General Chemistry CHG101 PDF

General Chemistry CHG101 (2025) Introduction to Chemistry (Terms 1 & 2) by Dr Richard Betz 2025 General Chemistry CHG101 (2025) Introduction to Chemistry (Terms 1 & 2) by Dr...

General Chemistry CHG101 (2025) Introduction to Chemistry (Terms 1 & 2) by Dr Richard Betz 2025 General Chemistry CHG101 (2025) Introduction to Chemistry (Terms 1 & 2) by Dr Richard Betz 2025 Contents List of Figures v List of Tables vi 1 Chemistry – A Science in Context 1 1.1 General Remarks................................ 1 1.1.1 Chemistry and its Brothers and Sisters................ 1 1.1.2 Chemistry – From Past to Present.................. 2 1.2 Some Basic Concepts.............................. 4 1.2.1 Accuracy and Precision........................ 4 1.2.2 The SI System............................. 4 1.2.3 Logarithms and Significant Figures.................. 7 2 Getting to the Core – Atomic Theory 11 2.1 A Closer Look at Matter............................ 11 2.1.1 States of Aggregation and Phase Transitions............ 11 2.1.2 Mixtures and Separation Techniques................. 12 2.1.3 Properties of Pure Compounds.................... 15 2.2 Atomic Theory................................. 16 2.2.1 An Old Concept – Historic Roots................... 16 2.2.2 Masses, Masses, Masses........................ 17 2.2.2.1 Conservation of Mass.................... 17 2.2.2.2 Law of Constant Proportions................ 18 2.2.2.3 Law of Multiple Proportions................ 19 2.2.3 Dalton’s and Rutherford’s Atomic Model.............. 22 2.2.4 The Basic Construction Plan of an Atom.............. 23 2.2.5 Atomic Masses............................. 25 iv List of Figures 1.1 The Arabian alchemist Jabir ibn Hayyan................. 2 1.2 The alchemists Henning Brand, Andreas Libavius and the chemist Antoine Lavoisier and his wife Marie-Anne............... 2 1.3 Accuracy and precision explained graphically................ 4 2.2 Phase transitions................................ 11 2.1 The phase diagrams of carbondioxide and water............... 12 2.3 The family tree of substances......................... 16 2.4 Demokrit, Isaac Newton and Aristotle................ 17 2.5 John Dalton, George Stoney, Joseph Thomson and Ernest Ruther- ford....................................... 22 2.6 James Chadwick............................... 23 v List of Tables 1.1 The si base units................................ 5 1.2 Conversion factors used in the metric system................ 6 1.3 Roman numerals................................ 6 2.1 Heterogeneous mixtures............................ 13 2.2 Mercuryoxide decomposition.......................... 18 2.3 Silveroxide decomposition........................... 18 2.4 Carbonoxide decomposition (Series A).................... 19 2.5 Carbonoxide decomposition (Series B).................... 19 2.6 Sulfurfluoride decomposition (Series A)................... 20 2.7 Sulfurfluoride decomposition (Series B).................... 20 2.8 Sulfurfluoride decomposition (Series C)................... 20 2.9 Relation between the two series of carbonoxide decomposition....... 21 2.10 Relation between the three series of sulfurfluoride decomposition..... 21 2.11 Important properties of atomic building blocks............... 23 2.12 Relative atomic masses based on hydrogen.................. 25 2.13 Selected relative atomic masses based on different reference points.... 26 2.14 Masses of atomic building blocks....................... 26 vi 1 Chemistry – A Science in Context 1.1 General Remarks 1.1.1 Chemistry and its Brothers and Sisters ince the dawn of mankind, humans have been observers of natural phenomena. S Born out of necessity or plain curiosity, simple observation quickly gave rise to identifying patterns, explaining occurrences and, best of all, making – more or less – reliable predictions into the future. Over time, this process allowed for the development of a series of natural sciences that – starting in the 17th , 18th and 19th century – have evolved as separate subjects such as botany, zoology, physiology, the geo-sciences, physics and – as the late-comer – chemistry. While all of these disciplines deal with nature, their key focus varies, for the few mentioned examples we would concentrate on: botany — looks at plants and vegetation, their systematics and characterization zoology — deals with animals, their systematics and characterization physiology — studies the human body, its set-up and processes enabling its exi- stence geo-sciences – focuses on the construction plan of our planet and the mechanisms that shape it physics — is concerned with the quantitative changes matter experiences, specifi- cally mass and energy chemistry — places its emphasis on the synthesis of new materials and the identi- fication of structure-property relationships Although most natural scientific fields are traditionally dealt with separately in a tertiary education context, significant overlaps