Introduction to General Chemistry 2024-25 PDF

Introduction to General Chemistry (345.240) 1 Courses in English and German Written exam 90 min In German and English Identical exams for both languages 3 options to write the exam: Tuesday, November 12th 2024...

Introduction to General Chemistry (345.240) 1 Courses in English and German Written exam 90 min In German and English Identical exams for both languages 3 options to write the exam: Tuesday, November 12th 2024 Thursday December 13th 2024 Friday January 10th 2025 For the exam you will have to register (or unregister if you change your mind) in KUSSS 2 There will be 48 points in the exam for the Introduction to General Chemistry Lecture. By asking (meaningful) questions about the previous lecture you can get up to 2 bonus points for the exams taking place in winter semester 2024/25. Note: It's no problem to pass the exam without bonus points. If you ask meaningful questions in min. 2/3 of all given lectures (most likely 10 out of 14), you will get 2 bonus points. If you ask meaningful questions in min. 1/3 of all given lectures (most likely 5 out of 14), you will get 1 bonus point. A meaningful question...... needs to be about the topic of the given lecture (and not about something from a week before or something which will be discussed a week later).... needs to be specific (just "I didn't understand anything, Can you explain it again?" is not specific!) You can ask your questions either in English or German. These questions should encourage you to think about the topics of the lecture directly after the lecture and not only on the evening before the exam (when it will be too late to ask questions). We will discuss these questions (or the most relevant/frequent ones) in the following lectures. So make use of this possibility to get your questions discussed again. Submitting your questions is only possible (1) after the current lecture and (2) on the very same day of the current lecture. If you have another question later on, you can of course ask it directly in the lecture, however a submission in Moodle is not possible after the deadline. Note: Other students are not able to see your question. 3 Main aim of this lecture is: Providing an introduction to the field of General Chemistry Priming students for the Introductory/General Lab Course These slides should be seen as a helpful tool for learning the different topics in “Introduction to General Chemistry“. Anyway, they do not replace the physical attendance at the corresponding lecture. Additional information can/will be given during the lecture that might not be fully included in these slides. It should be stated that these slides cannot be seen as a „full text script“ – so explanations given during the lecture will be needed for a proper understanding of the topics discussed within this lecture! Books (see next page) may be used as supportive tools while studying for the exam. 4 Recommended books (English): T.E. Brown, H.E. Lemay et al., Chemistry the Central Science, Pearson R. Lewis, W. Evans, Chemistry, Palgrave P. Atkins, L. Jones, Chemical Principles, MacMillan Lecture slides as pdf via MOODLE. 5  Content Chapter# Topic 1 General Introduction and Fundamentals 2 Atoms and Elements 3 The Periodic Table 4 The Chemical Bond 5 Properties of Gases, Liquids and Solids 6 Redox Reactions 7 Chemical Equilibrium and Properties of Solutions 8 Acids and Bases 9 Short Introduction to Coordination Chemistry 10 Complex/Coupled Equilibria 11 A Short Introduction into Thermochemistry and Chemical Thermodynamics 6 1 General Introduction and Fundamentals General Introduction and Fundamentals 7 1 General Introduction and Fundamentals Chemistry is the scientific discipline involved with elements and compounds composed of atoms, molecules and ions: their composition, structure, properties, behavior and the changes they undergo during a reaction with other substances Main aspect: Chemical Reaction Difference Chemistry/Physics: Chemical process: new substances with different properties are formed → Sodium and Chlorine form Sodiumchloride; Bromine and Benzene form Bromobenzene Physical process: substances are preserved but may (for example) change their state (for example from solid to liquid) → Ice melts to Water; solid sugar dissolves in Water; 8 1 General Introduction and Fundamentals Traditionally Chemistry has been sub-divided by classial subjects such as: Organic Chemistry (mainly dealing with carbon chemistry) Inorganic Chemistry (dealing with the chemistry of all the other elements) Physical Chemistry (dealing with the basic principles of chemistry) Analytical Chemistry (investigating the qualitative and quantitative composition of samples) Biochemistry (exploring the role of chemistry in biology) ………………………….. 9 1 General Introduction and Fundamentals Nowadays new (more application oriented) and interconnected categories have evolved such as: Figure from: https://www.utoledo.edu/nsm/chemistry/research_areas/ (3.3.2020). 10 1 General Introduction and Fundamentals The three levels of Chemistry Macroscopic level Magnesium metal reacts with gaseous oxygen under release of light and heat forming a white powder with completely different properties than the educts employed Microscopic level Upon reaction of magnesium with oxygen the bonds in the educts break to form a new ionic compound namely magnesiumoxide Symbolic level The language of chemists using elemental symbols and arrows 11  1 General Introduction and Fundamentals Matter: everything with a mass and a volume (e.g. water, iron, a leaf, air, sand, …) Substance (=pure substance): matter that has distinct properties and a composition that does not vary from sample to sample (e. g. water, quartz, gold,…….) Pure substances may be Elements or Compounds Elements: 118 are known (see Periodic Table of Elements) – 80 of them are stable - cannot be decomposed into simpler substances Atoms of an element Molecules of an element Compounds: substances composed from more than one element Molecules of a compound Mixture of elements and a compound 13  1 General Introduction and Fundamentals Mixtures Most of the matter consists of mixtures of different pure substances. In a mixture each compound retains its identity and