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chemical bonding Flashcards

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1
Q

properties of ionic compounds

A

high melting and boiling points, hard but brittle, good conductor of electricity in molten and aqueous states, soluble in polar solvents but insoluble in non-polar solvents

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2
Q

high melting point and boiling point (ionic compound)

A

the melting and boiling points of ionic compounds are high because a large amount of thermal energy is required to overcome and separate the strong electrostatic attractive forces between oppositely-charged ions

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3
Q

hard (ionic compound)

A

most ionic compounds are hard, the surfaces of their crystals are not easily scratched. this is because the ions are bounded strongly to the lattice and are not easily displaced.

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4
Q

brittle (ionic compound)

A

brittleness – ease of shattering or fracturing under stress. under sharp blows (high stress force), slight distortion can occur along a plane in the ionic solid. this happens because distortion causes ions of like charges to come close together and they sharply repel.

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5
Q

electrical conductor in molten and aqueous state (ionic compound)

A

there are free mobile ions present to carry the electrical charges

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6
Q

electrical non-conductor in solid state (ionic compound)

A

strong electrostatic attractive forces prevent the movement of charged ions. since the ions are unable the move, they cannot carry the electrical charges.

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7
Q

solubility in water and polar solvents (ionic compound)

A

the formation of ion-dipole attraction results in the release of energy that causes the detachment of ions from the crystal lattice for salvation. some ionic crystals dissolve readily in water. since water is a polar molecule, it attacks an ionic lattice and pulls it apart. once removed from the lattice, an ion is quickly surrounded by the water molecules. these water molecules are strongly attracted to the ions because of the electrostatic ion-dipole attraction.

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8
Q

solubility in non-polar solvents (ionic compound)

A

non-polar solvents like hexane and benzene form weak forces of attraction with ions which do not pull the ions away from the lattice structure and as such the compound does not dissolve in such solvents.

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9
Q

covalent bond

A

electrostatic attraction between a shared pair of electrons and the nuclei of the atoms being bonded

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10
Q

electron sharing in covalent bonds

A

usually between atoms of non-metals (elements with an electronegativity difference <1.8)
achieve noble gas configuration
probability distribution of the shared electrons is relatively high in the region between the two nuclei

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11
Q

overlapping of atomic orbitals

A

in order to form a covalent bond, two atoms must come close enough for their atomic orbitals with unpaired electrons to overlap. too large an overlap results in a strong repulsion between the bonding nuclei. Most stable situation achieved by partial overlapping of two atomic orbitals.

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12
Q

two ways of electron sharing

A

normal covalent bonds & coordinate (dative) covalent bonds

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13
Q

normal covalent bonds

A

electrons from the shared pair originate from the two atoms involved in bonding

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14
Q

coordinate (dative) covalent bonds

A

both electrons from the shared pair originate from the same atom

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15
Q

bond pairs

A

shared pairs of electrons between two atoms

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16
Q

lone pairs

A

pairs of electrons not shared

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17
Q

octet rule

A

in forming chemical bonds, atoms tend to achieve the stable noble gas electronic configuration with 8 electrons in the valence shell. this can be done by gaining, losing or sharing of electrons.

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18
Q

exceptions of octet rule

A

incomplete octet structure and expanded octet structure

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19
Q

incomplete octet structure (exception)

A

molecules whose central atoms have fewer than 8 atoms after bonding (BF3 & AlCl3, usually group 2 and group 13 elements)
molecules with an odd number of electrons (NO, NO2 etc. N has 5 valence electrons)

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20
Q

expended octet structure (exception)

A

molecules whose atoms have more than 8 electrons after bonding (SF6 & PF5, usually compounds of period 3 elements onwards)

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21
Q

BeCl2 & BF3 (molecules without noble gas configuration)

A

after bonding, Be only has 4 assigned electrons. after bonding, B only has 6 assigned electrons. Be and B are period 2 elements with an n=2 shell which can hold a maximum of 2n^2=8 electrons.
Be and B form covalent compounds due to the high ionisation energy and involved in forming Be2+ and Be3+ respectively. They often have incomplete valence shells (< 8e-) in their compounds.
These electron-deficient compounds are very reactive.