in between the various subjects exist and necessitate a basic understanding across the board. As an example, the life-sustaining processes in plants, animals and the human body cannot be understood without chemi- stry, the forces behind earthquakes cannot be explained without physics, the motion of planets and stars could not be predicted without mathematics, etc. As a consequence, your first year in a science program will expose you to a multitude of disciplines – so please fasten your seatbelt! , 1  1 Chemistry – A Science in Context 1.1.2 Chemistry – From Past to Present ithout knowing it consciously, people have always "done chemistry". Burning W wood and coal, producing metals, cooking, baking, brewing beer – all this dates back until ancient Egypt and even neolithic times. The first people to really venture into chemistry – or better say its predecessor alchemy – were Arabian scholars in the 7th and 8th century (cf. Figure 1.1). Important basic chemical operations such as distilla- tions were first described in this period. Many terms used up to this day have their roots in the Arabic world such as ’alkaline’, ’alcohol’ or ’elixir’. Even the term ’alchemy’ itself has partially Arabic roots and roughly translates either to ’changing metals’ or ’black soil’. Although alchemy was very closely related to modern-day chemistry, it was not a real natural science as it entailed many elements of magic and witchcraft that sought to solve the alchemists’ main quest – the Philosopher’s Stone. The latter is an elusive compound shrouded in legend that was said to have the power Figure 1.1: The Arabian al- to dissolve all compounds it comes into contact with, chemist Jabir ibn Hayyan. turn all metals into gold and, if applied to humans, pro- Source of photo: Wikipedia vide never-ending youthfulness and eternal life. Irrespec- (user: Halfdan). tive of the contradictory properties of the Philosopher’s Stone it is needless to point out that alchemists never succeeded in finding it. But on their way, they discovered many working techniques, synthesis methods and described the properties of chemical compounds (cf. Figure 1.2). The transition of alchemy to chemistry as a real natural science occurred during the second half of the 18th century by the introduction of quantitative measurements by Antoine Lavoisier (cf. Figure 1.2). Figure 1.2: The alchemist Henning Brand (left) and Andreas Libavius (centre) in their laboratories. Antoine Lavoisier and his wife Marie-Anne (right). Source of photos: Wikipedia (users: File Upload Bot (Magnus Manske), Janericloebe, Soerfm). 2  1 Chemistry – A Science in Context Over the decades and few centuries chemistry has been around as an independent na- tural science it has – as all other major sciences – experienced significant diversifications and split-ups in smaller sub-disciplines whose names you will have encountered many times already and are prone to bump into over and over again throughout your studies, professional life as well as in news stories and the media in general. Just as an appetizer – a short look at what is out there: 1. General Chemistry — the core of this module. It constitutes a basic but com- prehensive outlook about topics ranging from inorganic, organic and analytical to physical chemistry 2. Inorganic Chemistry — deals with all compounds except hydrocarbons, i.e. com- pounds featuring bonds between carbon and hydrogen (few exceptions) 3. Organic Chemistry — deals with hydrocarbons and compounds derived thereof 4. Physical Chemistry — looks at the quantitative and energetic changes entailed in the wake of chemical reactions as well as the physical principles guiding the "construction plan" of matter 5. Analytical Chemistry — studies the quantitative and qualitative composition of compounds and their spectroscopic characteristics 6. Biochemistry — focuses on the processes that are found in living creatures 7. Radiochemistry — deals with the disintegration and transmutation of elements as well as compounds featuring radioactive atoms ("Hot Chemistry") 8. Macromolecular Chemistry — is concerned with molecules of high molecular mass, natural and artificial (in the latter case also often referred to as ’Polymer Chemi- stry’) 9. Theoretical Chemistry — seeks to describe matter and chemistry "as such" by means of mathematical methods 10. Computational Chemistry — applies established mathematical methods to predict the properties of compounds by means of simulations 11. Technical Chemistry — focuses on the upscaling of reactions, procedures and equip- ment from a laboratory scale to bigger industrial dimensions ("Industrial Chem- istry") 12. Pharmaceutical Chemistry — specializes in the synthesis of new drugs applying structure-property relationships While the second, third and fourth mentioned fields constitute the "classical split" in chemistry – that is also reflected by the Department of Chemistry at nmu – the boundaries are, by far, not