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Chemistry > 2: Bonding and Structure > Flashcards

2: Bonding and Structure Flashcards

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1
Q

Define ionic bonding

A

The electrostatic force of attraction between a positively charged metal ion and a negatively charged non-metal ion, involving the transfer of electrons from the metallic element to the non-metallic element

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2
Q

How are cations and anions bonded together

A
  • Cations and anions are oppositely charged, so are attracted to each other
  • Electrostatic forces of attraction are formed between the oppositely charged ions forming ionic compounds
  • This ionic bond is very strong and require lots of energy to overcome
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3
Q

How are ionic solids arranged

A
  • Ions form giant ionic lattice structures, which are evenly distributed crystalline structures
  • The ions are arranged in a regular repeating pattern, so positive charges cancel out negative charged, so the lattice is neutral
  • Non-directional bonding occurs, so each ion is attracted to all the oppositely charged ions around it
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4
Q

What are dot and cross diagrams, draw one using CaF2

A
  • They show the arrangement of the outer shell electrons in ionic or covalent compounds or elements
  • Each species electrons are represented differently (dot/cross)
  • Charge of the ion is shown in square brackets
  • The ions together always make a neutral compound
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5
Q

What is ionic radius, and its trend in -ve and +ve ions

A
  • The radius of an ion
  • Ionic radii increase with increasing negative charge
  • Ionic radii decrease with increasing positive charge
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6
Q

Why does ionic radii increase with increasing negative charge

A
  • Anions are formed by gaining electrons
  • This makes the valance electrons be further away from the nucleus, having a weak hold on it
  • Causing the ionic radii to increase
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7
Q

Why does ionic radii decrease with increasing positive charge

A
  • Cations are formed by losing electrons
  • There are fewer electrons which undergo a greater electrostatic force of attraction to the nucleus
  • Decreasing the ionic radii
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8
Q

What are isoelectric ions, and describe their trend in ionic radii

A

Ions that have the same electronic configuration, however the number of protons in the nucleus stays the same as the original atom, so the electrons get pulled in more, and the ionic radii decreases with isoelectric ions that have larger proton numbers

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9
Q

What does the type of giant ionic lattice depend on

A
  • Depends on the sizes of the positive and negative ions, arranged in an alternating fashion
  • E.g. MgO is cubic
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10
Q

Why are most ionic compounds solid at room temperature

A
  • Due to the strong electrostatic forces of attraction between the oppositely charged ions in the lattice requiring a lot of energy to overcome
  • Giving them very high melting/boiling points
  • Lattices with ions that have higher charge also have high melting/boiling temperatures due to the stronger attraction between the ions
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11
Q

Can ionic compounds conduct electricity

A
  • Electrical current can only floe is there are freely moving charged particles
  • Ionic compounds can conduct electricity in molten states or in solutions, as their ions can move and carry charge
  • Ionic compounds can’t conduct electricity in solid states
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12
Q

Are ionic compounds soluble

A
  • Solubility is dependant on breaking the ionic lattice, and the polar molecules of the solution attracting and surrounding the ions
  • Polar solvents (e.g. water) can break down and surround the ionic lattice and ions, with the +ve end of the polar molecule attracting the -ve ions and visa versa
  • Solubiltiy if ionic compounds depends on the relative strength of the electrostatic forces of attraction within the ionic lattice, and the attractions between the ions and the polar molecules
  • The greater the ionic charge, the less soluble the compound
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13
Q

Describe how to fine evidence for the presence of ions

A
  • Usuing electrolysis to see the positive ions in a solution being attracted to the negative cathode, and the negative ions being attracted to the positive anode
  • The separation can be viewed using copper(II) chromate
  • The blue Cu+2 ions are attracted to the cathode, and the yellow chromate ions attracted to the anode
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14
Q

Define covalent bonding

A

The electrostatic attraction between the nuclei of two non-metal atoms, and their shared pair of electrons

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15
Q

What is double covalent bonding

A

Where two atoms share 4 electrons to form a double covalent bond (e.g. O2/CO2/C2H4)

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16
Q

What is triple covalent bonding

A

Where each atom shares 3 electrons to form a triple bond (e.g. N2)

