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Chemistry > Topic 2 - Bonding and Structure > Flashcards

Topic 2 - Bonding and Structure Flashcards

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1
Q

what is an ionic bond

A

ionic bonding refers to the strong electrostatic forces between oppositely charged ions in a lattice

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2
Q

what is an ion

A

a charged particle or molecule formed from gain or loss of electrons

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3
Q

what does ionic bonding occur between

A

occurs usually between a metal and a non metal.

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4
Q

describe the relationship between the ionic radius and the strength of the ionic bond

A
  • size increases, strength decreases
  • larger ions have a greater ionic radius
  • so have a weaker attraction to the oppositely charged ion as the forces have to act over a greater distance
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5
Q

how does the ionic radius change across a period and why

A

across a period (from the left to right) the atomic radius decreases
due to the nuclear charge increasing

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6
Q

how does the ionic radius change down a group and why

A

the ionic radius increases down the group
due to an increasing number of shells

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7
Q

Describe the relationship between the ionic charge and the strength of the ionic bond

A
  • charge increases, strength increases
  • ions with a greater charge will have a greater attraction to the other ions
  • resulting in stronger forces of attraction and therefore stronger ionic bonding
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8
Q

describe the formation of ion in terms of electron loss or gain

A
  • generally metals lose outer electrons to become positively charged ion and non-metals gain electrons to complete their outer shell and become negatively charged
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9
Q

what is x-ray diffraction used for

A

It is a method used to study ionic compounds and spacing between ions.
which makes it possible to calculate the radii of the ions

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10
Q

describe the trends for a set of isoelectronic ions
N^3- to Al^3+

A
  • as the atomic number increases the ionic radii decreases because there are more protons.
  • so there is a greater electrostatic force between the nucleus and the outer shell.

order of increasing radii:
Al3+, —————-> N3-

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11
Q

what are isoelectric ions

A

different atoms that have the same number of electrons

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12
Q

what happens to the ionic radius in isoelectric ions

A
  • the ionic radius decreases in isoelectric ions as the atomic number increase.
  • the shielding and the electrons in the outer shell stay the same
  • so the radius will decrease as the attractive force from the nucleus increases
  • as the nuclear charge increases.
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13
Q

explain how the physical properties of ionic compounds provide evidence for the existence of ions

A
  • have high mp and bp due to strong electrostatic forces of attraction which are from oppositely charged ions
  • they do not conduct electricity unless when aqueous or liquid as ions are free to move. Charged particles must be present to carry charge
  • soluble in water and other polar solvents as they are polar and are insoluble in non-polar solvents
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14
Q

explain how the migration of ions provide evidence for the existence of ions

A
  • migration of ions can be seen in electrolysis
  • in copper chromate there are blue Cu 2+ and yellow CrO4 2- ions
  • at the cathode blue cu 2+ ions will appear and at the anode yellow CrO4 2- ions will appear

this provides evidence for ions due to its attraction to electrodes

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15
Q

what are intermolecular forces

A

weak attractive forces between molecules

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16
Q

what are covalent bonds

A

refers to the strong electrostatic attraction between two nuclei and the shared pairs of electrons between them

generally occurs between two or more non metals

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17
Q

what is a dative covalent bond

A

a bond in which the shared pair of electrons comes from one atom (lone pair)
(the other atom doesn’t share any electrons)

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18
Q

what is a lone pair

A

a pair of electrons in the outer shell which comes from one atom to form a bond with a molecule or ion

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19
Q

how is a dative bond shown

A

instead of a line there is an arrow
H
H–N–>H
H

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20
Q

how are dative covalent bonds different to covalent bonds

A

it is treated the same as a standard covalent bond and reacts in the exact same way

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21
Q

what is meant by the bond length

A
  • the distance between the nuclei of two bonded atoms in a molecule
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22
Q

what is meant by bond enthalpy

A

the strength of a bond

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23
Q

how is bond length investigated

A

by using x-ray diffraction and microwave spectroscopy

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24
Q

what is the relationship between bond lengths and bond strengths for covalent bonds

A
  • bond length depends on the size of the atoms involved and the number of pairs of electrons shared.
  • larger atoms form longer bonds since there is more shielding ( requires less energy to overcome)
  • the strength of the bond varies inversely with its length.
  • a short bond is stronger with a greater bond energy
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25
Q

how does electron density effect bond length and enthalpy

A

The greater the electron density the stronger the attractive force
this means the atoms are pulled in further towards each other
making the bond length short and the bong enthalpy high

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26
Q

what is the length and strength of bonds in relation to single, double and triple covalent bonds

