6. Bonding Flashcards
ionic bonding
- the sum of all electrostatic attractions between oppositely charged ions
- resulting from the transfer of electrons from one bonding atom to another
lattice energy
energy given out when ionic bond is formed
polyatomic ions
- ammonium (NH4+)
- hydroxide (OH-)
- nitrate (NO3-)
- hydrogencarbonate (HCO3-)
- carbonate (CO32-)
- sulfate (SO42-)
- phosphate (PO43-)
characteristics of ionic compounds
- crystalline at room temperature
- have high melting and boiling points
- conduct electrical current in molten or solution state (not in solid)
- polar bonds (dissolve in polar solvents)
metallic bond
- electrostatic attraction between a lattice of positive metal ions and delocalized electrons
(e- in random movement except when connected to a source)
strength of metallic bonds depends on
- number of valance electrons that become delocalized
- charge of the metal ion
- ionic radius of the metal ion
alloys
- homogeneous solution of a metal in another metal
- steel, brass, pewter
characteristics of metallic compunds
- melting point decreases as ionic size increases
- malleability (can be beaten into shape)
- ductility (can be drawn into a wire)
- electrical conductivity
covalent bonds
- chemical bond resulting from the sharing of electrons between two bonding atoms (formed between non-metals, except beryllium)
- electrostatic attraction between a shared electron pair and positively charged nuclei
- obeying the octet rule and sharing of electrons
- electrons in the bond move back and forth so that each atom has a stable outer E level for some time
ionic
polar covalent
non-polar/pure covalent
x>1.8
0>x>1.8
x=0
polar vs non-polar molecules
non-polar molecules have an even distribution of charge in the molecule due to an equal sharing of bonding electrons (no electronegativity difference) - the opposite for polar
coordinate bonding
or dative bonding
- when the electrons in a shared electron pair come from only one atom
VSEPR theory
electrons are organized in electron domains on the furthest possible stable distance from one another
LP-LP>LP-BP>BP-BP
electron domain
one direction in space/region of electron density
lewis structures (list, state the bonding angle, LP, BP and draw each)
1) Linear (1)
2) Trigonal planar
3) Angular/bent (1)
4) Tetrahedral
5) Trigonal pyramidal
6) Angular/bent (2)
7) Trigonal bipyramidal
8) Seesaw
9) T-shaped (1)
10) Linear (2)
11) Octahedral
12) Square pyramidal
13) Square planar
14) T-shaped (2)
15) Linear (3)
covalent bonding leads to two kinds of molecules
1)giant structures (strong covalent bonds forming a network, high MP and BP)
2) simple molecular structures (few covalent bonds, low MP and BP, exist as gases and low boiling liquids)
characteristics of covalent compounds
- have a definite and predictable shape
- very strong but weaker than metallic and ionic
- low MP and BP
molecular substances
- subgroup of covalent bonding
- a substance that has atoms held together by weak covalent bonds
- low MP and BP
- fullerene (allotropic modification of C, conductor)
network solids
- subgroup of covalent bonding
- a solid that has covalently bonded atoms linked in one big network/macromolecule/lattice structure
- graphite (stacked graphene sheets of linked hexagonal rings, conductor), diamond (tetrahedral structure, strong bonds insulator), graphene (single layer of hexagonal C-atom lattice), SiO2 (giant covalent structure, doesn’t conduct)
intermolecular forces
attractive forced between molecules (much weaker than intramolecular)
Van der Waals
also called London forces or dispersion forces
- due to random movement of electrons leading to the formation of (temporary) instantaneous dipole and hence induced dipoles in molecules
- strength depends on molar mass
- effective only over a short range
- depends on the SA of the molecule
dipole-dipole
- due to electrostatic attraction between molecules with permanent dipoles
- significantly stronger than VdW
hydrogen bonding
- in molecules where H is bonded to an atom of high electronegativity value
- the strongest of all intermolecular forces
- important for DNA base pairing, secondary structure of proteins (alpha helix and beta pleated sheets)
ion-dipole forces
- attraction forces between an ion and a polar molecule (dipole)
resonance hybrids (list and draw)
1) carbonate ion
2) nitrate ion
3) sulfur dioxide
4) ozone
5) benzene
6) ethanoate ion
sigma bonds vs pi bonds
sigma bonds - formed when two atomic orbitals overlap in one small area (s-s, s-p, p-p)
pi bonds - formed as a result of sideways overlapping of parallel p orbitals in two areas (p-p)
delocalization of electrons
happens whenever double and single bonds alternate between atoms - resonance bonding (1.5 bonds) - resonance structures (conductors)
ozone layer
- stratospheric ozone is in dynamic equilibrium with oxygen (continuously being formed and decomposed by UV radiation - protecting the Earth’s surface from it)
- formed oxygen atoms are called radicals and they react with O2 molecules to form ozone:
O=O -> 2O*
O2 + O* -> O3
and vice versa:
O3 -> O2 + O*
O3 + O* -> 2O2 - steady-state - rate of O3 production is equal to the rate of its decomposition
how is the steady state altered and by whom
by O3-depleting pollutants - CFCs and oxides of nitrogen
1) CFCs - weak C-Cl bonds broken by UV
CF2CL2 -> CF2Cl* + Cl* (radical initiation)
Cl* + O3 -> ClO* + O2
2) nitrogen oxides
NO + O3 -> NO2 + O2
free radicals
have an unpaired electron(s) - very reactive
wavelength =
(h*c)/E
hybridization
merging of s and p orbitals (they meet on the same energy level): sp3, sp2, sp
formal charge
charge of every atom in the structure
valance electrons (theoretical) - actual number of electrons surrounding it (actual value)
how are all bonds ranked by strength
metallic and bonding > covalent > ion-dipole > H-bonds > dipole-dipole > VdW
electron domain vs molecular shape
electron domain shape - lone electrons are regarded as bonds (not the case in molecular shape)