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17
Q

What is dative covalent bonding

A
  • In some molecules which have a lone pair of electrons, it can be donated to form a bond with an electron deficient atom
  • Both of the electrons are from the same atom (e.g. ammonium)
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18
Q

How is aluminium chloride formed with dative covalent bonds

A
  • At high temperatures chloride can exist as the monomer AlCl3
  • At low temperatures, two molecules of AlCl3 from to make a dimmer Al2Cl6, forming two dative bonds
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19
Q

Define bond energy

A

The energy required to break one mole of a particular covalent bond in the gaseous state. The larger the bond Enthalpy, the stronger the covalent bond

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20
Q

What is bond length

A

The internuclear distance of two covalently bonded atoms

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21
Q

How is bond length decreased

A

It is decreased by the electrons and the nuclei in the atoms having greater forces of attraction, decreasing bond length and increasing strength of the covalent bond

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22
Q

Why are triple bonds the shortest and strongest covalent bonds

A
  • Due to the large electron density between the nuclei of the two atoms
  • This increases the forces of attraction between the electrons and nuclei, causing them to be pulled closer together
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23
Q

What is the valence shell electron pair repulsion theory

A
  • It predicts the shape and bond angles of molecules
  • It considers how negatively charged electrons will repel each other, and how bonding pairs of electrons will repel other electrons around the central atom
  • It predicts shapes for molecules that minimises the repulsion forces
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24
Q

What VSEPR rules are followed when determining shape and bond angles of a molecule

A
  • Valance shell electrons are the outer shell electrons
  • Electron pairs repel each other as they have the same charge
  • Lone pair electrons repel each other more than bonded pairs
  • Repulsion between multiple and single bonds is treated the same as repulsion between single bonds
  • Repulsion between pairs of doubled bonds are greater
  • The most stable shape is adopted to minimise repulsion forces and maximise attraction
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25
Q

What are lone pairs of electrons more repulsive

A
  • Have more concentrated electron charge clouds than bonding pairs
  • Electron charge clouds are wider and closer to the central atoms nucleus
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26
Q

What is the linear bond shape and angle

A
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27
Q

What is the trigonal bond shape and angle

A
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28
Q

What is the non-linear bond shape and angle

A
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29
Q

What is the pyramidal bond shape and angle

A
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30
Q

What is the tetrahedral bond shape and angle

A
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31
Q

What is the trigonal bipyramidal bond shape and angle

A
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32
Q

What is the octahedral bond shape and angle

A
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33
Q

What is the bond shape and angle in BF3 and why

A
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34
Q

What is the bond shape and angle in CO2 and why

A
35
Q

What is the bond shape and angle in CH4 and why

A
36
Q

What is the bond shape and angle in NH3 and why

A
37
Q

What is the bond shape and angle in H2O and why

A
38
Q

What is the bond shape and angle in SF6 and why

A
39
Q

What is the bond shape and angle in PF5 and why

A
40
Q

Define electronegativity

A

The power of an atom to attract the pair of electrons in a covalent bond towards itself

41
Q

Why does electronegativity occur

A

The positive nucleus has the ability to attract the negatively charged electrons in the outer shell, towards itself. With highly electronegative atoms, they can pull both of the electrons in the covalent bond towards their nucleus

42
Q

What is the Pauling scale

A

The scale used to assign a value of electronegativity for each atom, with Fluorine being the most electronegative

43
Q

How does nuclear charge affect electronegativity

A
  • Attraction exists between the positively charged protons in nucleus, and negatively charged electrons in the shells
  • Increased number of protons means an increase in nuclear attraction for the valance electrons
  • Therefore an increased nuclear charge results in increased electronegativity
44
Q

How does atomic radius affect electronegativity

A
  • Atomic radius is the distance from the outer shell to the nucleus
  • Valance electrons closer to the nucleus are more strongly attatched to it
  • Therefore increased atomic radius results in decreased electronegativity
45
Q

How does shielding affect electronegativity

A
  • Filled energy shields can shield the affect of the nuclear charge, causing valance electrons to be less attracted to the nucleus
  • The addition of extra shells and sub shells in an atom causes the valance electrons to experience less attractive force to the nucleus
  • Therefore increased number of inner shells and sub shells, increases shielding and decreases electronegativity
46
Q

