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Chapter 6

Chapter 6. Chemical Bonds. Stable Electron Configurations. In Chapter 4, we talked about this… How electrons are arranged in the atom We also mentioned ground and excited state Ground: Stable Excited: Not stable There is another type of stableness that we can discuss in an atom.

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Chapter 6

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  1. Chapter 6 Chemical Bonds

  2. Stable Electron Configurations • In Chapter 4, we talked about this… • How electrons are arranged in the atom • We also mentioned ground and excited state • Ground: Stable Excited: Not stable • There is another type of stableness that we can discuss in an atom

  3. Stable as a Table • Energy levels are stable when the outer ring is completely filled up • Think back to valence electrons and it would be when they have 8 electrons in their ring • Which group of elements did we mention previously that already has 8 v electrons

  4. Connect the dots • Chemical properties are based on valence e- • Useful to have a model with valence electrons • Electron dot diagram: Model of an atom that focuses only on valence electrons • Symbol: Represents nucleus • Dots: Represent electrons

  5. Electron Dot diagrams • A way of keeping track of valence electrons. • How to write them • Write the symbol. • Put one dot for each valence electron • Don’t pair up until they have to X

  6. The Electron Dot diagram for Nitrogen • Nitrogen has 5 valence electrons. • First we write the symbol. N • Then add 1 electron at a time to each side. • Until they are forced to pair up.

  7. Write the electron dot diagram for Na Mg C O F Ne He

  8. Why bond together? • Elements that do not have full valence electrons usually are involved in reactions • By reacting, they achieve a goal to have their electron configurations to be the most stable • Two types: Ionic and Covalent Bonds

  9. Ionic Bonding • Let’s look at Chlorine and Sodium EDD • What would make it become stable? • Losing or gaining an electron • So what do you think happens when we react Cl with Na?

  10. Ionic Bonding Na Cl

  11. Ionic Bonding Na+ Cl-

  12. Ions • Gains or loses an electron to reach that stable electron configuration • When this happens, # of protons does not equal the # of electrons • Ion: An atom that has a new positive or negative charge • This charge is represented by a + and -

  13. What is the atomic number of chlorine? How many electrons does it have? What is its overall charge? What happens to its charge if it gains an electron? An ion with a negative charge is called an anion. 17 17 Neutral Becomes negative 1 ( Cl 1-) Formation of Ions

  14. What is the atomic number of sodium? How many electrons does it have? What is its overall charge? What happens to its charge if it loses an electron? An ion with a positive charge is called an cation. 11 11 Neutral Becomes positive (Na 1+)

  15. Chemical Bonds • A force that holds atoms or ions together as one • Remember when we said opposites attract • Well that works with cations and anions too • Ionic bond: force that holds cations and anions together

  16. Get Energized • We talked about when electrons move to a higher level they absorb energy • Ionization Energy: The energy required to allow an electron to completely escape the protons attraction • Varies from element to element. Lower the energy the easier it is to remove an electron

  17. Ionization energy trends • Two trends exist on the periodic table when it comes to ionization energy • Increases as you move from left to right • Decreases as you move from top to bottom • Examples

  18. Ionic Compounds • We have talked that two or more elements coming together is a compound • Well, if they contain ionic bonds they are ionic compounds • Chemical Formula: notation that shows what elements are in a compound and the ratio of atoms or ions in the compound

  19. Examples • NaCl- 1 Na and 1 Cl • Look at Mg and Cl • Mg has 2 valence while Cl each has 7 valence • Mg would need two Cl to be stable • We would right it like this : MgCl2 • The number needs to be a subscript and after the element it is referring to • If there is only 1 atom present then no subscript

  20. How are they arranged? • Chemical formula tells the ratio but not the structure of the compound • NaCl: 6 Na surround a Cl and vice versa • Keeps a fixed position in a rigid framework (lattice) • Crystal Lattice: Replacement of ions in a regular repeating arrangement in an ionic compound

  21. Covalent Bonds • Covalent Bond: Chemical bond in which two atoms share a pair of valence electrons • Occur between nonmetals since they have high ionization energy • When they share 1 pair it is a single bond • When they share 2 pair it is a double bond • When they share 3 pair it is a triple bond

  22. Diatomic molecules • Usually most nonmetals exist as a diatomic molecules • Diatomic means two atoms • Makes since because they have high valence electron #’s

  23. Attraction • In general… • Elements on the right have a greater attraction for electrons • Elements at the top of a group have greater attraction for electrons • Fluorine: Most reactive and has biggest attraction • Francium: Least

  24. Molecules of Elements • Molecule: a neutral group of atoms that are joined together by one or more covalent bonds • The attraction between the shared electrons and the protons in each nucleus holds the atoms together in a covalent bond • We can write the formula similar to how we did it earlier

  25. Unequal Sharing… • Polar Covalent Bond: a covalent bond where the electrons are not shared equally • Example: HCl or H2O • When these bonds form, the atom with the greater attraction will be slightly negative while the lesser attraction will be slightly positive

  26. Polar and Nonpolar • Just because elements aren’t sharing equally though doesn’t mean they are all polar. • Nonpolar Covalent Bond: Unequal sharing but ends up with none of the elements being more negative/positive than the other.