strict. Effectively, even a hardcore organic chemist has to 3  1 Chemistry – A Science in Context Figure 1.3: Accuracy and precision explained graphically. From left to right: 1. accurate (right) = correct and precise; 2. not correct but precise; 3. correct but not precise; 4. not accurate (not right) = neither correct nor precise. understand and use principles and techniques from inorganic, physical and analytical chemistry. Therefore, stereotyped thinking within Chemistry is as meaningful as it is in Science in general – not at all. 1.2 Some Basic Concepts Each natural science is characterized by the fact that it can be approached in a quan- titative manner, i.e. one can make numerical statements about amounts – e.g. masses, energy, etc. – based on (reproducible) measurements. As mentioned in Section 1.1.2 on page 2, it was this quantifying step taken by Lavoisier – who introduced the concise use of a balance in the laboratory – that uplifted non-scientific alchemy to the exact natural science of chemistry. Therefore, a short overview over a couple of important basic terms shall be given. Very likely, you will be familiar with most of them. 1.2.1 Accuracy and Precision ll measurements conducted in natural sciences and real life bear a small, ran- A dom error caused by factors beyond control or systematic errors. Repeating a measurement multiple times and averaging the results can be a way to cancel out the random errors. To judge the quality of a series of such measurements one can use the terms ’accurate’ and ’precise’. Looking at the shooting results on a Bull’s Eye as shown in Figure 1.3 will explain the two terms graphically. You will encounter and discuss these two concepts in greater detail in Analytical Chemistry in Second Year. 1.2.2 The SI System ver time, natural scientists have looked at a multitude of phenomena and de- O duced the rules that allowed for their quantification. Although there is a seem- ingly endless number of properties that can be observed and treated in that way, it was found that there is only a limited number of properties one needs to really know and measure as all other properties can be derived from these. A set of seven basic properties – known as the si Units – was identified and is listed in Table 1.1 on page 5 4  1 Chemistry – A Science in Context Table 1.1: The si base units. Property Unit name Unit symbol length metre m mass kilogram kg time second s temperature Kelvin K electric current Ampère A amount of substance mole mol luminous intensity candela cd To handle larger spans in between big and small values, the metric system provides for suitable prefixes that can be used to address smaller and larger values with regards to the base unit. The most common ones are listed in Table 1.2 on page 6. Some care has to be taken when dealing with the conversion of derived properties as the example of length (in meters), area (in square meters) and volume (in cubic meters) shall illustrate: ×10 ×10 ×10 ×1000 mm −−→ cm −−→ dm −−→ m −−−−→ km ×100 ×100 ×100 ×1000000 mm2 −−−→ cm2 −−−→ dm2 −−−→ m2 −−−−−−→ km2 ×1000 ×1000 ×1000 ×1000000000 mm3 −−−−→ cm3 −−−−→ dm3 −−−−→ m3 −−−−−−−−→ km3 Reasoning behind this: mm2 = mm × mm. So if there are 10 mm in one cm, there are (10 × 10) mm2 in one cm2. Mind: 1 mL = 1 cm3 1 L = 1 dm3 Question: How many liters of water are there in 2 km3 of water? Nelson Mandela University grounds often have a speed limit of 30 km h−1. A student that needs to rush for a study session is really working the pedals and reaches a speed of 8 m s−1 at the control room on South Campus where a friendly po- lice officer is checking how fast vehicles are passing by. Would the student be in trouble? Convert the speed to a common unit (e.g. m s−1 to km h−1 ): 1 1 1 1000 km 1 × 3600 km 1m= km and 1 s = h =⇒ 1 m s−1 = 1 = (1) 1000 3600 3600 h 1 × 1000 h Filling in the known values allows for calculating the correct answer: 8 m s−1 = 28.8 km h−1 (2) 5  1 Chemistry – A Science in Context Table 1.2: Conversion factors used in the metric system. Prefix Symbol Multiplication factor Yotta Y 1024 = 1 000 000 000 000 000 000 000 000 Zetta Z 1021 = 1 000 000 000 000 000 000 000 Exa E 1018 = 1 000 000 000 000 000 000 Peta P 1015 = 1 000 000 000 000 000 Tera T 1012 = 1 000 000 000 000 Giga G 109 = 1 000 000 000 Mega M 106 = 1 000 000 kilo k 103 = 1 000 hecto h 102 = 100 deca da 101 = 10 Base 100 = 1 deci d 10−1 = 0.1 centi c 10−2 = 0.01 milli m 10−3 = 0.001 micro µ 10−6 = 0.000 001 nano n 10−9 = 0.000 000 001 pico p 10−12 = 0.000 000 000 001 femto f 10−15 = 0.000 000 000 000 001 atto a 10−18 = 0.000 000 000 000 000 001 zepto z 10−21 = 0.000 000 000 000 000 000 001 yocto y 10−24 = 0.000 000 000 000 000 000 000 001 Table 1.3: Roman numerals. 