Atoms, Molecules and Stoichiometry

Atoms, Molecules and Stoichiometry Define and use the terms relative atomic, isotopic, molecular and formula masses, based on the 12 C scale oRelative atomic mass: The weighted average mass of naturally occurring atoms of an element on a scale where an atom of carbon-12 has a mass of exactly 12 units.oRelative isotopic mass: The mass of a particular isotope of an element on a scale in which an atom of the carbon-12 isotope has a mass of exactly 12 units.oRelative molecular mass: The weighted average mass of a molecule measured on a scale in which an atom of the carbon-12 isotope has a mass of exactly 12 units.oRelative formula mass: The weighted average mass of a molecule’s formula measured on a scale in which an atom of the carbon-12 isotope has a mass of exactly 12 units.Define and use the term mole in terms of the Avogadro constant oOne mole of a substance is the amount of that substance that has the same number of specific particles (atoms, molecules or ions) as there are atoms in exactly 12g of the carbon-12 isotope which is equal to the Avogadro constant (6.02x10 23 ) Analyse mass spectra in terms of isotopic abundances (knowledge of the working of the mass spectrometer is not required) oThe isotopic abundance shows how much of each isotope exists in a sample of an element.Calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum oRelative Atomic Mass = Isotopic Mass * Percentage Abundance / 100 Define and use the terms empirical and molecular formula

oEmpirical Formula: The simplest whole number ratio of atoms in a compound

oMolecular Formula: The total number of atoms of each element present in one molecule or one formula unit of the compound Calculate empirical and molecular formulae, using combustion data or composition by mass Write and construct balanced equations

Perform calculations, including use of the mole concept, involving:

o(i) reacting masses (from formulae and equations) o(ii) volumes of gases (e.g. in the burning of hydrocarbons) o(iii) volumes and concentrations of solutions oWhen performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question. When rounding up or down, candidates should ensure that significant figures are neither lost unnecessarily nor used beyond what is justified (see also Practical Assessment, Paper 3, Display of calculation and reasoning on page 51).on = m/Mr oMoles = Mass / Relative Molecular/Formula Mass on = V/Vr oMoles = Volume / Molar Volume on = cV 1 / 4

oMoles = Concentration * Volume Deduce stoichiometric relationships from calculations such as those above Atomic Structure Identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses oProtons have mass 1, charge +1 oNeutrons have mass 1, charge 0 oElectrons have mass 1/1836, charge -1 Deduce the behaviour of beams of protons, neutrons and electrons in electric fields oNothing happens to neutrons.oProtons and electrons experience the same magnitude force but opposite directions.oTherefore electrons have greater acceleration due to less mass.Describe the distribution of mass and charge within an atom oMajority of the mass is in the nucleus.oAll the positive charge is in the nucleus oAll the negative charge is outside the nucleus Deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (atomic and mass numbers) and charge Describe the contribution of protons and neutrons to atomic nuclei in terms of proton (atomic) number and nucleon (mass) number oAtomic number is number of protons oMass number is number of nuclides (protons + neutrons) Distinguish between isotopes on the basis of different numbers of neutrons present Recognise and use the symbolism A y x for isotopes, where x is the nucleon (mass) number and y is the proton (atomic) number Describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals oElectrons are arranged in energy levels called shells which are describe by a principal quantum number. As the number increases, the energy of the shell increases.os subshell has 1 orbital op subshell has 3 orbitals od subshell has 5 orbitals Describe and sketch the shapes of s and p orbitals 2 / 4

o State the electronic configuration of atoms and ions given the proton (atomic) number and charge, using the convention 1s 2 2s 2 2p 6 , etc.os is for groups 1 and 2 (and helium) op is for groups 14-18 (except helium) od is for groups 3-12 oThe s shell is filled before the preceding d shell. e.g. 4s before 3d oHowever, an electron will be moved from 4s to 3d to get a full or half full 3d shell as it is more stable due to reduced repulsion.oDuring ionisation, electrons are removed from the last shell first with the exception of s shell removed before the d shell. e.g. 4s before 3d.Explain and use the term ionisation energy oIonisation energy is the energy needed to remove 1 mole of electrons from 1 mole of atoms of an element in the gaseous state to form 1 mole of gaseous ions.Explain the factors influencing the ionisation energies of elements oIn general, ionisation energy increases as the proton number increases as the nuclear charge increases, causing the attraction force between nucleus and electrons to increase.

oExceptions:

The further the valence electron shell is from the nucleus, the lower the ionisation energy as the number of full electron shells between the valence electrons and the nucleus increases, causing electron shielding.Spin pair repulsion – less energy to remove electron from pair Explain the trends in ionisation energies across a period and down a group of the Periodic Table (see also Section 9.1) oGeneral increase in ΔH across period as nuclear charge increases, while distance and shielding remains reasonably constant Anomalies occur between the different subshells and spin-pair repulsion.oDecreases in ΔH down a group as distance and shielding increases which outweigh the increased nuclear charge Deduce the electronic configurations of elements from successive ionisation energy data Interpret successive ionisation energy data of an element in terms of the position of that element within the Periodic Table oMassive difference in ΔH between removing electron “1” and electron “8”. 3 / 4

Explain and use the term electron affinity oElectron affinity is the enthalpy change when 1 mole of electrons is added to 1 mole of gaseous atoms to form 1 mole of gaseous 1- ions under standard conditions Chemical Bonding Describe ionic bonding, using the examples of sodium chloride, magnesium oxide and calcium fluoride, including the use of ‘dot-and-cross’ diagrams oIonic bonding consists of cations and anions electrostatically bonded.

Describe, including the use of ‘dot-and-cross’ diagrams:

ocovalent bonding, in molecules such as hydrogen, oxygen, chlorine, hydrogen chloride, carbon dioxide, methane, ethene Covalent bonding consists of the sharing of pairs of electrons between two atoms oco-ordinate (dative covalent) bonding, such as in the formation of the ammonium ion and in the Al2Cl6 molecule Co-ordinate bonding / dative covalent bonding is a covalent bond where the pair of electrons come from the same atom Describe covalent bonding in terms of orbital overlap, giving σ and π bonds, including the concept of hybridisation to form sp, sp 2 and sp 3 orbitals (see also Section 14.3) oEvery single bond is a σ bond. All bonds except one on multi bonds are π bonds with the one being σ.oHybridisation occurs when orbitals overlap. One electron from the s orbital moves up into the empty p orbital to have four half filled subshells.oWhen an s orbital overlaps with three p orbitals, an sp 3 hybrid is formed where each orbital is ¼s character and ¾p character.oWhen an s orbital overlaps with two p orbitals, an sp 2 hybrid is formed where each orbital is ⅓s character and ⅔p character.oWhen an s orbital overlaps with one p orbital, an sp hybrid is formed where each orbital is ½s character and ½p character.osp, sp 2 , and sp 3 are all σ bonds.oA π bond is formed when two p orbitals overlap.Explain the shapes of, and bond angles in, molecules by using the qualitative model of electron- pair repulsion (including lone pairs), using as simple examples BF3 (trigonal planar), CO2 (linear), CH4 (tetrahedral), NH3 (pyramidal), H2O (non-linear), SF6 (octahedral), PF5 (trigonal bipyramidal) olone pair-lone pair > lone pair-bond pair > bond pair-bond pair oTrigonal Planar - 120˚ oLinear - 180˚ oTetrahedral - 109.5˚ oTrigonal Pyramidal - 107˚ oNon-linear - 104.5˚ oOctahedral - 90˚ oTrigonal Bipyramidal - 120˚, 90˚

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