properties. Whereas pure substances have a fixed composition, the composition within a mixture can vary. Typical examples for mixtures varying in texture and appearance from sample to sample are rocks or wood. These mixtures are called heterogeneous mixtures. Examples for homogeneous mixtures are solutions of sugar or salt in water Heterogenous mixture Homogenous mixture 14 1 General Introduction and Fundamentals Cassification of matter Figure from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 15  1 General Introduction and Fundamentals Pysical properties - States of matter Solid (s): particles (atoms, molecules or ions) are packed densely; a solid is incompressible, and rigid; particles move with higher temperature but cannot swap places. Crystalline: built from regular units (e.g. crystal lattice in NaCl) Amorphous: non-regular (soot) Liquid (l): densly packed, (almost) incompressible; particles move with higher temperature and may swap places (flow); (s) (l): melting point Solid Liquid Gaseous Gaseous (g): particles can move freely (e.g. H2 @ One substance – three phases 20°: 1760 m/s); particles collide, but interact only marginally; gases take any shape Definition: A phase is a region where all physical properties are essentially uniform. Phase boundaries are boundaries between different phases Figure from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 16 1 General Introduction and Fundamentals Transition between states of matter: Gaseous melting Liquid Solid freezing 17 1 General Introduction and Fundamentals Physical and chemical properties of matter Intensive properties: do not depend on the amount of substance (temperature, density, melting point, …..) Extensive properties: do depend on the amount of substance (mass, volume, energy…..) Physical properties: can be determined without changing the identity or composition of the substance (colour, odour, hardness, solubility…..) Chemical properties: describe how matter can be changed/react (flammability, reaction with specific reagents, oxidizability…..) 18 1 General Introduction and Fundamentals Units of measurement Unit / Definition Symbol SI – Base units Mass m Kilogramm kg Length l Meter m Time t Second s Temperature T Kelvin K Amount n Mole mol Electric current I Ampere A Luminous intensity Iv Candela cd Derived units Volume V Cubicmeter m3, 1 m3=1000 l Density d Mass /Volume g/cm3 Molar mass M Mass / amount g/mol Force F Mass. acceleration kg m s-2=N (Newton) Pressure P Force/ Area N/m2 Charge q current. time A s=C (Coulomb) 19 More about units - see „Chemical Calculations“! 1 General Introduction and Fundamentals Amount of substance, Mole, Molar Mass The mole is one of the most important concepts in chemsitry! Definition: 1 mole of objects contains the same number of objects as there are atoms in exactly 12.000 g of 12C. The number of atoms present in exactely 12.000 g of 12C equal Avogadro´s constant NA: NA = 6.0221. 1023 Therefore 1 mole of an object equals a number of 6.0221. 1023 of this object. The amount of this substance equals 1 mole. The molar mass M [in g/mol] of an element is the mass m (in gramm) of one mole of its atoms. m M= n The molar mass of molecules is composed from the molar masses of the elements combined in the molecule. 20  1 General Introduction and Fundamentals 1 mole … equals: Carbon: 12 g (3,4-5,3 cm³) Sulfur: 32 g (~16 cm³) Copper: 64 g (~7 cm³) Lead: 207 g (~18 cm³) Mercury: 201 g (~15 cm³) Oxygen: 32 g (~22400 cm³) Water: 18 g (~18 cm³) Sodiumchloride: 58,5 g (~27 cm³) = Na (23.0) + Cl (35.45) Figure from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 21 1 General Introduction and Fundamentals Chemical reactions and reaction equations A chemical reaction is a process, where either one or multiple chemical compounds are transformed into other compounds whereby energy is either released or taken up. Example: Hydrogen and Oxygen form Water No chemical reactions are: Physical processes Nuclear reactions 22 1 General Introduction and Fundamentals Chemical reactions and reaction equations Rules: Coefficients The number of atoms of each for reaction stochiometry element must be the same on both sides of the reaction arrow The number of charges must be the same on both sides of the reaction 2 H2 + O2 → 2 H2O arrow Coefficients should be as small as possible and integers Educt(s) Product(s) Reactant(s) Starting material(s) Reaction-arrow 23 1 General Introduction and Fundamentals Chemical reactions and reaction equations: Symbols and additional information Arrows Information on states Reaction-arrow → gaseous (g), ↑ Sequence of reactions → → liquid (l) Forward and back reaction ⇄ solid (s), ↓ Equilibrium reaction ⇌ dissolved in water (aq) Transfer of electrons Reaction conditions heat ∆ Arrows for mesomeric structures ↔ light hν (no reaction!) use of a catalyst … 24 1 General Introduction and Fundamentals Chemical reactions and reaction equations: 2S + 3 O2 → 2 SO3 react to 2 particles S and 3 particle O2 2 particles SO3 (form) 2 mol S and 3 mol O2 react to 2 mol SO3 → Coefficients describe the ratio between the particles involved! 2gS and 3 g O2 react to 5 g SO3 64 g S and 96 g O2 react to 160 g SO3 32 g S and 48 g O2 react to 80 g SO3 1S + 1½ O2 → 1 SO3 1 particle S and 1½ particle O2 react to 1 particle SO3 1 mol S and 1½ mol O2 react to 1 mol SO3 6,022∙1023 9,033∙1023 6,022∙1023 and react to particles S particles O2 particles SO3 25 2 Atoms and Elements Atoms and Elements 26 2 Atoms and Elements Development of the theory behind the concept of atoms Demokrit (460-370 b.c.): the world consist of impartible particles („atomos“) Platon, Aristoteles: there are no impartible particles 17. century, amongst others Isaac Newton (1642-1727): revival of Demokrits ideas; small impartible particles are used to describe properties of e.g. gases/air Graphics from http://www.artnet.com/artists/artus-wolfaerts/demokrit-der-lachende-philosoph-_nQLd0X-UJhyvZ3cGuhiGQ2 (20.2.2020) and https://de.wikipedia.org/wiki/Isaac_Newton#/media/Datei:GodfreyKneller-IsaacNewton-1689.jpg (20.2.2020). 