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Revision topic 1-3 Properties of gases: ● ● ● ● ● Gases have a small mass All gases respond in a similar way to changes in temperature, pressure and volume. They exert a pressure, that depends on the amount of gas and the temperature There is no bonding between molecules The molecules may move in all directions allowing the gas to expand throughout any container Pressure ● Pressure is the amount of force exerted on one unit of area. Avogadro´s law ● One mole of any gas will occupy the same volume, if the temperature and pressure are the same. = equal volumes of different gases at the same temperature and pressure contain the same number of particles. ● This volume is known as the molar volume of a gas, Vmolar (unit dm3/mol) n = V / Vmolar http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/gasesv6.swf Boyles law ● ● PV = k1 i.e. the relationship between volume and pressure for a gas The law describes how the volume of a given amount of gas at constant temperature varies inversely with the applied pressure: Gay Lussac's law ● ● i.e. the relationship between temperature and pressure for a gas states that the pressure of a given amount of gas held at constant volume is directly proportional to its temperature in kelvin = an increase in temperature increases the kinetic energy of the particles, which means they will move faster and collide with the walls with more energy and more frequency P = k 2T Charles' law ● ● i.e. the effect of temperature on the gas volume The law describes how the volume of a given amount of gas is directly proportional to its temperature in kelvins. V = k3T The combined gas law ● The three gas laws can be combined to one expression: where 1 refers to the initial conditions and 2 to the final conditions. Only ideal gases obey the gas laws perfectly. Real gases can be treated as ideal gases, unless dealing with extremely precise measurements. ● The ideal gas equation ● The ideal gas law relates pressure, volume, temperature and amount of substance: where R is the gas constant = 8,31451 J K-1 mol-1 ● Only ideal gases will follow this equation exactly. Isotopes ● Atoms of the same element always contain the same number of protons, but if they contain a different number of neutrons they are called isotopes. • Isotopes have: - The same number of protons = same atomic number (Z) - Different number of neutrons, thus different mass number (A) • Relative atomic mass • The mass of an atom depends on the number of protons and neutrons in the nucleus (the mass of electrons is so small that it can be ignored in chemistry). • The relative atomic mass is a weighted average mass (according to relative abundances) of all the naturally occuring isotopes of an element compared with an atom of the C-12 isoptope, which has a mass of exactly 12. • e.g. Ar (H) = 1,008 • Many isotopes of elements are radioactive, because their nuclei break down spontaneously and emit radiation. 2.2 The mass spectrometer • The mass of individual isotopes can be determined using a mass spectrometer: 1) sample is vaporized 2) ionized to positive ions: M (g) → M+ (g) + e3) the ions are accelerated and will pass through a charged slit 4) the ions are deflected by an external magnetic field 5) the ions are recorded on a detector (relative amounts of the ions, mass to charge ratio) • Electromagnetic radiation can be viewed as a stream of tiny packets (quanta) of energy. • The energy of electromagnetic radiation is expressed: E = hf where: h = Planck's constant ≈ 6,626 · 10-34 J s • The smaller the wavelength, thus the higher the frequency, the more energy the wave possesses. Line spectra ● ● ● ● When white light is passed through hydrogen gas, some of the light is absorbed. A continous spectrum is an emission spectrum that contains all the wavelengths or frequencies of visible light. An absorption spectrum is produced where some colours are missing ( those that are absorbed by hydrogen). A corresponding emission line spectrum has only certain wavelegths or frequencies of visible light.The lines correspond to the light of particular wavelengths given off (=emitted) by the element. Excitation • When an atom absorbs energy, an electron moves into an orbit of higher energy further from the nucleus. An unstable excited state is produced. • The electron soon falls back to a lower level and gives out (= emittes) the energy in the form of electromagnetic radiation with a specific frequency, f. • The lowest energy level of the electron is called the ground state. ∆Eelectron = hf Topic 3. Peridoicity ● Periodicity = repeating patterns of physical and chemical properties. 3.2 Physical properties ● ● Groups: the group number (1-7) gives the number of electrons in the outermost energy level (= valence electrons). All atoms in the same group (or column) have similar chemical properties. Periods: All elements in the same period have the same number of energy levels (electron shells). Effective nuclear charge ● ● ● In every atom there is a balance between the attraction of the positively charged nucleus for the negatively charged electrons and repulsion between the electrons. The outer electrons (valence electrons) do not experience the full attraction of the positive nucleus because of the presence of inner electrons. They are shielded from the nucleus and repelled by the inner electrons. The effective charge experienced by the outer electrons is less than the full nuclear charge. Atomic radius ● ● The effective charge increases (= atomic radii decreases) as a period is crossed from the left to the right: ● One proton is added to the nucleus and one electron to the outermost electron shell. ● There is no change in the number of inner electrons. The effective charge remains almost the same down a group, because: ● The increase in the nuclear charge is offset by the increase in the number of inner electrons. Ionic radius ● ● ● ● Positive ions are smaller than their parent atoms (because of loss of the outer shell). Negative ions are larger than their parent atoms (because of increased electron repulsion by addition of electrons). The ionic radii decrease as a period is crossed from the left to the right (because of increased attraction between the nucleus and the electrons). The ionic radii increase down a group as the number of electron shells increase. Electronegativity ● ● A measure of the tendency of an atom in a molecule to attrat a pair of shared electrons towards itself. Electronegativity values increase from left to right across the periodic table and decrease down a group. The first ionization energy • The energy required to remove one mole of electrons from one mole of isolated gaseous atoms of an element to form one mole of gaseous unipositive ions under standard conditions. X (g) → X+ (g) + e● The second and third ionization energies : X+ (g) → X2+ (g) + eX2+ (g) → X3+ (g) + e- ● The stronger the bonds, the higher the melting point. 