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Chemistry > 6. Bonding > Flashcards

6. Bonding Flashcards

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1
Q

ionic bonding

A
  • the sum of all electrostatic attractions between oppositely charged ions
  • resulting from the transfer of electrons from one bonding atom to another
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2
Q

lattice energy

A

energy given out when ionic bond is formed

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3
Q

polyatomic ions

A
  • ammonium (NH4+)
  • hydroxide (OH-)
  • nitrate (NO3-)
  • hydrogencarbonate (HCO3-)
  • carbonate (CO32-)
  • sulfate (SO42-)
  • phosphate (PO43-)
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4
Q

characteristics of ionic compounds

A
  • 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)
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5
Q

metallic bond

A
  • electrostatic attraction between a lattice of positive metal ions and delocalized electrons
    (e- in random movement except when connected to a source)
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6
Q

strength of metallic bonds depends on

A
  • number of valance electrons that become delocalized
  • charge of the metal ion
  • ionic radius of the metal ion
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7
Q

alloys

A
  • homogeneous solution of a metal in another metal
  • steel, brass, pewter
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8
Q

characteristics of metallic compunds

A
  • melting point decreases as ionic size increases
  • malleability (can be beaten into shape)
  • ductility (can be drawn into a wire)
  • electrical conductivity
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9
Q

covalent bonds

A
  • 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
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10
Q

ionic
polar covalent
non-polar/pure covalent

A

x>1.8
0>x>1.8
x=0

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11
Q

polar vs non-polar molecules

A

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

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12
Q

coordinate bonding

A

or dative bonding
- when the electrons in a shared electron pair come from only one atom

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13
Q

VSEPR theory

A

electrons are organized in electron domains on the furthest possible stable distance from one another
LP-LP>LP-BP>BP-BP

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14
Q

electron domain

A

one direction in space/region of electron density

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15
Q

lewis structures (list, state the bonding angle, LP, BP and draw each)

A

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)

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16
Q

covalent bonding leads to two kinds of molecules

A

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)

17
Q

characteristics of covalent compounds

A
  • have a definite and predictable shape
  • very strong but weaker than metallic and ionic
  • low MP and BP
18
Q

molecular substances

A
  • 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)
19
Q

network solids

A
  • 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)
20
Q

intermolecular forces

A

attractive forced between molecules (much weaker than intramolecular)

21
Q

Van der Waals

A

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

22
Q

dipole-dipole

A
  • due to electrostatic attraction between molecules with permanent dipoles
  • significantly stronger than VdW
23
Q

hydrogen bonding

A
  • 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)
24
Q

ion-dipole forces

A
  • attraction forces between an ion and a polar molecule (dipole)
25
Q

resonance hybrids (list and draw)

A

1) carbonate ion
2) nitrate ion
3) sulfur dioxide
4) ozone
5) benzene
6) ethanoate ion

26
Q

sigma bonds vs pi bonds

A

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)

27
Q

delocalization of electrons

A

happens whenever double and single bonds alternate between atoms - resonance bonding (1.5 bonds) - resonance structures (conductors)

28
Q

ozone layer

A
  • 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
29
Q

how is the steady state altered and by whom

A

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

30
Q

free radicals

A

have an unpaired electron(s) - very reactive

31
Q

wavelength =

A

(h*c)/E

32
Q

hybridization

A

merging of s and p orbitals (they meet on the same energy level): sp3, sp2, sp

33
Q

formal charge

A

charge of every atom in the structure
valance electrons (theoretical) - actual number of electrons surrounding it (actual value)

34
Q

how are all bonds ranked by strength

A

metallic and bonding > covalent > ion-dipole > H-bonds > dipole-dipole > VdW

35
Q

electron domain vs molecular shape

A

electron domain shape - lone electrons are regarded as bonds (not the case in molecular shape)

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