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IB Chemistry HL > 4. Bonding > Flashcards

4. Bonding Flashcards

(46 cards)

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1
Q

Ionic Bonding

A

The electrostatic attraction between oppositely charged ions

  • Cations: positive charge, lose valence electrons (metals)
  • Anions: negative charge, gain valence electrons (non-metals)
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2
Q

Ionic Bond Properties

A
  • non directional bond - strength of bond eaqual in all directions
  • conducts electricity when molten/ in solution
  • high melting/ boiling points
  • hard solids
  • low volatility
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3
Q

Covalent Bonding

A

The electrostatic attraction between a pair of electrons and positively charged nuclei

  • bonding electron pairs
  • electrons not involved in bonding are called non-bonding e- pairs
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4
Q

Electronegativity Difference

A
  • Covalent - high electronegativity locks the electrons in covalent bonds
  • Metalic - Low electronegativity allows electrons to float in seas
  • Ionic - high electronegativity difference means uneven sharing

< 0.4 ——————————————— 1.8<

nonpolar polar ionic

covalent bond covalent bond

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5
Q

Polarity

A
  • Partial positive charge (positive/ negative)
  • The overall charge = Dipole movement
  • The molecular is polar if the electron densities are not symmetrical (eg H2O)
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6
Q

Ionic Bond strength

A
  • electrostatic attraction between positive and negative ions
  • the smaller the ions / the greater the charge on the ions, the stronger the attraction between the ions
    • greater charge density within the structure
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7
Q

Covalent Bond strength

A
  • shared pair of electrons between atoms
  • the shorter the bond, the stronger it is
    • triple bonds are shorter and stronger than double bonds
    • triple > double > single
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8
Q

Metallic Bond strength

A
  • attraction between lattice of positive metal ions and delocalized outer shell electrons
  • the smaller the metal ions, the greater the charge on ions
  • the more delocalised outer shell electrons there are, the stronger the attraction between the ions and electrons
    • greater charge density within the structure
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9
Q

Lewis Structures

A
  • representations of molecules showing all electrons, bonding and nonbinding
  • each bond contains 2 electrons
  • elements pair up according to the octet rule
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10
Q

Exceptions to the Octet rule

A
  • Boron - forms stable compounds with just 3 valence electrons
    • highly reactive
  • Beryllium - forms stable compounds with just 2 valence electrons
  • compounds that break the octet rule are often toxic and dangerous, ready to react so the octet rule is obeyed
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11
Q

VSEPR Theory

A

Valence Shell Electron Pair Repulsion

Electron domain - the number of bonds/ electron pairs there are on the central atom

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12
Q

Shapes w/ non-bonding electron pairs

A

If a molecule / ion has lone pairs on the central atom, the shapes are slightly distorted away from the regular shapes

because of extra repulsion caused by the lone pairs

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13
Q

Shape Angles - Linear

A

180

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14
Q

Shape Angles - Bent

A

105

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15
Q

Shape Angles - Trigonal Planar

A

120

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16
Q

Shape Angles - pyramidal

A

107

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17
Q

Shape Angles - Tetrahedral

A

109.5

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18
Q

Shape Angles - Octahedral

A

90

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19
Q

Resonance structures

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A

When a lewis structure allows for the same arrangement of atoms but a different but equally valid arrangement of electrons, resonance occurs

  • Electrons are delocalized and are spread over 2 bonding orbitals
  • Located in the p orbitals - overlapping p orbitals create pi bonds with delocalized electrons
20
Q

Allotropes

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A

Compounds of the same element that differ in structure

eg Carbon

21
Q

Diamond (structure & properties)

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A
  • Each carbon bonded to 4 other carbons by sp3 hybridization
  • High melting point
    • strong directional covalent bond
  • Extremely hard
    • difficult to break atoms apart/ move them in relation to one another
  • No electrical conductivity
    • electrons are localized in specific bonds
  • Insoluble in polar and non-polar solvents because molecular bonds are stronger than any IMF
22
Q

Graphite (structure)

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A
  • Carbon atoms bonded via sp2 hybridization
  • form sheets of 6 sided rings with p-orbitals perpendicular from plane of ringt
    • presence of p orbitals allows for stronger london forces that hold the sheets together
23
Q

Graphite (properties)

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A
  • Conducts electricity - delocalized electrons
  • Slippery - sheets can slip past each other
  • High melting point - delocalized electrons
  • Insoluble
24
Q

Fullerenes

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A

Buckyballs

  • carbon atoms bond in units of 60 atoms, with interlocking 6 sided and 5 sided rings
  • sp2 hybridization - extra p orbitals form pi bonds resulting in the properties