27 2 Atoms and Elements Development of the theory behind the concept of atoms John Dalton (1766-1844): Each element consists of small particles (atoms) All atoms within an element have the same mass and the same properties, atoms of different elements are different Upon chemical reactions atoms are not transformed (into other atoms), modified, created or destroyed By combining different atoms compounds are formed, whereby the ratio between these atoms within one compound is constant Graphics from https://geboren.am/person/john-dalton (20.2.2020). 28 2 Atoms and Elements Development of the theory behind the concept of atoms Daltons atomic theory allows explaining the following (already known) principles: Law of constant composition: In a given compound, the relative numbers and kinds of atoms are constant.1 Law of conservation of mass: The total mass of materials present after a chemical reaction is the same as the total mass present before the reaction.1 and predicts the following: Law of multiple proportions: If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers.1 1 from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 29  2 Atoms and Elements Development of the theory behind the concept of atoms Joseph J. Thomson (1856-1940): Explanation of cathode-rays as a current of negatively charged particles (electrons), allows to determine the relationship between charge and mass charge/mass = 1.76 x 108 C/g 30 Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson.  2 Atoms and Elements Development of the theory behind the concept of atoms Robert Millikan (1868-1953): oil-drop experiment → allows to determine the charge and subsequently the mass of electrons charge/mass: 1.76 x 108 C/g charge: 1.602 x 10-19 C mass: 9.10 x 10-28 g 31 Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 2 Atoms and Elements Development of the theory behind the concept of atoms Video von: https://infocenter.pearson.de/media/cwsfiles/3827371910/demovideos/ch02/RutherfordExpNuclearAtom.html (abgerufen am 20.2.2020). 32 2 Atoms and Elements Development of the theory behind the concept of atoms Ernest Rutherford (1871-1937): → The positive charge and most of the mass resides in a very small extremely dense region called the nucleus → Most of the volume of the atom is empty space – electrons move within this space Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 33 2 Atoms and Elements Structure of an atom Atoms consist of protons, neutrons and electrons protons and neutrons (=nucleons) form the nucleus (small, high density) Most of the volume of the atom is empty space – electrons move within this space Every atom has equal numbers of electrons and protons 34 2 Atoms and Elements Average molar mass © Christian G. Huber University of Salzburg 35 2 Atoms and Elements In chemical reactions only the electron shell is involved. The law of the conservation of matter is valid. Energies released are in the range of several eV. In nuclear reactions also the nucleus is affected. Energies involved are in the MeV range. Transformation of matter into energy must be considered. 36  2 Atoms and Elements Nuclear chemistry Radioavtive decay – an instable nucleus is transformed into a more stable one by emitting radiation Nuclear transmutation – a nucleus is bombarded with subatomic particles resulting in a transformation of the nucleus Nuclear fission – A heavy nucleus is split into two lighter nuclei (spontaneous or induced by bombardment with e.g. neutrons) – see: nuclear power plants Nuclear fusion – Two light nuclei fuse into a heavy one – see: sun 37 2 Atoms and Elements Radioactive radiation α-radiation (e.g. 210 84𝑃𝑜 → 206 82𝑃𝑏 + 42𝐻𝑒) – Emission of α-particles (mass = 4, atomic number = 2; v > 10000 km/s); equals Helium nucleus He2+ – Only for heavy nuclei mass number >209 – Shielding required: piece of paper β--radiation (e.g. 82 35𝐵𝑟 → 82 36𝐾𝑟 + −10𝑒) – Emission of an electron ( −10𝑒; v > 130000 km/s) – Mass number remains the same; atomic number +1 as a neutron is transformed into a proton – Shielding required: aluminium foil γ-radiation – Electromagnetic radiation with very low wavelength = high energy – Shielding required: massive walls of lead – Further decay reactions: β+-decay electron capture … 38 2 Atoms and Elements Use of radioactivity Radiometric dating – minerals 238 206 92𝑈 →→ 82𝑃𝑏 – organic materials → radiocarbon dating 146𝐶 → 14 7𝑁 + −10𝑒 Medicine – for example: diagnosis and therapy of tumors Maintenance-free sources of energy – for example: pace makers, satellites Sterilization (of food or pharmaceutical formulations) Chemistry – investigations on reaction mechanisms … 39 2 Atoms and Elements Electromagnetic radiation 40 2 Atoms and Elements Development of the theory behind the concept of atoms The wave depicted in (a) has double the wavelength and half the frequency of the one shown in (b) The wavelength λ (nm) is the distance between two successive wave crests or troughs The frequency ν (Hz): the number of waves passing a given point per second The amplitude: intensity of radiation Grafiken aus: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Basiswissen Chemie, Pearson. 41  2 Atoms and Elements The structure of the electron shell The understanding of the structure of the electron shell is closely related with the analysis of electromagnetic radiation (light) that is either emitted or absorbed by matter The following theoretical and experimental considerations which are discussed in the following slides are regarded essential for the understanding of the structure of the electronshell: Light emission from hot objects (black body radiation) Photoelectric effect Line spectra from atoms Grafics from wikipedia 42  2 Atoms and Elements Hot objects emitt radiation The wavelength/frequency/colour of the radiation depends on the temperature These observations could only be explained when assuming energy as a quantisized parameter E = hn = hcl-1 h = Planks constant = 6.626 x 10-34 Joule second [J s] c = speed of light= 2.997 x 108 m / second [m s-1] Potential energy increases continously Potential energy increases stepwise (quantisized) 43 Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson.  2 Atoms and Elements Light shining on clean metal surfaces can cause electrons to be emitted from the surface – this is called the photoelectric effect This only happens if the energy/frequency of the light exceeds a certain threshold Increasing the intensity of the light does not lead to emission of electrons only increasing the energy up to the threshold value results in the emission of an electron – but above the energy threshold increasing the intensity of light increases the number of electrons emitted This can be explained in a way that a certain amount of energy is needed to overcome the attractive forces holding the electron within the metal; if the energy of the photon striking the surface exceeds the value needed for the emission of the electron, excessive energy is converted to kinetic energy determining the speed of the electron 44 Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 2 Atoms and Elements Video von: https://infocenter.pearson.de/media/cwsfiles/3827371910/demovideos/ch06/PhotoelectricEffect.html (abgerufen am 20.2.2020). 