CHEMICAL BONDINGCHAPTERS 8-9(IONIC, COVALENT)

Chemistry

1

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!!!

2

FORMING CHEMICAL BONDS

Bonds may form in three ways:1. ionic bond: electrostatic force that holds oppositely charged particles together -called ionic compunds2. covalent bond: attractive force between

atoms due to the sharing of valence electrons -called molecules3. metallic bond: attraction of a metallic cation

for the delocalized electrons that surround it3

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

4

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.

5

-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 6

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

7

8

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

9

-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

10

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 11

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

12

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

13

TYPES OF COVALENT BONDS

Single Covalent-2 electrons shared between atoms

-represented by a single line C C

-sigma bond (): 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

14

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 (): parallel orbitals overlap -only occurs with multiple bonds

15

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

16

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

17

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. 18

4. Subtract the number of pairs you used in step 3 from the number of bonding pairs determined in step 2.

5. 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

19

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)

20

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

21

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.

22

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

23

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

24

25

VALENCE BOND THEORYvalence 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

BeCl2Orbital notation: 2p =>: 2p

2s 2sp -take one s orbital and one p orbital we create an equal

energy hybrid orbital known as ‘sp’BCl3CCl4

26

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--

27

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 28

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)1. dispersion force (induced dipole)2. dipole-dipole force3. hydrogen bonding

29

TYPES OF INTERMOLECULAR FORCESdispersion force (induced dipole) -occurs between nonpolar molecules -very weakdipole-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

30

TYPES OF INTERMOLECULAR FORCEShydrogen 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

31

HOMEWORK

Worksheet on Lewis Structures and Identifying Shapes of Molecules.

32