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Chemical Bonding Chapters 8-9 (Ionic, Covalent)

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Presentation on theme: "Chemical Bonding Chapters 8-9 (Ionic, Covalent)"— Presentation transcript:

1 Chemical Bonding Chapters 8-9 (Ionic, Covalent)

2 What is a chemical bond? chemical bond: force that holds two atoms together -determines the properties of compounds -creates stability in the atom ►nature tends to favor lower energy systems ►bonded atoms are lower energy Bond breaking is endergonic and bond formation is exergonic!!!

3 Forming Chemical Bonds
Bonds may form in three ways: 1. ionic bond: electrostatic force that holds oppositely charged particles together -called ionic compounds 2. covalent bond: attractive force between atoms due to the sharing of valence electrons -called molecules 3. metallic bond: attraction of a metallic cation for the delocalized electrons that surround it

4 Ionic Bonds -forms between metals and nonmetals ◊metals lose electrons, forms a cation ~cation: positive ion from loss of electrons ◊nonmetals gain electrons, forms an anion ~anion: negative ion formed from gain of electrons -most are binary, which means they contain 2 different elements, such as MgO, Al2O3

5 Properties of Ionic Compounds
-alternating positive and negative ions form an ionic crystal -the ratio of positive to negative ions is determined by the number of electrons transferred ◊due to high difference in electronegativity -strong attraction results in a crystal lattice, a 3-D arrangement of atoms.

6 -high melting and boiling points -hard, rigid,brittle solids at room temperature -electrolyte when dissolved in water or in molten state -formulas are in smallest whole number ratio of elements -creates very strong bonds

7 Metallic Bonds -similar to ionic bonds because they often form lattices in the solid state. ◊ outer orbitals overlap ~no sharing/transfer of electrons -electron sea model: all metal atoms in a metallic solid contribute their valence electrons to form a ‘sea’ of electrons around the metal atoms. -valence electrons are free to move from atom to atom (delocalized electrons), forming metallic cations


9 Properties of Metallic Bonds
-formula written as an atom -generally have high melting and boiling points, with especially high boiling points ~due to the amount of energy needed to separate the electrons from the group of cations ~varies due to # valence electrons -malleable & ductile ~mobile electrons can easily be pulled and pushed past each other

10 -durable ~though electrons move freely, they are strongly attracted to the metal cations and are not easily removed from the metal -good conductors ~free movement of the delocalized electrons, allowing heat and electricity to move from one place to another very quickly -luster ~interaction between light and delocalized electrons -forms alloys, a mixture of elements with metallic properties -properties differ from those of the individual elements

11 Covalent Bonds & Their Properties
-form between: -atoms with small difference in electronegativity ~2 or more nonmetal atoms ~metalloids and nonmetals -formulas give true ratio of atoms (molecular formula) -low melting and boiling points. -many vaporize readily at room temperature

12 More Properties of Covalent Bonds
-may exist as liquids, gases or relatively soft solids -some can form weak crystal lattices (sugar) -nonelectrolytes when dissolved in water -weakest of the three types ~low bond strength

13 Strength of Covalent Bonds
What affects bond strength? bond length: distance that separates the bonded nuclei -determined by the size of the atoms and how many electron pairs are shared ♦larger the atom, the longer the bond length, the weaker the bond ♦more shared electrons gives a shorter, stronger bond

14 Types of Covalent Bonds
Single Covalent -2 electrons shared between atoms -represented by a single line C C -sigma bond (s): single covalent bond formed when an electron pair is shared by the direct overlap of orbitals ♦can occur between s & s, s & p , or p & p orbitals

15 Multiple Bonds -two atoms share more than 2 electrons. ~double bond: 4 electrons shared ( 2 pairs) O = O ~triple bond: 6 electrons shared (3 pairs) N N -commonly formed by C, N, O, P, S pi bond (p): parallel orbitals overlap -only occurs with multiple bonds

16 Single vs Multiple Bonds
-the more electrons shared, the stronger the bond ~triple bond, shortest, strongest ~single bond, longest, weakest -due to increase in electron density between the 2 nuclei, which increases the attraction between the nuclei N N O O C C

17 Molecular Structures (Lewis Structures)
structural formula: uses letter symbols and bonds to show relative positions of atoms -can be predicted for many molecules by drawing Lewis structures (covalent only) -H is always an end (terminal) atom, never a central atom -less electronegative atom is the central atom -nature favors symmetry

18 Rules for Drawing Structural Formulas
Once you have the central atom: 1. Find the total number of valence electrons -for negative ions, add electrons -for positive ions, subtract electrons 2. Determine the number of bonding pairs by dividing the total number by 2 3. Place one bonding pair (single bond) between the central atom and each terminal atom.

