Chemical Bonding Part I Flashcards
(64 cards)
1
Q
Noble Gases
A
- chemically inert/inactive
- extremely unreactive
- stabilities are related to electronic configurations
2
Q
Number of electrons in the outer shell of a noble gas
A
- either 8 (ns^2np^6)
- 2 (1s^2) for helium
- stability gained in octet or duplet state
3
Q
Atoms achieve stability by
A
acquiring the nearest noble gas electronic configuration. Meaning possessing a completely filled valence shell.
4
Q
The acquisition of a noble gas electronic configuration can be achieved by
A
- Ionic (electrovalent) bonding
2. Covalent bonding
5
Q
Ionic (electrovalent) bonding
A
transfer of one or more electrons from one atom to another
6
Q
Covalent bonding
A
sharing of electrons between atoms
7
Q
Chemical bonds
A
attractive forces that hold atoms together in compounds
8
Q
Overall process of bond formation usually results in
A
energy loss from atoms
9
Q
The nucleic of atoms in a chemical bond
A
is unaffected. Usually only involves electrons in the outermost shell (valence electrons)
10
Q
All bonds are based on
A
electrostatic forces of attraction between two entities (positive and negative)
11
Q
Electronegativity
A
the tendency of an atom to attract bonding electrons to itself in a chemical bond
12
Q
Electronegativity increases
A
across a period
13
Q
Electrognegativity decreases
A
on descending down a group
14
Q
Most electronegative elements
A
Fluorine, Oxygen, Nitrogen
Full of Noncense
15
Q
Metallic bond
A
the strong electrostatic forces of attraction between the positive metal ions and mobile sea of delocalised electrons
16
Q
Electron-Sea model (metallic bonding)
A
- The metallic lattice is a regular array of positive ions (metallic atoms minus their valence electrons)
- In a metal, atoms are packed closely together, the outer shell orbital of one of the metal atoms can overlap with several outer shell orbitals of neighboring atoms
- Multiple overlapping of atomic orbitals = the valance electrons from each atom come under the influence of a very large number of atoms = the valence electrons can wander freely throughout the lattice (electrons are delocalised)
- Metallic bonds are non-directional. (Malleability can be mold into any shape because attraction is not in any particular direction)
17
Q
Strength of Metallic Bonds
A
- Number of valence electrons: as it increases the charge of the positive metal ions increases and the number of mobile electrons increases (stronger)
- Size of metal atom/ion: the smaller the closer the metal cations are to the mobile electrons, the stronger the forces of attraction between the electrons and nuclei, the stronger the bonds
18
Q
Metallic bonds: Melting and boiling points
A
- high
- large amount of energy required to overcome the strong electrostatic forces of attraction between the positive metal ions and mobile sea of delocalised electrons in a giant metallic lattice structure
- increase with the strength of metallic bond
19
Q
Metallic bonds: electrical conductivity
A
- high
- good conductors in solid or molten (liquid) states because delocalised electrons can carry charges throughout the metallic lattice
- if a potential difference is applied between the ends of a metal, the delocalised electron cloud will flow towards the positive potential
20
Q
Metallic bonds: Hardness
A
- fairly hard, depends on the strength of metallic bond
- made harder by alloying
- malleable and ductile because the layers in the lattice can slide over each other without breaking the strong metallic bond, the delocalised electrons can still hold the displace ions together
21
Q
Metallic bonds: Thermal conducitivy
A
- high
- heat travels through the metal lattice as a result of the rapid random movement of the mobile electrons
- when heat is supplied to one end, the kinetic energy of the electrons is increased. the increase is transmitted through the system of delocaslied electrons to other parts
22
Q
Ionic bonding
A
- the strong electrostatic attraction between two oppositely charged ions, formed from the transfer of electrons between the original atoms
23
Q
Formation of ionic bonds
A
Generally, metals form positive ions (cations) by loosing electrons and non-metals from negative ions (anions) by gaining electrons. Ions formed have stable noble gas configuration
24
Q
Lattice Energy
A
- A quantitative measure of the stability of any ionic solid
- required to break lattice down
- energy released when 1 mole of ionic compound is formed from its constituent gaseous ions under standard conditions
- the more exothermic the lattice energy (numerically larger), the greater the attraction between the oppositely charged ions, the stronger are the ionic bonds)
25
Strength of ionic bond/ factors that affect the lattice energy
1. Charge of ion: charge of positive ion q+ and charge of negative ion, q-
- the more highly charged the ions, the greater the magnitude of lattice energy
2. Size of ions: cationic radius r+ and anionic radius r-
- the smaller the ions, the smaller the inter-ionic distance (r+ + r-), and the greater the magnitude of lattice energy
26
Giant Ionic Lattice Structure of Ionic Compounds
- regular three dimensional regular packing of positive and negative ions
- positive and negative ions are surrounded and held strongly by a fixed number of oppositely charged ions
- strong electrostatic forces hold the ions together in an orderly manner
- exist as a crystalline solid
- non directional, strong in all directions
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Physical properties of Ionic Compounds