45  2 Atoms and Elements Line spectra: When applying very low pressures and high voltages, gases emit light with different colours (wavelengths) → using a suitable device (e.g. prism) this light may be split up into the different lines (colours) Line spectra from hydrogen lead to 4 distinct lines: 410 nm, 434 nm, 486 nm, 656 nm Already in 1885 Johann Balmer derived a simple formula: 𝑛2 𝜆=𝐴⋅ ,𝑛 = 3,4,5,6 𝑛2 −4 1 1 1 Finally with the Rydberg formula we get: = 𝑅𝐻 ⋅ ( − ) 𝜆 𝑛12 𝑛22 Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 46 2 Atoms and Elements Bohr´s model Combining the knowledge from Rutherfords „nuclear atom“ and the necessity to explain line spectra led to the formulation of Bohr´s model with the following three postulates: Only orbits with certain radii, corresponding to specific energies are permitted for the electron in the hydrogen atom; the first orbit (n=1) is closest to the nucleus; for n ∞ the energy of the electron approaches 0 An electron in a permitted orbit is in an „allowed“ energy state – for this reason it does not emit energy Energy is emitted or absorbed when the electron changes from one permitted energy state to another; this energy is quantisized and equals E=hn; 47 2 Atoms and Elements From this the energy states of the hydrogen atom can be calculated as follows: constant negative!!!! with n = principal quantum number; the first orbit (n=1) is closest to the nucleus – leading to the highest value for negative energy = highest stability if, for hydrogen, the electron is in the n=1 orbit the atom is in the ground state, if it is in a higher orbit (n=2,3,4….) it is in the excited state if the electron jumps from one state to another, the energy changes according to: i…initial f…final resulting in: 48 2 Atoms and Elements Photon emitted has a wavelength of 102 nm 49 2 Atoms and Elements Bohr´s model However Bohr´s model suffers from several shortcomings: it only works for one-electron and not for multi-electron atoms according to classical physics the electron should loose energy and fall into the positively charged nucleus the model violates the principle of Heisenbergs uncertainty Solution: As radiation can be described as wave-like or particle like, also matter should possess wave properties the concept of matter waves allows to solve the problems related to Bohr´s model 50 2 Atoms and Elements The concept of matter waves For connecting matter with wave-properties De Broglie formulated the following equation: Electrons orbiting around the nucleus can be described as waves (= matter waves). The wavelength of these matter waves depends on the mass and the velocity of the electron (= momentum). As can be understood from the equation above, reasonably small wavelengths (i.e. high energies) are only obtained for very small values of mv. Matter waves for larger particles have so tiny wavelengths that they are almost unobservable. So matter waves only play a role in the microscopic world 51 2 Atoms and Elements The concept of matter waves h = 6.626 10-34 J s Examples for calculation: J = [kg m2 s-2] - A golf ball: m = 45.9 g; v = 50 m s-1 6.626 𝑥 10−34 𝜆= 2.89 x 10-34 m 45.9 𝑥 10−3 5 𝑥 101 Add appropriate units and check for plausibility! - An electron: m= 9.11 10- 31 kg; v = 5.97 106 m s-1 6.626 𝑥 10−34 𝜆= 1.22 x 10-10 m = 0.122 nm 9.11 𝑥 10−31 5.97 𝑥 106 52 2 Atoms and Elements Quantum mechanic treatment of the atom The Schrödinger equation succeeded in describing both the wave-like and the particle-like behavior of the electron: Thereby: - the electron is seen as a „standing wave“ - the solution of the Schrödinger equation provides a wave function Y describing the electron - Heissenbergs uncertainty does not allow to determine the position of the electron exactely; only the probability density or electron density Y2 can be given 53 2 Atoms and Elements Quantum mechanic treatment of the atom standing wave wave function Y probability density Y2 Graphics modified from https://physikunterricht-online.de/jahrgang-12/linearer-potentialtopf/. 54 2 Atoms and Elements Quantum mechanic treatment of the atom 𝜓 Wave functions 𝜓: solutions of Schrödinger´s equation Probability/electron density/ 𝜓 2 2 Radial probability function 4𝜋𝑟 2 ⋅ 𝜓 2 𝜓 → leads to orbitals Schrödinger´s equation yields a set of wave functions → are described by 3 quantum numbers: 4𝜋𝑟 2 ⋅ 𝜓 2 Graphics from http://www2.physki.de/PhysKi/index.php/Wasserstoff-Atom#cite_note-FN5-9 (Wolfram CDF Player Animation, abgerufen am 20.2.2020) und Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Basiswissen Chemie, Pearson. 55 2 Atoms and Elements 56 2 Atoms and Elements The angular momentum quantum number (l) 57 2 Atoms and Elements The magnetic quantum number (ml) the magnetic quantum number characterizes the energy of an electron when an external magnetic field is applied Zeeman effect 58  2 Atoms and Elements s - Orbitals Contour plot Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 59  2 Atoms and Elements s - Orbitals X 4pr2 Contour plot Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 60  Solutions of the Schrödinger equation 61 2 Atoms and Elements p - Orbitals Dumbbell shaped (different lobes have opposite signs (+/-) for opposite directions - but identical probability! One node each through the nucleus 3 possible orientations (px, py, pz) Orbitals show identical (degenerate) energy 62 Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson.  2 Atoms and Elements d - Orbitals Leaf-clover shaped; the four orbital lobes have opposite signs (+/-) – the fith looks like a floating tyre Two nodes through the nucleus 5 possible orientations (dxy, dyz, dzx dx2-y2, dz2) d-Orbitals are particularely important in transition metals Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 63 2 Atoms and Elements f - Orbitals Complex geometry Three nodes through the nucleus 7 possible orientations f-Orbitals are particularely important in lanthanois and actinoids Grafik von https://commons.wikimedia.org/wiki/File:F_orbital.png (abgerufen am 20.2.2020). 