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Chapter 6

Chapter 6 . Chemical Bonds. 6.1 Ionic Bonding. Stable Electron Configurations When the highest occupied energy level of an atom is filled with electrons, the atom is stable and will NOT react. Ex: Noble Gases (8 valence electrons) and Helium (2 v.e .). Electron Dot Diagrams.

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Chapter 6

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  1. Chapter 6 Chemical Bonds

  2. 6.1 Ionic Bonding • Stable Electron Configurations • When the highest occupied energy level of an atom is filled with electrons, the atom is stable and will NOT react. • Ex: Noble Gases (8 valence electrons) and Helium (2 v.e.)

  3. Electron Dot Diagrams • Electron Dot Diagrams (Lewis Structures) is a model of an atom in which each dot represents a valence electron

  4. Ionic Bonds • Elements that do not have complete sets of valence electrons will be reactive. By reacting, they can achieve complete sets. • Some elements achieve stable electron configurations through the transfer of electrons between atoms.

  5. Formation of Ions • When an atom gains or loses an electron, the protons do not equal the electrons. The atom is no longer neutral. • Ion- Atom that has a net positive or negative charge. • Anion- Ion with a negative charge. • To name an anion, you change the suffix to –ide and add ion. • Ex: Chloride ion • Cation- Ion with a positive charge. • To name a cation, you add ion to the end. • Ex: Sodium ion

  6. Formation of Ionic Bonds • Chemical Bond- force that holds atoms or ions together. • Ionic Bond- force that holds anions and cations together. • An ionic bond forms when electrons are transferred from one atom to another • Ionization Energy- the amount of energy needed to remove an electron. • The lower the ionization energy, the easier it is to remove an electron. • It’s easier to remove an electron from a metal than from a nonmetal

  7. Ionic Compounds • A chemical formula is a notation that shows what elements a compound contains and the ratio of the atoms or ions of these elements in the compound. • Ex: MgCl2

  8. Ionic Compound Diagrams to try: • K and Cl Li and S • Sr and Cl Al and N • *Ga and S

  9. Crystal Lattices • Remember that orderly, 3-D structure of solids? • In NaCl, each chloride ion is surrounded by 6 sodium ions and each sodium ion is surrounded by 6 chloride ions. • The opposite charges keep ions in their fixed positions in a rigid framework, or lattice. • Crystals- Solids whose particles are arranged in a lattice structure.

  10. Properties of Ionic Compounds • The properties of an ionic compound can be explained by the strong attractions among ions. • high melting point • poor conductor of electricity when solid, good conductor when melted. • shatter when struck by a hammer

  11. 6.2 Covalent Bonds • Sharing Electrons • H has 1 electron. If it had two electrons, it would be happy. (It would have a full outer energy level.) • Two H atoms can share their electrons  • Covalent Bond- chemical bond in which 2 atoms share a pair of valence electrons. • When two atoms share 1 pair of electrons, the bond is a single bond.

  12. Molecules of Elements • Two hydrogen atoms bonded together form a unit called a molecule • Molecule- neutral group of atoms that are joined together by one or more covalent bonds • The attractions between the shared electrons and the protons in each nucleus hold the atoms together in a covalent bond.

  13. Diatomic Molecules & Multiple Covalent Bonds • Diatomic Molecules- occurs in many nonmetals. It’s two atoms of the same element. • Ex: Halogen Group • HONClBrIF • Multiple Covalent Bonds • Nitrogen has 5 valence electrons. It needs to share 3 pairs of electrons in order to be happy.

  14. Unequal Sharing of Electrons • In general, elements on the right of the periodic table have a greater attraction for electrons. • Fluorine has the strongest attraction.  • Polar Covalent Bonds • Electrons are not shared equally. Ex: H-Cl • Electrons will hang out by Cl, giving it a slight negative charge. The other atom will have a slight positive charge. Shared electrons in a hydrogen chloride molecule spend less time near the hydrogen atom than near the chlorine atom.

  15. Nonpolar Covalent Bonds • Electrons are shared equally. • Polar molecules have stronger attractions to each other.

  16. 6.3 Naming Compounds and Writing Formulas • Naming Ionic Compounds • Binary Ionic Compounds – Made from only 2 elements – 1 metal and 1 nonmetal • Cationfirst followed by anion. • Change the ending of the anion to -ide

  17. 6.3 Naming Compounds and Writing Formulas • Writing Formulas for Ionic Compounds • Place the symbol of the cation first, followed by the symbol for the anion. • Use subscripts to show the ratio of the ions in the compound (all compounds must be neutral) • Use criss-cross method

  18. 6.3 Naming Compounds and Writing Formulas • Naming Molecular Compounds • The more metallic element appears first. • The name of the second element changes its ending to –ide. • Ex: carbon dioxide • Use the Greek prefixes to reflect the number of atoms of each element. • Ex: dinitrogentetraoxide N2O4 • If mono appears in the first element, the prefix is dropped. • Ex: NO2 mononitrogen dioxide becomes nitrogen dioxide.

  19. 6.3 Naming Compounds and Writing Formulas • Writing Molecular Formulas • Write the symbols for the elements in the order the elements appear in the name • Prefixes indicate the number of atoms of each element – they become subscripts • If there is no prefix, there is only 1 atom • Ex: Diphosphorustetrafluoride P2F4

  20. 6.4 The Structure of Metals • Remember, metals like to lose electrons. • What happens if no nonmetals are present to accept the extra electrons? • In a metal, valence electrons are free to move along the atoms. • Metal atoms become cations surrounded by a “sea of electrons” • Metallic Bond- the attraction between a metal cation and the shared electrons that surround it. • The cations in a metal form a lattice that is held in place by the strong metallic bonds between cations and the surrounding valence electrons. • Overall, a metal is neutral because the total number of electrons does not change.

  21. The sea of valence electrons can explain the various properties of metals • 1. Alkali metals are very soft. The bonds between the cations and the electrons are very weak because each metal only contributes 1 valence electron. • 2. An electric current is a flow of charged particles. Metals have a built-in flow of charged particles. • 3. Metals are malleable. When metal is struck with a hammer, the metal ions shift their position and the shape changes In a metal, cations are surrounded by shared valence electrons. If a metal is struck, the ions move to new positions, but the ions are still surrounded by electrons.

  22. Alloys • Gold is a soft metal. It becomes harder when mixed with Ag, Cu, Ni, or Zn • Alloy – A mixture of two or more elements, at least one of which is a metal. • Alloys have the characteristic properties of metals. • EX: • Gold that is 100 percent pure is labeled 24-karat gold. • Gold jewelry that has a 12-karat label is only 50 percent gold. • Jewelry that has an 18-karat label is 75 percent gold.

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