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IB Chemistry HL > Chapter 4/14 > Flashcards

Chapter 4/14 Flashcards

Chemical bonding and structure

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1
Q

Ionic bonding

A

Involves transfer of one or more electrons from outer shell of one atom to outer shell of another atom
- transfer of these electrons results in formation of positive and negative ions

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2
Q

Cations

A

formed when metals lose valence electrons

  • hence, have more protons than electrons in the atom overall
  • this gives them a positive charge overall, forming a positive ion
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3
Q

Anions

A

formed when non-metals gain electrons

  • hence, have less protons than electrons in the atom overall
  • this gives them a negative charge overall, forming a negative ion
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4
Q

Ionic bond

A

Electrostatic attraction between oppositely charged ions

  • formed by elements w/ an electronegativity difference of approx. 1.8 units and greater
  • ionic bonding is non-directional, force of attraction occurs in all directions around the individual atoms
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5
Q

Ionic compounds

A

Composed of a metal and non-metal element (electrically neutral)

  • formation of positive and negative ions results in atoms achieving full outer shells of electrons
  • oppositely charged ions are attracted to each other, resulting in formation of an ionic bond
  • in accordance with the octet rule

Exception to metal and non-metal rule:
Ammonium chloride- has both types of bonding
- ionic bonding between the ions
- covalent bonding in the ammonium ion

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6
Q

Octet rule

A

States that atoms are more stable with the electron configuration of a noble gas

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7
Q

Structure of ionic compounds

A

Under normal conditions, ionic compounds are usually solids w/ lattice structures
- ions that make up an ionic compound are arranged in a regular crystalline structure, lattice structure

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8
Q

Polyatomic ions

A

Ions that consist of more than one atom

  • ions that make up the polyatomic ions are bonded by covalent bonds
  • but bonding between a polyatomic ion and another ion is ionic
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9
Q

Formula of an ionic compound

A
  • ratio of the ions that make up the compound, so , overall, the charges cancel out
  • referred to as the ‘formula unit’, combination of ions is repeated throughout the lattice structure
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10
Q

Physical properties of ionic compounds

A
  • MP and BP
  • solubility
  • electrical conductivity
  • brittleness
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11
Q

Melting point and boiling point

A
  • an ionic compound is held together by electrostatic attractions between oppositely charged ions
  • strong attraction between ions results in ionic compounds, relatively high MP and BP
  • hence, relatively large amounts of energy are needed to break these forces of attraction between the ions
  • hence, ionic compounds are solids under standard conditions
  • MP of an ionic compound depends on ionic charge and ionic radius of its component ions
  • greater the charge on ion and smaller its ionic radius, greater attraction between oppositely charged ions and higher the MP
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12
Q

Volatility

A

How easily a substance evaporates

- ionic compounds have very low volatility because of strong forces of attraction between ions in lattice structure

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13
Q

Solubility

A

Not all ionic compounds dissolve in water

  • depends on whether forces of attraction between water molecules and ions in lattice are strong enough to overcome attractions between ions
  • water molecules are polar due to difference in electronegativity between oxygen and hydrogen
  • at surface of ionic lattice, where it’s in contact with water, positive and negative dipoles of water molecule are attracted to oppositely charged ions in lattice structure
  • these ions break off from lattice and are surrounded by water molecules (hydration)
  • when this happens, the solid dissolves
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14
Q

Non-polar solvents

A

Non-polar solvents can’t disrupt lattice structure, hence solubility of ionic substances eg. hexane and propanone is limited

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15
Q

Solubility and solvents

A

Polar substances are soluble in polar solvents

Non-polar substances are soluble in non-polar solvents

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16
Q

Electrical conductivity

A

Depends on presence of mobile ions

  • ionic compounds don’t conduct electricity when solid because ions are held in fixed positions in lattice structure
  • when an ionic compound is molten or dissolved, ions are free to move and carry an electric current
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17
Q

Brittleness

A

Ionic compounds tend to shatter when a force is applied- are said to be brittle

  • fracture across a plane when layers of ions become incorrectly aligned
  • occurs because movement of ions within lattice when force is applied places ions of same charge next to each other
  • forces of repulsion between ions of same charge cause lattice structure to split and fracture
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18
Q

General properties of ionic compounds

A
  • generally highly soluble in water
  • good electrical conductivity only when molten or dissolved
  • solids at room temperature
  • high MP and BP
  • brittle
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19
Q

Why are ionic compounds soluble in water?

A

Water molecules are attracted to oppositely charged ions causing them to dissolve

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20
Q

Why are ionic compounds good electrical conductors?

A

When molten/in solution, ions are free to move about

- when solid, ions are held in fixed positions in lattice structure

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21
Q

Why are ionic compounds solids at room temperature?

A

Ions are held in fixed positions by strong electrostatic attractions in lattice structure

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22
Q

Why do ionic compounds have high MP and BP?

