Lecture V.Chemical Bondingjude.edu.sy/assets/uploads/lectures/Lecture V.Chemical Bonding.pdf ·...
Transcript of Lecture V.Chemical Bondingjude.edu.sy/assets/uploads/lectures/Lecture V.Chemical Bonding.pdf ·...
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Chemical Bonding
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What is Chemical Bonding?
Atoms want to share electrons so that their valence shells arefull and they are chemically stable
Chemical Bonding is two or more atoms combining bysharing electrons so that a new substance isproduced that has different physical andchemical properties than its componentelements.
Chemical bonding occurs when atoms share
valence electrons
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Electron Distribution in MoleculesElectron Distribution in Molecules
Gilbert. N. Lewis 1875 - 1946
• Atoms combine to achieve a more stable electronconfiguration.
•Maximum stability results when an atom isisoelectric with a nobel gas.
•Whin atoms interact to form a chemical bond, onlytheir outer regions are in contact( for this reason,when we study chemical bonding, we are concernedprimarily with the valence electrons of atoms).
•Chemists use a system of dots devised by lewisand called lewis dot symbols.
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Lewis Symbols
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• All noble gases except He has an s2p6 configuration.
• Octet rule: atoms tend to gain, lose, or share electronsuntil they are surrounded by 8 valence electrons (4electron pairs).
• Caution: there are many exceptions to the octet rule.
The Octet Rule
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Lewis Symbols of Atoms and Ions• Also known as electron dot symbols• Use symbol of element to represent nucleus and inner
electrons• Use dots around the symbol to represent valence electrons
– put one electron on each side first, then pair• Elements in the same group have the same Lewis symbol
– Because they have the same number of valence electrons• Cations have Lewis symbols without valence electrons• Anions have Lewis symbols with 8 valence electrons
Li• Be• •B• •C• •N• •O: :F: :Ne:• •
•
• • • •
•• •• •• ••
••
Li• Li+1 :F: [:F:]-1•
•• ••
••
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Writing Lewis Structures ofMolecules
• Count the total number of valence electrons from all theatoms
• Attach the atoms together with one pair of electrons– A line is often used as shorthand for a pair of electrons
that attach atoms together• Arrange the remaining electrons in pairs so that all hydrogen
atoms have 2 electrons (1 bond) and other atoms have 8electrons (combination of bonding and nonbonding)
• Occasionally atoms may violate this rule– Nonbonding pairs of electrons are also know as Lone
Pairs
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Structural Formulas- LewisStructural Formulas- LewisStructuresStructures
• Valence electrons are indicated aroundthe symbol for the element
Oxygen has 6 valenceelectrons
Nitrogen has 5valence electrons
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Drawing Lewis StructuresDrawing Lewis Structures• Imagine each side (top, bottom, left,
right) of the symbol of the element canhold 2 electrons for a total of 8electrons.
• Each side will hold one electron first,then will double up.
• In covalent bonding the number ofsingle electron sides (unpairedelectrons) indicates the number ofcovalent bonds the atom must haveto satisfy its octet.
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• Oxygen has 6valence electrons.
• Two unpairedelectrons meansthat oxygen mustform two bonds tosatisfy its octet.
• Draw the Lewisstructure for thefollowing:– Chlorine– Phosphorus– Carbon
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Lewis StructuresLewis Structures
• Atoms shareelectrons to fill theiroctets.
• A solid line indicatesa shared pair ofelectrons.
• Dots are used toindicate unsharedpairs of electrons.
Formation of a single covalent bond
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Double and Triple BondsDouble and Triple Bonds• A unique characteristic of covalent
compounds is their ability to form multiplebonds between two atoms.
• Refer back to the Lewis Structures fornitrogen and oxygen.
• Nitrogen needs to share three electrons• Oxygen needs to share two electrons.
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Technique for DrawingTechnique for DrawingLewis StructuresLewis Structures
1. Determine the number of valence electrons ineach atom making up the molecule
2. Add the valence electrons and divide by two3. Draw the “skeleton.” If carbon is present,
place it at the center of the molecule.4. Distribute the pairs of electrons around the
skeleton to satisfy each atoms octet.(Remember: Hydrogen only needs twoelectrons to fill its octet.)