A

single , double, triple
<————– bond length
bond strength————->

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27
Q

what is a bond angle

A

the angle between two covalent bonds in a molecule or giant covalent structure

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28
Q

what is the shape of a molecule determined by

A
  • the shape of a molecule is by the repulsion between electron pairs.
  • the electron-pair repulsion theory - electron pairs in the outer shell repel and get as far apart as possible
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29
Q

what bond pairs have the most repulsion. (lone pairs and bond pairs combinations)

A

lone pair- lone pair
has the most repulsion
lone pair - bond pair
bond pair - bond pair
the least repulsion

for every lone pair you reduce the remaining bong angles by 2.5 degrees

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30
Q

what are the different shapes for molecules
with no lone pairs only bond pairs

A
  • linear
  • trigonal planar
  • tetrahedral
  • trigonal bipyramid
  • octahedral
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31
Q

what is the reason for the shape and bond angles of the linear shape

A

2 bond pairs, each 180 degrees
- in order to get as far apart as possible electron pairs must be on opposite sides
EG.
BeCl2 Cl - Be - Cl
CO2 C = O = C

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32
Q

what is the reason for the shape and bond angles of the trigonal planar shape

A

3 bond pairs, 120 degrees between them
- to get as far apart as possible the electrons pairs occupy the corners of the triangle
eg BCl3

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33
Q

what is the reason for the shape and bond angles of the tetrahedral shape

A

4 bond pairs, 109.5 degrees between them
- electron pairs repel and occupy corners of the tetrahedron
eg CH4

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34
Q

what is the reason for the shape and bond angles of the trigonal bipyramid shape

A

5 bond pairs, 90 and 120 degrees between them
eg PCl

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35
Q

what is the reason for the shape and bond angles of the octahedral shape

A

6 bond pairs, 90 degrees between them
eg SF6

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36
Q

what are the different shapes of molecules with lone pairs and bond pairs

A
  • trigonal pyramidal
  • bent
  • distorted T
  • seesaw
  • square pyramidal
  • square planar
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37
Q

what is the reason for the shape and bond angles of the trigonal pyramidal shape

A

1 lone pair and 3 bond pairs, 107 degrees
would be 109.5 but due to the one lone pair it would make the bond angle 107 degrees.
eg NH3

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38
Q

what is the reason for the shape and bond angles of the bent shape

A

2 lone pairs 2 bond pairs, 104.5 degrees
two lone pairs so minus 5 degrees so the angle is 104.5
eg H20

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39
Q

what is the reason for the shape and bond angles of the T-shape

A

2 lone pairs 3 bond pairs, 87.5 degrees
eg ClF3

40
Q

what is the reason for the shape and bond angles of the seesaw shape

A

1 lone pair 4 bond pairs, 87 and 102 degrees
eg SF4

41
Q

what is the reason for the shape and bond angles of the square pyramidal shape

A

1 lone pair and 5 bond pairs, 81.9 and 90 degrees
eg IF5

42
Q

what is the reason for the shape and bond angles of the square planar shape

A

2 lone pairs and 4 bond pairs, 90 degrees
the bond angles remain the same as the 2 lone pair repel equally from both sides
eg XeF4

43
Q

how do you predict the shape and bond angles in a molecule

A
  • draw a dot and cross diagram of the molecule
  • identify how many bond pairs and lone pairs there are
  • match the amount of LP and BP to the shape.
  • eg 1 lone pair and 5 bond pairs will be a square pyramidal shape
44
Q

define the term electronegativity

A

the power /ability of an atom to attract the bonding electrons in a covalent bond

45
Q

how does electronegativity change across a period

A

across a period the electronegativity increases
- due to the atomic radius decreasing so the electrons are closer to the positive nucleus so stronger electrostatic forces of attraction

46
Q

how does electronegativity change up a group

A

up a group electronegativity increases as the shielding and atomic radius decreases up a group

47
Q

what is the most electronegative element

A

fluorine

48
Q

What does differences in electronegativity lead to

A

differences in electronegativity leads to bond polarity in bonds and molecules

49
Q

what does no differences in electronegativity between two atoms lead to

A

it leads to a pure, non - polar covalent bond

50
Q

explain the electronegativity in ionic bonding

A
  • the difference in electronegativity is so great that the atom with he high electronegativity takes the electron from the other atom

NaCl - Cl is more electronegative so takes the electron from the Na

51
Q

explain the electronegativity in pure covalent bonding

A
  • there is no difference in electronegativity
  • the molecule in electronically symmetrical

eg H2 - they are the same atom so have the same electronegativity.

