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Chemistry A-Level > Bonding and Structure > Flashcards

Bonding and Structure Flashcards

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1
Q

What is a pure substance?

A

Pure substances are made up of only one kind of particle - made up of only 1 element /compound

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2
Q

what are semi-metals/metalloids?

A

They have properties of both metals and non-metals.

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3
Q

How are ionic compounds formed?

A

There is a transfer of electrons from metal atoms which lose electrons from their outer energy levels to non-metal atoms which gain electrons to their outer energy levels.
The metal atoms lose electrons and non-metal atoms gain electrons which gives them noble gas electron configuration.

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4
Q

what is particular about d block elements and the transfer of electrons

A

d block element atoms do not always achieve noble gas electron configuration - the electrons are transferred to 4s subshell (in fourth energy level) before the 3d (in the third energy level)

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5
Q

How are ions formed

A

when atoms lose / gain electrons they are no longer electrically neutral and have acquired an overall charge

Metal atoms - lose outer electrons and become positively charged

Non-metal atoms - gain electrons to fill their outer energy level and become negatively charged

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6
Q

What is a positive ion called

A

Cation

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7
Q

What is a negative ion called

A

Anion

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8
Q

What is the ionic bond

A

The electrostatic attraction btwn oppostiely charged ions

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9
Q

Remember to find out the formula of ionic compounds by adjusting number of atoms in compound according to their relative charges

e.g. MgCl2

A
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10
Q

What is an ionic lattice

A

Ions are held in lattice - a regular, repeated 3 dimensional arrangement in a metal or other crystalline solid
e.g 6:6 configuration

Ionic solid formed has many strong electrostatic forces of attractions between oppositely charged ions

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11
Q

What makes ionic compounds crystalline

A

The regular pattern of ions within the structure.

  • decrepitation: when NaCl heated it makes a cracking sound caused by ionic crystalline structure breaking up
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12
Q

What are the properties of Ionic Compounds

A
  • High melting/boiling points - solid at room temperature
  • Usually soluble in water
  • Conduct electricity when molten/in aqueous solution
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13
Q

Why ionic compounds have high melting/boiling points - solid at room temperature

A

There is a large number of strong electrostatic attractions btwn pos and neg ions which means that lots of energy is required to melt an ionic solid

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14
Q

What are factors which affect the melting/boiling point of ionic compounds

A

Smaller the ions and the Higher the Charge on the ion = the stronger the ionic bond
- ionic radius decide the Size of the ion

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15
Q

Which are generally smaller positive or negative ions

A

Positive ions are smaller:
- the metal atoms lose electrons from outer energy level so one less energy level occupied.
- The effective nuclear charge increases so electrons are pulled close to the nucleus (higher ration of protons to electrons)

Negative ions are larger than the parent atom:

  • the repulsion btwn the electrons moves them farther apart from each other
  • the effective nuclear charge decrease as there are more electrons with the same number of protons
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16
Q

Why are ionic compounds usually soluble in water

A

Water is a polar molecule and surrounds charged ions.
When moving water molecules hit ionic lattice they knock ions off - then water molecules surround the ions.

If ELECTROSTATIC ATTRACTION btwn positive and negative ions is so strong that water can’t break up lattice - then compound insoluble in water.

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17
Q

Why do ionic compounds conduct electricity when molten/in aquoeus solution

A

Ionic compounds CANNOT conduct electricity in a solid state
In liquid/aqueous state, ions are free to move around + carry charge

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17
Q

What are the molecular ions and their charges

A

SO4 2-

NO3 -

OH -

CO3 2-

HCO3 -

NH4 +

Compounds containing molecular ions are Ionic and have same properties as ionic compounds

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18
Q

Define Covalent Bond

A

Shared pair of electrons between two atoms

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19
Q

Where are covalent bonds found

A
  • molecular elements and compounds - Cl2, H2O
  • macromolecular covalent elements/ compounds - C(graphite/diamond), SiO2
  • molecular ions
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20
Q

single bond
double bond
triple bond

A

single PAIR of shared electrons
two PAIRS of shared electrons
three PAIRS of shared electrons

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21
Q

Normally electrons shared will have existed as an unpaired electron in an orbital

A
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22
Q

How do atoms form covalent bonds

A
  • Atoms use unpaired electrons in orbitals in order to form covalent bonds
  • unpaired e- in orbital of one atom shared w unpaired e-in orbital of another atom
  • Atoms promote electrons into unoccupied orbitals in same energy level to form more covalent bonds but may not - so variety of compounds can be formed - PCl3, PCl5
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23
Q

