Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

IB Chemistry HL > 4. Bonding > Flashcards

4. Bonding Flashcards

You may prefer our related Brainscape-certified flashcards:
AP® Chemistry flashcards Periodic Table flashcards Chemistry 101 flashcards Organic Chemistry 101 flashcards
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)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

VSEPR Theory

A

Valence Shell Electron Pair Repulsion

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Shape Angles - Linear

A

180

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Shape Angles - Bent

A

105

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Shape Angles - Trigonal Planar

A

120

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Shape Angles - pyramidal

A

107

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Shape Angles - Tetrahedral

A

109.5

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Shape Angles - Octahedral

A

90

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Resonance structures

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

A

Compounds of the same element that differ in structure

eg Carbon

21
Q

Diamond (structure & properties)

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)

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)

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

Fullerenes

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
Q

Silicon structures

A
  • 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
Q

Intermolecular Forces - London Forces

A
  • 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
Q

Intermolecular Forces - Dipole Dipole Forces

A
  • 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
Q

Intermolecular Forces - Hydrogen Bonds

A
  • 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
Q

Effect of IMF on Boiling Points

A
  • 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
Q

Metallic Bonding

A

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
Q

Alloys

A

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
Q

Composition of Brass and Steel

A

Brass

Copper and Zinc

Steel

Iron and Carbon

33
Q

Formal Charge (FC)

A

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
Q

Overlapping Orbitals

A
  • 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
Q

Sigma Bonds

A
  • Head to head overlap
  • Cylindrical symmetry of electron density around the internuclear axis
36
Q

Pi Bonds

A
  • 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
Q

Importance of Ozone

A
  • 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
Q

Ozone Formation Equation

A

O2(g) + UV (242nm) –> O•(g) + O•(g)

O•(g) + O2(g) –> O3(g)

exothermic

39
Q

Ozone Depletion Equation

A

O3(g) + UV(330nm) –> O2(g) + O•(g)

O3(g) + O•(g) –> 2O2(g)

40
Q

Catalytic Ozone Destruction - Nitrogen Oxides

A

NO•(g) + O3(g) –> NO2•(g) + O2(g)

NO2•(g) + O• –> NO•(g) + O2(g)

41
Q

Catalytic Ozone Destruction - Chlorofluorocarbons

A

CCl2F2(g) –> CClF2•(g) + Cl•(g)

Cl•(g) + O3(g) –> O2(g) + ClO•(g)

ClO•(g) + O•(g) –> O2(g) + Cl•(g)

42
Q

Bond Order

A

The measurement of the number of e- involved in bonds between two atoms in a molecule (if bond order = 0, there is no bonding)

Bond Order = Total bonding pairs / total no. of positions

eg Ozone O3 = 3/2

43
Q

Hybridization (definition)

A

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
Q

Two sp orbitals

A
  • 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
Q

Three sp orbitals

A
  • sp2 orbital on carbon overlaps to form a sigma bond with the orbital in oxygen
  • the unhybridized p orbitals overlap in pi fashion
46
Q

Four sp orbitals

A
  • sp3 orbitals on carbon overlap to form a sigma bond with orbitals in H

IB Chemistry HL (19 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions