ib_chemistry_definitions[1]

IB Chemistry Definition Objectives (For Paper 1)
Topic 1: Quantitative Chemistry (SL)
1.2.1: Define the terms relative atomic mass (A r) and relative molecular mass (M r).
o relative atomic mass (A
r
): the weighted average of atomic masses of all naturally occurring
isotopes of an element.
o Ex: A
r of hydrogen = 1.01
o Ex: A
r of oxygen = 16.00
o Ex: A
r
of copper = 63.55
o relative atomic mass is the same thing as atomic weight and average atomic mass.
o relative molecular mass (M
r ): the sum of the relative atomic masses of a given molecule.
o Ex: M
r of H2O = 18.02
o Ex: M
r
of C2H4 = 28.06
o relative molecular mass is the same thing as molecular weight and average molecular mass and is equivalent to molar mass of a substance in grams.
o Ex: M
r
of O2 is 32.00; molar mass of O2 is 32.00 g/mol
1.5.1: Distinguish between the terms solute, solvent, solution and concentration (g⋅dm-3 and
mol⋅dm-3):
o solute: a substance dissolved in a liquid to form a solution.
o solvent: the dissolving medium in a solution (usually present in greatest proportion)
o solution: a homogeneous solution comprised of one or more solutes dissolved in a solvent.
o concentration:
o 1 cm3 = 1 mL
o1000 cm3 = 1 dm3 = 1 L
o g/L = g⋅dm-3
o molarity = mol/L = M = mol⋅dm-3
Topics 2 & 12: Atomic Structure (SL & HL)
2.1.3: Define the terms mass number(A), atomic number (Z) and isotopes of an element:
o mass number (A): the sum of an atom’s protons and neutrons.
o Ex: the mass number of a carbon atom with 7 neutrons is 13.
o atomic number (Z): the number of protons in an atom’s nucleus (determines
identity of atom).
o Ex: Z=79 for the element gold.
o isotope: Atoms of the same element with a different number of neutrons.
o Ex: hydrogen atoms exist in nature as one of three isotopes:
Topics 3 & 13: Periodicity (SL & HL)
3.2.1: Define the terms first ionization energy and electronegativity.
•first ionization energy: the energy required to remove one electron from an atom in its gaseous state (measured in kJ⋅mol-1):
X (g) → X+ (g) + e-
o electronegativity: a measure of an atom’s relative ability to attract a bonding pair of
electrons to itself.
13.2.4: Define the term ligand.
o ligand: neutral molecules or negative ions that contain a non-bonding pair of electrons.
O, NH3, CN1-
•Ex: H
2
Topics 4 & 14: Bonding (SL & HL)
4.1.1: Describe the ionic bond as the electrostatic attraction between oppositely charged ions:
•ionic bond: the electrostatic attraction (+,-) between oppositely charged ions.
4.2.1: Describe the covalent bond as the electrostatic attraction between a pair of electrons and
positively charged nuclei:
•covalent bond: the electrostatic attraction (+,-) between a pair of electrons and positively charged nuclei.
4.4.1: Describe the metallic bond as the electrostatic attraction between a lattice of positive ions
and delocalized electrons:
delocalized electrons: valence electrons that are detached from the atoms they
came from and are free to move throughout the metallic structure. Topics 5 & 15: Energetics (SL & HL)
5.1.1: Define the terms exothermic reaction, endothermic reaction and standard enthalpy of
reaction (∆Hө):
o exothermic reaction: a reaction that releases heat to the surroundings.
o endothermic reaction: a reaction that absorbs heat from the surroundings.
o enthalpy (H): heat energy converted from energy contained in chemical bonds.
o enthalpy change (∆H): H
products
– H reactants
o standard change in enthalpy of reaction (∆Hө): the enthalpy change when molar quantities
of reactants in their normal states react to form products
in their normal states under standard conditions of
temperature (298 K) and pressure (101.3 kPa or 1 atm).
5.4.1: Define the term average bond enthalpy:
o chemical reactions involve the breaking and making of bonds.
o it requires energy (+∆H) to break bonds.
o energy is released (-∆H) when new bonds are formed
o average bond enthalpy: the energy change that occurs when 1 mol of a covalent bond in
the gaseous state is formed from its gaseous atoms:
X (g) + Y (g) → X—Y (g)
15.1.1: Define and apply the terms standard state, standard enthalpy change of formation (∆Hөf)
and standard enthalpy change of combustion (∆Hөc):
o standard state: a precisely defined reference state as:
•temperature of a substance at exactly 25°C.