Bonding pair

A

Pair of electrons shared btwn two atoms

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24
Lone pair
unshared(non-bonding)pair of electrons
25
Structure of metals
- generally solids and have particles packed closely together - atoms packed in layers and outer shell electrons are delocalised and can move between the layers - movement of delocalised electrons explains why metals conduct electricity and heat - atoms in layers are without outer shell electrons = ions - when discussing bonding refer to metal particles as positive ions when discussing structure and reactivity refer as metal atoms Lattice of positive metal ions in a sea of delocalised electrons
26
what is a metallic bond
The electrostatic attraction btwn delocalised electrons and positive metal ions in the lattice
27
Properties of Metals
- Metals conduct electricity - Metals conduct heat - Metals are ductile and malleable - Metals have high densities - Metals have high melting points
28
Why do metals conduct electricity?
Delocalised electrons in the metal structure can move. The delocalised electrons carry the charge and their movement causes an electric current to flow.
29
Why do metals conduct heat?
Heat is conducted when particles can move and are close enough to pass on heat energy from one to another. The delocalised electrons enable heat energy to be passed through the metal.
30
Why are metals ductile/malleable?
ductile = can be drawn into wires malleable = can be hammered into shape malleable/ductile due to the layered structure of the lattice bc the layers can slide over each other WITHOUT disrupting bonding layers are still held together by delocalised electrons - the strong attraction btwn the positive ions and delocalised electrons holds the structure together force applied causes layers to move
31
why do metals have high densities?
Positive ions are packed tightly together = high density
32
why do most metals have high melting points?
Metals have large regular structures strong forces of attraction btwn the positive ions and delocalised electrons. Overcoming these strong electrostatic attractions requires a large amount of heat energy.
33
Why do transition metals have much higher melting points than main group metals?
There are a large number of d sub-shell electrons which can be delocalised to create a stronger metallic bond.
34
What are the two main types of covalent substances?
- Molecular/simple covalent - Macromolecular/giant covalent
35
Molecular covalent crystals
molecular covalent substances which exist in a crystalline form solid (low melting point) molecular covalent substances form crystalline structures
36
Iodine structure
Large iodine molecules pack together into regular arrangement causing crystalline form of iodine small units pack together in regular lattice -- solid iodine has a crystalline nature
37
Ice structure
Molecules of water arranged in regular arrangements and form crystalline structure.
38
Properties of molecular covalent crystals
- Low Melting Points - Brittle bc they don't have the strong bonds holding together like other crystalline substances(held together by relatively weak intermolecular forces) - Do not conduct electricity - bc no charged particles to carry charge
39
What is an allotrope?
Different forms of the same element in the same physical state.
40
Physical properties of macromolecular subtances mostly depend on the MANY strong covalent bonds within the structure
41
Diamond structure
- Very hard - many strong covalent bonds and a rigid 3D structure - High Melting Point - many strong covalent bonds - this req a lot of energy to break Each carbon atom is strongly bonded to 4 other carbon atoms in a tetrahedral arrangement - the bond angle is 109.5 - Does not conduct heat or electricity - no charged particles which can move
42
Graphite structure
- Conducts electricity - delocalised electrons btwn layer which move and carry charge - Very high melting point - many strong covalent bonds which req a lot of energy to break - Layered structure w weak forces of attraction btwn layers so the layers can slide over each other = soft and used as lubricant/pencil lead Each carbon atom is covalently bonded to three other carbon atoms in a hexagonal arrangement - the bond angle = 120 The fourth electron is delocalised between the layers and provides weak forces of attraction btwn the layers - Can conduct electricity bc delocalised electrons can move and carry charge
43
As substance changes from one state to another there is an energy change
44
Melting
Endothermic Energy taken in to overcome forces/bonds
45
Freezing
Exothermic temp at which substance freezes = melting point Energy released when substance freezes as bonds/forces are formed
46
Boiling
Endothermic Energy taken in when substance boils to overcome attraction btwn particles/molecules
47
Condensing
Exothermic condensing point = boiling point Energy released - bonds/forces formed
48
Subliming
solid to gas/gas to solid Solid iodine and Solid carbon sublime.
49
Shape of molecule/ion determined from:
- Total num of electron pairs around central atom - Number of bonding pairs - Number of lone pairs Molecule/Ion takes up a shape that MINIMISES REPULSIONS/ MAXIMISES SEPARATION. The shape depends on the arrangement of atoms around the central atom.
50
Lone pairs are held closer to central atom so have greater repulsive effect on other pairs of electrons
51
Order of strength of repulsions of electron pairs
LP <-> LP > LP <-> BP > BP <-> BP
52
2 bonding pairs of electrons
Bond Angle = 180 Shape = Linear Explanation: 2 bp repel each other equally - molecule takes up this shape to minimise repulsions
53
3 bonding pairs of electrons
Bond angle = 120 Shape = Trigonal Planar Explanation: 3 bp repel each other equally - molecule takes up this shape to minimise repulsions
