Date post: | 12-Jan-2016 |
Category: |
Documents |
Upload: | charleen-mills |
View: | 222 times |
Download: | 3 times |
Electrons in outermost shell are called valence electrons and the valence shell is the outermost
occupied shell
Electron Configuration of IonsElectron Configuration of Ions
CATIONS CATIONS are formed when a neutral atom loses one or more electrons:
np electrons are lost first
ns electrons are lost second
nd electrons of previous shell lost third (if there’s any)
Fe [Ar] 3d6 4s2
Fe3+ [Ar] 3d5
Anions: Anions: to form monoatomic anions, add electrons until the next noble gas configuration
has been reached
N [He] 2s2 2p3 (room for 3 more electrons)
N3- [He] 2s2 2p6
O [He] 2s2 2p4 (room for 2 more electrons)
O2- [He] 2s2 2p6
When you gain an electron, you gain a positive charge
When you lose an electron, you lose a negative charge
Periodic Table and Electronic Configuration
Periodic table is divided into s,p,d & f blocks which are named for the last subshell of the element
that’s occupied
Groups 1 & 2 – s block [no. of group]
Groups 13-18 – p block [no. of group minus 10]
Transition metals – d block
Lanthanides & actinides – f block
PeriodsPeriods (horizontal rows) are numbered according to the principle quantum number of the valence
shell
Exceptions: Exceptions:
Helium 1sHelium 1s2 2 put in with the noble gases in group 18put in with the noble gases in group 18
Hydrogen 1sHydrogen 1s11 ( can act like group 1 & 17 ) ( can act like group 1 & 17 )
Periodicity of Atomic Properties
Atomic RadiusAtomic Radius of an element is defined as half the distance between the centres of neighbouring atoms.
•Values generally increase down a group and decrease from left to right across a period
Ionic RadiusIonic Radius of an element is its share of the distance between neighbouring ions in an ionic solution
•Values generally increase down a group and decrease from left to right across a period
• Cations are smaller and anions larger than their parent atoms
Ionisation energyIonisation energy of an element is the energy needed to remove an electron from an atom of an element in
the gas-phase
• 1st & 2nd ionisation energies
• generally higher for elements close to He and lower for elements close to Cs
•Very high value if electron is expelled from a closed shell
Electron Affinity, EElectron Affinity, Eeaea,, of an element is the energy released when an electron is added to a gas-phase atom
• generally positive values - generally energetically favourable
•elements with highest values are those close to O, F, Cl.
Chemical Bonds
Chemical bondChemical bond is a link between atoms
•A bond forms if the resulting arrangement of atoms has a lower energy than the sum of the energies of the separate atoms.
•Chemical bonds result from changes in location of electrons in the outermost shells of atoms (valence shells)
Ionic Bonds
An ionic bond is the electrical attraction between the opposite charges of cations and anions.
Lewis Symbols for atoms and ions - keeps track of valence electrons
Lewis Symbol : Symbol of Element + dot for each valence electron
H 1s1
He 1s2
O 1s2 2s2 2px2 2py
1 2pz1
N 1s2 2s2 2px1 2py
1 2pz1
A single dot: represents an electron alone in an orbital
A pair of dots: represents two paired electrons in the same orbital
To work out the formula of an ionic compound:
1. remove dots from Lewis symbol for metal atom.
2. transfer the dots to the Lewis symbol for the nonmetal atom and complete its valence shell
3. adjust the no.s of atoms to ensure all dots removed from metals are accommodated by nonmetals.
4. add charges of ions.
Ca [Ar] 4s2 Ca Cl [Ne] 3s2 3p5 Cl
Cl + Ca + Cl Ca2+Cl
-Cl
-
Chemical formula for calcium chloride is CaCl2
•Octet of electrons formed in each case
OCTET RULE: atoms of the reactive representative elements tend to undergo those chemical reactions that
most directly give them the electronic configuration of the nearest noble gas
Energy is needed to produce ions and when they pack together there’s a net overall lowering of energy - lattice
enthalpy is a measure of this attraction
Properties of Ionic Solids
• ions stack together in regular crystalline structures (crystalline solids)
• High m.p.s and b.p.s
• Brittle
• Form electrolyte solutions when dissolved in water
Covalent Covalent BondsBonds
• Compounds of nonmetals are not ionic - ionic compounds need ions of both positive and negative
charge.
•Atoms of nonmetals don’t become cations - too many electrons have to be lost to achieve noble gas
configurations
• Nonmetals form covalent bonds to one another by sharing pairs of electrons
• A covalent bond is a pair of electrons between two atoms.
• In covalent bonds, atoms share electrons to reach a noble gas configuration.
