1 / 22

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.

colin
Télécharger la présentation

Chapter 6

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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.

More Related