54
4 bonding pairs of electrons
Bond angle = 109.5 Shape: Tetrahedral Explanation: 4 bp repel each other equally - molecule takes up this shape to minimise repulsions
55
5 bonding pairs of electrons
Bond angle = 90 and 120 Shape = Trigonal Bipyramidal Explanation: 5 bp repel each other equally - molecule takes up this shape to minimise repulsions
56
6 bonding pairs of electrons
Bond angle = 90 Shape = Octahedral/Square bipyramidal Explanation: 6 bp repel each other equally - molecule takes up this shape to minimise repulsions
57
1 lone pair, 3 bonding pairs
. . H - N ---H ^ H Bond angle = 107 Shape = Pyramidal Explanation: Lone pair of electrons has greater repulsion than bonding pairs of electrons - molecule takes up this shapes to minimise repulsions
58
2 lone pairs and 2 bonding pairs
.. .. H - O - H Bond angle = 104.5 Shape = Bent Explanation: Lone pair of electrons has greater repulsion than bonding pairs of electrons - molecule takes up this shapes to minimise repulsions
59
When a coordinate bond forms it converts a lone pair of electrons into a bonding pair of electrons
coordinate bond = bonding pair of electrons formation of a coordinate bond causes change in shape
60
3 bonding pairs and 2 lone pairs
Bond angle = 86 Shape = T-shaped Explanation: Basic shape is trigonal bipyramidal but the lone pairs have greater repulsion than bonding pairs - the molecule takes up this shape to minimise repulsions
61
4 bonding pairs and 2 lone pairs
Bond angle = 90 Shape = Square pyramidal Explanation: Lone pair of electrons has greater repulsion than bonding pairs of electrons - molecule takes up this shapes to minimise repulsions
62
2 bonding pairs and 3 lone pairs
Bond angle = 180 Shape = Linear Explanation: basic shape is trigonal pyramidal but lone pairs of electrons have greater repulsion than bonding pairs - the molecule takes up this shape to minimise repulsions
63
Electronegativity
A measure of how strongly an atom attracts electrons to itself when in a covalent bond
64
relationship btwn electronegativity and polarity of bond
the greater the electronegativity difference between two atoms in a covalent bond, the more polar the bond will be. electrons found closer to one atom than other = more electronegative
65
Which is the most electronegative element?
Fluorine(4)
66
Factors affecting electronegativity
- Distance of bonding electrons from the attractive power of the nucleus(same as atomic radius) - across period atomic radius decreases - Size of nuclear charge(atomic number) Shielding of attractive power of the nucleus by inner electrons - going down the group = more inner energy levels which shield the attractive power of the nucleus
67
Trends in electronegativity
- Electronegativity increases across a period BECAUSE: - atomic radius decreases across a period so stronger attraction btwn the positive nucleus and electrons in the covalent bond - Nuclear charge increases across a period = greater attraction btwn electrons in covalent bond and positively charged nucleus of more electronegative atom - Electronegativity decreases down a group BECAUSE: - Atomic radius increases down a group - weaker attraction btwn pos nucleus and electrons in covalent bond - Shielding of nuclear charge/attractive power of nucleus increases down groupbc there are more electrons in inner energy levels Change in electronegativity seen more across a period than down a group as down a group the nuclear charge will still be increasing
68
Delta -
higher electronegativity value
69
Delta +
lower electronegativity value
70
No difference/very small difference in electronegativity
Non polar molecule
71
Small difference in electronegativity
Polar molecule
72
Large difference in electronegativity
Ionic compound
73
Intermolecular forces
Attractive forces btwn covalent molecules.
74
What are the types of intermolecular forces
- induced dipole-dipole forces(van der waals) - permanent dipole-dipole forces - hydrogen bonds
75
Van der waals
Larger a molecule is, has a higher Mr, has more electrons. The more electrons there are the greater the induced dipoles so greater van der waals forces. In solids van der waals forces more imp than in gases bc molecules closer together and have stronger van der waals forces btwn molecules.
76
Permanent dipole-dipole forces
between molecules which have permanent dipole within the molecule permanent dipole cause molecules to be attracted to each other Delta - <---> Delta + Permanent dipole-dipole forces responsible for holding together molecules in polar substances - van der waals forces also present btwn these molecules
77
Hydrogen bonds
Between delta+ H atom (covalently bonded to O,N,F atom) and the LONE PAIR OF ELECTIONS ON AN O,N OR F atom of another molecule. Much weaker than covalent bond. H-O, H-N, H-F bonds most polar, H-F only found in hydrogen fluoride. hydrogen bonds can break and reform - water is fluid
78
strength of intermolecular forces
van der waals < permanent dipole-dipole < hydrogen bonds
79
Bonding
A force of attraction btwn two speicies
80
Ionic bond
The force of attraction btwn positive ions and negative ions
81
Metallic bond
The force of attraction btwn positive cations and a sea of mobile, delocalised valence electrons
82
Covalent bond
The force of attraction resulting from 2 atoms sharing electrons
83
Molecules
Groups of atoms held together by covalent bonds
84
Coordinate/Dative covalent bond
Force of attraction resulting from 2 atoms sharing a pair of electrons where both of the electrons come from the same atom
85
Van der waals forces def
caused by random fluctuations of the electron cloud - causing temporary dipoles
86

Chemistry A-Level (2 decks)

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