1 2 3 4 5 6 7 8 9 10 11 12 I II III IV V VI VII VIII IX X XI XII 6  1 Chemistry – A Science in Context Although all natural scientists are encouraged to exclusively use the si units as well as the prefixes provided by the metric system, several old and established units have survived up to the present day as "outdated but accepted" such as liters for volumes and Å for short distances. A similar situation is at hand for Roman numerals – although their use is discouraged in favour of their Arabic counterparts, many fields of application persist up to the present day. In the case of chemistry, Roman numerals as listed in Table 1.3 on page 6 are the preferred choice to indicate oxidation states. 1.2.3 Logarithms and Significant Figures iven the sometimes very large and sometimes miniscule amounts of properties G natural sciences have to deal with, a logarithmic way of expressing values is convenient. Just as a refresher from high school mathematics, a short glimpse on the most basic principles of this subject shall be given. loga (b) is the mathematical operation to answer the question: "Which number has to be used as the exponent for a to yield b?". An example could be: log2 (32) = 5 because 25 = 32. a is referred to as "the basis". As the logarithms for a basis of 10 are very important, the decadic logarithm is of- ten abbreviated as "lg" in many textbooks. Another important abbreviation is "ln" (natural logarithm or logarithmus naturalis) applying Euler’s figure – e – as the basis. For every type of logarithm, the following arithmetical rules apply (no specific basis was given as they apply to all(!) logarithms!): log(ab) = log(a) + log(b) log( ab ) = log(a) − log(b) log(ab ) = b × log(a) Although all natural sciences are exact sciences and strive for perfection in predicting qualitative and quantitative descriptions and predictions, they have retained their status as experimental sciences, i.e. measurements have to be made to obtain values. Different measurement methods offer different degrees of accuracy, therefore, calculations based on such values cannot be infinitely precise. It is common practice to limit numerical statements to the lowest number of significant figures given for any value in a specific problem. To determine the number of significant figures, the following rules apply: 1. all non-zero digits are significant 2. zeros between non-zero digits are significant 7  1 Chemistry – A Science in Context 3. zeros beyond the decimal point at the end of a number are significant 4. zeros preceding the first non-zero digit in a number are not significant Determine the number of significant figures (sf) in the following examples. 1. 43.7609 kg – 6 sf 2. 0.0000001 cm3 – 1 sf 3. 230 V – 3 sf (this one is a shitshow of epic proportions) 4. 0.01000 Å – 4 sf The rules as stated here are sufficient for this module, however, regardless of which science you pursue in the end, you will have to familiarize yourself with some additional rules surrounding these issues in greater depth. For chemistry, a more detailed descrip- tion of the rules to be used for calculations with significant figures as well as the statistics behind measurements will be presented to you in Analytical Chemistry in Second Year. You should have learned the following things from this section: 1. How chemistry fits the context of other natural sciences 2. What the historical roots of chemistry are 3. Which major sub-disciplines are dealt with in chemistry 4. How the quality of measurements is judged 5. Which (and how) si units are used 6. How to determine significant figures To practice you could answer the following questions: 1. Name several natural sciences and what their respective key focus of research is! 2. Why is the Philosopher’s Stone as described by the alchemists an impossible thing to obtain? 3. Name several sub-disciplines of chemistry and describe what they are dealing with! 4. Regarding the sub-disciplines of chemistry – would it be possible to re-group some of them together in a common field? 5. Name the seven basic si properties and their respective units! 8  1 Chemistry – A Science in Context 6. Determine the number of significant figures in each case: a) 12.560 kg b) 0.02 cm c) 0.10005 Å d) 2.000070 V e) 25.07 ◦ C f) 0.000000000000000000000000007 A 7. Convert the following quantities into the unit given in round brackets: a) 20 m ( pm) b) 12 g ( kg) c) 72 km h−1 ( m s−1 ) d) 75 g dL−1 ( kg mL−1 ) 8. Although the si system has been accepted as standard for scientific measurements in nearly all countries on this planet, old traditional units still persist in common day use, especially in former British colonies. Convert the following quantities into the correct si units: a) length of a piece of string: 15 inches (1 inch = 2.54 cm) b) amount of milk in a jug: 0.5 gallons (1 gallon = 3.785411784 L) c) mass of butter for a cake: 8 troy ounces (1 troy ounce = 31.1 g) d) pressure in a tire: 50 pound per square inch (1 pound = 0.45359237 kg) References: 1. Chemistry as Natural Science, Alchemy, History — A. Burrows, J. Holman, A. Parsons, G. Pilling, G. Price: Chemistry3 — introducing inorganic, organic and physical chemistry, Oxford University Press, 2009, 4–6. J. C. Kotz, P. M. Treichel, J. R. Townsend: Chemistry & Chemical Reac- tivity (International Edition), 8th Edition, Brooks/Cole, Cengage Learning, 2012, 335–343; 1059. http://en.wikipedia.org/wiki/Alchemy (last accessed: Jan 12th , 2025). http://en.wikipedia.org/wiki/Chemistry (last accessed: Jan 12th , 2025). 