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22
Q

factors affecting ionic bond strength

A

lattice energy, covalent character

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23
Q

covalent character (factor affecting ionic bond strength)

A

introduction of covalent character in ionic bond increases strength of ionic bond

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24
Q

properties of ionic compounds with covalent character

A

ionic compounds with covalent character exhibit lower melting point because ionic compounds with a high degree of covalent character may be soluble in organic solvent

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25
electronegativity
relative attraction atoms has for the shared pair of electrons in a covalent bond
26
pauline's scale most electronegative element + value
Flourine, 4.0
27
most electronegative elements
top right hand corner, F, O, N, Cl --> gain electrons from atoms of other elements and are powerful oxidising agents
28
low electronegative elements
metallic elements, electropositive
29
trends of electronegativities across a period (left to right)
atoms gets smaller, resulting in decreased distance between bonding electrons and nuclei nuclear charge increases but shielding effect remains relatively constant since the inner quantum shells of electrons remain the same effective nuclear charge increases as a result, electrostatic attraction between bonding electrons and nuclei increases as atoms get smaller hence, electronegativity increases
30
chemical bond
electrostatic force which holds two or more atoms or ions together
31
bond breaking
absorbs energy, endothermic
32
bond making
releases energy, exothermic
33
iconic bond
attraction between positive charged ions (cations) and negatively charged ions (anions)
34
covalent bond
attraction between nuclei and shared electrons
35
metallic bond
attraction between positively charged metal ions and delocalised electrons
36
valence electrons
electrons involved in chemical bond formation from the outermost shell of atoms, valence electrons rearrange to attain state of minimum energy (nuclei of atoms are unaffected)
37
octet configuration
possessing a completely filled valence shell with 8 outermost shell electrons (nearest noble gas electronic configuration)
38
exceptions of octet configuration
helium
39
types of chemical bonding
Ionic, covalent (including coordinate covalent bonds), metallic
40
structure of ionic bonding
giant ionic structure
41
structure of covalent bonding
simple molecular, giant covalent (macromolecular)
42
structure of metallic bonding
giant metallic structure
43
types of molecules in simple molecular structure
polar, non polar
44
hydrogen bonding
between polar molecules that contain H directly bonded to N, O, or F
45
examples of hydrogen bonding
HF, NH3, H2O, CH3COOH, CH3CH2OH, CH3CONH2
46
dipole-dipole forces
between polar molecules that do not necessarily contain H or are directly bonded to N, O or F
47
examples of dipole-dipole forces
CO, HCl, CH3Cl, SO2
48
non-polar molecules
instantaneous induced dipole-induced dipole forces aka London dispersion forces
49
examples of London forces
CCl4, CO2, H2, I2, PCl5, SF6
50
ionic/electrovalent bond
electrostatic attraction experienced between electric charges of a cation and an anion
51
cation
atom loses electrons
52
anion
atoms gains electrons
53
electronic configuration of ionic bonded ions
nearest noble gas electronic configuration
54
ionic compounds usually formed
between two elements of very different electronegativities, usually difference of >1.8. highly electronegative non-metal gains electrons to form anion & low electronegative metal loses electrons to form cation
55
nature of ionic bonding
electrostatic in nature usually strong, so ionic compounds generally have high melting and boiling points non-directional (equally strong in any direction)
56
what gives rise to electrostatic attraction
interaction between charged particles
57
lattice energy
enthalpy change when one mole of a solid ionic compound is separated into gaseous ions under standard conditions (indicator of strength of electrostatic attraction)
58
magnitude of lattice energy
directly proportional to the charges on the ions and inversely proportional to the distance separating them
59
when molecules have minimum potential energy
greatest electrostatic attraction between ions & existence of bonding
60
when ions come closer
electrostatic attraction between ions decreases
61
what prevents inter-nuclear distance from being smaller than the radius of the ions
repulsion of electron clouds