•pressure of gaseous substance are exactly 1 atm.
•liquids and solids are pure.
•aqueous solutions are exactly 1 mol⋅dm-3
o standard enthalpy change of formation (∆Hө
f ): the change in enthalpy that accompanies the
formation of one mole of a compound from its elements with all substances in standard states.
•Example: ½ N
2
(g) + O2 (g) NO2 (g) ∆Hөf = + 34 kJ/mol for NO2 (g)
o standard enthalpy change of combustion (∆Hө
c ): the change in enthalpy that accompanies the
formation of one mole of a compound from its complete reaction with oxygen gas under standard conditions.
•Example: C (s) + O
2
(g) → CO2 (g) ∆Hөc = -393 kJ/mol for CO2 (g)
15.2.1: Define and apply the terms lattice enthalpy and electron affinity:
o lattice enthalpy (∆Hө
latt ): the enthalpy change that results when a solid ionic salt is
formed from its gaseous ions under standard conditions.
•Ex: Na1+ (g) + Cl1- (g) → Na+Cl- (s) ∆Hөatt = -771 kJ⋅mol-1 for NaCl(s)
o electron affinity (∆Hө
at ): the enthalpy change associated with the addition of an
electron to a gaseous atom.
Ex: Cl (g) + e-→ Cl1- (g) ∆Hөat = -364 kJ⋅mol-1 for Cl atom
Topics 6 & 16: Kinetics (SL & HL)
6.1.1: Define the term rate of reaction:
o rate of reaction: the change in concentration of a reactant or product per unit time.
6.2.2: Define the term activation energy, E a:
o activation energy (E
a ): the minimum amount of energy required to bring about a chemical reaction.
Topics 8 & 18: Acids & Bases (SL & HL)
8.1.1: Define acids and bases according to the Brønsted-Lowry and Lewis theories:
•Brønsted-Lowry Definition
•acid: a proton (H+) donor.
•base: a proton (H+) acceptor.
o Lewis Acid-Base Model
o acid: a substance that is an electron-pair acceptor.
o base: a substance that is an electron-pair donor.
18.2.1: Describe the composition of a buffer solution and explain its action:
•buffer solution: an aqueous solution that resists changes in pH when small amounts of acid, base or water are added to it.
Topics 9 & 19: Oxidation and Reduction (SL & HL)
9.1.1: Define oxidation and reduction in terms of electron loss and gain.
o oxidation: the loss of one or more electrons from a substance.
o Zn (s) Zn2+ (aq) + 2e- (Zn is oxidized)
o reduction: the gain of one or more electrons.
o2H+ (aq) + 2e- H2 (g) (H+ is reduced)
9.2.3: Define the terms oxidizing agent and reducing agent.
•oxidizing agent: a substance that is able to oxidize other substances.
•the oxidizing agent itself undergoes reduction.
•reducing agent: a substance that is able to reduce other substances.
•the reducing agent itself undergoes oxidation.
19.1.1: Describe the standard hydrogen electrode.
o consists of a platinum electrode (inert) surrounded by hydrogen gas at 1 atm of pressure.
o the platinum electrode is immersed in an aqueous solution of acid in which the concentration of H1+ ions is exactly 1 M (1 mol⋅dm3).
o the temperature is maintained at 298 K.
19.1.2: Define the term standard electrode potential (Eθ)
•standard electrode potential (Eθ): the potential (“pull” on the electrons from the oxidizing agent) of a half-reaction under standard state conditions, as measured against the potential
of the standard hydrogen electrode.
Topics 10 & 20: Organic Chemistry (SL & HL)
10.1.1: Describe the features of a homologous series:
o have the same general formula (C
n
H2n+2 for alkanes)
•Example: the homologous series of alkanols: CH
3OH (methanol), CH3CH2OH (ethanol),
CH3CH2CH2OH (propanol), CH3CH2CH2CH2OH (butanol), etc
o neighboring members differ by a CH
2
o have similar chemical properties
o have gradually changing physical properties (vapor pressure, density, viscosity, solubility, etc).。