- usually possess giant crystal lattice structure
- oppositely charged ions are assembled so that the bonds extend throughout the structure
- properties came from both the structure and bonding
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Ionic Compounds: Melting and boiling points
- high
- large amount of energy needed to overcome the strong electrostatic forces
- large quantities of heat energy must be supplied so that the ions vibrate vigorously enough to overcome the forces of attraction
- in molten lonic liquid, still strong
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Ionic Compounds: Electrical conducitivity
- good in molten state of aqueous solution because mobile ions that can carry charges between the electrodes
- electrical currents carried by moving charges
- poor in solid state because the ions int he crystals are fixed in the lattice points, only type of movement is vibrational
30
Ionic Compounds: Hardness
- crystals are hard because strong electrostatic attraction must be partially overcome to deform the crystal
- but the crystals are brittle and have good cleavage planes
- under sharp blows (high stress forces) slight displacement can occur along a plane in the ionic solid, once this occurs, similarly charged ions are brought near each other and strong mutual repulsion between neighboring layers of ions occurs (very regular fashion)
31
Ionic Compounds: Solubility
- generally soluble in water
- in water, the strong electrostatic attraction between ions may be overcome by hydration energy released when ions are solvated (hydrated) by the water molecules
- there is attraction of water molecules to ions because of ion-dipole force
- the O end orients towards the cation, whereas an H atom orients towards the anion
- Solvation of ions (where ion-dipole attractions are formed) releases hydration energy to break down the ionic lattice structure
- insoluble in non-polar solvents
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Covalent Bonding
the strong electrostatic forces of attraction between the positive nuclei of the two atoms and the bonding electrons shared between them
33
Formation of Covalent bonds
- results from the sharing of a pair of electrons between two atoms
- each atom in sharing the electrons achieves the nearest noble gas electronic configurations, gains stability
- usually directional, has a preferred orientation in space due to the orientation of the atomic orbitals taking part in bond formation
- defined shapes
- usually occur between two atoms that do not easily lose or gain electrons from one another; usually for two non-metallic elements with similar electronegativity
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Bonding electrons
electrons which are shared between the atoms
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bond pair
two bonding electrons
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non-bonding electrons
valence electrons not involved in bonding
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lone pair
two non-bonding electrons
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Molecules have a central atom with less than 8 valence electrons
Covalent compounds with incomplete octet have a high tendency to reach the stable octet by forming dative covalent bonds
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Molecules have a central atom with more than 8 valence electrons
Elements in P3 onward can utilise their energetically accessible d orbitals in bonding.
P2 elements can only have a max of 8 electrons in the valence shell.
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Molecules have a central atom with unpaired electrons
Example: NO
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Nature of Covalent bonds
1. Normal: two combined atoms contribute an equal number of electrons for sharing
2. Dative: the two combined atoms shared a pair of electrons which is contributed by one of the two
(i) one must possess at least one lone pair (donor)
(ii) other must at least one empty orbital in its outer shell to accept electron pair (acceptor, likely to be electron deficient)
- electron deficient species have a tendency to accept electrons by forming dative covalent bond to form a dimmer or adduct
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Bond length
the distance (in nm) between the nuclei of the two atoms joined by a covalent bond
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bond energy/bond dissociation energy
energy required to break one mole of covalent bond between the two atoms in the gaseous sate (in kJ mol-1)
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Factors affecting the strength of covalent bonds
1. Size of atom
- the smaller the size, the larger the extent of overlapping between atomic orbitals, the stronger the covalent bond
- size of atom is larger down the group
2. Number of bonding electrons
- the more the stronger the forces of electrostatic forces of attraction the stronger the covalent bond (triple > double > single)
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Molecule is reactive when
energy required for bond breaking is very low, or when energy released for bond formation is very high
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simple molecular compounds
- made of covalent compounds
- H2O, CO2, CH4
- solid: simple discrete molecules at the lattice points
- atoms in the molecule held by strong covalent bonds.
