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Chapter 6 Sections 6.1 – 6.4.

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Presentation on theme: "Chapter 6 Sections 6.1 – 6.4."— Presentation transcript:

1 Chapter 6 Sections 6.1 – 6.4

2 6.1 – Chemical Bonds Chemical Bond = A link between atoms
Why does it occur? The nucleus of one atom is attracted to the electrons of another.

3 Types of Bonds (an overview) (You will see all of these again later in the chapter!)
Ionic Bond Ion = Atom which has gained or lost electron(s) Metal = -LEFT side of Periodic Table -Weak nucleus / Low Electronegativity -LOSERS of electrons -Become + charged ions

4 Types of Bonds (an overview)
Ionic Bond Nonmetal = -RIGHT side of Periodic Table -Strong nucleus / High Electronegativity -GRABBERS of electrons -Become - charged ions

5 Types of Bonds Ionic Bonds
Atoms gain or lose valence electrons to become a NOBLE GAS CONFIGURATION Examples: Na = Na ION = Cl = Cl ION =

6 Types of Bonds Ionic Bonds
Ionic bond = A chemical bond between a cation (+) and an anion (-). Caused by a TRANSFER of electron(s). Usually a metal + a nonmetal

7 Types of Bonds (an overview) (You will see all of these again later in the chapter!)
2. Covalent Bond = A bond caused by a SHARING of electrons Usually a nonmetal + a nonmetal Nonpolar Covalent = Equal sharing of the electrons. Atoms are close in strength Polar Covalent = Unequal sharing of the electrons. One atom is a little bit stronger than the other World of Chemistry; #8 Chemical Bonds; End at 16:25 – formation of Hydrogen molecule

8 Types of Bonds How do you tell which type of bond it is?
-By ELECTRONEGATIVITY A chart of electronegativity will be provided to you: -The greater the difference in electronegativity – the more ionic the bond. -Electrons spend more time closer to the element with higher electronegativity.

9 Types of Bonds If the ABSOLUTE VALUE of the electronegativity difference is: GREATER THAN 1.7 = IONIC Bond LESS THAN 0.3 = NONPOLAR COVALENT Bond 0.3 – 1.7 = POLAR COVALENT Bond Examples:

10 Types of Bonds (an overview)
3. Metallic bond Usually metals only -The metal gives up valence electrons. -Electrons are free to move about. Atom Electron Sea

11 More Detail on the Bond Types 6.2 – Covalent Bonds
Covalent Bond = A sharing of electrons Molecule = A group of atoms held by covalent bonds (ex – water) Diatomic Molecule = Molecule with only 2 atoms Molecular Compound = Compound made of molecules Molecular Formula = The type and number of atoms in a molecule (ex – H2O)

12 Formation of Covalent Bonds
Sharing electrons in a covalent bond makes the atoms more stable and decreases the energy of the atoms. Energy is released when a bond is FORMED. Overlapping of Orbitals – Example H2: + H H H2

13 The Octet Rule Atoms in a compound obtain the electron configuration of a NOBLE GAS to gain stability

14 Drawing Lewis Structures
-A picture of the covalent bonds in a molecule -Connect valence electrons with LINES -For Academic classes, atoms follow the octet rule unless stated Examples:



17 Single Bond = 1 pair of electrons (2 e-s total) shared between two atoms
Double Bond = 2 pairs of electrons (4 total e-s) shared between two atoms Triple Bond = 3 pairs of electrons (6 total e-s) shared between two atoms *Single bonds are the LONGEST in length; Triple are the SHORTEST *Single bonds have the LOWEST bond energy; Triple have the HIGHEST

18 More Detail on the Bond Types 6.3 – Ionic Bonds
Ionic Bond = Bond formed by the attraction of a cation (lost electrons) to an anion (gained electrons) Crystal Lattice = 3-Dimensional network of ions Formula Unit = Simplest ratio of ions NaCl

19 Dot structures for Ionic Compounds:
-Will reach noble gas configuration -Draw an ARROW to show the transfer of e- -Draw as many of each ion as needed Examples:

20 Comparison of Ionic and Molecular
Molecular Compounds IONIC Compounds Bond Type Covalent Bonds Ionic Structure Individual Molecules Crystal Lattice Strength of Bond Strong Bonds Very Strong Bonds mp/bp Low mp / bp High mp / bp Drawing Connect Dots Arrows with Charges Other Conduct electricity when melted or in water

21 Drawing the Pictures – when you’re not told the TYPE of substance
Do electronegativity difference first!! Examples:

22 6. 4- Metallic Bonding Metals have LOW electronegativity – Will LOSE electrons The steps: -Donate valence electrons to electron sea -Electrons free to move about -All electrons in sea are shared by all atoms

23 6.4 – Metallic Bonding Properties of Metals:
Good conductors of heat – e- sea shakes Good conductors of electricity – e- in sea can move Malleable – atoms can be pushed closer Ductile – atoms can be pushed closer Luster – light bounces off e- sea

24 Chapter 6 Section 6.5

25 Putting Partial Charges on Molecules

26 Properties of MOLECULAR Compounds
VSEPR = Valence Shell Electron Pair Repulsion Theory Valence electrons move as far away from each other as possible Draw Lewis Structure Look at Central Atom Count electron areas (bonds + lone pairs) Use chart info *Academic will be given a WORD BANK and the option to use Model Kits*

27 VSEPR Areas Bonds
Lone pairs Shape Bond Angles Structure Example 2 Linear 180o BeCl2;CO2 3 Trigonal Planar 120o BF3 1 Bent <120o SO2 4 Tetrahedral 109.5o CH4 Trigonal Pyramidal <109.5o NH3 H2O

28 VSEPR Examples:

29 VSEPR Examples:

30 VSEPR Examples:

31 Hybridization Combination of equal energy orbitals to form new orbitals which all have the same shape and energy Carbon: BECOMES C C 1s22s22p2 four sp3 hybrid

32 Types of Molecules FIRST – Draw Lewis Structure & Include Partial Charges! 1. Dipole = Molecule with overall charge 2. NonPolar With Polar Sites = Molecules with area of charge which cancel out 3. Nonpolar = Molecule with no areas of charge

33 Types of Molecules How do you tell the difference?
-Ask yourself these questions… 1. Is there charge on the molecule? Yes No = Nonpolar 2. Can it be sliced? YES = Dipole NO = NPWPS

34 Examples

35 Intermolecular Forces
AKA – EXTERNAL BONDS The attraction BETWEEN Molecules Types of External Bonds: Dipole-Dipole Interactions -Occur due to attraction between partial charges -Occur between two dipoles (Fix notes) – the strongest external bond Hydrogen Bond = External bond that involves a hydrogen atom d - d +

36 Intermolecular Forces
London Force -Occurs between Nonpolar (or Nonpolar With Polar Sites) molecules – CHANGE THIS IN THE NOTES!! -Very weak connection (nonpolar to nonpolar is the weakest) The Steps: (only needed for Honors) Electrons in one molecule shift instantaneously to one side Instantaneous charge results Electrons in another molecule are repelled Very weak attraction results

37 Properties Based on Number / Strength of External Bonds
1. State of Matter s>l>g 2. Evaporation (*volatility) slow>fast 3. Thickness (*viscosity) thick>thin 4. Wetness (*adhesion) To feel wet the substance must bond to your skin (to the Na+Cl-) 5. Dissolving LIKE DISSOLVES LIKE

38 Demonstrations

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