The valence of an element refers to the number of covalent bonds an atom of the element forms
H has one valence electron 1s1
H H+ H H H HLine represents shared pair of
electrons
H forms one bond, therefore its valence is 1
F has 7 valence electrons [He] 2s2 2p5
F F F F
F forms one bond, therefore its valence is 1
F2 has lone pairs of electrons - pairs of valence electrons not
involved in bonding. H2 has no lone pairs
Structure of Polyatomic Species
Methane CH4
HH H HC
Carbon is
tetravalent
Single bond: single shared pair of electrons (2)
Double bond: double shared pair of electrons (4)
Triple bond: triple shared pair of electrons (6)
Bond Order: the number of electron pair bonds that link 2 atoms
Single bond has a bond order of 1
C
H
H HH
• choose atom with lowest I.E. for central atom
• H is never central
•arrange atoms symmetrically around central atom
•central atom often written first in chemical formula (CH4)
1. Count total no. of valence electrons on each atom & divide by 2 to get no. of electron pairs
Ammonia NH3 N 1 x 5 = 5
H 3 x 1 = 3
Total = 8
8 2 = 4
4 valence e- pairs
2. Write chemical symbols of atoms to show layout in molecule
H
NH H
3. Place one electron pair between each pair of bonded atoms
( one pair remains)
H
NH H
4. Complete octet of each atom (or duplet for H) by placing
electron pairs as lone pairs around atoms
H
NH H
• If don’t have enough electron pairs to form octets, form multiple bonds
• Check each atom has an octet or duplet
Water H20 H 2 x 1 = 2
O 1 x 6 = 6
Total = 8
8 2 = 4
4 valence e- pairs
O
H H
O
H H
O
H H
Polyatomic Ions
A shared electron pair can originate from one atom
NH3 + H+ NH4+ (ammonium ion)
•If both shared electrons come from one atom- called a coordinate covalent bond
H
NH H + H+
H
NH H
H+
Resonance
Benzene C6H6
Kekulé structure
2 structures have exactly same energy - blend together as a resonance hybrid - electron
density is spread evenly around the ring
Resonance stabilises a molecule by lowering its total energy
Resonance occurs between structures with the same arrangement of atoms, but
different arrangement of electron pairs
C
CC
C
CC
H
H
H
H
H
H
Lewis Acids & Bases
Lewis Acid: electron pair acceptor
Lewis Base : electron pair donor
Acid + Base Complex
O
H H
H+ + O HH
H
Lewis acid Lewis base Complex
+
ElectronegativitElectronegativityy
• most bonds lie somewhere between pure ionic and pure covalent
ElectronegativityElectronegativity is the electron withdrawing power of an atom
A measure of the electron pulling power of an atom on an electron pair in a molecule
Highest at upper right hand corner of periodic table and lowest at bottom left hand corner
The greater the difference in electronegativities of 2 elements, the greater the extent of ionic character
The greater the difference in electronegativities of 2 elements, the greater the extent of ionic character
Molecular StructureMolecular Structure
Lewis structures : 2D diagrams & generally don’t show how molecules are arranged in space
VSEPR Model: Valence Shell Electron-Pair Repulsion Model
• molecules consist of central atom & attached atoms
• attached atoms lie to corners of different shapes - describe the shapes of the molecules
• Bond Angles: angles between the bonds
• regions of high electron concentration (found in bonds and lone pairs) repel one another and take up positions as
far away from each other as possible
PredictingPredicting Shapes of Molecules• write down Lewis structure & decide how electron pairs can be arranged around each “central” atom to minimise
repulsions
CASE A: Central atom with no lone pairs
1. CaCl2 CaCl Cl Ca ClCl
Linear with bond angle of 180o
2. BF3
B
F
FF
Trigonal Planar
120o
3. CH4
Tetrahedral
109.5o
5. SF6
Trigonal bipyramidal
120o & 90o
4. PCl5
Octahedral
90o
C
H
H HH
CASE B: Molecules with multiple bonds
• VESPR theory doesn’t distinguish between single & double bonds
• Electron pairs in a double bond act as a single unit of high electron concentration
CO2Linear
CASE C: Molecules with lone pairs on central atom
VESPR formula; A = central atom
X = atom bonded to central atom
E = lone pair of electrons on central atom
NB: lone pairs on attached atoms not included
CO O
e.g. BF3 = AX3 species (no lone pairs on central atom)
NH3 = AX3E species (1 lone pair on nitrogen)
Electron arrangement:Electron arrangement: 3D arrangement of all regions of high electron concentration (bonds and lone pairs) around
central atom.
