1 / 16

Chapter 6 Chemical Bonds

Chapter 6 Chemical Bonds. Warm-up Question: 1. What are valence electrons? 2. List the number of valence electrons for the following elements: Na ________K ________ Cl __________Al _________ ***Hint: You need your periodic table***. Chemical Bonds….

rane
Télécharger la présentation

Chapter 6 Chemical Bonds

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 Warm-up Question: 1. What are valence electrons? 2. List the number of valence electrons for the following elements: Na ________K ________ Cl __________Al _________ ***Hint: You need your periodic table***

  2. Chemical Bonds… • Chemical Bond – force that holds atoms together • Bonding occurs between the valence electrons • These chemical bonds are constantly being made and broken down in our environment. • Your body breaks the chemical bonds in the food we eat to get energy so our body can do work. • Chemical bonding forms COMPOUNDS • Reading a Chemical compound formula C6H12O6

  3. Chemical Bonds… (cont) • There are two types of chemical bonds. • The difference is what they do with their valence electrons. • 1. Ionic Bonds - One atom steals electrons from another atom. • Ionic is always a metal with a nonmetal • 2. Covalent Bonds – Electrons are shared between the atoms. • Covalent is always two nonmetals

  4. Representing Valence Electrons Lewis Dot Structure – model of an atom in which each dot represents a valence electron * symbol in the center is the nucleus and all the inner electrons in the atom C = S = He =

  5. IONIC COMPOUNDS • Ion -an atom that has a charge. (+ or -) • 2 types of ions 1. anion –negatively charged ion (must gain electrons) ex: Cl-1 O-2 N-3 2. cation - positively charged ion (must lose electrons) ex: Ca+2 K+1 Al+3 ***You can determine charge by where they are located on the periodic table.

  6. Stable Electron Configurations When the highest energy level is filled with electrons, the atom is stable and not likely to react. Octet Rule – the tendency of atoms to gain or lose electrons in order to fill the outer most electron shell.(a full outer shell has 8 valence electrons) What group on the periodic table already has a full outer most shell?

  7. Drawing Lewis Dot Structures of Ions • Cations – they are + charge so they must have ________ electrons. • Examples: • Anions – They are – charge so they must have _______ electrons. • Examples:

  8. Dot Structures for Ionic Compounds • Examples: • BaO • CaCl2 • Mg3 N2 • AlBr3 • *** Now for the homework assignment***

  9. II. Ionic Bonds 1. ion – an atom that has a positive or negative charge a. When an atom gains or loses an electron, the # of protons no longer = the # of electrons b. 2 types of ions 1. anion – negatively charged ion (gains a electron) 2. cation - positively charged ion (loses a electron) → Na+Cl- → Mg+2S-2

  10. 2. Ionic Bond – Form between metals and nonmetal (transfer of electrons) ex. NaCl, MgO, CaBr2 3. Chemical Bond – force that holds atoms together 4. Ionization energy – amount of energy used to remove an electron • a. Lower the ionization energy, the easier it is to remove an electron from an atom • b. Ionization energies increase from left to right in a period • Ex. More energy is required to remove an electron from a nonmetal than a metal. • c. Energy increases from the bottom of a group to the top • Ex. Easier to remove an electron from K than Na

  11. III.Covalent Bonds – form between two non- metals (sharing of electrons) a. Molecule – neutral group of atoms that are joined together by one or more covalent bonds IV. Metallic bonds 1. Metal in a sea of electrons Ex. Cu, Al, Fe

  12. V. Unequal Sharing of Electrons 1. Electronegativity – ability of an atom attract electrons to itself, greater this number the greater its ability to attract electrons a. Non-Polar Covalent Bonding- an equal sharing of electrons ex. F2 b. Polar Covalent Bonding - an unequal sharing of electrons ex. HF attraction between polar molecules are stronger

  13. VI. Naming Compounds and Writing Formulas a. The subscripts on H2 = 2 atoms in 1 molecule of hydrogen b. determining the number of atoms Ex. P2F4 = 2 phosphorus atoms 4 fluorine atoms Ex. (NaSO4)3 = 3 sodium 3 sulfate 12 oxygen

  14. 1. Writing Formulas for ionic compounds a. list the symbols and charges for the cation and anion b. use the criss cross method ex. 2. Naming Ionic Compounds (TYPE 1) A. Cations (metals) and Anion (nonmetal) • 1st cation before the anion (or metal before the nonmetal) • 2nd drop the end of the second element and add “ide” Ex. NaCl = BP = = CaCl2 Sodium Chloride Boron Phosphide

  15. 3. Naming TYPE (2) A. If the metal is a transition metal Roman Numerals in parentheses are used 1 = (I) 2 = (II) 3 = (III) 4 = (IV) 5 = (V) ex. Cu+2 = copper (II) ion Fe+3 = iron (III) ion CuCl2= Iron (III) oxide = Copper (II) Chloride Fe2O3

  16. 3. Naming Covalent Compounds (TYPE 3) A. Name between two nonmetals 1st – most metallic element first (element farthest to the left) 2nd – then add “ide” on the end of the second element 3rd – then add a prefix (mono is never used with the 1st element) Ex. N2O = Ex. diphosphorus tetraflouride = DiNitrogen Monoxide P2F4

More Related