- only weak intermoelcular forces between the molecules
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Giant covalnet compounds/ giant molecular (macromolecular) structures
- infinite assembly of atoms
- diamond, graphite, silicon, silica
- high melting points because melting points involves the breaking of covalent bonds between atoms in the giant network of atoms
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Physical properties of simple molecular compounds
- low boiling and melting points because they involve overcoming the weaker intermoelcular forces (van der Waals forces/hydrogen bonds)
- non conductors of electricity in the fused state or solid state because of absence of mobile ions/mobile delocalised electrons
- non-polar are soluble in non-polar solvents (hexane,benzene, methylbenzene)
- polar are soluble in polar solvents (water)
- soft because breaking involves weak intermolecular forces
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Properties and uses of graphite
1. Soft and Slippery: layers are held by van der Waals' forces which are relatively weak and layers can slide over one another
2. High melting point (3730 degrees C): large amount of energy required to break the strong covalent bonds within the layers
3. Good conductor of electricity: delocationsed mobile pi electrons (3 out of 4 valence are used in bonding).
- good in the direction parallel to the planes containing hexagonal rings go carbon
- poor perpendicular to these planes
4. Insoluble in all solvents: the solvent molecules cannot penetrate the graphite lattice due to the presence of strong covalent bonds within layers
5. Graphite is rather unreactive, unless at high temps: due to strong covalnet bonds
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Structure of Diamond
- giant molecular structure
- three dimensional covalent network of interlocking hexagons
- every carbon is covalently bonded to 4 other carbon atoms in a tetrahedral arrangement
- carbon-carbon bond is extremely strong (0.154 nm)
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Properties of diamond
1. Very high melting point (4500 degrees C): large amount of energy required to break covalent bonds
2. Non-conductor of heat and electricity: no mobile electrons to carry current, all valence are in bonding
3. Insoluble in all solvents: solvent molecules cannot penetrate, strong covalent bonds
4. Very hard: mechanical strong and rigid due to C-C bonds and tetrahedral arrangement. hardest natural substance, used as abrasive and to cut concrete and other hard substances
5. high refractive index
52
Effect of electronegativity on types on bonding
- pure covalent bond: bonding electrons equally shared, needs to have same electronegativity
- pure ionic bond: electrons completely transferred to form positive and negative ions: two atoms to have very large difference in electronegativity
- but no substances is purely covalent or ionic
- most bonds have some of the other character due to polarisation of bond
- % of ionic charter of a bond increases with increasing difference in electronegativtiy of the two atoms
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Polarisation of ionic bonds
- the positive charge on the cation attract electrons towards it form the anion -> distortion of the electron charge cloud of anion
- polarisation = distortion of electron cloud on the anion by the cation
- the spherical charge cloud in an anion is fairly easily distorted (polarised) due to the nuclear charge insufficiency to hold it firmly -> degree of overlap of charge clouds -> degree of covalent character
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polarising power
ability of a cation to polarise an anion
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polarisability
tendency for an anion to be polarised
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Factors affecting th extent of polarisation
- positive ions: when the charge is high, the size is small, the greater the polarising power
- charge density proportional to charge/ionic radius
- negative ions: when the charge is high, the isze is large, the negative ions are more easily polarised, higher
57
An ionic bond will have a high degree of covalency when
- cation is small
- anion is big
- both ions are highly charged
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Non-polar covlaent bonds
a covalent bond between two atoms of the same electronegativity, there is no charge separation in these molecules and hence, no dipole moment
(H-H, Cl-Cl)
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Polar covalent bonds/polar bond
- covalent bond between two atoms of different electronegativity
- the more electronegative the atom (with greater attractive power) polarises the bonding electrons of the covalent bond to itself, acquiring a partial-negative charge (delta-) and the less electronegative atom acquiares a partial negative charge (delta+)
- polarisation of covalent bond leads to some ionic character -> additional attractions to strengthen the bond
- the more polar the bond, the greater the ionic character, the stronger the bond
- produces dipole which has a dipole moment
60
Dipole moment
a quantitative measure of the polarity of the bond
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polar molecule
one that the overall dipole moment of the molecule is not zero
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non-polar molecule
one that the overall dipole moment of the molecule is zero
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A covalent bond will be polar if
the difference in electronegativity between the two bonding atoms is approximately between 0.5 to 2.0
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polarity of the bond increases as
the difference in electronegativity increases