AX3E species has 4 regions of high electron density
Strengths of repulsions are in the order:
lone pair-lone pair > lone-pair-bonding pair
> bonding pair - bonding pair
Water H2O
H O H AX2E2
• 4 electron pairs adopt a tetrahedral arrangement
• only 2 positions are occupied by atoms - classified as angular or “bent”
• lone pairs push away from each other
• bonding atoms forced closer together
O
H H
Bond angle is less than that of tetrahedron
( 104.5o compared to 109o)
Ammonia NH3Ammonia NH3
H
NH H AX3E
• 4 pairs adopt a tetrahedral arrangement
• 3 positions occupied by atoms
• H atoms have moved slightly towards each other
H
NH HTrigonal pyramidal
( 107o compared to 109.5o )
VSEPR typeNo. of e - pairs in bonds
No. of lone pairs
Structure
Example
AX2
AX2E2
AX3
AX3E
AX4
AX5
AX6
2
2
3
3
4
5
6
0
2
0
1
0
0
0
linear
nonlinear bent
planar triangular
trigonal pyramidal
tetrahedral
trigonal bipyramida
l
octahedral
BeCl2
CaCl2H2O
H2SBCl3
BF3
NH3
CH4
PCl5
SF6
Polar Polar MoleculesMolecules
Polar Covalent Bond: electron pair shared unequally between 2 atoms resulting in some partial ionic character
• arises from differences in electronegativities of atoms
O H Polar bondO more electronegative than H &
gains greater share in the bonding e- pair
O- H+- = O has a partial negative charge
+ = H has a partial positive charge
Two atoms in a polar bond with partial charges give rise to an electric dipole moment
When bonded atoms have different electronegativities, the bond is polar
Nonpolar molecule has zero dipole moment e.g. H2 & Cl2
O OC
+- - Polar bonds, but overall a non-polar molecule
O
H H
+
+
-
Polar bonds and polar molecule
AX2, AX3, AX4, AX5 & AX6 - non-polar (where X = same element)
Bond Strengths
• strength of bonds is measured by bond enthalpy, HB
• bond breaking requires energy, so all HB values are positive
• HB typically increase as bond order increases and decrease as atomic radius and no. of lone pairs increases
Bond Lengths
• Bond length: distance between the centres of 2 atoms joined by a chemical bond
• for bonds between same atoms, the shorter the bond, the stronger it is
e.g. C C is shorter & stronger than C C
• covalent radii- added to estimate bond lengths in molecules
1s 1s
Molecular Orbitals - another view of the covalent bond
H2
bond
(sigma bond)
Looking along internuclear axis,
e- distribution resembles that of an s orbital
Molecular Orbitals - another view of the covalent bond
• a shared e- pair resides in a molecular orbital formed by the partial overlapping of two atomic orbitals
• space created by the overlapping of 2 atomic orbitals is called a molecular orbital
• M.O. Can hold a maximum of 2 e- with opposite spins (like an A.O.)
• A.O. That overlap are generally those of valence shell electrons
• p-orbitals can also form sigma bonds
N2
N 1s2 2s2 2px1 2py
1 2pz1
Two 2pz orbitals form a “head-on” sigma bond
Two 2px orbitals form a pi bondLooking along internuclear
axis,
e- distribution resembles that of an p orbital
• remaining two 2py orbitals form another bond
• N2 has one sigma bond ( from both 2pz
orbitals) and two pi bonds (from both 2px and 2py orbitals)
A single bond is a bond
A double bond is a & a bond
A triple bond is a & two bonds
Promotion & HybridisationBonding in polyatomic orbitals more complex to explain e.g. CH4
• C 1s2 2s2 2px1 2py
1 (looks like C can only form 2 bonds)
• Know carbon can nearly always form 4 bonds. HOW?HOW?
• NB NB carbon has one empty p orbital (2pcarbon has one empty p orbital (2pzz))
• If we promote an eIf we promote an e-- from the 2s orbital into this empty from the 2s orbital into this empty 2p2pzz orbital, get 4 unpaired electrons. orbital, get 4 unpaired electrons.
2s
2p
1s
2s
2p
1s
Before promotion
(can only form 2 bonds)
After promotion
(can form 4 bonds)
BUT now looks like there are 2 different bond types in CH4
Type 1 H 1s C 2s
H 1s C 2px
H 1s C 2py
H 1s C 2pz
Type 2
•In fact, four sp3 hybrid (mixture) orbitals formed from combination
of s orbital and 3 p orbitals
• These orbitals have equal energies lying between the energies of s and p orbitals
sp3 orbital
Bonding in Methane (CH4)
• One unpaired electron occupies each of carbon’s sp3 hybrid orbitals
• Each of these 4 electrons can pair with an electron in a H 1s orbital
• 4 sigma () bonds are formed
• Structure of methane is tetrahedral