Brainscape's Knowledge GenomeTM

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


See full index

IB > Chem > Flashcards

Chem Flashcards

1
Q

How can random uncertainties be reduced?

A

Taking repeated readings and calculating the average value

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

How can systematic errors be reduced?

A

Subtracting the error or replacing faulty equipment

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

Random uncertainties lead to a loss of ___ and are reflected in ___

A
  • Precision
  • Percentage uncertainty
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Systematic errors lead to a loss of ___ and are reflected in ___

A
  • Accuracy
  • Percentage error
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Define precision and accuracy

A
  • Precision: How close repeated readings are to one another
  • Accuracy: How close readings are to true value
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

If calculations involve addition and subtraction, final answer should be quoted to same ___ as piece of original data with fewest ___

A

Decimal places

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

If calculations involve multiplication and division, final answer should be quoted to same ___ as piece of original data with fewest ___

A

Significant figures

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Define an element

A

Pure compound containing only 1 type of atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Define an atom

A

Smallest part of an element that can still be recognised at that element

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Define a compound

A

Pure substance formed when 2 or more elements combine chemically in a fixed ratio

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

Define a mixture

A

2 or more substances physically mixed together in an unfixed ratio

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Give the formula for number of molecules

A

Moles x 6.02 x 10^23

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

Give properties of a mixture

A
  • Retain individual properties
  • Can be separated by physical means
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Give the formula for number of atoms

A

Moles x No. of atoms x 6.02 x 10^23

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

Give the formula for mass of 1 molecule

A

Molar mass / 6.02 x 10^23

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

Give the formula for volume of gas (dm^3) at STP

A

Amt of gas (mol) x 22.7

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

Give the formula for the relationship between the concentration, volume and moles of two substances

A

C1V1/N1 = C2V2/N2

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

Give the formula for % composition

A

(no. of atoms x relative atomic mass)/relative molecular mass * 100

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

Define empirical and molecular formula

A
  • Empirical formula: Simplest ratio of different types of element in a compound
  • Molecular formula: Actual number of each elements present in a molecule of the compound
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Give the formula for % yield

A

Theoretical yield/Actual yield * 100

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Give two assumptions in an ideal gas

A
  • Volume of gas is negligible compared to volume of gas container
  • Negligible intermolecular forces of attraction between gas particles
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Give two conditions for achieving an ideal gas

A
  • Low pressure –> Gas particles are far apart, resulting in insignificant volume of gas compared to volume of gas container
  • High temperature –> Gas particles possess high kinetic energy to overcome intermolecular forces of attraction between them
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Give the ideal gas law

A

pV = nRT
- P = Pressure (Pa)
- V = Volume (m^3)
- n = Amt (mol)
- R = 8.31 J k^-1 mol^-1
- T = Kelvin

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

Give the combined gas law

A

P1V1/T1 = P2V2/T2
- Units for P and V: Any as long as they are consistent
- Units for T: Kelvin

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

Give the formula for concentration in ppm

A
  • (Mass of solute * 10^6)/Mass of solution
  • (Volume of gas * 10^6)/Volume of air
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Define hydrated and anhydrous salt

A
  • Hydrated: Salt with water
  • Anhydrous: Salt without water
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

Give the formula for total concentration of ions when ionic substances dissolve in water

A

Original concentration * No. of ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

Define mass number

A

Number of protons and neutrons in the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

Define atomic number

A
  • Number of protons in the nucleus and thus the number of electrons
  • Unique to an element
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

Define an isotope

A

Different atoms of the same element with the same number of protons but different number of neutrons in the nucleus
- Same chemical properties as chemical properties depend on the number and arrangement of electrons
- Different physical properties because their mass is different e.g. mass, speed of atoms

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

Define relative atomic mass

A

Weighted mean mass of all naturally-occurring atoms of an element, relative to 1/12 of the mass of a carbon-12 atom
- (% abundance * mass) + … + (% abundance * mass)/100

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

Why are isotopes radioactive

A
  • Nuclei break down spontaneously
  • When they break down, they emit radiation which is dangerous to living things
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

Energy levels/Shells ___ in energy and are ___ in distance as they get further from the nucleus.

A
  • Increase
  • Closer
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

Give the formula for the number of electrons in each energy level/shell

A

2n^2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

Give the order of the electromagnetic spectrum

A

RMIVUXG
- For Visible region: ROYGBIV

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

Longer wavelength, ___ frequency, ___ energy

A

Lower

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
37
Q

Explain the emission spectra

A
  • When energy is supplied to an atom, electrons are excited from their ground state to an excited state
  • (Electrons can only exist in certain fixed energy levels) –> Electrons will be unable to exist in the higher level and fall to a lower energy level
  • When electrons drop to a lower energy level they emit energy which corresponds to a particular wavelength and shows up as a line in the spectrum
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
38
Q

When electrons return to __, they show up as a line in which region?
- n = 1
- n = 2
- n = 3

A
  • n = 1: Ultraviolet region (largest energy change)
  • n = 2: Visible region
  • n = 3: Infrared region
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
39
Q

Why do lines in the electromagnetic spectrum converge?

