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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. 11 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. 31 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