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Chemical Bonding Bonds form in 2 main ways atoms share electrons electrons are transferred between atoms Type of bond depends on the atom’s electronegativity.

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Presentation on theme: "Chemical Bonding Bonds form in 2 main ways atoms share electrons electrons are transferred between atoms Type of bond depends on the atom’s electronegativity."— Presentation transcript:


2 Chemical Bonding

3 Bonds form in 2 main ways atoms share electrons electrons are transferred between atoms Type of bond depends on the atom’s electronegativity and electron configuration

4 3 Main Types of Bonds 1.Ionic Bonds electrostatic force atoms transfer e - to become ions usually between a metal and a nonmetal 2.Covalent Bonds electrons are shared by atoms usually between two nonmetals

5 3.Metallic Bonds forces that hold metals together metals have many freely moving electrons that attract positive metal ions Ex. What type of bonding would exist in solid aluminum?

6 Ionic Bonding valence electrons: outermost s and p electrons of an atom Dot Diagrams: show valence electrons Examples: isoelectronic: having the same electron configuration

7 Characteristics of Ionic Compounds high melting point able to conduct electricity in molten state tend to be water soluble crystallize in definite patterns (crystal lattice)

8 A closer look at ionic bonding

9 Naming BINARY Ionic Compounds name the metal first – do not change ending name nonmetal second – change ending to –ide Examples:

10 Writing Formulas for Ionic Compounds Sum of all ion charges MUST equal ZERO! Use the “criss-cross” method Examples:

11 Covalent Bonding molecule: name for a covalently bonded particle

12 Characteristics of Covalent Compounds low melting point do not conduct electricity usually brittle solids, liquids, or gases

13 A closer look at covalent bonding

14 Naming Covalent Molecules make sure the bond is covalent (usually 2 nonmetals) first element’s name does not change second element’s ending becomes –ide Use Greek prefixes to indicate the # of atoms of each element mono = 1 di = 2 tri = 3 tetra = 4 penta = 5 hexa = 6 hepta = 7 octa = 8

15 Writing Formulas for Covalent Molecules Prefixes tell the # of atoms of each element Examples:

16 Molecular Geometry VSEPR Theory

17 Lewis Structures Show arrangement of atoms in molecules Show shared (bonding) and free electrons

18 Drawing Lewis Structures Used for covalently bonded molecules ONLY! 1.Determine the atoms in the molecule 2.Count valence electrons for each atom. 3.Find total # of valence electrons 4.Arrange atoms in skeleton structure. *Least electronegative atom in center!* 5.Add electrons to structure.

19 The number of covalent bonds normally formed by an atom in a Lewis structure depends on its group in the periodic table. H is expected to form one bond. F, Cl, Br, I, all in group 17 are expected to form one bond each. O, S, Se, in group 16, are expected to form two bonds each. N, P, As, in group 15, are expected to form three bonds each. C, Si, Ge, in group 14 are expected to form four bonds each.

20 Octet Rule Atoms try to achieve Noble Gas configuration (8 outer e - ) Hydrogen – forms “duet” instead Some atoms exceed octet – more than 8 bonding e -

21 VSEPR Theory Valence Shell Electron Pair Repulsion Theory electron groups arranged to minimize repulsion

22 Molecular Shapes FILL IN SHAPE CHART! Show relative positions of atomic nuclei MUST determine Lewis structure to determine shape!

23 Resonance Equivalent Lewis structures Shows possible locations of double bonds

24 Resonance

25 Polarity of Bonds polar: having opposite ends polar bond: e - shared unequally caused by difference in electronegativity nonpolar bond: e - shared equally

26 ALL COVALENT BONDS ARE POLAR EXCEPT: 1. C – H 2. any atom bonded to itself

27 A closer look at polar bonds

28 Polarity of Molecules Bonds must be polar for molecule to be polar. Molecule must have a definite top and bottom with opposite charges in order to be polar.

29 Intermolecular Forces (Weak Bonds) Three main types 1. Dispersion forces (London, van der Waals) 2. Dipole-Dipole Interactions 3. Hydrogen bonding

30 Dispersion Forces (van der Waals) Very weak Between nonpolar molecules Induces momentary (temporary) dipole Ex. – occurs in Cl 2, CO 2, CH 4, etc.

31 Dipole-Dipole Interactions Stronger than dispersion Occur between molecules with permanent dipoles (aka – polar) Partially + end of one molecule attracted to partially – end of another

32 Hydrogen bonding Stronger type of dipole-dipole interactions Results from H being covalently bonded to either F, O, or N Stronger because… H is so small F, O, & N are very EN Partial +/- charges are more intense


34 H-bonds > Dipole-dipole > Dispersion Affect BP, MP, solubility More E required to boil/melt substances w/ stronger intermolecular forces Why? Intermolecular Forces

35 Pop Quiz 1.Name CaCO 3 2.Write a formula for sodium sulfite. 3.Draw a dot diagram for boron (B). 4.How many valence electrons does carbon have? 5.What is the oxidation number of potassium?

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