A

Ions are attracted to each other by strong electrostatic attractions
- large amounts of energy are needed to separate them

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23
Q

Why are ionic compounds brittle?

A

Repulsion between ions of same charge causes lattice structure to split and fracture

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24
Q

Covalent bonding

A

Occurs between non-metal elements and results in formation of molecules and giant covalent structures
- occurs between elements with a difference in electronegativity of fewer than 1.8 units

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25
Q

Hydrogen gas, H2

A
  • exists as a diatomic molecule w/ a single covalent bond between two hydrogen atoms
  • atoms are held together by electrostatic attraction that exists between two nuclei and shared pair of electrons
  • a hydrogen atom needs two electrons in its valence shell to achieve noble gas structure
  • two hydrogen atoms in a hydrogen molecule share pair of bonding electrons, so each atom now has two electrons in its valence shell
  • electrons are shared
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26
Q

Covalent bond

A

Electrostatic attraction between a shared pair of electrons and the positively charged nuclei

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27
Q

Halogens as diatomic molecules

A
  • all form diatomic molecules
  • atoms are bonded via single covalent bond
  • atomic radius increases down a group, so length of bond between atoms also increases
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28
Q

Different covalent bonds

A

Single: 2 shared bonding electrons
Double: 4 shared bonding electrons
Triple: 6 shared bonding electrons

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29
Q

Bond length

A

Distance between the nuclei of the bonded atoms

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30
Q

Bond strength

A

Amount of energy needed to break the bond between the atoms

  • described in terms of bond enthalpy
  • can affect reactivity of a compound
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31
Q

Long bonds vs. short bonds

A
  • longer bonds are weaker than shorter bonds due to increased distance between nuclei and shared pairs of bonding electrons
  • results in a weaker electrostatic attraction between atoms
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32
Q

Multiple bonds

A
  • involve more shared electrons between atoms
  • electrostatic attraction between bonded nuclei is greater
  • this greater attraction brings nuclei closer together, so multiple bonds are shorter and stronger than single bonds
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33
Q

Bond order

A

Number of bonds between a pair of atoms

  • single bonds = bond order of 1
  • double bonds = bond order of 2
  • triple bonds = bond order of 3

The higher the bond order, the greater the strength of the bond

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34
Q

Calculating bond order

A

sum of individual bond orders/ number of bonding groups

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35
Q

Electronegativity and bonding

A

Measure of attraction of an atom for a shared pair of bonding electrons

  • difference in electronegativity determines type of bonding that takes place between atoms
  • large difference in electronegativity = ionic bond
  • small difference in electronegativity = covalent bond
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36
Q

Electronegativity values and bonding

A

Ionic = greater than or equal to 1.8
Polar covalent = 0.5 - 1.7
Non-polar covalent = 0.0 - 0.4

The greater the electronegativity difference between the atoms, greater the polarity of the bond, greater its ionic character

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37
Q

Non-polar covalent bonds

A

Occur between atoms w/ a small difference in electronegativity (0.0-0.4 units)

  • in these molecules, bonding electrons are equally shared
  • present in bonds in diatomic molecules composed of the same atom e.g. oxygen, chlorine gas, nitrogen gas
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38
Q

Polar covalent bonds

A

Occurs between atoms w/ a electronegativity difference between 0.5-1.7 units

  • in these molecules, bonding electrons aren’t shared equally between atoms
  • bond w/ greatest polarity = H-F bond (biggest difference in electronegativity between the two atoms)
  • H-F bond is still a covalent bond, despite large difference in electronegativity
  • bond w/ least polarity = H-I bond (smallest difference in electronegativity between the two atoms)
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39
Q

Polarity of hydrogen halides

A

Polarity of H-X bond decreases as you descend group 7

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40
Q

Bond polarity symbols

A

Bond polarity= results from the difference in electronegativities of the bonded atoms

Bond polarity is designated by the SIGMA+ and SIGMA- signs placed on the molecules

  • refers to partial charges called dipoles
  • SIGMA- is assigned to the more electronegative element in the bond
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41
Q

Factors to consider for polarity of a molecule

A
  1. Presence of polar bonds within the molecule

2. Molecular geometry of the molecule

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42
Q

Non-polar molecules

A
  • Molecules is symmetrical, despite containing polar bonds
  • because these molecules are symmetrical, so, bond polarities cancel each other out
  • molecule has no net dipole movement
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43
Q

Polar molecules

A
  • have a net dipole movement
  • presence of polar bonds within molecules, together w/ molecular geometry, means bond polarities don’t cancel out

NB/presence of a net dipole movement determines type of intermolecular force that exists between molecules

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44
Q

Octet rule

A

States that the most stable arrangement for an atom is to have 8 electrons in its outermost energy level w/ electron configuration of a noble gas
- it’s the tendency of atoms to gain a valence shell w/ a total of 8 electrons

