioxide CO2http://dbhs.wvusd.k12.ca.us

49. Triple Covalent Bonds • When 3 pairs of electrons are shared between atoms, a triple bond is formed • Eg, nitrogen gas, N2 http://www.webchem.net/notes/chemical_bonding/covalent_bonding.htm

50. BOND LENGTH and BOND STRENGTH • The more pairs of electrons shared between atoms, the stronger and shorter the bond • Therefore single bonds tend to be longer and weaker than double bonds and triple bonds are shorter and stronger than corresponding double bonds, as shown in Table 1 below