A

Because energy levels themselves converge. Beyond the continuum, electrons can have any energy and are free from the influence of the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
40
Q

Why do transition metals show different coloured flames when burnt?

A
  • Heat causes electrons to move to a higher energy level
  • Electrons are unstable in the higher energy level and fall back down
  • When they fall from a higher to lower energy level, energy is emitted in the form of visible light energy (with the wavelength of the observed energy/colour)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
41
Q

Give the difference between a line spectrum and continuous spectrum

A
  • Line spectrum: Only has specific wavelengths
  • Continuous spectrum: Has all wavelengths
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
42
Q

Give the number of orbitals in the s, p, d, f subshells

A

s:1, p:3, d: 5, f:7

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
43
Q

Is the electronic configuration of the chromium ion (transition metal) [Ar]4s13d5 or [Ar]4s23d6

A

[Ar]4s13d5

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
44
Q

When transition metals form ions, the electrons are removed from the ___ orbital first.

A

S orbital

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
45
Q

Give the formula for effective nuclear charge

A

Number of protons - Shielding/inner electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
46
Q

Describe and explain the change in atomic radius across a period

A

Decreases because
- Electrons are added to the same valence shell
- But number of protons increase so effective nuclear charge increases
- Valence shells are more strongly attracted to the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
47
Q

Describe and explain the change in atomic radius down the group

A

Increases because
- Outermost electron is in a valence shell further from the nucleus
- Attraction of nucleus for electrons decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
48
Q

Describe the difference in radius between a parent atom and its cation

A

The radius of the cation will be smaller because
- The cation has less electrons but the same number of protons
- Valence electrons are more strongly attracted to the nucleus
- Number of electron shells also decrease

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
49
Q

Describe the difference in radius between a parent atom and its anion

A

The radius of the anion will be larger because
- The anion has more electrons but the same number of protons
- Valence electrons are less strongly attracted to the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
50
Q

Given that they have the same number of electrons, rank them from the smallest to largest radius and explain why: Noble gas atom, cation, anion

A

Cation < Noble gas < Anion
- Same number of electrons but an increasing number of protons
- Effective nuclear charge increases
- Valence electrons are more strongly attracted to the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
51
Q

Define first ionisation energy and give a chemical equation representing it

A
  • Energy required to remove an electron from an atom in a gaseous state (kJ mol^-1)
  • X (g) –> X+ (g) + e-
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
52
Q

Describe and explain the change in first ionisation energy across a period

A

Increases because
- Number of protons increase but number of inner electrons remain the same
- Effective nuclear charge increases
- Valence shell is more strongly attracted to the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
53
Q

Describe and explain the change in first ionisation energy down the group

A

Decreases because
- Electron to be removed is in a valence shell further from the nucleus
- Attraction of nucleus for electrons decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
54
Q

Explain why boron (1s2,2s2,2p1) has a lower ionisation energy than beryllium (1s2,2s2) even though boron has a higher effective nuclear charge

A
  • Electron to be removed from boron is in the 2p-sublevel while electron to be removed from beryllium is in the 2s-sublevel
  • The p-sublevel is higher in energy than the s-sublevel
  • Less energy will be required to remove an electron from boron
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
55
Q

Explain why oxygen (1s2,2s2,2p4) has a lower ionisation energy than nitrogen (1s2,2s2,2p3) even though oxygen has a higher effective nuclear charge

A
  • Oxygen has 2 electrons paired up in the same p-orbital but nitrogen does not
  • It will be easier to remove one of the paired electron from oxygen due to interelectronic repulsion
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
56
Q

Define electron affinity and give a chemical equation representing it

A
  • Energy change/enthalpy change when one electron is added to each atom in gaseous state
  • X (g) + e- –> X- (g)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
57
Q

Describe and explain the change in electron affinity across a period

A

Increases because
- Increase in effective nuclear charge and decrease in atomic radius
- Electron will be more strongly attracted when brought into the valence shell

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
58
Q

Describe and explain the change in electron affinity down Group 17

A

Generally decreases because
- Electron is brought into the valence shell
- As size of atom gets bigger, there is weaker attraction between added electron and nucleus as it is brought into a position further from the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
59
Q

Define electronegativity

A

Measure of the attraction an atom has for a shared pair of electrons when it is covalently bonded to another atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
60
Q

Describe and explain the change in electronegativity across a period

A

Increases because
- Number of protons increase but number of inner electrons remain the same
- Effective nuclear charge increases
- Stronger attraction of bonding electrons to nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
61
Q

Describe and explain the change in electronegativity down the group

A

Decreases because
- Increase in number of valence shells
- Bonding electrons are further from the nucleus
- Attraction of nucleus for bonding electrons decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
62
Q

Explain why metals have low electronegativity

A

Metals have few valence electrons and increase their stability by losing electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
63
Q

Describe and explain the change in melting and boiling points down Group 1

A

Decreases because
- Ions get larger
- Weaker electrostatic forces of attraction between positive ions in lattice and delocalised electrons
- Less energy required to overcome electrostatic forces of attraction

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
64
Q

Describe and explain the change in reactivity down Group 1

A

Reactivity increases/ Reactions become more vigorous as
- Size of atom increases
- Ionisation energy decreases
- Outer electrons are lost more easily

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
65
Q

Describe and give the equation for the reaction between a Group 1 element and oxygen

A
  • React vigorously with oxygen and tarnish rapidly in air
  • 4M (s) + O2 (g) –> 2M2O (s)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
66
Q

Give the equation for the reaction between a Group 1 element and water

A

2M (s) + H2O (l) –> 2MOH (aq) + H2 (g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
67
Q

Give some observations for the reaction between a Group 1 element and water

A
  • React rapidly
  • Produces strong base
  • Effervescence
  • Solution gets hot
  • Solution becomes basic
  • Metal hydroxide/hydrogen gas produced
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
68
Q

Describe the change in melting and boiling points down Group 17

A

Increases because
- Molecular mass increases
- London dispersion forces of attraction get stronger
- More energy required to overcome stronger intermolecular forces of attraction

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
69
Q

Describe the change in reactivity down Group 17

A

Decreases because electron affinity decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
70
Q

When a reactive halogen reacts with a halide ion of a less reactive halogen, what happens?

A

More reactive halogen displaces the halide ion

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
71
Q

Describe change in metallic properties across Period 3 oxides (sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon)

A
  • Sodium, Magnesium, Aluminium: Metal
  • Silicon: Metalloid
  • Phosphorus, Sulfur, Chlorine, Argon: Non-metal
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
72
Q

Describe the solubility of the Period 3 oxides (sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon)

A
  • Sodium, Magnesium: Soluble
  • Aluminium, Silicon: Insoluble
  • Phosphorus, Sulfur: Soluble
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
73
Q

Describe change in acid-base properties across Period 3 oxides (sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon)

A
  • Sodium, Magnesium: Basic
  • Aluminium: Amphoteric
  • Silicon, Phosphorus, Sulfur: Acidic
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
74
Q

Describe change in melting points across Period 3 (sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine, and argon)

A

Sodium - Aluminium: Increase
- Metallic bonding
- Increase in number of valence electrons and decrease in size

Silicon: High melting point
- Giant covalent structure

Phosphorus - Chlorine: S8 > P4 > Cl2
- P4, S8, Cl2
- Higher molecular mass
- Stronger intermolecular forces of attraction

Argon: Low melting point
- Weak intermolecular forces of attraction

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
75
Q

Define metallic bonding

A

Electrostatic forces of attraction between a lattice of positive metal ions and sea of delocalised electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
76
Q

Describe structure of metallic bonding

A

Giant metallic structure consisting of lattice of positive metal ions and sea of delocalised electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
77
Q

Describe physical properties of metallic bonding

A
  • Lustrous
  • Good conductor of electricity because of presence of delocalised electrons
  • Good conductors of heat
  • Ductile
  • Malleable
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
78
Q

Describe and explain the melting and boiling points of metallic bonding

A

High as large amount of energy is required to overcome strong electrostatic forces of attraction between positive ions and sea of delocalised electrons
- Increases with size of cation, charge of cation, number of delocalised electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
79
Q

Describe the solubility of metallic bonding

A

Insoluble in both polar and non-polar solvent

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
80
Q

Define alloys and explain why alloys are stronger and stiffer than pure metals

A
  • Homogenous mixtures of two or more metals, or a metal and non-metal
  • Introduction of a larger atom makes it harder for planes of atoms/ions to slide over one another
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
81
Q

Define ionic bonding

A
  • Electrostatic forces of attraction between oppositely charged ions
  • Formed when electrons are transferred from one atom to another to form ions with noble gas configuration
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
82
Q

Describe the structure of ionic bonding

A
  • Giant crystal lattice
  • Consists of oppositely charged ions held together by ionic bonds. Ions are arranged such that attraction between oppositely charged ions is maximum and repulsion between similarly charged ions is minimum
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
83
Q

Describe physical properties of ionic bonding

A
  • High melting and boiling point (Strong electrostatic forces of attraction)
  • Hard and brittle (Ionic bonds are strong and non-directional; Ordered arrangement of ions can be easily dislocated when sheer stress is applied)
  • Non-electrical conductor in solid state but good conductor when molten or dislocated
  • Soluble in polar solvent but insoluble in non-polar solvent
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
84
Q

Explain why ionic compounds are good conductors when molten or dissolved in aqueous solution but are non-conductors in solid state

A
  • In solid state, there is the absence of mobile charge carriers as ions are held in fixed positions by strong electrostatic forces of attraction in giant ionic lattice
  • In molten or aqueous state, there are mobile ions as lattice is broken down and ions are free to move
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
85
Q

Define covalent bonding

A
  • Electrostatic forces of attraction between the nuclei and the shared pair/s of electrons between them
  • Formed when atoms share one or more pairs of electrons so each atom can achieve a noble gas configuration
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
86
Q

Describe a simple molecular structure

A
  • Atoms are held by strong covalent bonds but molecules are held together by weak intermolecular forces of attraction
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
87
Q

Describe a giant molecular structure

A
  • Atoms are bonded together by strong covalent bonds to give a giant three dimensional structure
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
88
Q

Describe the melting and boiling points of covalent compounds

A
  • Simple molecular structure: Low as low amount of energy is required to overcome weak intermolecular forces
  • Giant molecular structure: High as large amount of energy is required to overcome strong intermolecular forces
89
Q

Describe the solubility of covalent compounds

A
  • Simple molecular structure: Polar molecules soluble in polar solvent, non-polar molecules soluble in non-polar solvent
  • Giant molecular structure: Insoluble in both polar and non-polar solvents
90
Q

Describe the electrical conductivity of covalent compounds

A

Simple molecular structure:
- Generally do not conduct electricity in any physical state
- Absence of mobile charge carriers as valence electrons are localised in covalent bonds

Giant molecular structure:
- Insoluble in both polar and non-polar solvents
- Exception: Graphite due to presence of delocalised electrons within layers

91
Q

Describe the structure and bonding of silicone dioxide (SiO2)

A
  • Tetrahedral
  • Each silicon atom is covalently bonded to 4 oxygen atoms and each oxygen atom is covalently bonded to 2 silicon atoms
92
Q

Describe the structure and bonding of diamond (C)

A
  • Tetrahedral
  • Each carbon atom is covalently bonded to 4 other carbon atoms
93
Q

Describe the structure and bonding of silicon (C)

A
  • Tetrahedral
  • Each carbon atom is covalently bonded to 4 other carbon atoms
94
Q

Describe the structure and bonding of graphite

A
  • Trigonal planar
  • Each carbon atom is bonded to 3 carbon atoms
  • Honeycomb structure within each layer
  • Continuous layers of hexagons
95
Q

Explain why graphite is soft and slippery

A

Two-dimensional layers are held by weak intermolecular forces of attraction and layers can slide over one another when force is applied

96
Q

Describe the physical properties of graphite

A
  • Soft and slippery
  • Conductor of electricity
97
Q

Explain why graphite is a good conductor of electricity

A

Each carbon bonds to 3 other atoms. One valence electron per carbon atom can delocalise within the layer. The delocalised electrons can act as charge carriers.

98
Q

Describe the bonding of graphene

A
  • Each carbon atom is covalently bonded to 3 other carbon atoms
99
Q

Describe the structure and bonding of buckminsterfullerene

A
  • Each carbon atom is bonded covalently to 3 others
  • Weak intermolecular forces of attraction between molecules that require little energy to overcome
  • Slippery
100
Q

Describe the direction of the dipole moment in a polar molecule

A

From less electronegative to more electronegative

101
Q

Explain why diatomic molecules are non-polar

A

Both atoms have the same electronegativity value. Both atoms will exert identical attraction thus the electron pair is shared equally.

102
Q

Give the bond angles from
- Linear
- Trigonal planar
- Bent
- Tetrahedral
- Trigonal pyramidal
- Bent

A
  • 180
  • 120
  • 117.5
  • 109.5
  • 107
  • 105
103
Q

Explain London dispersion forces

A
  • Electrons in an atom are in constant motion and at any one time will not be symmetrically distributed about nucleus
  • This results in a temporary instantaneous dipole in the atom which will induce an opposite dipole in a neighbouring atom
  • Dipoles attract each other so that there is an attractive force between atoms
104
Q

Explain dipole-dipole forces

A
  • Electrostatic forces of attraction between polar molecules
105
Q

Explain hydrogen bonding

A
  • Occurs when hydrogen is bonded directly to a small highly electronegative element (e.g. O, N, F)
  • As the electron pair is drawn away from hydrogen atom by the electronegative element, all that remains is the proton in the nucleus as there are no inner electrons
  • Proton attracts a non-bonding pair of electrons from O, N or F resulting in a much stronger dipole-dipole attraction
106
Q

Explain why overall force between atoms are always attractive if dipoles constantly disappear and reappear

A

A dipole always induce an opposite one

107
Q

In an exothermic reaction, heat is __ surroundings.

A

Released to

108
Q

In an endothermic reaction, heat is __ surroundings

A

Absorbed from

109
Q

In an exothermic reaction, bonds in product are __ than bonds in reactant and the product is more __ than reactants

A

Stronger
Stable

110
Q

Define enthalpy change

A

Amount of heat energy absorbed/evolved in a reaction

111
Q

Define standard enthalpy change of reaction. Give the formula to find enthalpy change of reaction if the enthalpy change of formation of products and reactants are known

A
  • Enthalpy change when molar amounts of reactants react together to form products under standard conditions
  • Enthalpy change of formation of products - Enthalpy change of formation of reactants
112
Q

Define standard enthalpy change of combustion

A

Enthalpy change when 1 mole of substance is completely burnt in oxygen under standard conditions

113
Q

Define standard enthalpy change of formation

A

Enthalpy change when 1 mole of compound is formed from its elements under standard conditions

114
Q

Define specific heat capacity (c)

A

Energy required to raise temperature of 1kg of substance by 1K

115
Q

Define heat capacity (C)

A

Energy required to raise temperature by 1K

116
Q

Give some sources of error in a calorimeter and explain how to minimise these errors

A
  • Heat loss to surroundings. Insulate reaction vessel or use a lid
  • Calorimeter could absorb some of the heat evolved. Determine the specific heat capacity of the calorimeter and add the heat absorbed by the calorimeter to heat absorbed by water to find total heat evolved by burning of fuel
  • Incomplete combustion (formation of soot and carbon monoxide)
117
Q

Give possible assumptions when calculating enthalpy change

A
  • Density of all aqueous solutions is assumed to be 1 g cm-1
  • Specific heat capacity is assumed to be the same as water
118
Q

Define bond enthalpy and give the formula for calculating enthalpy change from the bond enthalpy values of the product and reactant

A
  • Energy required to break one mole of a covalent bond in a gaseous molecule averaged over similar compounds
  • BE (bonds broken) - BE (bonds formed)
119
Q

Bond breaking is ___ (exothermic/endothermic).
Bond forming is ___ (exothermic/endothermic)

A
  • Bond breaking is endothermic
  • Bond forming is exothermic
120
Q

Give the limitations of using bond enthalpy values

A
  • Values are average values and are not specific to the bond broken for that reaction
  • Values only apply to bonds in gaseous state
121
Q

Define rate of reaction

A

Increase in concentration of one of the products per unit time or decrease in concentration of one of the reactants per unit time

122
Q

Define the collision theory

A

For a reaction to occur between two particles, particles must
- Collide with the appropriate geometry so the reactive parts of the particle come into contact with each other
- Must collide with sufficient energy to bring about the reaction

123
Q

Define activation energy

A

Minimum amount of energy required for a reaction to take place

124
Q

Describe the effect of increasing temperature on rate of reaction

A
  • Increase in rate
  • As temperature increases, kinetic energy of particles increase
  • Greater number of particles with energy greater than or equal to the activation energy
  • Frequency of effective collisions increase
125
Q

Describe the effect of increasing surface area on rate of reaction

A
  • Increase in rate
  • Larger exposed surface area for reaction
  • Higher frequency of effective collisions
126
Q

Describe the effect of increasing concentration on rate of reaction

A
  • Increase in rate
  • Number of reacting particles per unit volume increases
  • Frequency of effective collisions increases
127
Q

Describe the effect of increasing pressure on rate of reaction

A
  • Increase in rate
  • Reactant particles are closer
  • Frequency of effective collisions increases
128
Q

Describe the effect of a catalyst on a reaction

A
  • Provides an alternative pathway with lower activation energy
  • Greater proportion of reacting particles colliding with energy greater than or equal to activation energy
  • Frequency of effective collisions increases
  • Greater rate of reaction
129
Q

Define dynamic equilibrium

A

Equilibrium where
- The concentration of all the reactants and products are constant
- Rate of forward reaction is equal to the rate of the reverse reaction

130
Q

Define closed system

A

One where neither matter nor energy can be lost or gained

131
Q

If equilibrium constant is a large number, reaction will proceed to the ___

A

Right

132
Q

If the reaction quotient is larger than the equilibrium constant, reaction will proceed to the ___

A

Left

133
Q

Equilibrium constant can be affected by ___

A

Temperature
- If forward reaction is favoured, equilibrium constant increases

134
Q

Define Le Chatelier’s principle

A
  • Provided temperature remains constant, value of equilibrium constant must remain constant
  • If a system at equilibrium is subjected to a small change, equilibrium will shift to minimise the effect of the change
135
Q

Describe the effect of temperature on the position of equilibrium

A
  • Increase in temperature favours endothermic reactions
  • Decrease in temperature favours exothermic reactions
136
Q

Describe the effect of pressure on the position of equilibrium

A
  • Increase in pressure will favour the reaction pathway leading to less volume (to minimise pressure)
  • Reduction in pressure will favour the reaction pathway leading to greater volume
137
Q

Describe the effect of concentration on the position of equilibrium

A
  • Increasing the concentration of reactants will favour the forward reaction
138
Q

Describe the effect of a catalyst on the position of equilibrium

A
  • No effect on position of equilibrium as it lowers the activation energy for both reactions equally and increases the rate of both reactions equally
139
Q

Define an acid

A

Substance that produces hydrogen ions, H+ (aq), in aqueous solution

140
Q

Define a base

A

Substance that can neutralise an acid

141
Q

Define an alkali

A

Base that is soluble in water

142
Q

Define a Bronsted-Lowry acid

A

Substance that can donate a proton

143
Q

Define a Bronsted-Lowry base

A

Substance that can accept a proton

144
Q

Define a conjugate base

A

Species remaining after acid has lost a proton

145
Q

Define a conjugate acid

A

Species remaining after base has accepted a proton

146
Q

Define an amphoteric species

A

Substance that can donate or accept a proton

147
Q

Give the equation for a reaction between acid and base

A

Acid + Base –> Salt + Water

148
Q

Give the equation for a reaction between acid and reactive metal

A

Acid + Reactive Metal –> Salt + Hydrogen
* Reactive metal must be above copper

149
Q

Give the equation for a reaction between acid and carbonate

A

Acid + Carbonate –> Salt + Carbon dioxide + Water

150
Q

Give the equation for a reaction between acid and hydrogencarbonate

A

Acid + Hydrogencarbonate –> Salt + Carbon dioxide + Water

151
Q

Define pH

A
  • Measure of the concentration of H+ (aq) ions in a solution
  • pH - log10 [H+ (aq)]
152
Q

Define a strong acid and give examples

A
  • Acid that completely dissociates into its ions in aqueous solution
  • E.g. HCl, HNO3, H2SO4
153
Q

Define a strong base and give examples

A
  • Base that completely dissociates into its ions in aqueous solution
  • E.g. NaOH, KOH, Ba(OH)2
154
Q

Define a weak acid and give examples

A
  • Acid that partially dissociates into its ions in aqueous solution
  • Only 1 molecule in 100 dissociates
  • E.g. CH3COOH, CO2 in H2CO3
155
Q

Define a weak base and give examples

A
  • Base that only partially dissociates into its ions in aqueous solution
  • E.g. NH3, C2H5NH2
156
Q

Give possible solutions to differentiate between a strong and weak acid/base

A
  1. Measure pH using pH meter
  2. Measure conductivity using a conductivity meter. Strong acids and bases will give higher readings because they contain more ions in solution
  3. For acids: React with metals or carbonates. Strong acids will give a greater rate of reaction because the concentration of hydrogen ions is greater
157
Q

Define VSEPR theory

A

Pairs of electrons arrange themselves around the central atom so they are as far apart as possible in order to minimise repulsion and obtain maximum stability

158
Q

Give the order of repulsion for: Lone pair-lone pair, lone pair-bond pair, bond pair-bond pair

A

Lone pair-lone pair, lone pair-bond pair, bond pair-bond pair

159
Q

Sulfur dioxide is produced from

A
  • Volcanoes
  • Combustion of sulfur-containing fossil fuels
  • Smelting of sulfide ions
160
Q

Give the reactions of sulfur that produce acid deposition

A

Formation of sulfur dioxide: S (s) + O2 (g) –> SO2 (g)

Oxidation of sulfur dioxide: SO2 (g) + 1/2 O2 (g) –> SO3

Reaction between sulfur dioxide and water in air:
- SO2 (g) + H2O (l) –> H2SO3 (aq) Sulfurous acid
- SO3 (g) + H2O (l) –> H2SO4 (aq) Sulfuric acid

161
Q

Nitrogen oxides are produced from

A
  • Electrical storms
  • Internal combustion engine
  • Jet engine
162
Q

Give the reactions of nitrogen that produce acid deposition

A

Formation of nitrogen oxide: N2 (g) + O2 (g) –> 2NO (g)

Oxidation of nitrogen dioxide: 2NO (g) + O2 (g) –> 2NO2 (g)

Reaction between nitrogen dioxide and water: 2NO2 (g) + H2O (l) –> HNO3 (aq) + HNO2 (aq)

OR: Reaction between nitrogen dioxide and oxygen: 4NO2 (g) + O2 (g) + 2H2O (l) –> 4NO3 (aq)

163
Q

Define acid deposition

A

Process by which acidic particles, gases and precipitation leave the atmosphere

164
Q

Define acid rain

A

Rain with a pH less than 5.6

165
Q

Give some effects of acid deposition

A
  • Increased acidity in the soil can leech important nutrients such as Ca2+, Mg2+, K+
  • Can kill fish
  • Stone that contains calcium carbonate is eroded by acid rain and can dissolve away exposing a fresh surface to react with more acid
  • Can irritate mucous membranes and increase the risk of respiratory diseases such as asthma
  • Dissolve heavy metal compounds and release poisonous ions such as Cu2+ and Pb2+
166
Q