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45
Q

Exceptions to the octet rule

A
  • Hydrogen is stable w/ only 2 electrons in its outer shell
  • Atoms eg. boron, beryllium and aluminium (in compounds) are stable w/ fewer than 8 electrons in their outer shell
  • atoms in period 3 and higher, e.g. sulphur, can form expanded octets w/ up to 12 electrons in their valence shell
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46
Q

Incomplete octets

A

Atoms with less than 8 electrons in their outer shells- said to have incomplete octets or be electron-deficient

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47
Q

Electron-deficient molecules

A

Do not have a full outer shell of electrons

  • atoms in molecules that are electron-deficient are able to form coordinate covalent bonds
  • eg. Aluminium chloride
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48
Q

Example of an incomplete octet

A

Boron trifluoride (BF3)

  • a boron atom only has 3 electrons in its outer energy level- no possibility of it reaching a noble gas structure w/ 8 electrons simply by sharing electrons
  • hence, in BF3, boron atom has formed max. no. of bonds that it can under the circumstances- and is stable
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49
Q

Forming a molecule

A

An atom will tend to make as many covalent bonds as possible

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50
Q

Example of an electron-deficient molecule

A

Atoms in molecules that are electron-deficient are able to form coordinate covalent bonds

  • example of a molecule that has coordinate covalent bonds is AlCl3 (aluminium chloride)
  • at high temp. = aluminium chloride exists as molecules w/ formula AlCl3 (molecule is electron-deficient, still needs 2 electrons to complete octet in outer shell of aluminium atom)
  • at lower temp. = 2 molecules of AlCl3 combine to form a dimer w/ formula Al2Cl6
  • AlCl3 molecules are able to combine because lone pairs of electrons on 2 of the chlorine atoms form coordinate covalent bonds w/ aluminium atoms

NB/ coordinate covalent bond is represented by use of an arrow
- dots and crosses represent valence electrons of aluminium and chlorine

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51
Q

Expanded octets

A

Elements in period 3 and beyond can have expanded octets- accommodate more than 8 electrons in their valence shell
- molecules w/ central atoms from elements in period 3 can accommodate up to 18 electrons in valence shell

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52
Q

Lewis structures

A

Represent bonding in a molecule, they show both bonding electrons and non-bonding electrons
- also called electron dot diagrams

NB/ remember to draw the lone pairs of electrons so each atom has an octet of electrons
- remember to include square brackets and charge when drawing structure for an ion

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53
Q

Coordinate covalent bond

A

Covalent bond = formed between atoms that both contribute electrons to the bond

A coordinate covalent bond = differs from a ‘regular’ covalent bond
- both bonding electrons come from one atom

NB/ once formed, a coordinate covalent bond is identical to a ‘regular’ covalent bond

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54
Q

Resonance structures

A

is one of two or more alternative Lewis structures for a molecule or ion that can’t be described fully w/ one Lewis structure alone

  • no. of possible resonance structures is equal to no. of different positions for multiple bond
  • occurs when molecules contain multiple bonds (double or triple), there is more than one possible Lewis structure that can be drawn
  • but none of the resonance structures represents the actual structure of the ion
  • the actual structure = resonance hybrid structure
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55
Q

Resonance hybrid structure

A

Intermediate between the forms of resonance structures

- all bonds are identical, and are intermediate in strength and length between a single and double bond

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56
Q

Why do covalent molecules have defined shapes?

A

Electrons in a covalent bond are located in a specific position within a molecule
- hence, covalent molecules have a defined shape, determined by the orientation of these ‘fixed’ or ‘localised’ bonds in space

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57
Q

Delocalised electrons

A

involves e- that are shared by/between all atoms in a molecule or ion as opposed to being localised between a pair of atoms

  • give greater stability to a molecule or polyatomic ion
  • exist in any molecule where there is more than one possible Lewis structure- hence, more than one position for a multiple bond
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58
Q

VSEPR theory

A

Valence shell electron pair repulsion theory

  • theory that electron pairs in molecules repel each other and orientate themselves as far away from each other as possible
  • molecule adopts shape that minimises repulsion between bonding and non-bonding electrons (lone pairs)
  • enables us to predict shapes of molecules
  • theory proposes that molecular shape adopted is that where strengths of repulsion between electron pairs in valence shell of central atom are minimised
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59
Q

Bonding domains

A

Bonding electrons in single, double or triple bonds are known as bonding domains

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60
Q

Non-bonding domains

A

Non-bonding electrons are called non-bonding domains

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61
Q

Electron domains

A

Both bonding domains and non-bonding domains are collectively known as electron domains
- single bonds, double bonds and triple bonds each count as one electron domain, as do non-bonding pairs of electrons

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62
Q

Molecular geometry

A

Shape of a molecule

  • deduced by looking at Lewis structure for molecule
  • count no. of electron domains around central atom
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63
Q