62
strength of ionic attraction varies when
decreasing the size and increasing the charge of the ion will increase the strength of ionic attraction, leads to higher melting and boiling points
63
crystalline structure/lattice structure
3-dimensional arrangement of alternating cations and anions. ions in ionic compound attract one another to form giant ionic lattices where positive ions are surrounded by the negative ions and vice versa
64
other ways strength of ionic bond can be affected (in crystal structure)
arrangement of ions in crystal structure, degree of covalent character in ionic bond
65
trends of electronegativities down a group (top to bottom)
atoms get larger, resulting in increased distance between bonding electrons and the nuclei increase in distance results in decrease in electrostatic attraction between the bonding electrons and the nuclei of the atoms hence, electronegativity decreases
66
assumptions of covalent bonding
model of covalent bonding assumes that the pair of electrons in a covalent bond is shared equally between two atoms, the electron density lies in an equidistant region from both nuclei
67
non-polar covalent bond
in a bond between identical atoms, eg H2 or Cl2, the bonding electrons are shared equally thus the electron density is symmetrically distributed between the bonded atoms
68
polar covalent bond
in a polar bond, electrons are shared unequally because of the difference in the electronegative values of the 2 atoms in the molecule. when these two atoms with different electronegativities form a covalent bond, the valence electron density distorts towards the atom with the higher electronegativity this distortion of the valence electron cloud is called polarisation. covalent bond that is polarised is a polar bond. polar bond will have partial electric charges on the two ends of the bonds due to the uneven distribution of electrons
69
dipole moment
the more electronegative atom with a high electron density will have a partial negative charge (δ-) the less electronegative atom with a lower electron density will have a partial positive charge (δ+) separation of charges creates a dipole (two equal and opposite charges separated over a distance) polar bond will therefore possess a small dipole moment (measures polarity of bond -- vector quantity)
70
polar bond and polar molecule
many molecules with polar bonds are polar themselves and. have a permanent dipole. (molecular polarity is a vector). if there is a net dipole, the molecule is polar. BUT some molecules with polar bonds are non polar as their bond polarities cancel each other out so net dipole =0.
71
for a covalent bond to be polar
there must be at least one polar bond the dipole moment of polar bonds must not cancel out so that there will be a net dipole moment
72
relationship between the type of bond and difference in electronegativity between two bonded atoms
when difference in electronegativity of two atoms increases, the polarity of the covalent bond between them increases until a stage when the electron pair essentially reside with the more electronegative atom, giving rise to an ionic bond.
73
large electronegativity difference
pure ionic bond, complete transfer of electrons
74
average electronegativity difference
either ionic bond with covalent character, partial transfer of electrons OR covalent bond with ionic character, unequal sharing of electrons
75
no electronegativity difference
covalent bond, equal sharing of electrons
76
what does difference in electronegativity show
can be used to predict whether bond is ionic of covalent
77
when electronegativity difference = 0
non polar covalent bond
78
when electronegativity difference < 0.4
mostly covalent bond
79
when electronegativity difference 0.4 to 1.8
polar covalent bond
80
when electronegativity difference > 1.8
mostly ionic bond
81
factors affecting covalent bond strength
bond length, bond multiplicity, bond polarity
82
bond length (factors affecting covalent bond strength)
(bond length is measured from the centre of one atom to the centre of another atom) the longer the bond length, the further from the shared pair of electrons the nuclei are. hence, larger bond length translates to a lower bond energy and strength (exception) bond in F2, molecule is relatively weak even though it is shorter than the bond lengths of the heavier dihalogen molecules. This is a result of the relatively large electron and internuclear repulsions and the relatively small overlap of bonding orbitals arising from the small size of the atoms.
83

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