What is the trend in electronegativity down a group

A

Electronegativity decreases down a group

47
Q

Why does electronegativity decrease down the group

A
  • Nuclear charge increases as more protons are added to the nucleus
  • However a new shell is added down the group, increasing shielding
  • This increases atomic radii, causing an overall decrease in the attraction between the valance electrons and the nucleus
48
Q

What is the trend in electronegativity across a period

A

Electronegativity increases across a period

49
Q

Why does electronegativity increase across a period

A
  • The nuclear charge increases without the addition of protons to the nucleus
  • Shielding remains constant as no new shells are being added to the toms
  • The nucleus has an increasingly strong attraction for the valance electrons, decreasing the atomic radii
50
Q

What is a nonpolar molecule

A

When two atoms in a covalent bond have the same electronegativity

51
Q

What determines bond polarity

A
  • The difference in electronegativity between atoms
  • Different electronegativities (more than 1.7) cause ionic bonds
  • Different electronegativities of 0.3-1.7 cause a covalent bond to form, and the bond is polar
52
Q

What does a polar molecule look like

A
  • The negative charge centre and positive charge centre don’t coincide
  • Therefore the electron distribution is asymmetric
  • The less electronegative atom gets a partial charge of delta positive
  • The more electronegative atom gets a partial charge of delta negative
  • The greater the difference in electronegativity, the more polar the bond becomes
53
Q

How is polarity assigned in molecules with multiple atoms

A
  • Polarity of each bond, and arrangement of bonds is considered
  • Molecules can have polar bonds, but their arrangement,ent means the molecule is no polar (e.g. CCl4/CO2/BF3)
  • Molecules can also have polar bonds that are arranged in a way that they don’t cancel out a the molecule is polar (e.g. CH3Cl)
54
Q

What is a helpful clue about molecular polarity

A

Molecules which are symmetrical are unlikely to be polar

55
Q

What are intramolecular forces

A

Forces within a molecule (usually covalent bonds), and drawn with full straight lines

56
Q

What are intermolecular forces, and the types

A
  • Weak forces between molecules, drawn with dashed straight lines
  • Induced dipole-dipole (Van der Waals/London dispersion) forces
  • Permanent dipole dipole forces
  • Hydrogen bonding (special type of permanent dipole force)
57
Q

List the intermolecular forces from strongest to weakest

A
58
Q

What are induced dipole-dipole forces

A
  • Exist between all atoms or molecules
  • Electron clouds in non-polar molecules are always moving unevenly around the molecule
  • This can cause a temporary dipole to arise
  • The temporary dipole can induce a dipole on a neighbouring molecule, causing the +ve and -ve side of the two molecules to be attracted to one another
  • Only temporary as electron clouds move constantly
59
Q

What are permanent dipole-dipole forces

A
  • Polar molecules have permanent dipoles, so always have a negative and positively charged end
  • The +ve end of one molecule can be attracted to the -ve end of another molecule, forming a permanent dipole
60
Q

What is the relative strength between permanent and induced dipole-dipole forces

A

For small molecules with the same number of electrons, permanent dipole forces are stronger, and require more energy to break

61
Q

What is hydrogen bonding

A
  • The strongest form of intermolecular bonding
  • They are a type of permanent dipole-dipole force
  • Takes place only between species that has an O/N/F (very electronegative) atom bonded to a hydrogen
  • Causing the bond to become highly polarised, allowing the H to become delta positive so it can form a bond with the lone pair on O/N/F
  • E.g. water
62
Q

What anomalous properties does water have due to hydrogen bonding

A
  • High melting/boiling points
  • High surface tension
  • Higher density in liquid form rather than solid
63
Q

Why does water have high melting/boiling points

A
  • Due to the strong intermolecular forces of hydrogen bonding between the molecules, which require a lot of energy to overcome
  • The graph comparing Enthalpy of vapourisation of different hydrides shows an increase from H2S to H2Te due to increased electrons and therefore instantaneous dipole dipole forces
  • But it shows H2O having a much higher Enthalpy due to hydrogen bonds
64
Q

Why does water have high surface tension

A
  • Water molecules at the surface are bonded together by hydrogen bonds
  • These molecules pull downwards on the surface molecules causing the surface of them to become compressed and more tightly together at the surface, increasing its tension
65
Q