52
Q

explain what happen in polar covalent bonding

A
  • the difference in electron negativity between the atoms are small.
  • the electrons are still shared between the atoms
  • however the shared pairs of electrons move closer to the more electronegative atom
  • results in the bond being polarised
53
Q

explain the link between electronegativity and the polarity of a bond

A

the larger the difference in electronegativity between the two atoms, the more polar the bond

HCl is more polar than H2
as HCl has a high difference in electronegativity between atoms

54
Q

what is the difference between a non-polar bond and a polar bond

A
  • polar bond result in the more electronegative atom to be slightly negative δ−
  • the less electronegative atom becomes slightly positive δ+
  • this is because the electrons move closer to the more electronegative atom
  • in a non-polar bond the electronegativity between the two atoms are the same
  • so the electrons are equally distributed between the bonding atoms.
    the bond is considered perfectly convalent
55
Q

in a polar bond, which atom does the pair of electrons move towards

A
  • the electrons move towards the more electronegative atoms.
  • this is because the atom has a higher ability to attract electrons
56
Q

what are ionic and covalent bonding

A

ionic and covalent bonding are extremes of a continuum of bonding

57
Q

do polar bonds lead to a polar molecule

A

molecules with polar bonds may not always be polar molecules

58
Q

what is a dipole

A

a bond or molecule which has oppositely charges
HCl - Cl is partially negative, H is partially positive so has a dipole
H2 has no dipole due to being the same atom so is a pure covalent bond

59
Q

what are intermolecular forces

A

weak attractive forces between molecules

60
Q

How are intermolecular forces affected when changing states of matter

A
  • intermolecular forces are partially broken when changing from solid to liquid
  • and fully broken down when turning liquid into gas via evaporation
61
Q

what are the three types of intermolecular forces

A
  • London forces (Van der Waals or induced dipole dipole interaction)
  • permanent dipole interactions
  • hydrogen bonds
62
Q

How are Permanent dipoles produced

A
  • a polar bond is formed when two atoms that are bonded have sufficiently different electronegativities
  • the more electronegative atom draws the negative charge towards itself producing a ∂- region and a ∂+ region.
  • this produces a permanent dipole
    eg HF
63
Q

How do permanent dipoles lead to a lattice -like structure

A

The ∂+ and ∂- regions of neighbouring polar molecules attract each other and hold the molecules together in a lattice-like structure

∂+ ∂- ∂+ ∂- ∂+ ∂-
H - Cl H - Cl H - Cl

64
Q

what atoms do hydrogen bonds act between

A

hydrogen bonds only form between hydrogen and the three most reactive atoms: nitrogen, oxygen and fluorine

H2O, NH3, HF

65
Q

How are Hydrogen bonds formed

A

nitrogen, oxygen and fluorine’s lone pair will form a bond with a partially positive H (∂+) from another molecule.

( the hydrogen is from another molecule)
H (∂+) - O ….. H (∂+) - O
l l
H(∂+) H (∂+)
( water bonding with the hydrogen from another water molecule to show a hydrogen bond shown through A DOTTED LINE)

66
Q

what are the properties of hydrogen bonds

A
  • the strongest intermolecular force
  • molecules with a hydrogen bond have a much higher bp/mp
67
Q

How do London forces act and behave

A
  • London forces act as an induced dipole between molecules
  • they are the weakest type of intermolecular force
68
Q

Describe the link between Mr ( relative formula mass) and London forces

A
  • ## the greater the Mr of the molecule, the stronger the intermolecular forces.
69
Q

Explain why straight chain molecules have a higher mp/bp compared to branched chain molecules

A
  • straight chain molecules experience stronger London forces as they are packed more tightly compared to branched chain molecules.
  • this means that the distance over which the intermolecular forces have to act is decreased, which strengthens the london forces.
  • meaning branched chain molecules have a lower bp/mp.
70
Q

How does the chain length affect London forces

A
  • the longer the chain, the stronger the london forces
  • as the chain length of the molecule increases, so does the relative formula mass (Mr)
  • this results in stronger intermolecular forces (london) between the chains.
  • so the compound has a higher bp.
71
Q

explain why water has a high mp/bp refering to waters structure and bonding

A
  • water has a simple molecular structure but has a high mp/bp
  • this is due to the hydrogen bonds which require a lot of thermal energy/enthalpy to overcome
  • there are alsop lots of bonds as each molecule is bonded to 4 others in a tetrahedral structure
72
Q

explain why ice is less dense than liquid water referring to waters bonding and structure

A
  • water is in a tetrahedral structure
  • when solid the hydrogen bonds hold the molecules in a rigid structure with lots of air gaps.
  • which makes ice less dense than water
73
Q

when will there be a hydrogen bond in molecules

A
  • there will only be hydrogen bonds present if there is a ∂+ hydrogen and a ∂- oxygen, nitrogen or fluorine atom with a lone pair of electrons in a molecule
74
Q

Describe the trends in Boiling point of alkanes with increasing chain lengths

A
  • as the chain length increases so does the relative formula mass,
  • meaning more London forces which means it requires more thermal energy/enthalpy to overcome.
  • so longer the chain length the bp would increase
75
Q

