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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 compunds 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 a 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 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

27 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--

28 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

29 Intermolecular & Intramolecular Forces
These properties are due as a result of differences in attractive forces -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 Types of Intermolecular Forces (van der Walls forces) dispersion force (induced dipole) dipole-dipole force hydrogen bonding

30 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

31 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

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

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