19 4. Subtract the number of pairs you used in step 3 from the number of bonding pairs determined in step Take the remaining electron pairs and place them around the terminal atoms so each satisfies the octet rule. -place any remaining pairs on the central atom

20 6. If the central atom is not surrounded by 4 electron
pairs, it does not have an octet -convert one or two of the lone pairs on a terminal atom to a double or triple bond between that terminal atom and the central atom (remember which can form multiple bonds) 7. Exceptions: -reduced octet (H & B can have less than 8) -expanded octet (period 3-7 central atoms)

21 Resonance Structures (& an example)
-when one or more valid Lewis structure can be written for a molecule, resonance occurs ~let’s look at NO3-1 -each molecule/ion that undergoes resonance behaves as if it only has one Lewis structure

22 Shape & Hybridization 1. Count areas of electron density around the central atom -multiple bonds count as 1 area 2. Count the number of lone pairs on the central atom 3. Identify the shape & hybridization 4. Identify the polarity: -polar molecules have uneven electron forces, caused by the presence of lone pairs on the central atom or different terminal atoms.

23 Molecular Shape & Hybridization
The shape of molecules determines if two or more molecules can get close enough for a reaction to occur. VSEPR (Valence Shell Electron Pair Repulsion) model: atoms in a molecule are arranged so that the pairs of electrons (bonded and lone) minimize repulsion. -unshared electron pairs have greater repulsive force than shared electron pairs

24 VSEPR model The repulsion between electron pairs result in fixed angles between atoms -bond angle: angle formed by any two terminal atoms and the central atom ♦lone pairs take up slightly more space than bonded pairs (greater repulsive forces) ♦multiple bonds have no affect on the geometry because they exist in the same region as single bonds -example: H2O


26 Electronegativity and Polarity
Remember that atoms have different attractions for electrons (electronegativity). -electronegativity increases left to right and decreases down a period The character and type of bond can be predicted using the difference in electronegativities between bonded atoms. -pure covalent bond: equal sharing of electrons

27 Most atoms do not have equal sharing of electrons, producing a purely covalent bond. -polar covalent bond: unequal sharing of electrons When a polar bond forms the shared electrons are pulled more strongly toward one atom. -this creates partial charges at opposite ends of the molecule, which is called a dipole ♦ d- indicates a partial negative d+ indicates a partial positive

28 Polar molecule or not? A molecule can have individual polar bonds, but make a nonpolar molecule. How? We look at the shape of the molecule and the terminal atoms. Example: H2O vs CCl4

29 - “symmetric” molecules like CCl4 are nonpolar because the polar bonds (electron forces) cancel each other out. CCl4 - “asymmetric” molecules like H2O are polar because the electron forces do not cancel each other out. H2O

30 If water is polar, why will oil not dissolve in it
If water is polar, why will oil not dissolve in it? Oil must be nonpolar because A substance is only soluble (dissolvable) when combined with a like molecule. “Like Dissolves Like” hydrophobic- “fear of water” hydrophilic- “likes water”

31 Valence Bond Theory valence bond theory (VB theory): explains which atomic orbitals must have overlapped in order to obtain a particular geometry where all bonds are created equal. -explains why an atom with a full valence shell can bond BeCl2 Orbital notation: 2p =>: 2p 2s 2sp -take one s orbital and one p orbital we create an equal energy hybrid orbital known as ‘sp’ BCl3 CCl4

32 Self Checks #1 Predict the bond type found in the following: 1. NaCl 2. H2O 3. Ca #2 Predict the number of valence electrons for the following: 1. Li 2. Ba 3. B 4. Si 5. N 6. S 7. Br 8. Ne #3 Draw Lewis structures and identify the shapes for the following: 1. CCl4 2. BF3 3. OH--

33 Intermolecular & Intramolecular Forces
Properties, such a melting points & boiling points, are due as a result of differences in attractive forces -strong forces = strong bonds = higher mp/bp -attraction between atoms within a molecules is strong ~called intramolecular forces -attraction between different molecules is weak ~called intermolecular forces or van der Walls forces ~not bonds

34 Types of Intermolecular Forces
dispersion force (induced dipole) -occurs between nonpolar molecules -very weak dipole-induced dipole force -occurs between a polar molecule and a nonpolar molecule dipole-dipole force -occurs between polar molecules -the more polar the molecule, the stronger the force

35 Types of Intermolecular Forces
hydrogen bonding -strong intermolecular force between the hydrogen end of one dipole and the lone pairs of a fluorine, oxygen or nitrogen atom on another molecule’s dipole -special case of dipole-dipole

36 Homework Worksheet on Lewis Structures and Identifying Shapes of Molecules.

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