64 2 Atoms and Elements Atomic Orbitals Grafik von: https://www.researchgate.net/figure/Wavefunctions-of-some-s-p-d-and-f-orbitals_fig3_226412903 (abgerufen am 20.2.2020). 65 2 Atoms and Elements Orbitals and quantum numbers The shell with the principal quantum number (n) consists of exactely n subshells Each subshell consists of a specific number of orbitals whereby each orbital corresponds to a different allowed value of (ml). There are (2l+1) values for ml ranging from -l to +l. The total number of orbitals in a shell is n2. 66 2 Atoms and Elements Orbitals and quantum numbers 67 2 Atoms and Elements One-electron vs. many-electron atoms Hydrogen Atom with more than one electron needs additional quantum number: ms 68 2 Atoms and Elements Orbitals and quantum numbers Each orbital can hold a maximum of two electrons These electrons behave as they would spin around their own axis This results in an intrinsic angular momentum called spin Electrons within an orbital must have different spins Possible values for this spin are +1/2 and -1/2 often symbolized as and So each electron is unambiguosly characterized by its quantum numbers n, l, ml and ms (Pauli principle). 69 2 Atoms and Elements Orbitals and quantum numbers 70 2 Atoms and Elements 71 2 Atoms and Elements Ne 72  2 Atoms and Elements Occupation of orbitals Orbitals are filled with increasing energy – so for example 4s is filled before 3d ! Figures from: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson. 73 2 Atoms and Elements Occupation of orbitals There are exceptions to the Aufbau-principle, as half- or fully filled shells are particularly stable. So for example in the case of Zn2+ the two 4s electrons are removed to maintain the full d shell - so we have [Ar] 3d10 instead of [Ar] 4s2 3d8 74 3 The Periodic Table The Periodic Table of the Elements (PTE) 75 3 The Periodic Table 76 3 The Periodic Table The Periodic Table of the Elements Development of the PTE 77 3 The Periodic Table The Periodic Table of the Elements Dmitri Mendeleev (1834-1907) & Lothar Meyer (1830-1895): 1869: both (independently) presented the first „PTE´s“ Elements were listed according to increasing atomic mass Chemical and physical properties are repeated periodically Meyer Mendeleev 78 3 The Periodic Table The Periodic Table of the Elements 79 3 The Periodic Table The Periodic Table of the Elements Henry Mosley (1887-1915): Measured X-rays emitted by atoms after bombardment with high- energy electrons Observed that the frequency (energy) of the X-rays increased with increasing atomic mass Arranged X-ray frequencies in order by assigning a unique number called atomic number Atomic number was identified as number of protons in the nucleus The effective nuclear charge Zeff is defined as: Zeff = Z- S Wherey S is the so called „shielding constant“ 80 3 The Periodic Table The Periodic Table of the Elements 81 3 The Periodic Table The Periodic Table of the Elements 82 „blocks“ based on the valence electrons 3 The Periodic Table Electron configuration For an element its electron configuration can be written as: for example Sulfur (S): 1s22s22p63s23p4 For chemical reactions primarily the electron configuration of the outer shell (valence electrons) is of importance. As they have identical configuration in the outer shell, elements of the same group show similar chemical properties. For this reason it is common to use a shortened notation by adding the preceeding noble gas element in square brackets For our example this would be Neon (Ne) with an electron configuration of 1s22s22p6 So we can write for Sulfur (S): [Ne]3s23p4 83 3 The Periodic Table Periodic properties of the elements Anions are generally bigger than neutral atoms as the additional negative charge leads to expansion of the electron shell (increased electron-electron repulsion) Cations are distinctly smaller than neutral atoms as the additional positive charge leads to compression of the electron shell 84 3 The Periodic Table Periodic properties of the elements 85 3 The Periodic Table Periodic properties of the elements The first ionization energy of an element is the minimum energy (in kJ/mol) required to remove the first electron from an isolated gaseous atom thereby forming a cation: In the same way also a second electron can be removed (whereby the second ionization energy is usually higher than the first one: Noble gases are hard to ionize Generally elements at the bottom left of the periodic table (e.g. Cs) easily release an electron, elements at the top right are hard to ionize (F) 86 3 The Periodic Table Periodic properties of the elements Ionization energies show an opposite trend compared to atomic radii – the larger the radius, the lower the first ionization energy Ionization energies decrease within a group as the electrons are farther away from the nucleus Ionization energies increase within a period 87 3 The Periodic Table Periodic properties of the elements Differences between the n and (n+1) ionization energy for selected elements: Na - removal of second electron requires a lot more energy Na+ Mg - removal of third electron requires a lot more energy Mg++ Al - removal of fourth electron requires a lot more energy Al+++ 88 3 The Periodic Table Periodic properties of the elements The electron affinity specifies the energy change observed when an electron is added to an isolated gaseous atom forming an anion. There is not such a clear trend as for ionization energies but in general elements situated in the upper right corner take up electrons easier than those on the lower left corner of the PTE In contrast to the situation for ionization energies, for