Download Topics 1, 2, 3, 4, 5, 10, 11, 12, 13, 14, 15.1-2

IB Chemistry HL
Define and Distinguish Assessment Statements
for IB Topics Covered in Grade 11
(Topics 1, 2, 3, 4, 5, 10, 11, 12, 13, 14, 15.1-2, and 20)
1
Quantitative Chemistry
Topic 1.2.1
Define relative atomic
mass (Ar).
2
Quantitative Chemistry
Topic 1.2.1
The average mass of an atom, taking into
account the relative abundances of all
naturally occurring isotopes, relative to
one atom of carbon-12.
Note: There are no units for relative atomic mass.
3
Quantitative Chemistry
Topic 1.2.1
Define relative molecular
mass (Mr).
4
Quantitative Chemistry
Topic 1.2.1
The average mass of a molecule,
calculated by adding the relative atomic
masses of its atoms.
Note: There are no units for relative molecular mass.
Example: Mr for water = 2(1.01) + 16.00 = 18.02
5
Quantitative Chemistry
Topic 1.2.2
Define molar mass (M).
6
Quantitative Chemistry
Topic 1.2.2
The mass of one mole of a substance
expressed in grams per mole.
Example: M of water = 2(1.01) + 16.00 = 18.02 g mol-1
7
Quantitative Chemistry
Topic 1.2.4
Distinguish between
empirical formula and
molecular formula.
8
Quantitative Chemistry
Topic 1.2.4
The empirical formula of a compound is
the simplest whole number ratio of the
atoms it contains.
The molecular formula of a compound
gives the actual number of atoms of each
element.
example: butane
molecular formula = C4H10
empirical formula = C2H5
9
Quantitative Chemistry
Topic 1.5.1
Distinguish between solute,
solvent and solution.
10
Quantitative Chemistry
Topic 1.5.1
The solute is the substance dissolved
when a solution is formed. It is present
in the smallest quantity.
The solvent is the liquid that dissolves
another substance. It is present in the
greatest quantity.
A solution is a homogenous mixture of a
two or more substances.
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Quantitative Chemistry
Topic 1.5.1
Define concentration.
12
Quantitative Chemistry
Topic 1.5.1
Concentration is the amount of solute in
a known volume of solution.
We usually use molar concentrations.
The unit used is mol dm-3 .
Volume conversions:
1 dm3 = 1L
1 cm3 = 1 mL
1 dm3 = 1000 cm3 = 1000 mL
13
Atomic Structure
Topic 2.1.3
Define atomic number (Z).
14
Atomic Structure
Topic 2.3.1
Atomic number is the number of
protons in the nucleus of an atom.
The atomic number can also be used to find the
number of electrons in a neutral atom only.
atomic number >
15
Atomic Structure
Topic 2.1.3
Define mass number (A).
16
Atomic Structure
Topic 2.1.3
Mass number is the total number of
protons and neutrons
in the nucleus of an atom.
mass number >
lithium-7
number of neutrons in an atom = mass number - atomic number
17
Atomic Structure
Topic 2.1.3
Define isotopes.
18
Atomic Structure
Topic 2.1.3
Isotopes are atoms of the same element
but with different numbers of neutrons.
Isotopes are atoms of with the same atomic
number but different mass numbers.
Examples:
35Cl, 37Cl
uranium-235, uranium-238, uranium-239
19
Atomic Structure
Topic 2.3.2
Distinguish between a
continuous spectrum and
a line spectrum.
20
Atomic Structure
Topic 2.3.2
A continuous spectrum shows an
unbroken sequence of wavelengths.
A discontinuous spectrum shows only certain
wavelengths of light which show as discrete
lines or bands.
(This is also known as a line or an emission spectrum.)
21
Periodicity
Topic 3.1.2
Distinguish between the
terms group and period.
22
Periodicity
Topic 3.1.2
A group is a vertical column in the
Periodic Table.
A period is a horizontal row of elements in
the Periodic Table.
In a group, atoms have the same number of valence electrons.
In a period, atoms have the same number of energy levels.
23
Periodicity
Topic 3.2.1
Define first ionization energy.
24
Periodicity
Topic 3.1.2
First ionization energy is the energy
required to remove one mole of
electrons from one mole of gaseous
atoms.
This is also the enthalpy change for the reaction
X(g)
X+(g) + e-
25
Periodicity
Topic 3.2.1
Define electronegativity.
26
Periodicity
Topic 3.1.2
Electronegativity is a measure of the
tendency of an atom in a molecule to
attract a shared pair of electrons
towards itself.
Electronegativity is measured on a scale of 0 to 4.
27
Periodicity
Topic 13.2.4
Define ligand.
28
Periodicity
Topic 13.2.4
A ligand is an ion or molecule that
donates a pair of electrons to a metal
atom or ion in the formation of a
coordination complex.
Common ligands include H2O, NH3,
OH- ions, CN- ions and Cl- ions.
Dative covalent bonds form when the ligand bonds to
the metal atom or ion.
[ CuCl4 ]229
Energetics
Topic 5.1.1
Define exothermic reaction.
30
Energetics
Topic 5.1.1
An exothermic reaction releases heat to
the surroundings.
The temperature increases in an exothermic reaction.
The products have stronger bonds than the reactants.
Exothermic reactions have a -∆H value.
The enthalpy of the products is lower than the reactants.
The stability of the products is greater than the reactants.
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Energetics
Topic 5.1.1
Define endothermic reaction.