Describe the effect of branching in the carbon chain on the boiling temperatures of alkanes

A
  • branched chain hydrocarbons are less able to pack tightly together
  • branched alkanes also have a smaller surface area
  • this means there are less points of contact.
  • and the distance over which the intermolecular forces at upon increases, which reduces the attraction.
  • meaning less energy is needed to break the bonds apart
76
Q

explain why alcohol have a higher BP than alkanes

A
  • they have a similar relative formula mass, but alcohols have a lone electron pair.
  • this lone pair on oxygen is able to form a hydrogen bond with hydrogen.
  • meaning there are more intermolecular forces that are needed to be overcome
77
Q

describe the trends in boiling temperature of the hydrogen halides, HF to HI

A
  • hydrogen fluoride has the highest BP as it forms hydrogen bonds which is the strongest intermolecular force
  • however after hydrogen fluoride, the BP increases down the group as there number of electron increase so there is more london forces and more energy is needed to separate them
78
Q

explain how water is able to dissolve ionic compounds

A

-ionic compounds dissolve in water as the ionic compound is strongly hydrated by the water molecules due to their polarity.
-the partially negative oxygen would be attracted to the positive ions
-the partially positive hydrogen would attract negative ions
-this forms a new compound

79
Q

explain how water is able to dissolve simple alcohols

A

-alcohols form hydrogen bonds through the oxygen and hydrogen
- alcohols are able to dissolve with water because they form hydrogen bonds with water.
- alcohols with longer hydrocarbon chains do not mix easily with water as they have a hydrophobic tail that is insoluble.
- the longer the hydrocarbon the less miscible

80
Q

explain why water is a poor solvent for compounds

A

polar organic compounds are insoluble in water as they are unable to form hydrogen bonds.

81
Q

why are non aqueous solvents used for compounds with similar intermolecular bonds

A

-compounds with similar intermolecular forces to those in the solvent will generally dissolve
- non polar solutes will dissolve in non - polar solvents

82
Q

define metallic bonding

A
  • it is the strong electrostatic attraction between metal ions and the delocalised electrons
83
Q

How are the ions held together in a metallic bond

A
  • there are strong electrostatic forces of attraction between these oppositely charged particles
  • the greater the charge, the stronger the force of attraction
  • the larger the ion the weaker the attraction due to a greater atomic radii.
84
Q

describe the arrangement of ionic solds

A
  • they are arranged in giant lattices with thousands of ions
85
Q

describe the arrangement of covalently bonded solids

A
  • Macromolecular covalent substances are covalently bonded into a giant lattice, each has multiple covalent bonds.
  • diamond is an allotrope of carbon that forms 4 bonds per carbon atom
  • graphite is an allotrope of carbon that forms 3 bonds per carbon atom
  • silicon (IV) oxide is also a giant covalent compound
86
Q

describe the structure of solid metals

A
  • metal solids contain layers of thousands of ions in a sea of delocalised electrons which forms giant metallic lattices
87
Q

How is iodine, I2 and ice H2O bonded

A
  • iodine and water is a simple covalent molecular structure
88
Q

what are some of the structures formed from carbon atoms

A
  • graphite
  • diamond
  • graphene
89
Q

what is an allotrope

A
  • different forms of the same element, which have different properties
90
Q

describe the structure of diamond

A
  • it is a macromolecular structure
  • each carbon atom is bonded to 4 other atoms
  • forms a rigid tetrahedral structure
91
Q

describe the structure of graphite

A
  • it is a macromolecular structure
  • each carbon atom is bonded to 3 others
  • forms a flat hexagonal sheet
  • this means there is a delocalised electron per carbon which allows it to be able to conduct electricity
  • the intermolecular forces between the layers are weak allowing layer to easily slide over each other, so graphite can act as a lubricant
92
Q

describe the structure of graphene

A
  • consists of 2D sheets of graphite that are one atom thick.
  • light weight, strong and able to conduct electricity
93
Q

what are the physical properties of ionic substances

A
  • electrostatic forces of attraction
  • no intermolecular forces
  • giant lattice
  • can conduct electricity
  • high BP/MP
94
Q

what are the physical properties of giant covalent structures

A
  • covalently bonded
  • weak intermolecular forces
  • giant structure
  • mostly cannot conduct electricity
  • high MP/BP
95
Q

what are the physical properties of simple covalent structures

A
  • covalently bonded
  • weak or strong intermolecular forces can be both
  • simple structure
  • cannot conduct electricity
  • low MP/BP
96
Q

what are the physical properties of metallic structures

A
  • electrostatic forces of attraction
  • no intermolecular forces
  • giant metallic structure
  • can conduct electricity
  • high MP/BP

Chemistry (3 decks)

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