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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
Q

electronegativity

A

relative attraction atoms has for the shared pair of electrons in a covalent bond

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26
Q

pauline’s scale most electronegative element + value

A

Flourine, 4.0

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27
Q

most electronegative elements

A

top right hand corner, F, O, N, Cl –> gain electrons from atoms of other elements and are powerful oxidising agents

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28
Q

low electronegative elements

A

metallic elements, electropositive

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29
Q

trends of electronegativities across a period (left to right)

A

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

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30
Q

chemical bond

A

electrostatic force which holds two or more atoms or ions together

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31
Q

bond breaking

A

absorbs energy, endothermic

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32
Q

bond making

A

releases energy, exothermic

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33
Q

iconic bond

A

attraction between positive charged ions (cations) and negatively charged ions (anions)

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34
Q

covalent bond

A

attraction between nuclei and shared electrons

35
Q

metallic bond

A

attraction between positively charged metal ions and delocalised electrons

36
Q

valence electrons

A

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
Q

octet configuration

A

possessing a completely filled valence shell with 8 outermost shell electrons (nearest noble gas electronic configuration)

38
Q

exceptions of octet configuration

A

helium

39
Q

types of chemical bonding

A

Ionic, covalent (including coordinate covalent bonds), metallic

40
Q

structure of ionic bonding

A

giant ionic structure

41
Q

structure of covalent bonding

A

simple molecular, giant covalent (macromolecular)

42
Q

structure of metallic bonding

A

giant metallic structure

43
Q

types of molecules in simple molecular structure

A

polar, non polar

44
Q

hydrogen bonding

A

between polar molecules that contain H directly bonded to N, O, or F

45
Q

examples of hydrogen bonding

A

HF, NH3, H2O, CH3COOH, CH3CH2OH, CH3CONH2

46
Q

dipole-dipole forces

A

between polar molecules that do not necessarily contain H or are directly bonded to N, O or F

47
Q

examples of dipole-dipole forces

A

CO, HCl, CH3Cl, SO2

48
Q

non-polar molecules

A

instantaneous induced dipole-induced dipole forces aka London dispersion forces

49
Q

examples of London forces

A

CCl4, CO2, H2, I2, PCl5, SF6

50
Q

ionic/electrovalent bond

A

electrostatic attraction experienced between electric charges of a cation and an anion

51
Q

cation

A

atom loses electrons

52
Q

anion

A

atoms gains electrons

53
Q

electronic configuration of ionic bonded ions

A

nearest noble gas electronic configuration

54
Q

ionic compounds usually formed

A

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
Q

nature of ionic bonding

A

electrostatic in nature
usually strong, so ionic compounds generally have high melting and boiling points
non-directional (equally strong in any direction)

56
Q

what gives rise to electrostatic attraction

A

interaction between charged particles

57
Q

lattice energy

A

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
Q

magnitude of lattice energy

A

directly proportional to the charges on the ions and inversely proportional to the distance separating them

59
Q

when molecules have minimum potential energy

A

greatest electrostatic attraction between ions & existence of bonding

60
Q

when ions come closer

A

electrostatic attraction between ions decreases

61
Q

what prevents inter-nuclear distance from being smaller than the radius of the ions

A

repulsion of electron clouds

62
Q

strength of ionic attraction varies when

A

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
Q

crystalline structure/lattice structure

A

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
Q

other ways strength of ionic bond can be affected (in crystal structure)

A

arrangement of ions in crystal structure, degree of covalent character in ionic bond

65
Q

trends of electronegativities down a group (top to bottom)

A

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
Q

assumptions of covalent bonding

A

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
Q

non-polar covalent bond

A

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
Q

polar covalent bond

A

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
Q

dipole moment

A

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
Q

polar bond and polar molecule

A

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
Q

for a covalent bond to be polar

A

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
Q

relationship between the type of bond and difference in electronegativity between two bonded atoms

A

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
Q

large electronegativity difference

A

pure ionic bond, complete transfer of electrons

74
Q

average electronegativity difference

A

either ionic bond with covalent character, partial transfer of electrons OR
covalent bond with ionic character, unequal sharing of electrons

75
Q

no electronegativity difference

A

covalent bond, equal sharing of electrons

76
Q

what does difference in electronegativity show

A

can be used to predict whether bond is ionic of covalent

77
Q

when electronegativity difference = 0

A

non polar covalent bond

78
Q

when electronegativity difference < 0.4

A

mostly covalent bond

79
Q

when electronegativity difference 0.4 to 1.8

A

polar covalent bond

80
Q

when electronegativity difference > 1.8

A

mostly ionic bond

81
Q

factors affecting covalent bond strength

A

bond length, bond multiplicity, bond polarity

82
Q

bond length (factors affecting covalent bond strength)

A

(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
Q
A

ib chemistry hl (1 decks)

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