Electron domain geometry

A

Total no. of electron domains (both bonding and non-bonding) around the central atom
- for some atoms, molecular geometry is different to electron domain geometry

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64
Q

Electron domain geometry vs. molecular geometry

A

Electron domain geometry:
- total no. of electron domains around the central atom

Molecular geometry:

  • shape of a molecule
  • no. of electron domains around the central atom
  • ALSO takes into account the extra repulsion between bonding and non-bonding domains
65
Q

Linear geometry

A
  • molecules w/ two bonding domains around the central atom- linear molecular geometry
  • bond angle in a linear molecule = 180 degrees
  • electron domain geometry is also linear
66
Q

Trigonal planar geometry

A
  • molecules/ polyatomic ions w/ 3 bonding domains
  • bonding domains are at 120 degrees to each other in the same plane
  • 3 bonding domains around the central atom
  • electron domain geometry is the same as molecular geometry

2 bonding domains + 1 non-bonding = shape is bent or V-shaped

67
Q

Tetrahedral geometry

A
  • molecules w/ 4 bonding domains around central atom
  • bond angle = 109.5 degrees
  • electron domain geometry is same as molecular geometry- because, all 4 electron domains are bonding domains
68
Q

Non-bonding electrons

A

Cause slightly more repulsion than bonding pairs of electrons

69
Q

Order of repulsion between non-bonding domains (NBD) and bonding domains (BD)

A

NBD-NBD > NBD-BD > BD-BD

Greatest repulsion = between non-bonding domains
Least repulsion = between bonding domains

70
Q

Purpose of Lewis structures

A

show all the valence electrons in a covalently bonded species

71
Q

Giant covalent structures

A

Substances that don’t exist as discrete molecules

- instead, exist as many atoms bonded together in a regular arrangement

72
Q

Diamond

A
  • each carbon atom in diamond is covalently bonded to 4 other carbon atoms in a tetrahedral arrangement w/ a bond angle of 109.5 degrees
  • strong covalent bonds between carbon atoms mean that it’s a very hard substance w/ a high MP and BP
  • precious gemstone, used in jewellery
  • poor electrical conductor- no delocalised electrons within its structure
73
Q

Silicon

A
  • group 14 element
  • each silicon atom is bonded to 4 other silicon atoms in a tetrahedral arrangement, bond angle of 109.5 degrees
  • bonds between silicon atoms are strong covalent bonds- electrons are held in fixed positions, can’t move within the structure- silicon is a poor conductor of electricity at low temp.
  • electrical conductivity can be improved by doping (addition of small amounts of elements eg. phosphorus or boron to pure silicon)- process used in production of photovoltaic cells
74
Q

Silicon dioxide

A
  • each silicon atom is bonded by strong covalent bonds to 4 oxygen atoms in a tetrahedral arrangement w/ bond angle of 109.5 degrees
  • Si - O - Si bonds have a bent arrangement caused by lone pairs of electrons on oxygen atoms
  • strong covalent bonds between atoms are responsible for properties of silicon dioxide
  • a very hard substance, a poor conductor of electricity, high MP and BP
75
Q

Allotropes of carbon

A

Carbon exists as 3 allotropes:

  • diamond
  • graphite
  • fullerenes

These allotropes have different properties eg. electrical conductivity and hardness
- because of different bonding within structures

76
Q

Allotropes

A

Different forms of the same element in the same physical state

77
Q

Graphite

A
  • each carbon atom is covalently bonded to 3 other carbon atoms
  • has a layered structure consisting of carbon atoms arranged in fused hexagonal rings
  • a graphite crystal is composed of many such planar sheets of hexagonally arranged carbon atoms, stacked on top of one another
  • layers are held together by relatively weak London dispersion forces
  • each carbon atom has an electron which becomes delocalised across the plane
  • delocalised electrons explains why graphite can conduct electricity along the plane of the crystal when a voltage is applied
  • graphite is soft, hence, used as ‘lead’ in pencils (layers are able to slide over one another due to weak intermolecular forces between layers, making graphite soft and slippery)
78
Q

Properties and uses of diamond

A
  1. colourless transparent crystals that sparkle in light - used in jewellery and ornamental objects
  2. Hardest natural substance - used in drill bits and diamond saws
  3. Doesn’t conduct electricity
79
Q

Properties and uses of graphite

A
  1. Dark grey shiny solid
  2. Soft- used in pencils and as a lubricant
  3. Good electrical conductor- used as electrodes and in electric motors
80
Q

C60 fullerene

A
  • a simple molecular substance even though it contains so many bonded carbon atoms
  • structure is made of carbon atoms bonded together in 20 hexagons (6 carbon rings) and 12 pentagons (5 carbon rings), truncated icosahedron
  • each carbon atoms is covalently bonded to 3 other carbon atoms
  • presence of delocalised electrons within structure means that it does conduct electricity, poorer electrical conductor than graphite
81
Q