most elements this is an exothermic reaction (value 10-6 mol∙l-1) → 𝐊 𝐚 ⋅ 𝐊 𝐖 can be neglected [𝐇𝟑 𝐎+ ]𝟐 +𝐊 𝐚 ⋅ 𝐇𝟑 𝐎+ − 𝐊 𝐚 ⋅ 𝐇𝐀 𝟎 + 𝐊𝐖 = 𝟎 𝐊𝐚 𝐊 𝐇𝟑 𝐎+ ≈ − + 𝐊 𝐖 + 𝐊 𝐚 ⋅ [𝐇𝐀]𝟎 +( 𝐚 )𝟐 𝟐 𝟐 Autoprotolysis can be neglected (Guideline Ka∙[HA]0 > 100∙KW) + 𝐊𝐚 𝐊𝐚 𝟐 𝐇𝟑 𝐎 ≈ − + 𝐊𝐚 ⋅ [𝐇𝐀]𝟎 +( ) 𝟐 𝟐 240  8 Acids and Bases Calculating pH for mono- and polyprotic acids Monoprotic acids are capable to donate one proton. Polyprotic acids are capable to donate more than one proton. The dissociation takes place stepwise, every deprotonation step has its own Ka-value. H2CO3 + H2O ⇌ H3O+ + HCO3- Ka1 = 4.3 ∙ 10-7 HCO3- + H2O ⇌ H3O+ + CO32- Ka2 = 4.7 ∙ 10-11 H2SO4 + H2O ⇌ H3O+ + HSO4- Ka1 = 103 HSO4- + H2O ⇌ H3O+ + SO42- Ka2 = 1.2 ∙ 10-2 For the calculation of the pH-value it is often possible to neglect the dissociation of the second, third,… proton and the pH-value can be approximated by the dissociation of the first proton (min. factor of 1000 between the Ka-values). 241 8 Acids and Bases Strong bases and weak bases All considerations made for acids can also be applied to bases Examples for strong bases are: NaOH or KOH In water they are fully protonated so: B + H 20 BH+ + OH- Examples for weak bases are: acetate, pyridine and formate In water they are only partly protonated The dissociation constant Kb gives us an idea about the strength of a base: 242  8 Acids and Bases Kb values for some weak bases Often instead of the Kb the pKb value (= -log Kb) is used to describe the strength of a base! 243 From: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson  8 Acids and Bases Relationship between Ka and Kb Between the Ka and Kb (or pKa and pKb) values of conjugated acid/base pairs there is a relationship: Ka x Kb = Kw = 10-14 pKa + pKb = pKw = 14 244 From: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson  8 Acids and Bases Buffers Solutions containing relatively high concentrations (> 10-3 M) of a weak conjugate acid-base pair are called buffered solutions or simply buffers. They resist drastic changes in pH when small amounts of a strong acid or a strong base are added. A typical buffer solution can be obtained from a weak acid HA and a salt with the anion A-. HA + H2O D A- + H3O+ Typical buffers are made by mixing acetic acid with sodiumacetate or by mixing ammoniumchloride with aqueous ammonia solution. 245  8 Acids and Bases Buffers Coming back to the weak acid HA. The equilibrium in this case is further shifted to the left (undissociated form of HA) as additional A- is added (in the form of a salt). This means that [HA] equals approximately c°(HA) which is the initial concentration of HA. In the same way [A-] may be replaced by c°(A-). Of course we can write the acidiy constant for HA: K sA = A H O  − 3 + HA  246  8 Acids and Bases Buffers K sA = A H O  − 3 + HA  By substituting [HA] and [A-] with c°(HA) and c°(A-) respectively, followed by taking the logarithms we receive: c(A − ) pH = pK s(HA) + log A(HA) c(HA) c(korrespon dierende corresponding base Base) pH = pK s(HA) + log A(HA) c(Säure) acid This equation is called the Henderson-Hasselbalch equation 247  8 Acids and Bases Buffers From the Henderson-Hasselbalch equation we can deduct: Diluting the buffer should not lead to pH changes – as long as our assumption [HA] = c°(HA) and [A-] = c°(A-) is valid. The change in pH observed for a buffer that is diluted 1:1 with pure water is called „dilution effect“ pH changes due to the addition of strong acid or base should be much less pronounced as in the case of an addition of strong acid or base to pure water 1000 ml H2O + 100 ml HCl (1M): 0.1 moles of H3O+ in 1100ml pH = 1.04 DpH = 7 – 1.04 = 5.96 248  8 Acids and Bases Buffers If on the other hand we have 1000 ml of a buffered solution being 1M in acetic acid and 1M in sodiumacetate, the pH before addition of the acid (base) can be calculated using the Henderson-Hasselbalch equation: 1 pH = 4.75 + log = 4.75 1 If we add 100 ml of HCl (1M) the protons from the acid will convert acetate to acetic acid. So instead o a 1M acetate solution we will have a 0.9M acetate solution (if we consider the dilution to 1.1 L we get 0.818M). On the other hand the concentration of acetic acid is increased from 1 M to 1.1M (if we consider the dilution to 1.1 L we get 1M). Calculating the pH using the Henderson- Hasselbalch equation we receive: 0.818 𝑝𝐻 = 4.75 + log = 4.66 1 249  8 Acids and Bases Buffers So whereas in the case of water we had a DpH of 5.96 after adding 100 ml of a strong acid (1M) for the buffered solution it is only 0.09. The buffering capacity  of a buffer gives the number of moles of a strong acid that have to be added to a buffer to induce a pH change of 1. d c((acid) d c( Base) = − Säure ) = d pH d pH 250  8 Acids and Bases Buffers 251 From: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson  8 Acids and Bases Buffers The buffering capacity of a buffer is determined by: the initial concentration of acid and salt (the higher the better) the ratio of HA/A- whereby 1:1 is the optimum 252  8 Acids and Bases "We are so habituated to the use of water as a solvent, and our data are so frequently limited to those obtained in aqueous solutions, that we frequently define an acid or a base as a substance whose aqueous solution gives, respectively, a higher concentration of hydrogen ion or of hydroxide ion than that furnished by pure water. This is a very one sided definition …“ [Lewis, 1923] More details on the Lewis-concept will be provided in the Inorganic Chemistry lecture! 253  8 Acids and Bases Acid/base concepts discussed so far are based on proton transfer. Lewis was looking for a more general definition of acids and bases! A Lewis acid is an e- - pair acceptor A Lewis base is an e- -pair donor More details on the Lewis-concept will be provided in the Inorganic Chemistry lecture! 254  8 Acids and Bases Reaction of a B(OH)3 with water Upon reaction with water Lewis acids generate Bronsted acids! 255  8 Acids and Bases Reaction of oxides with water: Acidic: Basic: Amphoteric: 256  8 Acids and Bases Hard and Soft Acids and Bases (based on the Lewis-acid-base concept) Allows to estimate stability and reactivity of chemical compounds Separated in hard and soft acids/bases Not equivalent with strong and weak acids/bases!  