Properties

  • Super strong
  • conduct electricity and heat with low resistance
  • free radical scavenger
25
Silicon structures
* SiO2 repeating in a tetrahedral structure * Si is larger than C so the Si-Si bond length is greater * the greater the bond length, the lower the bond enthalpy * easier to break * more reactive than diamond
26
Intermolecular Forces - London Forces
* between **2 non-polar molecules** * caused by the motion of electrons around a nucleus * **creation of temporary dipoles / instantaneous induced dipoles** * increases with atomic radii
27
Intermolecular Forces - Dipole Dipole Forces
* between 2 polar molecules * **permanent dipole** * the partially charged ends of the molecules attract and repel each other * LDFs are always acting as well * **usually asymmetrical molecules** * except for bent structured molecules eg water
28
Intermolecular Forces - Hydrogen Bonds
* between an electronegative atom from one molecule and a hydrogen atom from another * **N, O, F** * super strong dipole dipole bond * hydrogen bond 5% stronger than an average covalent bond due to lack of underlying layer of electrons
29
Effect of IMF on Boiling Points
* **phase change is when the IMF are overcome** * liquid turns into a gas when the attractive forces between the particles are completely broken * Covalent macromolecular structure - extremely high melting / boiling points * Metals/ Ionic compounds - high boiling points due to ionic attractions * **The weaker the attractive forces, the more volatile the substance**
30
Metallic Bonding
The electrostatic attraction between a lattice of positive ions and delocalized electrons * strength of bond increases with charge of ion * conducts electricty - delocalized electrons * malleable - close-packed layers of positive ions can slide over each other without breaking more bonds than are made
31
Alloys
Mixtures of metals and one or more other elements Produces increased * strength * durability * hardness * resistance to corrosion * magnetic properties * due to the different elements being of different sizes preventing the easy sliding between 2 atoms
32
Composition of Brass and Steel
Brass Copper and Zinc Steel Iron and Carbon
33
Formal Charge (FC)
**FC = valence e- - 0.5bonding e- - non-binding e-** Most favored Lewis Structure: Structure with the _FC closest to zero_ The structure with the _negative FC on the most electronegative atom_
34
Overlapping Orbitals
* Sigma bonds result from the axial overlap of orbitals * Pi bonds result from the sideways overlap of parallel p orbitals * double bonds fomed by 1 sigma and 1 pi bond * triple bonds formed by 1 sigma and 2 pi bonds
35
Sigma Bonds
* Head to head overlap * Cylindrical symmetry of electron density around the internuclear axis
36
Pi Bonds
* Side-to-side overlap * Electron density above and below the internuclear axis * electron overlap is weaker, yet there is an extra bond --\> energy required to break them is more
37
Importance of Ozone
* At ground level **reacts with chemicals to form smog**, harms respiratory systems and degrades materials * At atmospheric levels is essential for life * **absorbs UV radiation** * exothermic reactions cause a temp. inversion in the stratosphere --\> **prevents convection keeping the layers of atmosphere stable**
38
Ozone Formation Equation
**O2(g) + UV (242nm) --\> O•(g) + O•(g) ** **O•(g) + O2(g) --\> O3(g)** exothermic
39
Ozone Depletion Equation
**O3(g) + UV(330nm) --\> O2(g) + O•(g)** **O3(g) + O•(g) --\> 2O2(g)**
40
Catalytic Ozone Destruction - Nitrogen Oxides
**NO•(g) + O3(g) --\> NO2•(g) + O2(g)** **NO2•(g) + O• --\> NO•(g) + O2(g)**
41
Catalytic Ozone Destruction - Chlorofluorocarbons
**CCl2F2(g) --\> CClF2•(g) + Cl•(g)** Cl•(g) + O3(g) --\> O2(g) + ClO•(g) ClO•(g) + O•(g) --\> O2(g) + Cl•(g)
42
Bond Order
The measurement of the number of e- involved in bonds between two atoms in a molecule (if bond order = 0, there is no bonding) ## Footnote Bond Order = Total bonding pairs / total no. of positions eg Ozone O3 = 3/2
43
Hybridization (definition)
The mixing of different types of orbitals to produce new types of orbitals --\> most hybrid orbitals are combinations of s and p orbitals that then form sigma bonds number of electron domains = amount of hybridization there is 2: sp // 3: sp2 // 4: sp3
44
Two sp orbitals
* Triple bonds * 2 sp orbitals form a sigma bond between the carbons, and two pairs of p orbitals overlap to form the 2 pi bonds
45
Three sp orbitals
* sp2 orbital on carbon overlaps to form a sigma bond with the orbital in oxygen * the unhybridized p orbitals overlap in pi fashion
46
Four sp orbitals
* sp3 orbitals on carbon overlap to form a sigma bond with orbitals in H

IB Chemistry HL (19 decks)

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