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Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
F N F
F
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
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Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
O C O
O
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-
4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
2 single bonds (2x2) = 41 double bond = 4
8 lone pairs (8x2) = 16Total = 24
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PracticePractice
• Draw Lewis Structures for the followingcompounds:– Ammonia
– Ethyne- C2H2
– Carbon Dioxide
– HCN
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Exceptions to the Octet RuleExceptions to the Octet Rule
• Atoms with more than an octet– SF4
• Molecules with an odd number ofelectrons– NO
– Generally short lived, unstable molecules
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SO• Lewis symbol- shows the valence e-s,
represened by dots, around the symbolof the element
• Octet rule - atoms tend to lose, gain orshare e-s until they are surrounded by 8valence e-s
– Consists of full s & p sublevels– Achieve the noble gas config closest to
them in the P-table
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Resonance: DelocalizedBonding
Resonance theory states that whenever amolecule or ion can be represented by two ormore plausible Lewis structures that differ onlyin the distribution of electrons, the truestructure is a composite, or hybrid, of them
Resonance structures
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Molecules that Don’t Followthe Octet Rule
Molecules with an odd number of valence electronshave at least one of them unpaired and are calledfree radicals
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Molecules that Don’t Followthe Octet Rule
Some molecules have incomplete octets. These areusually compounds of Be, B, and Al, generally havesome unusual bonding characteristics, and are oftenquite reactive
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Molecules that Don’t Followthe Octet Rule
Some compounds have expanded valence shells,which means that the central atom has more thaneight electrons around it
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Molecules that Don’t Followthe Octet Rule
An expanded valence shell may also need toaccommodate lone-pair electrons as well asbonding pairs
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Odd ElectronOdd electron - if there is an odd number of
valence electrons, it is not possible togive every atom eight electrons• Let’s look at NO, nitric oxide
• It is impossible to pair all electrons as thecompound contains an ODD number ofvalence electrons
N - ON - O
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Lewis Dot Symbols forRepresentative Elements
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What is a Bond?
• A force that holds atoms together.
• Why?
• We will look at it in terms of energy.
• Bond energy the energy required to breaka bond.
• Why are compounds formed?
• Because it gives the system the lowestenergy.
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See if you can define the following words before starting thelesson…
• Anion- negative ion• Cation-positive ion• Octet Rule- rule that states that atoms tend to gain,
lose, or share electrons so that each atom has fulloutermost energy level which is typically 8 electrons.
• Polyatomic Ion- charged group of covalently boundatoms
• Monatomic Ion- ion formed from a single atom• Molecule-neutral group of atoms united by covalent
bonds
• Unshared Pair- pair of electrons that is not involvedin bonding but instead is held exclusively by oneatom.
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Principal Types of Chemical Bonds: Ionic and Covalent
• Ionic bond - a transfer of one or moreelectrons from one atom to another
• Forms attractions due to the oppositecharges of the atoms
• Covalent bond - attractive force due to thesharing of electrons between atoms
• Some bonds have characteristics of bothtypes and not easily identified as one or theother
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Ionic Bonds
• metal to nonmetal, in ionic compounds
• metal loses electrons to form cation
• nonmetal gains electrons to form anion
• ionic bond results from + to - attraction– larger charge = stronger attraction
– smaller ion = stronger attraction
• Lewis Theory allow us to predict thecorrect formulas of ionic compounds
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electron transfer
IONIC BONDING
Lewis Symbol
1s22s22p63s1
Ne coreimplied insymbol
Na
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Lewis Symbol
1s22s22p63s1 1s22s22p63s23p5
Ne coreimplied insymbol
ClNa
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Na+→
1s22s22p6 1s22s22p63s23p6
the loss or gain of electrons(dots) until
IONIC BONDING
the formation of ionic bonds is representedin terms of Lewis symbols
both species have reached an octet of electrons
ClNa Cl
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Cl Cl
[Ne] 3s23p6
represents one orbital
(Pauli: 2 electrons)
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ions stack together in regular crystalline structures
ionic solids typically
1. high melting andboiling points
2. brittle
3. form electrolytesolutions if theydissolve in water
electrostatic interaction
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Li(s) + ½ F2(g) → LiF(s)
lattice energy
(up to few 1000 kJmol-1)
enthalpy of formation
Hess’s Law
Li+(g) + F-(g) → LiF(s)
Born-Haber Cycle
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The Covalent Bond
• Covalent bonds are formed when 2 atoms share 1 ormore pairs of electrons achieving the lowestpossible energy.
• Atoms which form covalent bonds are very similar intheir tendency to gain or lose electron. In general,covalent bonds are formed between nonmetalatoms.