Graphene

A
  • single- layered material made of individual sheets of graphite
  • composed of single layers of graphite w/ a bond angle of 120 degrees between carbon atoms in a trigonal planar arrangement
  • high tensile strength
  • high electrical and thermal conductivity
  • useful in energy and biomedical industries
  • behaves as a semi-metal, it’s suitable for electronic devices

NB/ of the different allotropes of carbon, graphene is the most chemically reactive
- because of the reactive edges of the structure, where there are carbon atoms with unoccupied bonds

82
Q

Intramolecular forces

A

Atoms within a molecule are bonded by covalent bonds- bonds that are within the molecule

83
Q

Intermolecular forces

A

Forces that exist between molecules

- responsible for physical properties eg. BP, MP, solubility of a substance

84
Q

3 types of intermolecular forces

A
  • London dispersion forces
  • dipole-dipole forces
  • hydrogen bonding

NB/ London dispersion, dipole- induced dipole and dipole-dipole forces are collectively known as van der Waal’s forces

85
Q

London dispersion forces

A

refers to instantaneous induced induced dipole-induced dipole forces that exist between any atoms or groups of atoms

  • should be used for non-polar entities
  • weakest type of intermolecular force

NB/ have significant effects on physical properties of molecular covalent substances

86
Q

Induced dipoles

A

A molecule with a temporary dipole can induce a dipole in a neighbouring molecule- an induced dipole
- temporary dipole in one atom induces a dipole in adjacent atoms

87
Q

Temporary dipoles

A

Temporary dipole- caused by changes in electron density within an atom or molecule

  • at a certain moment in time, electrons may be concentrated on one side of an atom or molecule- gives this side a slight -ve charge, and opp. side a slight +ve charge
  • these opposite charges are known as dipoles
88
Q

The strength of London dispersion forces depends on:

A
  • ease with which the electrons in an atom or molecule form a temporary or induced dipole (their polarizability)
  • surface area of the molecule
89
Q

Strength of London dispersion forces and polarisability

A

In general, as molar mass of a molecule increases, so does its polarizability

  • results in a larger temporary or induced dipole being formed within the molecule, leads to stronger London dispersion forces between the molecules
  • as London forces between molecules increase, BP of substance also increases
90
Q

Halogens and BP

A
  • as we move down group, molar mass of molecules increases
  • results in an increase in strength of London dispersion forces between molecules
  • increase in BP
91
Q

Strength of London dispersion forces and surface area

A
  • molecules that have a greater area between them have stronger London dispersion forces and higher BP
92
Q

Dipole-dipole forces

A
  • only exist between polar molecules that have a permanent dipole eg. HCl
93
Q

Hydrogen chloride as a permanent dipole

A
  • permanent dipole arises due to difference in electronegativity between hydrogen and chlorine
  • hydrogen, w/ lower electronegativity value, has a positive dipole while chlorine, w/ higher electronegativity value has a negative dipole
  • force of attraction is between -ve dipole of chlorine atom of one molecule and +ve dipole of hydrogen atom on another
94
Q

Hydrogen bonding

A

Occurs between molecules that have an electronegative nitrogen, oxygen or fluorine atom directly bonded to a hydrogen atom (H-N, H-O, H-F)
- stronger type of dipole-dipole attraction, responsible for high BP of water

95
Q

Hydrogen bonding in water

A
  • Hydrogen bonding is responsible for high BP of water

- hydrogen bonds occur between lone pairs of electrons on oxygen atom and hydrogen atom on a nearby water molecule

96
Q

Relative strength of intermolecular forces

A

Hydrogen bonding > dipole-dipole forces > London dispersion forces

97
Q

Solubility

A
  • water is the universal solvent- able to dissolve many different substances
  • many ionic compounds are soluble in water due to its polar nature
  • certain covalent substances are also soluble, able to form hydrogen bonds w/ water molecules
  • polar substances are soluble in polar solvents
  • non-polar substances are soluble in non-polar solvents
98
Q

Metallic structure

A

In metals:

  • valence electrons are free to move throughout the metallic structure
  • electrons are delocalised (‘sea of delocalised electrons’)
  • metal atoms are ionised, forming positive ions that are arranged in a lattice structure
  • bonding in metals = metallic bond
99
Q

Properties of metals

A
  • good conductors of heat and electricity
  • ductile (can be made into wires)
  • malleable (can be bent into shape)
  • shiny when polished
100
Q

Metallic bond

A

The electrostatic attraction between lattice of positive metal ions and ‘sea’ of delocalised electrons
- non-directional- force of attraction occurs in all directions between positive ions and delocalised electrons within lattice structure

101
Q

Metals’ malleability and ductileness

A
  • if sufficient force is applied to metal, one layer of metal ions can slide over each other without disrupting metallic bond- bond remains intact
  • explains why metals are malleable and ductile
  • bending a metal into shape doesn’t break the metallic bond, instead, layers of metal ions slide over each other without disrupting metallic bond
102
Q

Why can metals conduct electricity?