9  1 Chemistry – A Science in Context 2. si Units and Measurements — A. Burrows, J. Holman, A. Parsons, G. Pilling, G. Price: Chemistry3 — introducing inorganic, organic and physical chemistry, Oxford University Press, 2009, 7–10. J. C. Kotz, P. M. Treichel, J. R. Townsend: Chemistry & Chemical Reac- tivity (International Edition), 8th Edition, Brooks/Cole, Cengage Learning, 2012, 25–31; 35–39. http://en.wikipedia.org/wiki/International_System_Of_Units (last accessed: Jan 12th , 2025). 10 2 Getting to the Core – Atomic Theory 2.1 A Closer Look at Matter 2.1.1 States of Aggregation and Phase Transitions hemists deal with compounds, i.e. various types of substances. All substances C are comprised of matter , i.e. something that has mass and a volume. Despite the very large number of compounds known, they can be grouped into three major categories: 1. solids 2. liquids 3. gases For completeness, it should be pointed out that – several decades ago – a potential fourth state of aggregation has been described: the super-fluid one. Theoretically, each compound can exist in all three states of aggregation, depending on the tempe- rature and pressure the compound under consideration is maintained. A plot of the phase a compound is in versus the temperature as well as the pressure is referred to as phase diagram. As an example, the phase diagrams of carbondioxide and water are given in Figure 2.1 on page 12. For each pair of data for tempera- ture and pressure one can now quickly gas see in what phase the substance will be in. While increasing the pressure on co n solid water at constant temperature nd tio n tio en will result in turning it into liquid wa- ev ma s ma ap at bli ter the opposite holds true for carbon- ion or bli su at su dioxide. Heating both compounds at ion re melting a given pressure will eventually melt solid liquid the solid and, later at even higher tem- solidi!cation peratures, vaporize the liquid. At one specific set of pressure and tempera- Figure 2.2: The transitions in between the phases. ture, all three phases co-exist. This point is known as triple point. At high temperatures and pressures, the phase diagrams stop. This point is known as critical point and is the point from which one onwards there is no more difference between "liquid" and "gas". 11  2 Getting to the Core – Atomic Theory p / bar p / bar solid liquid solid liquid gas gas T / °C T / °C Figure 2.1: The phase diagrams of carbondioxide (left) and water (right). The green points mark the triple point and the blue points mark the critical point each. Question: How can there be a situation where there is no more difference between "liquid" and "gas"? The phase diagrams also indicate that a substance can change from each phase it melting H2 O is in to each other phase. The corresponding terms for these transitions are given in boiling H2 O Figure 2.2 on page 11. They are important and you should familiarize yourself with condensing H2 O them! 2.1.2 Mixtures and Separation Techniques t is important to point out that many substances encountered in nature constitute I mixtures of multiple compounds – which is not necessarily a bad thing as biological processes often necessitate such mixtures. Mixtures of compounds, in general, can occur between substances in all states of aggregation and can be subdivided in two large categories: 1. homogeneous mixtures — one cannot tell that the substance under investigation salt + H2 O is a mixture because it appears uniform 2. heterogeneous mixtures — one can tell that the substance under investigation is chalk + H2 O a mixture as different components are clearly visible While all homogeneous mixtures are called solutions – irrespective of which phases have been combined – a set of specific expressions has been created for heterogeneous mixtures. Table 2.1 on page 13 provides an important overview. Question: Is there not one combination missing in Table 2.1 on page 13? However, in order to identify the typical properties of specific compounds it is ne- cessary to obtain them in pure form. As a consequence, chemists have developed a vast 12  2 Getting to the Core – Atomic Theory Table 2.1: The technical terms used to identify heterogeneous mixtures. Phases mixed Name of mixture Practical example solid–solid mixture beach sand solid–liquid suspension mud solid–gas aerosol smoke liquid–liquid emulsion milk liquid–gas aerosol foam range of techniques exclusively aimed at separating mixtures of compounds into their respective constituents. If a heterogeneous mixture is at hand, one can simply resort to the following methods: 1. solid–solid — to separate