32
Energetics
Topic 5.1.1
An endothermic reaction absorbs heat
from the surroundings.
The temperature decreases in an endothermic reaction.
The products have weaker bonds than the reactants.
Endothermic reactions have a +∆H value.
The enthalpy of the products is higher than the reactants.
The stability of the products is less than the reactants.
33
Energetics
Topic 5.1.1
Define standard enthalpy
change of reaction (∆H˚).
34
Energetics
Topic 5.1.1
Standard enthalpy change is the heat
transferred during a reaction carried out
under standard conditions.
Standard conditions for thermodynamics:
temperature = 25˚C = 298 K
pressure = 1 atm = 101.3 kPa = 1.01 x 105 Pa
concentration of solutions = 1.0 mol dm-3
35
Energetics
Topic 5.4.1
Define average bond
enthalpy.
36
Energetics
Topic 5.4.1
Average bond enthalpy is the energy
required to break one mole of a covalent
bond, where all reactants and products
are in the gaseous state. It is the average
energy used to break that bond in
different molecules.
The greater the average bond enthalpy, the stronger the bond.
Bond enthalpies are used to find the enthalpy change of a reaction:
∆H = ∑ BE reactants - ∑ BE products
37
Energetics
Topic 15.1
Define standard state.
38
Energetics
Topic 15.1
The standard state of an element or
compound is its most stable form
under the standard conditions
of 298 K and 101.3 kPa.
The standard state of most metals is solid (except mercury).
The standard state of most elemental gases is diatomic (except noble gases).
The standard state of carbon is a solid, as graphite.
39
Energetics
Topic 15.1
Define standard
enthalpy of formation.
40
Energetics
Topic 15.1
Standard enthalpy of formation is the
enthalpy change when elements in their
standard states combine to form one
mole of a compound.
Symbol: ∆Hºf
Units: kJ mol-1
Sample thermochemical equations for formation reactions:
Na(s) + 1/2 Cl2(g)
NaCl(s)
∆Hºf = -411.2 kJ mol-1
N2(g) + 3/2 H2(g)
NH3(g)
∆Hºf = -45.9 kJ mol-1
2 C(graphite) + 3 H2(g) + 1/2 O2(g)
C2H5OH(l)
∆Hºf = -235.5 kJ mol-1
1/2 N2(g) + O2(g)
NO2(g)
∆Hºf = +33.2 kJ mol-1
NOTE: The ∆Hºf for any element in its standard state is ZERO.
41
Energetics
Topic 15.1
Define standard
enthalpy of combustion.
42
Energetics
Topic 15.1
Standard enthalpy of combustion is the
enthalpy change when one mole of a
compound in its standard state undergoes
complete combustion in excess oxygen
under standard conditions.
Symbol: ∆Hºc
Units: kJ mol-1
Sample thermochemical equations for standard combustion reactions:
CH4(g) + 2 O2(g)
CO2(g) + 2 H2O(l)
∆Hºc = -890 kJ mol-1
C6H12O6(s) + 6 O2(g)
CO2(g) + 6 H2O(l)
∆Hºc = -2830 kJ mol-1
C(graphite) + O2(g)
CO2(g)
∆Hºc = -394 kJ mol-1
H2(g) + 1/2 O2(g)
H2O(l)
∆Hºc = -286 kJ mol-1
NOTE: The ∆Hºc for oxygen gas (O2) is defined as ZERO.
43
Energetics
Topic 15.2
Define lattice enthalpy.
44
Energetics
Topic 15.2
Lattice enthalpy is the enthalpy change that
occurs when 1 mole of an ionic compound
is separated into its gaseous ions.
NaCl(s)
CaF2(s)
Na+(g) + Cl-(g)
Ca2+(g) + 2 F-(g)
∆H = +790 kJ mol-1
∆H = +2651 kJ mol-1
NOTE: Lattice enthalpy MAY be defined as the enthalpy change when 1 mole of an ionic
compound forms from its gaseous ions. In this the enthalpy change is exothermic.
45
Energetics
Topic 15.2
Define electron affinity.
46
Energetics
Topic 15.2
Electron affinity is the energy change that
occurs when one mole of gaseous atoms
gains a mole of electrons to form a
negative ion.
Cl(g) + 1e-
Cl-(g)
∆Hº = -349 kJ mol-1
47
Organic Chemistry
Topic 10.1.3
Distinguish between empirical,
molecular and structural
formulas.
48
Organic Chemistry
Topic 10.1.3
An empirical formula shows the lowest whole number ratio of
the atoms in a compound.
A molecular formula shows the actual numbers of atoms of
each element in a compound.
A structural formula unambiguously shows the arrangement of
atoms in a compound.
Example: butane
EF
C2H5
MF
C4H10
SF
49
Organic Chemistry
Topic 10.3.4
Distinguish between alkanes
and alkenes using bromine
water.
50
Organic Chemistry
Topic 10.3.4
Alkanes do not react with bromine water, and
the colour of the bromine remains
reddish-brown.
Alkenes do react with bromine water, and the
colour of the bromine changes from reddish
brown to colourless .
51
Measurement and Uncertainty
Topic 11.1
Distinguish between precision
and accuracy.
52
Measurement and Uncertainty
Topic 11.1
Accuracy refers to how close an experimental
value is to the actual value for a quantity. (The
mean value of the experimental values is used.)
Precision refers to the how close together
multiple measurements for a value are.
53