A
  • when a potential difference is applied across a metal, a direction is imposed on movement of delocalised electrons
  • they’re repelled from negative electrode and move to positive electrode
  • this orderly flow of delocalised electrons in a given direction constitutes flow of an electric current
  • delocalised electrons can also conduct heat,
103
Q

Why can metals conduct heat?

A
  • delocalised electrons can conduct heat
  • electrons move through the metal, carrying kinetic energy (in form of vibrations) from hotter part of metal to colder part
  • these delocalised electrons are also responsible for shininess of metals, because they reflect wavelengths of visible light
104
Q

The strength of the metallic bond depends on:

A
  • charge on the metal ion
  • ionic radius of the metal ion

The MP can be taken as approx. measure of strength of metallic bond
- stronger the bond, higher the MP

105
Q

Alloys

A

Homogeneous mixtures composed of two or more metals or a metal and a non-metal

  • have different properties to the metal they’re made from
  • they’re often stronger and more resistant to corrosion than their component elements
106
Q

How are alloys made?

A
  • alloys are produced by adding one or more metal or non-metal elements to another metal in its molten state so different atoms can mix
  • as mixture solidifies, atoms of different metals or non-metals are scattered through the lattice structure
107
Q

Why are metal malleable?

A

Due to layers in metallic lattice being able to slide over each other without breaking the metallic bond

108
Q

Why are alloys harder than component metals?

A

Addition of different sized atoms in alloy means that layers can’t slide over each other as easily

  • results in alloys being harder than its component metals
  • because of different packing of atoms in metallic lattice, alloys have properties that are different from their component elements
  • alloys are often stronger, more chemically stronger and more resistant to corrosion
109
Q

Properties of Brass

A

Composition: 70% copper, 30% zinc

  • harder than pure copper
  • used to make musical instruments
110
Q

Properties of Bronze

A

Composition: 90% copper, 10% tin

  • harder than pure copper
  • used to make statues
111
Q

Properties of Mild steel

A

Composition: 99.7% iron, 0.3% carbon

- stronger and harder than pure iron

112
Q

Properties of Stainless steel

A

Composition: 74% iron, 18% chromium, 8% nickel

  • increased resistance to corrosion
  • used to make cutlery
113
Q

Properties of Solder

A

Lead solder: 50% tin, 50% lead
Lead-free solder: tin w/ various metals
- lower MP than either tin or lead
- used in electrical circuit boards

114
Q

Resonance structures

A

exist for molecules that have more than one possible Lewis structure

115
Q

Resonance hybrid structures

A
  • a hybrid of all the possible resonance structures of that compound
  • has bonds of intermediate length and strength
  • each resonance structure contributes to the hybrid structure depending on its energy
  • resonance structure w/ lowest energy contributes the most to the hybrid structure

NB/ resonance hybrid is always more stable than any of the possible resonance structures

Dashed lines indicate that each bond is identical in terms of length and strength

116
Q

Equivalent resonance structures

A
  • when all resonance structures are of equal energy

- thus, make equal contributions to resonance hybrid

117
Q

Resonance energy

A

difference in energy between the hybrid structure and that of the most stable resonance structure

118
Q

Equivalent Lewis structures

A

each resonance structure contains same no. of single and double bonds as another

119
Q

Non-equivalent Lewis structure

A

each resonance structure contains different no. of multiple bonds

120
Q

Formal charge

A

used to determine which Lewis structure is the preferred one when there is more than one possibility

  • the charge an atom would have if all atoms in the molecule had the same electronegativity
  • assumes that a bond is 100% covalent, w/ no ionic character
121
Q

Equation used to calculate formal charge

A

FC = (no. of valence e-) - 1/2(no. of bonding e-) - (no. of non-bonding e-)

Use:
V= valence e-
B= bonding e-
L= non-bonding e-

FC = V - 1/2B - L

NB/ preferred Lewis structure is the one w/ the formal charge closes to 0 on each atom

  • for neutral compounds, the sum of the formal charges must be equal to 0 eg. CO2
  • sum of formal charges must = charge on the ion
122
Q

What happens when there’s more than one Lewis structure w/ formal charge closest to 0?