Give some solutions for acid deposition

A
  • Lower amounts of sulfur dioxides and nitrogen oxides by improving engine design
  • Switch to alternate forms of energy such as wind and solar power
  • Reduce amount of fuel by using public teansport
  • Liming of lakes: Adding calcium oxide or calcium hydroxide to neutralise acidity
167
Q

Define reducing agent

A

Species that reduces another while itself getting oxidised

168
Q

Give some examples of reducing agents

A
  • KI: Colourless I- –> Brown I2
  • Iron (II) Salts: Pale green Fe2+ –> Yellowish-brown Fe3+
169
Q

Define oxidising agent

A

Species that oxidises another while itself getting reduced

170
Q

Give some examples of oxidising agents

A
  • KMnO4: Purple MnO4- –> Colourless Mn2+
  • K2CrO7: Orange Cr2O7^2- –> Green Cr3+
171
Q

Explain the function of a salt bridge in a voltaic cell

A

Provides an electrical connection between half cells to complete the circuit and keep half-cells electrically neutral

172
Q

Explain the movement of ions in a salt bridge in a voltaic cell

A

Anions move towards anode and cathodes move towards cathode

173
Q

In a voltaic cell, the cathode is __ (negative/positive) and the anode is ___ (negative/positive)

A

Cathode is positive and the anode is negative

174
Q

In a voltaic cell, at the cathode/positive cell, ___ occurs and ___ are produced

A

Reduction occurs and pure elements are produced

175
Q

In a voltaic cell, at the anode/negative cell, __ occurs and ___ are produced

A

Oxidation occurs and cations are produced

176
Q

Define electrolysis

A

Breaking down of a substance in molten state by the passage of electricity through it

177
Q

Explain how electric current is produced in an electrolytic cell

A

Current is produced due to movement of electrons through wires and movement of ions in the electrolyte

178
Q

In an electrolytic cell, the cathode is ___ (negative/positive) and the anode is ___ (negative/positive)

A

Cathode is negative and anode is positive

179
Q

In an electrolytic cell, at the cathode/negative cell, ___ occurs and ___ are produced

A

Reduction occurs and pure elements are formed

180
Q

In an electrolytic cell, at the anode/positive cell, ___ occurs and __ are produced

A

Oxidation occurs and pure elements are formed

181
Q

Define a homologous series

A
  • Series of compounds with the same functional group
  • Each member has the same general formula
  • Neighbouring members differ by -CH2-
182
Q

Define a functional group

A

Atom/group of atoms in a molecule giving it its characteristic chemical properties

183
Q

Give evidence for the resonance structure of benzene

A
  • C-C bond lengths are all the same and have a value of 0.140nm which lies between the values for C-C and C=C
  • Enthalpy of hydrogenation of cyclohexene is -120kJ mol-1. If benzene simply had the cyclohexa-1,3,5-triene structure with three double bonds the enthalpy change of hydrogenation of benzene would be 3 times the enthalpy change of hydrogenation of cyclohexene. The difference is the extra energy associated with the delocalisation
  • Only one isomer exists for 1,2-disubstituted benzene compounds. If there were simply alternate bonds two isomers would exist
  • If benzene had 3 normal double bonds it would be expected to readily undergo addition reactions. But it only undergoes addition reactions with difficulty and most commonly undergoes substitution reactions
184
Q

Define structural isomers

A

Compounds with the same molecular formula but different structural formulas

185
Q

Describe boiling points down a homologous series

A
  • Down the homologous series, boiling points increase because carbon chain increases and London dispersion forces of attraction increase
  • Increases steeply at first but as successive CH2 groups are added, the rate of increase decreases
186
Q

Describe change in melting and boiling point when branching occurs

A
  • Molecules become more spherical in shape and contact surface area between molecules decreases
  • Melting point increases but boiling point decreases
187
Q

What are the intermolecular forces of attraction in alkanes?

A

Only London dispersion forces of attraction

188
Q

Explain the volatility of alkanes

A

Only have London dispersion forces of attraction

189
Q

What are the conditions, reagents and products for a substitution reaction?