8 Acids and Bases Hard acids/bases – small atomic/ionic-radius – high oxidation number – high charge density – low polarizability – but high tendency to induce polarization in other molecules Soft acids/bases – large atomic/ionic-radius – low oxidation number – low charge density – high polarizability – but low tendency to induce polarization on other molecules 8 Acids and Bases from: Nanoscale, 2018, 10, 5035  8 Acids and Bases HSAB-concept – reactivity Hard acids prefer to react with hard bases. They prefer to form ionic bonds. Soft acids react with soft bases. They prefer covalent bonding. If there are different possibilities for reactions not only the strong/weak character of an acid / base has to be considered but also the fact wether we have hard or soft acids/bases. Additionally steric factors may play a role.  HSAB-concept – application Stability – Solubility of salts Fluoride Chloride Bromide Iodide Natrium Sodium 1,0 5,4 7,2 8,4 Silver Silber 14,3 1,3∙10-5 7,2∙10-7 9,1∙10-9 – Acidity HF (3) < HCl (-6) < HBr (-9) < HI (-10) (pKa in parantheses) Faster than Reactivity Faster than 9 Short Introduction to Coordination Chemistry Short Introduction to Coordination Chemistry 262  9 Short Introduction to Coordination Chemistry Definition: Species that are assemblies of a central (transition-metal) ion bonded to a group of surrounding molecules or ions, such as [Ag(NH3)2]+ and [Fe(H2O)6]3+, are called metal complexes, or merely complexes. If the complex carries a net charge, it is generally called a complex ion. 1. Lutz H. Gade (1998): Koordinationschemie, Weinheim: Wiley-VCH 263  9 Short Introduction to Coordination Chemistry Central atom (ion): Lewis-acid; often transition metal, lanthanoides, actinoides; cationic or neutral. Ligand: Lewis-base; either atom(s) or molecules/molecular ions; mostly negatively charged or neutral, rarely cationic. Coordinative bond: also donor-acceptor-bond; in this case both electrons used for bonding come from one side - here mostly from the ligand that possesses free electron pairs. Coordination number: number of ligands coordinated by a central atom. 264  9 Short Introduction to Coordination Chemistry Donor-atom: the atom of the ligand binding to the central atom – so as an example for [Ag(NH3)2]+ this would be the N. Bridging-ligand: ligand connecting more than one central atom. Mono/bi/poly-dentate ligands: ligands that have one/two/several donor atoms that can bind to the central atom. 265 From: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson 9 Short Introduction to Coordination Chemistry Complex with a single central atom Multi-nuclear complex with more than one central atom which are either connected by metal-metal bonds or bridging ligands 266  9 Short Introduction to Coordination Chemistry A complex with poly-dentate ligands that binds with at least two donor atoms to the central atom shows the so called chelate – effect. This means it has a higher constant of formation and shows increased stability. Reasons for the increased stability of these complexes are: ◦ The less pronounced change in entropy during complex formation ◦ The fact that all bonds need to be released before the ligand is set free for any alternative reactions 267  9 Short Introduction to Coordination Chemistry Important complexes with chelate effect: EDTA Enzymes ◦ Six-dentate ligand ◦ Forms very stable 1:1 complexes with metal-cations having a charge of +2 or higher ◦ Consumption: ~40.000 t/year in EU & Norway ◦ Use: CO2 + H2O ⇌ HCO3- + H3O+ ▪ Textiles- and paper-industry, cleaning, conservation, analytical chemistry, detergents…. 268 10 Complex/Coupled Equilibria Complex/Coupled Equilibria 269  10 Complex/Coupled Equilibria Reaction with complexing agents: M (aq) + n L (aq) ⇌ MLn (aq) [Cu(H2O)4]2+ (aq) + 4 NH3 (aq) ⇌ [Cu(NH3)4]2+ (aq) + 4 H2O (l) 270 From: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson  10 Complex/Coupled Equilibria Ligands are added step by - step: [𝑀𝐿] M (aq) + L (aq) ⇌ ML (aq) 𝐾1 = 𝑀 ⋅[𝐿] [𝑀𝐿2 ] ML (aq) + L (aq) ⇌ ML2 (aq) 𝐾2 = 𝑀𝐿 ⋅[𝐿] … [𝑀𝐿𝑛 ] MLn-1 (aq) + L (aq) ⇌ MLn (aq) 𝐾𝑛 = 𝑀𝐿𝑛−1 ⋅[𝐿] For every step we can give the complex formation constant 𝑲𝒊. In general we have: 𝐾1 > 𝐾2 > … > 𝐾𝑛 271  10 Complex/Coupled Equilibria [𝑀𝐿] M (aq) + L (aq) ⇌ ML (aq) 𝐾1 = 𝑀 ⋅[𝐿] [𝑀𝐿2 ] ML (aq) + L (aq) ⇌ ML2 (aq) 𝐾2 = 𝑀𝐿 ⋅[𝐿] … [𝑀𝐿𝑛 ] MLn-1 (aq) + L (aq) ⇌ MLn (aq) 𝐾𝑛 = 𝑀𝐿𝑛−1 ⋅[𝐿] The different steps can be added – so we get the cumulative constant 𝜷𝒏 : [𝑴𝑳𝒏 ] M (aq) + n L (aq) ⇌ MLn (aq) 𝜷𝒏 = 𝑴 ⋅[𝑳]𝒏 Whereby: 𝛽𝑛 = 𝐾1 ⋅ 𝐾2 ⋅ … ⋅ 𝐾𝑛 272  10 Complex/Coupled Equilibria We can also formulate a complex-dissociation 𝑲𝒅 : 𝑀 ⋅[𝐿] ML (aq) ⇌ M (aq) + L (aq) 𝐾𝑑 = [𝑀𝐿] 1 Thereby we can write: 𝐾𝑑 = = 𝐾𝑖−1 𝐾𝑖 Compare with the analogous constant for acids 𝑲𝒂 : 𝐻 + ⋅[𝐴− ] HA (aq) ⇌ H+ (aq) + A- (aq) 𝐾𝑎 = [𝐻𝐴] 273  10 Complex/Coupled Equilibria Equilibria interacting with solubility For more information on solutions see chapter 7 on „Solutions“ Solubility C: The maximum amount of a substance that can be dissolved in a certain volume of solvent at given conditions (temperature…). Saturated solution: A solution where the solute forms an equilibrium with undissolved solute. Solubility equilibria are heterogeneous systems where a solid precipitate forms an equilibrium with a saturated solution. Sparingly soluble salt: show only very limited solubility in water. The non- dissolved part forms a solid precipitate whereby the dissolved part is fully dissociated 274  10 Complex/Coupled Equilibria The Equilibrium constant 𝑲𝒔𝒑 (𝐾𝑠𝑝 ) is called solubility product constant or simply solubility product. The solubility product informs us how well a salt is soluble in water. In general for a salt AmBn (z = 1,2,3,4): AmBn (s) ⇌ m Azn+ (aq) + n Bzm- (aq) 𝑨𝒛𝒏+ 𝒎 𝑩𝒛𝒎− 𝒏 𝑲= [𝑨𝒎 𝑩𝒏 ] 𝑨𝒎 𝑩𝒏 = 𝒄𝒐𝒏𝒔𝒕. 