  27. Polyatomic Ions • Groups of covalently bonded atoms that act as a single atom when combining with other atoms. • Act like a gang. • Examples: • OH -1 Hydroxide • NO3-1 Nitrate • CO3-2 carbonate • SO4-2 Sulfate

  28. Binary Compound • Compound made from two elements is a binary compound • Naming these are easy • Cation is always 1st followed by anion • Cation is not going to change Ex: Sodium • Anion uses part of the name plus the suffix “ide” • Ex: Chloride, Fluoride, Iodide

  29. Here are examples of common roots: • Cl: chlor- F: fluor- Br: brom- O: ox- I: iod- N: nitr-

  30. Metals with Multiple Ions • Metals that have a group # has an ion with that amount of positive charge • Example: K+, Ca2+, Al3+ • Many transition metals have more than 1 ion • When this happens, Roman numerals are added to indicate the charge on the atom • Example: Copper

  31. Molecular Compounds • The name and formula of a molecular compound describes the type and number of atoms in a molecule

  32. Naming Molecular Compounds • General rule: Most metallic element is 1st • Name in the second element changed to “ide” • The number of atoms is reflected in the name by adding a prefix to it • Prefixes are as follows

  33. Writing Formulas • If you know the name, you know the formula • Cation/Most Metallic is 1st, anion/Nonmetallic is 2nd • Use subscripts to show ratio of ions • Since atoms are neutral, charges must be 0

  34. Example 1: Write the name of the following formula: H2S • Step #1 - Look at first element and name it. Result of this step = hydrogen. • Step #2 - Look at second element. Use root of its full name ( which is sulf-) plus the ending "-ide." Result of this step = sulfide. • These two steps give the full name of H2S.

  35. Example 2: Write the name of the following formula: NaCl • Step #1 - Look at first element and name it. Result of this step = sodium. • Step #2 - Look at second element. Use root of its full name ( which is chlor-) plus the ending "-ide." Result of this step = chloride.

  36. Br1- Br1- Br Br K e- e- Br1- Br1- bromine atom potassium atom e- O2- potassium ion bromide ion Mg2+ K1+ K1+ K1+ K1+ K bromine atom potassium atom bromide ion potassium ion potassium bromide KBr magnesium bromide potassium oxide MgBr2 K2O

  37. Br1- Br1- Br1- PO43- O2- N3- Al3+ Ca2+ Pb4+ Mg2+ S2- K1+ K1+ K1+ OH1- OH1- NH41+ NO31- ? Cu2+

  38. OH1- OH1- Na1+ OH1- N3- N3- N3- N3- Pb4+ Pb4+ Pb4+ (metal) N2- N3- N3- Al3+ (metal) (metal) Ca2+ Pb4+ M2+ Mg2+ ? M1+ M1+ Chemical Bonding Activity (metal) (nonmetal) Pb4+ N3- Pb3N4 lead (IV) nitride or plumbic nitride

  39. Br1- Br1- Br1- N3- N3- N3- N3- Pb4+ Pb4+ Pb4+ O2- N3- Al3+ Mg2+ K1+ K1+ K1+ NH41+ NO31- Key http://www.unit5.org/christjs/4bondingact.doc 4. 5. 1. KBr 2. AlN 6. OH1- Cu2+ K2O OH1- 3. Cu(OH)2 7. Pb3N4 MgBr2 NH4NO3

  40. Ca2+ Ca2+ Ca2+ PO43- PO43- PO43- O2- O2- O2- O2- Al3+ Al3+ Fe2+ NH41+ NH41+ NH41+ Key 8. 9. 10. (NH4)3PO4 11. Ca3(PO4)2 Al2O3 FeO

  41. 13. Fe3+ Fe3+ S2- S2- S2- S2- O2- O2- O2- O2- O2- S2- Pb4+ Pb4+ Cu2+ Pb2+ Cu1+ Cu1+ Key 14. 12. PbS 15. CuO 16. Fe2O3 Cu2O Pb2S4 PbS2 Pb2S3

  42. Writing molecular formula • Write symbols in the order they appear in name • Prefixes indicated the number of atoms • If no prefix assume that it has only 1 atom

  43. Metallic Bonds • Metal atoms achieve stable electron configuration by losing electrons • Metal valence electrons are free to move among the atoms, thus creating cations surrounded by shared electrons • This is called a metallic bond.

  44. Properties of these bonds • The cations form a lattice that is held in place by strong metallic bonds between cations and the valence electrons • Metallic bonds stronger than other metals • Metallic bonds are harder and melt higher

  45. + + + + + + + + + + + + Properties of Metals • The ability of the electrons to move is the reason for some of the properties • Can carry an electric current because of the “flow” of the electrons • Lattice is flexible, which allows it to be shaped

  46. + + + + + + + + + + + + Malleable

  47. + + + + + + + + + + + + Malleable • Electrons allow atoms to slide by.

  48. + - + - - + - + + - + - - + - + Ionic solids are brittle

  49. - + - + - + - + + - + - - + - + Ionic solids are brittle • Strong Repulsion breaks crystal apart.

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