chemical bonding chapters 8-9 (ionic, covalent)
DESCRIPTION
Chemical Bonding Chapters 8-9 (Ionic, Covalent). Chemistry. W h at 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 - PowerPoint PPT PresentationTRANSCRIPT
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CHEMICAL BONDINGCHAPTERS 8-9(IONIC, COVALENT)
Chemistry
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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!!!
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FORMING CHEMICAL BONDS
Bonds may form in three ways:1. ionic bond: electrostatic force that holds oppositely charged particles together -called ionic compounds2. 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
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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
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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.
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-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
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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
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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
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-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
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COVALENT BONDS & THEIR PROPERTIES
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-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
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MORE PROPERTIES OF COVALENT BONDS
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-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
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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
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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
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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
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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
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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
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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
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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
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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)
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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
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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.
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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
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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
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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
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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
♦ - indicates a partial negative + indicates a partial positive
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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
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- “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
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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”
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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
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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-- 32
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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 33
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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
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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
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HOMEWORK
Worksheet on Lewis Structures and Identifying Shapes of Molecules.
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