𝒃𝒆𝒄𝒂𝒖𝒔𝒆 𝒊𝒕 𝒊𝒔 𝒂 𝒔𝒐𝒍𝒊𝒅 ⇒ [𝑨𝒎 𝑩𝒏 ] =1 𝑲𝒔𝒑 𝑨𝒎 𝑩𝒏 = 𝑨𝒛𝒏+ 𝒎 𝑩𝒛𝒎− 𝒏 𝒂𝒏𝒅 𝒑𝑲𝒔𝒑 = −𝒍𝒐𝒈(𝑲𝒔𝒑 ) 275  10 Complex/Coupled Equilibria From the solubility product we can calculate the solubility (C): AmBn (s) ⇌ m Azn+ (aq) + n Bzm- (aq) 𝑲𝒔𝒑 (𝑨𝒎 𝑩𝒏 ) = [𝑨𝒛𝒏+ ]𝒎 [𝑩𝒛𝒎− ]𝒏 𝑨𝒛𝒏+ = 𝒎 ⋅ 𝑪 𝑨𝒎 𝑩𝒏 and 𝑩𝒛𝒎− = 𝒏 ⋅ 𝑪 𝑨𝒎 𝑩𝒏 𝒏 𝑲𝒔𝒑 𝑨𝒎 𝑩𝒏 = 𝑨𝒛𝒏+ 𝒎 𝑩𝒛𝒎− 𝒏 = (𝒎 ⋅ 𝑪 𝑨𝒎 𝑩𝒏 )𝒎 ⋅ 𝒏 ⋅ 𝑪 𝑨𝒎 𝑩𝒏 𝑲𝒔𝒑 𝑨𝒎 𝑩𝒏 = 𝒎𝒎 ⋅ 𝒏𝒏 ⋅ 𝑪 𝑨𝒎 𝑩𝒏 𝒎+𝒏 𝒎+𝒏 𝑲𝒔𝒑 𝑨𝒎 𝑩𝒏 𝑪 𝑨𝒎 𝑩𝒏 = 𝒎𝒎 ⋅ 𝒏𝒏 276  10 Complex/Coupled Equilibria Example: which salt is better soluble: AgI (𝐾𝑠𝑝 = 8,5 ⋅ 10−17 ) or BiI3 (𝐾𝑠𝑝 = 7,7 ⋅ 10−22 )? AgI (s) ⇌ Ag+ (aq) + I- (aq) 𝐾𝑠𝑝 𝐴𝑔𝐼 = 𝐴𝑔+ ⋅ [𝐼 − ] 𝐴𝑔+ = [𝐼 − ] 𝐶 𝐴𝑔𝐼 = 𝐾𝑠𝑝 𝐴𝑔𝐼 = 9,2 ⋅ 10−9 𝑚𝑜𝑙 ⋅ 𝑙 −1 277 10 Complex/Coupled Equilibria Example: which salt is better soluble: AgI (𝐾𝑠𝑝 = 8,5 ⋅ 10−17 ) or BiI3 (𝐾𝑠𝑝 = 7,7 ⋅ 10−22 )? BiI3 (s) ⇌ Bi3+ (aq) + 3 I- (aq) 𝐾𝑠𝑝 𝐵𝑖𝐼3 = 𝐵𝑖 3+ ⋅ [𝐼− ]3 3 ⋅ 𝐵𝑖 3+ = [𝐼 − ] 4 𝐾𝑠𝑝 𝐵𝑖𝐼3 𝐶 𝐵𝑖𝐼3 = = 1,3 ⋅ 10−5 𝑚𝑜𝑙 ⋅ 𝑙 −1 27 278  10 Complex/Coupled Equilibria Example: which salt is better soluble: AgI (𝐾𝑠𝑝 = 8,5 ⋅ 10−17 ) or BiI3 (𝐾𝑠𝑝 = 7,7 ⋅ 10−22 )? AgI (s) ⇌ Ag+ (aq) + I- (aq) 𝐾𝑠𝑝 𝐴𝑔𝐼 = 𝐴𝑔+ ⋅ [𝐼 − ] 𝐴𝑔+ = [𝐼 − ] 𝐶 𝐴𝑔𝐼 = 𝐾𝑠𝑝 𝐴𝑔𝐼 = 9,2 ⋅ 10−9 𝑚𝑜𝑙 ⋅ 𝑙 −1 BiI3 (s) ⇌ Bi3+ (aq) + 3 I- (aq) 𝐾𝑠𝑝 𝐵𝑖𝐼3 = 𝐵𝑖 3+ ⋅ [𝐼− ]3 3 ⋅ 𝐵𝑖 3+ = [𝐼 − ] 4 𝐾𝑠𝑝 𝐵𝑖𝐼3 𝐶 𝐵𝑖𝐼3 = = 1,3 ⋅ 10−5 𝑚𝑜𝑙 ⋅ 𝑙 −1 27 279 10 Complex/Coupled Equilibria Common Ion effect Addition of a well soluble salt ▪ That has a common ion with the sparingly soluble salt ▪ Reduces the concentration of the counter ion in solution! ▪ Example: we add 𝐼 − = 0,1 𝑚𝑜𝑙 ⋅ 𝑙 −1 (e.g. as NaI) to a saturated solution of AgI or BiI3 (compare results with previous slides) 𝐾𝑠𝑝 𝐴𝑔𝐼 𝐾𝑠𝑝 𝐴𝑔𝐼 = 𝐴𝑔+ ⋅ 𝐼− ⇒ 𝐴𝑔+ = [𝐼 − ] as the amount of I- set free by the dissolution of AgI or BiI3 is far less than 0.1 mol L-1 we can set the I- concentration to 0.1 mol L-1 𝐶 𝐴𝑔𝐼 = 𝐴𝑔+ = 8,5 ⋅ 10−16 𝑚𝑜𝑙 ⋅ 𝑙 −1 3+ − 3 𝐾𝑆𝑃 𝐵𝑖𝐼3 3+ 𝐾𝑆𝑃 𝐵𝑖𝐼3 = 𝐵𝑖 ⋅ [𝐼 ] ⇒ 𝐵𝑖 = [𝐼 − ]3 𝐶 𝐵𝑖𝐼3 = 𝐵𝑖 3+ = 7,7 ⋅ 10−19 𝑚𝑜𝑙 ⋅ 𝑙 −1 280  10 Complex/Coupled Equilibria What happens if further equilibria are of importance? 2 sparingly soluble salts with a common ion ▪ AgBr (s) ⇌ Ag+ (aq) + Br- (aq) 𝑲𝒔𝒑,𝟏 ▪ AgSCN (s) ⇌ Ag+ (aq) + SCN- (aq) 𝑲𝒔𝒑,𝟐 Sparingly soluble salt & complex formation ▪ AgCl (s) ⇌ Ag+ (aq) + Cl- (aq) 𝑲𝒔𝒑 ▪ Ag+ (aq) + 2 NH3 (aq) ⇌ [Ag(NH3)2]+ (aq) 𝜷𝟐 Sparingly soluble salt & acid/base-equilibrium ▪ PbF2 (s) ⇌ Pb2+ (aq) + 2 F- (aq) 𝑲𝒔𝒑 ▪ HF (aq) + H2O ⇌ H3O+ (aq) + F- (aq) 𝑲𝒂 … → all equilibria have to be considered! → coupled equilibria 281  10 Complex/Coupled Equilibria Common Ion effect – two sparingly soluble salts: AgBr (s) ⇌ Ag+ (aq) + Br- (aq) 𝐾𝑠𝑝 𝐴𝑔𝐵𝑟 = 5 ⋅ 10−13 = 𝐴𝑔+ ⋅ 𝐵𝑟 − AgSCN (s) ⇌ Ag+ (aq) + SCN- (aq) 𝐾𝑠𝑝 𝐴𝑔𝑆𝐶𝑁 = 1 ⋅ 10−12 = 𝐴𝑔+ ⋅ 𝑆𝐶𝑁 − 𝐴𝑔+ = 𝐵𝑟 − + 𝑆𝐶𝑁 − 𝐾𝑠𝑝 𝐴𝑔𝐵𝑟 𝐴𝑔+ ⋅ 𝐵𝑟 − = + − = 0.5 ⇒ 𝑆𝐶𝑁 − = 2 ⋅ 𝐵𝑟 − ⇒ 𝐴𝑔+ = 𝐵𝑟 − + 2 ⋅ 𝐵𝑟 − 𝐾𝑠𝑝 𝐴𝑔𝑆𝐶𝑁 𝐴𝑔 ⋅ 𝑆𝐶𝑁 𝐾𝑠𝑝 𝐴𝑔𝐵𝑟 = 5 ⋅ 10−13 = ( 𝐵𝑟 − + 2 ⋅ 𝐵𝑟 − ) ⋅ 𝐵𝑟 − ⇒ 𝐵𝑟 − = 4.1 ⋅ 10−7 𝑚𝑜𝑙 ⋅ 𝑙−1 ⇒ 𝐴𝑔+ = 12.2 ⋅ 10−7 𝑚𝑜𝑙 ⋅ 𝑙−1 and 𝑆𝐶𝑁 − = 8.2 ⋅ 10−7 𝑚𝑜𝑙 ⋅ 𝑙−1 ⇒ 𝐶(𝐴𝑔𝐵𝑟) = 4.1 ⋅ 10−7 𝑚𝑜𝑙 ⋅ 𝑙−1 and 𝐶(𝐴𝑔𝑆𝐶𝑁) = 8.2 ⋅ 10−7 𝑚𝑜𝑙 ⋅ 𝑙−1 S𝑜𝑙𝑢𝑏𝑖𝑙𝑖𝑡𝑦 𝑜𝑓 𝑡ℎ𝑒 𝑠𝑎𝑙𝑡𝑠: 𝐶 𝐴𝑔𝐵𝑟 = 7.1 ⋅ 10−7 𝑚𝑜𝑙 ⋅ 𝑙−1 , 𝐶(𝐴𝑔𝑆𝐶𝑁) = 10.0 ⋅ 10−7 𝑚𝑜𝑙 ⋅ 𝑙−1 → Solubility is reduced for both – but the effect is more pronounced for the less soluble salt! Effect: 58% for AgBr and 82% for AgSCN 282  10 Complex/Coupled Equilibria Common Ion effect – two sparingly soluble salts: AgBr (s) ⇌ Ag+ (aq) + Br- (aq) 𝐾𝑠𝑝 𝐴𝑔𝐵𝑟 = 5 ⋅ 10−13 = 𝐴𝑔+ ⋅ 𝐵𝑟 − 1:200 AgCl (s) ⇌ Ag+ (aq) + SCN- (aq) 𝐾𝑠𝑝 𝐴𝑔𝐶𝑙 = 1 ⋅ 10−10 = 𝐴𝑔+ ⋅ 𝐶𝑙− 𝐴𝑔+ = 𝐵𝑟 − + 𝐶𝑙− 𝐾𝑠𝑝 𝐴𝑔𝐵𝑟 𝐴𝑔+ ⋅ 𝐵𝑟 − = + − = 0.005 ⇒ 𝐶𝑙− = 200 ⋅ 𝐵𝑟 − ⇒ 𝐴𝑔+ = 𝐵𝑟 − + 200 ⋅ 𝐵𝑟 − 𝐾𝑠𝑝 𝐴𝑔𝐶𝑙 𝐴𝑔 ⋅ 𝐶𝑙 𝐾𝑠𝑝 𝐴𝑔𝐵𝑟 = 5 ⋅ 10−13 = ( 𝐵𝑟 − + 200 ⋅ 𝐵𝑟 − ) ⋅ 𝐵𝑟 − ⇒ 𝐵𝑟 − = 5.0 ⋅ 10−8 𝑚𝑜𝑙 ⋅ 𝑙−1 ⇒ 𝐴𝑔+ = 1.0 ⋅ 10−5 𝑚𝑜𝑙 ⋅ 𝑙−1 and 𝐶𝑙− = 1.0 ⋅ 10−6 𝑚𝑜𝑙 ⋅ 𝑙−1 ⇒ 𝐶(𝐴𝑔𝐵𝑟) = 5.0 ⋅ 10−8 𝑚𝑜𝑙 ⋅ 𝑙−1 and 𝐶(𝐴𝑔𝐶𝑙) = 1.0 ⋅ 10−5 𝑚𝑜𝑙 ⋅ 𝑙−1 S𝑜𝑙𝑢𝑏𝑖𝑙𝑖𝑡𝑦 𝑜𝑓 𝑡ℎ𝑒 𝑠𝑎𝑙𝑡𝑠: 𝐶 𝐴𝑔𝐵𝑟 = 7.1 ⋅ 10−7 𝑚𝑜𝑙 ⋅ 𝑙−1 , 𝐶(𝐴𝑔𝐶𝑙) = 1.0 ⋅ 10−5 𝑚𝑜𝑙 ⋅ 𝑙−1 → Solubility is substantially reduced for the less soluble salt! Effect: 7% for AgBr and 100% for AgCl 283 10 Complex/Coupled Equilibria Solubility influenced by a complexation reaction: Sparingly soluble salt plus ligand forming a complex with at least one of the ions: AgCl (s) ⇌ Ag+ (aq) + Cl- (aq) 𝑲𝒔𝒑 = 𝑨𝒈+ ⋅ [𝑪𝒍− ] [𝑨𝒈(𝑵𝑯𝟑 )+ 𝟐] Ag+ (aq) + 2 NH3 (aq) ⇌ [Ag(NH3)2]+ (aq) 𝜷𝟐 = 𝑨𝒈 ⋅[𝑵𝑯𝟑 ]𝟐 + ◦ Ag+ (from dissolved AgCl) is present as Ag+ (aq), [Ag(NH3)]+ (aq) and [Ag(NH3)2]+ (aq) and because of electroneutrality: [Cl-] = [Ag+] + [[Ag(NH3)]+] + [[Ag(NH3)2]+] ◦ but only Ag+ (aq) is relevant for the solubility equilibrium ◦ The concentration of Ag+ (aq) will be decreased upon addition of NH3 so more of the AgCl can dissolve → solubility is increased 284  10 Complex/Coupled Equilibria 285 From: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson  10 Complex/Coupled Equilibria In general, the solubility of a compound containing a basic anion (that is, the anion of a weak acid) increases as the solution becomes more acidic. Solubility influenced by pH If at least one of the ions is involved in an acid/base equilibrium (ammonium, hydroxide, carbonate, oxalate,…) PbF2 (s) ⇌ Pb2+ (aq) + 2 F- (aq) 𝑲𝒔𝒑 𝑷𝒃𝑭𝟐 = 𝑷𝒃𝟐+ ⋅ 𝑭− 𝟐 𝑯𝟑 𝑶+ ⋅[𝑭− ] HF (aq) + H2O ⇌ H3O+ (aq) + F- (aq) 𝑲𝒂 (𝑯𝑭) = [𝑯𝑭] ◦ F- (dissolved from PbF2) is present as F- (aq) and HF (aq) and due to electroneutrality we have [Pb2+] = ½ · ([F-] + [HF]) ◦ only F- (aq) is relevant for the solubility equilibrium ◦ The concentration of F- is decreased upon decrease of the pH (as more F- will be protonated) ◦ So more PbF2 will be dissolved → the solubility increases The solubility of salts containing a (moderate-weakly) basic anion increases with decreasing pH! The solubility of salts containing a (moderate-weakly) acidic cation 286 increases with increasing pH!  10 Complex/Coupled Equilibria From: Theodore L. Brown, H. Eugene LeMay, Bruce E. Bursten, Paula Y. Bruice (2014): Chemistry the central science, Pearson 287 11 Thermochemistry and Fundamentals of Chemical Thermodynamics Thermochemistry and Fundamentals of Thermodynamics 288 11 Thermochemistry and Fundamentals of Chemical Thermodynamics The „0“ Law of Thermodynamics When A is in thermal equilibrium with B and B is in thermal equilibrium with C also A is in thermal equilibrium with C 289 11 Thermochemistry and Fundamentals of Chemical Thermodynamics The first Law of Thermodynamics Energy can be converted from one form to another – it can neither be created nor destroyed. In any process in an isolated system, the total energy remains the same Isolated system: no exchange of energy or matter with the environment Closed system: only energy can be exchanged with the environment Open system: both energy and matter can be exchanged with the environment 290  11 Thermochemistry and Fundamentals of Chemical Thermodynamics DE