• Atoms may share more than one pair of electrons– Double bond –– Triple bond –– Multiple bonds are
•
• Covalent bonding between identical atoms meanselectrons are
• Covalent bonding between different atoms leads to
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Covalent Bonding
• Shared electron bonds
• Due to overlap of atomic orbitals– (Valence Bond Theory)
• Allows each atom to fill valence shell withelectrons
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Covalent Bonds
• Form between two or more nonmetals, inmolecular or covalent compounds
• Sharing of electrons– Equally shared---nonpolar covalent bond
– Unequally shared----polar covalent bond
• Bond order: single, double and triple
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Single Covalent Bonds
Nonpolar Polar
H
F••
••
•• ••
••F•• ••
H O•• ••••
••F F
F••
••
•• • F•• ••••• H•H• O
••• •
••
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Double Covalent Bonds
• two atoms sharing two pairs of electrons– 4 electrons
• shorter and stronger than single bond
O•••• O••
••••••
O O
O••
• •••
O••
• •
••
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Triple Covalent Bonds
• two atoms sharing 3 pairs of electrons– 6 electrons
• shorter and stronger than single or doublebond
N•••••••••• N
N N
N••
• ••
N••
• •
•
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BondsIonic Bonds
• Composed of ions thathave gained or lostelectrons to achieve afull outer shell
• Electrostatic attractiveforces
• Crystalline solids – nodiscrete molecules -formula units
• Identified by empiricalformulas
• Metal + non-metal
Covalent Bonds
• Composed of atomsthat are sharingelectrons to achieve afull outer shell
• Shared electron bonds
• Discrete molecules,forms gases, liquids,and solids
• Identified by molecularformulas
• Non-metal + non-metal
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electronegativity
difference
between two atoms
involved in the bond
IONIC OR COVALENT
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Covalent Bonding in Hydrogen
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2. ELECTRON SHARING
1. ELECTRON FULLY TRANSFERED
IONIC BONDING
COVALENT BONDING
A + B → AB
2 Na(s) + Cl2(g) → 2 NaCl(s)
2 H2(g) + O2(g) → 2 H2O(l)
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Electronegativity difference
Bond Type
Zero
Intermediate
Large
Covalent
Polar Covalent
Ionic
Co
valent C
haracter
decreases
Ion
ic Ch
aracter increases
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Bonding spectrum100% covalent 100% ionic
A+ B-A B A B
Increasing ΔEN
Increasing polarity Transfer
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• Bond order is the number ofshared electron pairs in a bond.
• A single bond has BO = 1, adouble bond has BO = 2, etc.
• Bond length is the distancebetween the nuclei of two atomsjoined by a covalent bond.
• Bond length depends on theparticular atoms in the bond andon the bond order.
Bond Order and Bond Length
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Electronegativity
• measure of the pull an atom has on bonding electrons
• increases across period (left to right)
• decreases down group (top to bottom)
• larger difference in electronegativity, more polar thebond
– negative end toward more electronegative atom
δ+ H — F δ-
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Bond Polarity
• .
0 0.4 2.0 4.0Electronegativity Difference
covalent ionic
polarnon
polar
3.0-3.0= 0.0
4.0-2.1= 1.9
3.0-0.9= 2.1
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Linus Pauling
1901–1994American chemistOne of the few recipientsof two Nobel prizes. His
paper “The Nature of the Chemical Bondand the Structure of Molecules andCrystals” won him the 1954 Nobel Prize inchemistry. Won the peace prize in 1962.
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ElectronegativityPauling (1901-1994)
�Electronegativity is the relative tendency of anatom in a molecule to attract a shared pair ofelectrons in a bond to itself.
�The most electronegative element is fluorine andit is given a value of 4.0.
�The higher the electronegativity value of anatom, the greater is the ability of an atom of thatelement to attract electrons to itself.
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Electronegativity
Electronegativity (EN) is a measure of the ability ofan atom to attract bonding electrons to itself
The greater theelectronegativity ofan atom in a molecule,the more strongly itattracts the electronsin a covalent bond
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Electronegativity
• Measure of the ability of an atom to attractshared electrons– Larger electronegativity means atom attracts more
strongly– Values 0.7 to 4.0
• Increases across period (left to right) on PeriodicTable
• Decreases down group (top to bottom) onPeriodic Table
• Larger difference in electronegativities meansmore polar bond– negative end toward more electronegative atom
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Electronegativity
• Polarity is determined by difference inelectronegativity– Nonpolar covalent
– Polar covalent
– Ionic compound
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•Electronegativity differences greater thanzero and less than two usually give a polarcovalent bond. (Some books use up to a 0.5difference as still being nonpolar because C-Hbonds are generally considered nonpolar.)
•Electronegativity differences of 2 or greaterare associated with ionic bonds.
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Electronegativity Differenceand Bond Type
Two identical atoms have the same electronegativityand share a bonding electron pair equally. This iscalled a nonpolar covalent bond
Example: chlorine gas
All homonuclear diatomic molecules have nonpolarcovalent bonds:
H2, N2, O2, F2, Cl2, Br2, I2
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Electronegativity Differenceand Bond Type
In covalent bonds between atoms with somewhatlarger electronegativity differences, electron pairs areshared unequally. This is called a polar covalentbond
Example: hydrogen chloride gas, HCl
The electrons are drawn closer to the atom of higherelectronegativity, Cl
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Electronegativity Differenceand Bond Type
With still larger differences in electronegativity,electrons may be completely transferred from metalto nonmetal atoms to form ionic bonds
Example: sodium chloride, NaCl
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Electronegativity Differences
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Figure 11.3: Electronegativity values for selectedelements
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Electronegativity