3.3 Chemical properties Group 1: the alkali metals ● ● Shiny, silvery, soft Reactivity increases down the group, because the first ionization energy decrease down a group. https://www.youtube.com/watch?v=uixxJtJPVXk Group 7: the halogens ● ● ● Reactivity decreases down a group as the atomic radius increases and the attraction for outer electrons decreases. Halide formation: 2 Na (s) + Cl2 (g) → 2NaCl (s) Displacement reactions: a more reactive halogen can displace the ions of a less reactive halogen (= take the electron from the less reactive ions) . Ex. Which of the following chemical reactions are possible? d Chemical properties of elements in the Period 3 ● ● ● Na, Mg and Al are metals. They are shiny and good conductors of heat and electricity. Si is a semi-conductor and is called a metalloid since it has some of the properties of a metal and some of a non-metal. P, S, Cl and Ar are all non-metals and do not conduct electricity. The Period 3 oxides ● ● Metals can also be distinguished from non-metals by their chemical properties. Oxides are products of reactions in which an element reacts with oxygen. Metal oxides Metal oxides (Na2O, MgO) are ionic compounds: - solids in room temperature - high mp & bp - conduct electricity when molten (or in aqueous solutions) - basic in aqueous solutions: Na2O (s) + H2O (l) → 2 NaOH (aq) MgO (s) + H2O (l) → Mg(OH)2 (s) Aluminium oxide ● ● Al2O3 is an ionic oxide with some covalent character. It is amphoteric as it acts as a base when it reacts with acids and acts as an acid when it reacts with bases: Silicon oxide ● ● ● Silicon oxide, SiO2, has a giant covalent structure with very high melting and boiling points. It is insoluble in water. It is classified as an an acidic oxide, because it reacts with NaOH at temperatures above 350° C. Non-metallic oxides ● Nonmetal oxides are covalently bonded because of the small difference in the elements' electronegativity values. ● sulfur: SO2, SO3 ● chlorine: Cl2O, Cl2O7 ● phosphorus: P4O6, P4O10 ● They have low mp and bp and do not conduct electricity. ● Non-metallic oxides are acidic in aqueous solutions: P4O10 + 6 H2O (l) → 4 H3PO4 (aq) phosphoric(V)acid SO3 (g) + H2O (l) → H2SO4 (aq) sulfuric(VI)acid Acidic rain • Rain is naturally acidic, because the water molecules react with the CO2 in the air and form the weak acid H2CO3. ● ● Acid rain is therefore: precipitation (rain, snow) with pH lower than 5.6. The main acids present in acid rain are sulfuric acid nitric acid. ● The main acids present in acid rain are sulfuric acid nitric acid. ● The sulfuric acid in the rain reacts with calcium carbonate (in limestone or marble) to create calcium sulfate, which then flakes off. CaCO3(s) + H2SO4(aq) → CaSO4(aq) + CO2(g) + H2O(l) TOPIC 2 (HL) ● Graph of successive ionization energies for sodium : ● Sub-levels of electrons (HL) • 1. Aufbau Principle: electrons are placed into orbitals of lowest energy first. • 2. Pauli exclusion principle: no two electrons in an atom can have exactly the same four quantum numbers no more than two electrons can occupy any one orbital, and if two electrons are in the same orbital they must have the opposite spin. 1s • 3. Hund´s rule: If more than one orbital in a sub-level is available, for example px, py, pz , electrons occupy the orbitals single with parallel spin before they are paired up. /watch?v=sMt5Dcex0kg The shapes of s, p and d orbitals The shapes of s, p and d orbitals Halv-filled and filled sub-levels are more stable: Cr Cu Electron configuration ● ● ● Electron arrangement: the number of electrons in each main level (e.g. Na: 2, 8,1) Electron configuration: the number of electrons in each sub-level (e.g. Na: 1s22s22p63s1) When writing electronic configurations, the sum of the superscripts must always total the number of electrons in the atom (or ion). 13.2 First row d-block elements (HL) Physical properties ● Small atomic radii compared to the neighbouring sblock elements. ● Only a small increase in atomic radii across the period. ● Metallic bonding ● High density Chemical properties Transition metals ● An element that contain an incomplete d sub-level in one or more of its oxidation states = forms ions with partially filled d sub-levels ● Neither Zn nor Sc are considered to be transition elements: Transition metals Complexes ● ● ● Because of their small size and relatively high charge, the transition metal ions have a high charge density. They attract species that are rich in electrons: ligands. A ligand is a molecule or negative ion that donates a pair of electrons to a central metal ion to form a dative (or coordinat) covalent bond, e.g. H2O, NH3, Cl-, CN- [Fe(H2O)6]3+ Ligand exchange (replacement) ● In aqueous solutions the water molecule usually act as a ligand, but it can be replaced with another ligand. Transition metals absorb visible light ● ● ● ● In an isolated gaseous transition metal atom, the five 3 d orbitals are degenerate since they all have the same energy. However, since the 3 d sub-shells all have different orientations in space they will be orientated differently relative to the ligands in a complex ion. The 3d electrons close to a ligand will experience repulsion and be raised in energy. The 3d electrons further from the ligand will be reduced in energy. → The 3d sub/shell splits into two energy levels. ● ● ● If all light is absorbed, the substance appears black. If only certain wavelengths are absorbed, the compound appears coloured. If all light is reflected, the compound appears white. Heterogeneous catalysts ● The catalyst is in a different state from the reactants. ● e.g. Fe in the Haber process (ammonia is produced) Palladium and platinum in converters in cars Homogenous catalysts ● ● ● Homogenous catalysts are in the same phase as the reactants and products. The two reacting species bond chemically to the transition metal for ing an intermediate, and then leave. Transition metals can be relatively easily oxidized and reduced due to their variable oxidation states.