A

Conditions: Presence of sunlight or ultraviolet light to provide extra energy required to overcome activation energy

Reagents: Alkane, Halogen

Products: Hydrogen halide, substituted alkane

190
Q

Give equations for the steps in a substitution reaction between methane and chlorine

A

Initiation:
Covalent bond breaks and each of the two atoms forming the bond retains one of the shared electrons, resulting in the formation of two free radicals
Cl2 –> 2Cl radical

Propagation:
CH4 + Cl radical –> HCl + CH3 radical
CH3 radical + Cl2 –> CH3Cl + Cl radical

Termination:
Cl radical + Cl radical –> Cl2
CH3 radical + Cl radical –> CH3Cl
CH3 radical + CH3 radical –> C2H6

191
Q

Arrange in order of bond strength and bond length: Single bond, Double bond, Triple Bond

A

Single bond: Weakest strength, longest bond length
Double bond: Medium, medium bond length
Triple bond: Strongest, shortest bond length

Stronger attraction between nuclei and bonding electrons

192
Q

Define an unsaturated compound

A

Contains carbon double or triple bonds

193
Q

Describe what happens when an alkane and an alkene react with bromine solution

A
  • Alkane and bromine solution: No reaction
  • Alkene and bromine solution: Bromine solution changes colour from orange to colourless
194
Q

Give the conditions for a reaction between an alkene and halogen

A

Conditions: Room temperature

195
Q

Give the conditions and products for a reaction between an alkene and hydrogen

A

Conditions: Heat, finely divided nickel (catalyst)
Products: Alkanes

196
Q

Give the conditions and products for a reaction between an alkene and halogen halides

A

Conditions: Heat

197
Q

Give the conditions and products for a reaction between an alkene and water

A

Conditions: Heat, concentrated sulfuric acid/phosphoric acid (catalyst)
Products: Alcohols

198
Q

Explain the process by which primary alcohols are oxidised

A
  • Oxidised with acidified K2Cr2O7
  • First to aldehydes and then to carboxylic acids
  • To stop the reaction at the aldehyde stage, the ethanal is distilled from the reaction mixture as soon as it is formed
  • If complete oxidation to carboxylic acid is required the mixture can be heated under reflux so none of the ethanal can escape
199
Q

How can primary alcohols be oxidised to aldehydes?

A

To stop the reaction at the aldehyde stage, the ethanal is distilled from the reaction mixture as soon as it is formed

200
Q

How can primary alcohols be oxidised fully to carboxylic acids?

A

If complete oxidation to carboxylic acid is required the mixture can be heated under reflux so none of the ethanal can escape

201
Q

Explain the process by which secondary alcohols are oxidised

A

Oxidised to ketones and cannot undergo further oxidation

202
Q

Explain the process by which tertiary alcohols are oxidised

A

Cannot be oxidised as they have no hydrogen atoms attached directly to the carbon atom containing the -OH- group

203
Q

What reactions do halogenoalkanes undergo?

A

Nucleophilic substitution

204
Q

Give the products for the nucleophilic substitution reaction between ammonia and C2H5Br

A

NH3 + C2H5Br → C2H5NH2 + HBr

205
Q

What reactions does benzene undergo?

A

Electrophilic substitution
- High electron density so it can undergo substitution reactions with electrophiles

206
Q

Explain why benzene does not readily undergo addition reactions

A

Delocalisation of electrons in benzene ring give it extra stability

207
Q

Give the products for the electrophilic substitution reaction between C6H6 and Cl2

A

C6H6 + Cl2 –> C6H6Cl + HCl

208
Q

Give the reaction name, conditions and products for a reaction between an alcohol and carboxylic acid

A

Reaction: Condensation

Conditions: Presence of small amount of concentrated sulfuric acid catalyst

Products: Ester, water

209
Q

Give the IHD values for:
- A double bond
- A triple bond
- A ring

A
  • Double bond and ring: 1 IHD
  • Trilpe bond: 2 IHD
210
Q

Explain how infrared spectroscopy works

A
  • When molecules absorb energy in the infrared region of the electromagnetic spectrum they vibrate
  • When infrared radiation is passed through a sample the spectrum shows the characteristic absorption
  • Confirms that a particular bond is present
211
Q

Infrared spectroscopy only works for ___ compounds

A

Polar covalent compounds

212
Q

What region of the electromagnetic spectrum is used in 1H NMR

A

Radio waves

213
Q

How can 1H NMR help to identify compounds?

A
  • Provides information on the chemical environment of all the hydrogen atoms in a molecule
  • Peaks/signals indicates number of chemical environments
214
Q

Determine the bond angle in benzene

A

109.5

215
Q

The specific heat capacity of X is twice that of Y. If the same amount of heat is supplied to X and Y, the temperature change of Y will be ___ of X.

A

Four times that

216
Q

In a voltaic cell, negative ions flow through the salt bridge from the __ to the ___.

A

Cathode/positive half-cell to the anode/negative half-cell

217
Q

Differentiate between a homogenous and heterogenous mixture

A
  • Homogenous: Uniform composition and can be separated by filtration
  • Heterogenous: Non-uniform composition and can be separated by chromatography, distillation, crystallisation, etc.
218
Q

In a voltaic cell, electrons flow through the wires from the ___ to the ___.

A

Anode to the cathode

219
Q

What is the difference between an amphiprotic and amphoteric substance?

A

Amphiprotic: Can donate or accept a proton
Amphoteric: Can act as an acid or a base

IB (2 decks)

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