Why does water have high surface tension

A
  • Water molecules at the surface are bonded to other molecules via hydrogen bonds
  • These molecules pull down on the surface molecules, causing the surface of them to become compressed and tight, increasing tension
66
Q

Why is water more dense as a liquid than a solid

A
  • The particles in solids are more closely packed than in liquids
  • The packing of the molecules and the length of the hydrogen bonds means that the water molecules are slightly further apart than in liquid state
  • Therefore ice has a lower density by about 9%
67
Q

How does branching of a molecule affect intermolecular forces

A
  • The larger the surface area of the molecule, the more contact it will have with adjacent molecules
  • Surface area is reduced by branching
  • Branching affects the ability of the dipole forces inducing adjacent molecules
  • Causing a decrease on melting/boiling temperature
68
Q

How do number of electrons affect intermolecular forces

A
  • The greater the number of electrons, the greater the distortion, and the greater the frequency and magnitude of temporary dipoles
  • Increasing the melting/boiling point
69
Q

Why is the boiling point of Alkenes lower than that of alcohols

A
  • Alkenes and alcohols contain induced dipole-dipole bonds, but only alcohols contain hydrogen bonds due to the O-H bonds
  • Hydrogen bonds require more energy to overcome than induced dipole bonds
70
Q

Why does the boiling point of hydrogen halides increase down the group, and why is HF an exception

A
  • Boiling points increase down the group as the molecules get larger and the extra electrons allow greater temporary dipoles
  • Hydrogen fluoride has hydrogen bonding, with the H-F bond being very polar, causing it to have a higher boiling point
71
Q

What is the general principle of solubility

A

Like dissolves like

72
Q

How is solubility affected by intermolecular forces

A
  • Polar covalent substances generally dissolve in polar solvents as a result of dipole-dipole interactions/the formation of hydrogen bonds between the solute and solvent
  • As molecules increase in size, solubility can decrease due to the polar part of the molecule being the small part
  • Giant covalent substances usually don’t dissolve due to lots of energy needed to overcome the strong covalent bonds in the lattice
  • Polar substances can’t dissolve well in non polar solvents due to the dipole-dipole interactions being unable to react well
73
Q

How does metallic bonding work

A
  • Metal atoms are packed tightly in lattice structures
  • In these structures, the electrons are allowed to freely move, becoming delocalised electrons
  • When electrons are delocalised, the metal atoms become positively charged ions, which repel each other and keep the shape of the lattice
  • Strong forces are present between the positive metal centres and the sea of delocalised electrons
74
Q

What is the structure of ionic lattices

A
  • Ions form giant ionic lattice structures, as evenly distributed crystalline structures
  • Ions are arranged in regular repeating patterns so positive charges cancel negative charges
  • Attraction between cations and anions occurs in all directions
75
Q

What is the structure of metallic lattices

A
  • Metals form giant metallic lattices where metal ions are surrounded by a sea of delocalised electrons
  • Metal ions are often packed in hexagonal layers or in cubic arrangement
76
Q

What is the structure of simple covalent lattices

A
  • Iodine/Ice/Buckminsterfullerene
77
Q

What is the structure of giant covalent lattices

A
  • Graphite/diamond/silicon oxide
78
Q

What are simple molecular compounds

A

Covalent substances with smaller molecular structures

79
Q

Why do giant covalent lattices have high melting/boiling points

A
  • Large number of covalent bonds
  • Lots of energy required to break bonds in lattice
80
Q

Describe the structure and bonding in graphite

A
  • Each carbon atom is bonded to three others, forming layers of hexagons, leaving one free electron per carbon atom
  • Free electrons freely move and can conduct electricity
  • Covalent bonds between layers are very strong
  • Layers are attracted to one another by weak intermolecular forces, so they can slide over each other, making graphite soft
81
Q

Describe the structure and bonding in diamond

A
  • Each carbon atom bonds to four other carbon atoms, forming a tetrahedron
  • No intermolecular forces
  • Very hard
82
Q

Describe the structure and bonding in graphene

A
  • Consists of a single layer of graphite (sheet of carbon atoms)
  • 2D molecule
  • Has unusual properties
83
Q

How many degrees do lone pairs take away from the bond angle

A

2.5

Chemistry (19 decks)

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