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Chapter 3

Chapter 3. Chemical Bonds. The Noble Gases. Noble Gases: do not participate in chemical reactions – they are inert, or not reactive.

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Chapter 3

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  1. Chapter 3 Chemical Bonds

  2. The Noble Gases • Noble Gases: do not participate in chemical reactions – they are inert, or not reactive. • As an attempt to understand the inert properties of noble gases, chemists looked at the number of valence electrons of noble gases. All noble gases have 8 valence electrons.

  3. The Noble Gases • Chemists predicted that a configuration of 8 valence electrons must be a particularly stable configuration, and this stable electron configuration led to increased stability for the entire atom.

  4. The Octet Rule • Chemists then looked closer at other atoms which did not have 8 valence electrons. • Those atoms would gain or lose electrons in order to have 8 valence electrons, just like the stable noble gases. • When atoms gained or lost electron in this way, they became much more stable and much less reactive.

  5. The Octet Rule • The octet rule is the tendency for an atom to gain or lose electrons in order to achieve a stable configuration of 8 valence electrons, sometimes known as the “Noble gas configuration”. • Atoms gain or lose electrons in order to create a valence electron configuration identical to the nearest noble gas on the periodic table.

  6. Ions • When an atom gains or loses electrons, it is no longer charge-neutral and is called an ion. • The charge of the ion is written in the upper right corner of the symbol. • Positive ions are called cations and are named by using the regular elemental name and adding the word “ion”. • Negative ions are called anion and are named by giving the elemental name an “-ide” ending.

  7. Monatomic Cations • Metal name + ion • Na+ = sodium ion • Ba2+ = barium ion • Cr3+ = chromium III ion (transition metals) • Ag+ = silver ion • Zn2+ = zinc ion • Cd2+ = cadmium ion

  8. Monatomic Anions • Subsititute suffix -ide • Cl- = chloride • S2-= sulfide • Se2- = selenide • N3- = nitride

  9. Common Ions

  10. Examples-Fluorine • Draw electron configuration of fluorine • Which is the nearest noble gas? • Draw electron configuration of neon • Does fluorine gain or lose electrons to match neon? How many? • S, N, Ca, K • Problem 3.18-3.26 page 100

  11. Polyatomic Ions • Ions that contain more than one atom • They have there own “rules” for naming • You must know the polyatomic ions. • Table 3.4 page 73

  12. Ionic Compounds • Positive and negative ions come together in a specific formula to form an ionic compound which is charge-neutral. Every positive charge is balanced by a negative charge. • Chemical formulas are written using the symbols of the elements and subscript numbers to indicate the number of each ion. Charges are not shown in a chemical formula.

  13. In Class Examples • Predict the compound formed by the ions of Li and F. • Predict the compound formed by the ions of Mg and Cl. • Problem 3.3 page 77

  14. Binary Ionic Compounds • Contain only 2 elements and both elements present in the compound are ions • To name a simple ionic compound, you only need to name the ions that make up the compound. You do not need to indicate the number of each type of ion (because the formula of an ionic compound is definite). • Naming: • Name the metal from which the cation was formed • Follow it with the name of the anion that was formed.

  15. Naming Compounds • Monatomic Binary Ionic Compounds • NaCl = sodium chloride (Na+ & Cl-) • MgS = magnesium sulfide (Mg2+ & S2-) • Al2O3 = aluminum oxide (Al3+ & O2-) • Ni3P2 = nickel(II) phosphide (Ni2+ & P3-) • Problem 3.4 page 78

  16. Ionic Compounds • Writing the formula of an ionic compound from its name is a little more complex. Since the name does not communicate the number of each type of ion in the formula, you will have to figure it out yourself. • The strategy involves writing the symbols for the two elements, figuring out the charges on the ions they form, and figuring out how the ions can come together to create a neutral compound. • Problem 3.5 page 78 • The trickiest part about writing the formulas of ionic compounds is trying to figure out how many of each type of atom. We have to figure out the charges and figure out how to balance those charges.

  17. Ionic Compounds with Polyatomic Ions • Name the positive ion first and then the negative ion. • Problem 3.7 page 79 • Worksheet from 121 Packet

  18. Covalent Compounds • Covalent compounds are formed between elements that either do not form ions or form ions with the same charge (for example, two negatively charged ions). • Covalent compounds do not contain ions. Therefore, covalent compounds are not held together by attractive forces.

  19. Covalent Compounds • The atoms in a covalent compound achieve their stable octet of 8 valence electrons by sharing electrons with other atoms (in contrast to gaining or losing electrons, like an ion). The atoms are held together by the sharing of electrons.

  20. Molecules • Covalent compounds are also known as molecules. Because the bonding in a molecule is not determined by definite positive and negative charges, the formula of a molecule is not definite. • Examples: CO and CO2; NO and NO2; SO2 and SO3; CH4 and C2H6 and C3H8 and C4H10

  21. Types of Bonding • Ionic Bonds • Covalent Bonds • Nonpolar Covalent Bonds • Polar Covalent Bonds • Table 3.6 page 80

  22. Nonpolar Covalent Bonds • A covalent bond between atoms that have the exact or very similar electronegativities • Equal sharing of electrons • Example: • CH4 • O2

  23. Polar Covalent Bonds • Electrons are shared unequally in a covalent bond • Will have a dipole • Most covalent bonds are polar covalent bonds • Table 3.5 page 74: Electronegativities

  24. Practice Problems • Problem 3.8 page 81 • The easiest way to differentiate between an ionic bond and a covalent bond is to look for a metal. Why?

  25. Lewis Dot Structures • Show where electrons are being shared between two atoms • Show electrons that are not involved in bonding. • Step-by-step procedure on page 82 • Problem 3.10 page 85

  26. Exceptions to the octet Rule • Any atom in period 3 or lower can “expand” its octet because of d orbitals • Examples: • Phosphoric acid • Phosphorus pentachloride

  27. Naming Covalent Compounds • Covalent compounds are named in a way that is very similar to ionic compounds. You still name the elements in the order that they are written, and you give the last element an “ide” ending. • However, because covalent compounds do not have a definite formula, we must indicate the numbers of each type of atom in the name. To indicate the numbers, we use prefixes.

  28. Prefixes for Naming Covalent Compounds • 1:Mono 6: hexa • 2: Di 7: hepta • 3: Tri 8: octa • 4: tetra 9: nona • 5: penta 10: deca • Problem 3.12 page 88

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