A

Assign -ve formal charge to the more electronegative atom

123
Q

Sum of formal charges

  • neutral molecule
  • polyatomic ion
A

Neutral molecule: sum of formal charges must = 0

Polyatomic ion: sum of formal charges must = overall charge on the ion

124
Q

Expanded octets

A
  • elements in period 3 and beyond can accommodate more than 8 e- in their valence shells
  • due to availability of d orbitals which can be used for bonding
  • expanded octets results in formation of molecules w/ 5 and 6 electron domains around the central atom
125
Q

Electron domains

A

repel each other to be as far apart as possible

126
Q

Electron domain geometries for 5 and 6 electron domains

A

5 electron domains: trigonal bipyramidal

6 electron domains: octahedral

127
Q

Electron domain geometry

A

total no. of electron domains (both bonding and non-bonding) around the central atom

128
Q

Molecular geometry

A
  • total no. of electron domains (both bonding and non-bonding) around the central atom
  • takes into account the extra repulsion between bonding and non-bonding e-

NB/ for some molecules, molecular geometry is different to e- domain geometry

129
Q

Molecules w/ 6 electron domains

A

Electron domain geometry: Octahedral

Molecular geometry:

  1. Octahedral:
    - 6 bonding domain, 0 non-bonding domains
    - bond angles are 90 degrees and 180 degrees
  2. Square pyramidal:
    - 5 bonding domains and 1 non-bonding domain
    - bond angle of less than 90 degrees
    - reduced bond angle is due to extra repulsion between bonding and non-bonding domains around central atom
  3. Square planar:
    - 4 bonding domains and 2 non-bonding domains
    - bond angle of 90 degrees
130
Q

Molecules w/ 5 electron domains

A

Electron geometry: Trigonal bipyramidal

Molecular geometry:

  1. Trigonal bipyramidal:
    - 5 bonding domains, 0 non-bonding
  2. See-saw:
    - 4 bonding domains, 1 non-bonding domain
    - bond angles are 90 degrees and less than 120 degrees
  3. T-shaped:
    - 3 bonding domains and 2 non-bonding domains
    - bond angle is less than 90 degrees
  4. Linear:
    - 2 bonding domains and 3 non-bonding domains
    - bond angle is 180 degrees

NB/ non-bonding domains occupy equatorial positions to minimise effect of repulsive forces between non-bonding and bonding domains

131
Q

Covalent bonding

  • single
  • double
A

Covalent bond: formed by the sharing of valence e- between atoms

Single: involves sharing of 2 e-
Double: involves sharing of 4 e-
Triple bond: involves sharing of 6 e-

132
Q

Delocalised pi electrons

A

Electrons that are shared between more than 2 nuclei

  • they are present in molecules and polyatomic ions for which there’s more than one possible Lewis structure
  • originate from overlap of pi bonds in a molecule or ion
133
Q

Single covalent bond

A

composed of 1 sigma bond

134
Q

Double covalent bond

A

composed of 1 sigma and 1 pi bond

135
Q

Triple covalent bond

A

composed of 1 sigma and 2 pi bonds

136
Q

Bond strengths between pi and sigma bonds

A
  • pi bond isn’t as strong as a sigma bond
  • extra strength of sigma bonds comes from the greater overlap of atomic orbitals in the bond
  • in a pi bond, atomic orbitals can’t overlap as much, results in a weaker bond
137
Q

Properties of the sigma bond

A
  • formed by head-on axial overlap of atomic orbitals
  • electron density is concentrated between the nuclei of the bonding atoms
  • a sigma bond can be present on its own, or in combinate w/ pi bonds
  • there’s free rotation of atoms around a sigma bond

Formed by:

  • head-on overlap of s and p orbitals
  • head-on overlap of two p orbitals
138
Q

Properties of pi bond

A
  • formed by sideways overlap of atomic orbitals
  • electron density is concentrated above and below the plane of the nuclei of the bonding atoms
  • it isn’t formed independently, it’s only formed in the presence of a sigma bond
  • there is no free rotation around a pi bond because it would involve breaking the bond
139
Q

Ozone

A
  • also called trioxygen
  • composed of 3 oxygen atoms bonded together w/ a V- shaped or bent molecular geometry
  • central oxygen atom has 3 electron domains- 2 bonding, 1 non-bonding
  • extra repulsion from non-bonding pair of e- gives a bond angle of approx. 116 degrees
  • ozone molecule has delocalised pi e-
  • these e- are spread over both bonding positions in the molecule- results in both bonds being identical in both bond length and bond strength
  • these intermediate bonds are represented by the dashed lines, resonance hybrid structure
  • ozone molecule is polar (permanent dipole) due to different formal charges on each oxygen atom
140
Q

Calculation of wavelength of light required to dissociate oxygen and ozone

A
  1. Divide oxygen’s bond enthalpy by Avogadro
  2. Use wavelength = (hc)/E
  3. Convert to nm by multiplying by 10^9
141
Q

Chain reaction

A
  • CFC’s are highly stable compounds
  • means that CFC molecules released into lower atmosphere could remain intact and reach upper atmosphere
  • when exposed to UV radiation, compounds such as trichlorofluromethane decompose to produce chlorine free radicals
  • chlorine free radicals act as catalysts- break down ozone to form diatomic oxygen molecules
  • free oxygen atoms are also present due to action of powerful UV radiation on molecular oxygen
  • they’re part of the normal generation of ozone that takes place in upper atmosphere
142
Q

Ozone depletion

A
  • other compounds that cause ozone depletion: nitrogen oxides eg. nitrogen monoxide and nitrogen dioxide
  • both molecules have an unpaired e- and are free radicals
  • nitrogen oxides are produced in internal combustion engines- high temp. allow atmospheric nitrogen and oxygen to react
  • in this reaction nitrogen oxide acts as a catalyst in the reaction
143
Q

Carbon in the ground state

A
  • each carbon atom forms 4 covalent bonds in many compounds
  • its electron config. shows that it only has 2 half-filled atomic orbitals available for bonding (2pz orbital is unoccupied)
  • in its excited state, an e- has been promoted from a doubly occupied 2s orbital to empty 2pz orbital- excitation
  • energy required to promote an e- to a higher energy orbital is made up for when carbon atom forms 4 new bonds

Electron config. of excited carbon atom:

  • has 4 half-filled orbitals (one 2s and three 2p orbitals)
  • but, if carbon were to form 4 covalent bonds w/ this electron config., bonds wouldn’t be equal
  • when carbon forms 4 bonds, all bonds are equal
144
Q

Hydbridisation

A

involves mixing of atomic orbitals to form hybrid orbitals to be used for bonding

  • orbitals of the same atom, which are of different energy, can overlap to form hybrid orbitals of equal energy
  • no. of hybrid orbitals formed is equal to that no. of atomic orbitals involved in the process of hybridisation that has taken place

3 main types:

  • sp
  • sp^2
  • sp^3
145
Q

sp^3 hybridisation

A
  • involves mixing of one s orbital and three p orbitals to form four sp^3 orbitals
  • these adopt a tetrahedral arrangement
  • the hybrid orbitals repel each other and arrange themselves in a config. that minimises e- repulsion by spreading themselves as far apart as possible
146
Q

sp^2 hybridisation

A
  • one 2s orbital mixes w/ tow of the 2p orbitals to form three sp^2 hybrid orbitals
  • one 2p orbital doesn’t participate in the hybridisation process, leaving one unhybridised 2p orbital
147
Q

sp hybridisation

A
  • one 2s orbital mixes w/ a 2p orbital to form two sp hybrid orbitals
  • one of the 2p orbitals is involved in the hybridisation, so there are two unhybridised 2p orbitals remaining
148
Q

2 electron domains

A

Electron domain geometry: linear

Molecular geometry: linear

Hybridisation: sp

149
Q

3 electron domains

A
  1. Electron domain geometry: trigonal planar
    Molecular geometry: trigonal planar
    Hybridisation: sp^2
  2. Electron domain geometry: trigonal planar
    Molecular geometry: V-shaped
    Hybridisation: sp^2
150
Q

4 electron domains

A
  1. Electron domain geometry: tetrahedral
    Molecular geometry: tetrahedral
    Hybridisation: sp^3
  2. Electron domain geometry: tetrahedral
    Molecular geometry: trigonal pyramidal
    Hybridisation: sp^3
  3. Electron domain geometry: tetrahedral
    Molecular geometry: V-shaped
    Hybridisation: sp^3
151
Q

What determines how many electrons are lost or gained?

A

The no. of e- lost or gained is determined by the electron configuration of the atom

152
Q

Increasing the bond length and strength

A

Bond length decreases and bond strength increases as the no. of shared electrons increases

153
Q

Deducing the polar nature of a covalent bond

A
  • polar bonds result from unequal sharing of electrons
  • the more electronegative atom exerts a greater pulling power on the shared electrons
  • hence this holds the electrons more than the other
  • hence, electron distribution is unsymmetrical
  • this is because one end of the molecule has more electrons than the other
  • the more electronegative atom has the greatest electron density when bonded
154
Q

Molecular polarity

A
  • If the dipoles in the BONDS cancel out then the MOLECULE will be non-polar
  • If the net BOND dipoles are non-zero then the MOLECULE will be polar - electronegativity values are needed to calculate bond dipoles and molecular geometry to see if these cancel
155
Q

Non-directional bonding in alloys

A

production of alloys is possible because of the non-directional nature of the delocalised electrons
- the lattice can accomodate ions of different sizes

156
Q

Alloys vs. regular metals

A

Alloys are usually more stronger than regular metals

  • because if different atoms are present, the regular network of +ve ions is disturbed
  • atoms of a different size make it harder for layers of +ve ions to slide over each other
  • thus prevent bending or denting of the metal
157
Q

Covalent bonding in terms of pi and sigma bonds

A
  • a covalent bond results from the overlap of atomic orbitals
158
Q

Resonance

A

using two or more Lewis structures to represent a particular molecule or ion

IB Chemistry HL (9 decks)

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