mixtures of solids one can take advantage of... manual sorting – can be done if there are clearly visible differences. This chalk sticks method is very tedious but has been applied successfully by Pasteur who sorted crystals of tartaric acid manually under a microscope sifting – if the solid materials differ sufficiently in size one can simply use a beach sand sieve to separate them from one another magnets – if one compound is magnetic, one can simply use a magnet for sand + Fe separation electrostatics – solids can be charged electrostatically and then passed in wax paper + rod between plates bearing a positive or a negative charge upon which separation can be effected winnowing – heavy particles can be separated from light particles by the confetti + pebbles action of streaming air which will carry the lighter particles further than the heavier particles flotation – if the solids possess different densities one can use a liquid of saw dust + Fe median density that will make the lighter solids float on top of it while the heavier solids will stay at the bottom sublimation — sometimes one of the solid components can be evaporated sand + I2 easily while the other cannot. In this case, heating of the mixture can result in the desired separation extraction — if one of the solid compounds is soluble in a specific solvent – sugar + I2 while the other one is not – one can simply use this solvent to transfer one of the two components into solution and, subsequently, use one of the techniques described for the separation of suspensions 2. solid–liquid — to separate the components of a suspension one can resort to... 13  2 Getting to the Core – Atomic Theory sedimentation and subsequent decantation — gravity will, over time, make mud the solid particles suspended in a liquid settle down at the bottom of a vessel. Carefully pouring off the supernatant liquid will povide this phase in pure form centrifugation — gravitational settling of particles can take a long time but clay + H2 O the process can be sped up significantly be means of a centrifuge filtration — a filter will hold back solid material while it allows the liquid sand + H2 O to pass through. The process can be made faster if a vacuum is applied. Meanwhile, filters are available that can even hold back bacteria 3. liquid–liquid — to separate the liquid components in emulsions it is possible to succeed by... sedimentation and subsequent decantation — gravity will, over time, make milk the liquids separate according to their density centrifugation — the liquids are sorted according to their density oil + H2 O 4. solid–gas & liquid–gas — separating aerosols can be effected by... filtration — an aerosol can be passed through a filter that holds back every- cigarette thing but the gas washing — an aerosol can be bubbled through a liquid that will hold back hairspray everything but the gas In case of homogeneous mixtures, the separation techniques are aimed at changing the physical conditions of the mixture in such a way that the homogeneous mixture is turned into a heterogeneous one. The most common operations in that aspect are: 1. evaporation — the oldest and easiest way to isolate a dissolved solid from solu- salt water tion is by simply evaporating the solvent. As an alternative, one can also simply NH4 Cl in ice decrease the temperature or add a different solvent that is miscible with the first NH4 Cl + ether one but is not a good solvent for the dissolved compound 2. extraction — a dissolved substance can be transferred from one solvent to another I2 water solvent in which it is more soluble. If the two solvents as such are not miscible, a separatory funnel can then effect the separation 3. distillation — a mixture of two liquids can be separated by heating, vaporization red wine and subsequent condensation of the lower boiling liquid if their boiling points are sufficiently different 4. chromatography — materials attract each other adhesively, i.e. "they stick to ink each other". But as different substances are attracted to the same material to a different extent, one can expose a homogeneous mixture to a system of two other materials of which the one is kept steady – the stationary phase – while the other 14  2 Getting to the Core – Atomic Theory one is allowed to travel – the mobile phase. Typically, the stationary phase is a solid such as silicondioxide or sugar while the mobile phase is a solvent or a gas. The differing adhesions will then gradually separate the two components of the homogeneous mixture. This technique is among the most important ones in Organic Chemistry – column chromatography, thin-layer chromatography and gas chromatography might cross your way later in your studies 2.1.3 Properties of Pure Compounds nce a substance has been isolated in pure form – sometimes one needs to repeat O separation steps multiple times or combine various techniques to reach this goal – one can measure and describe its properties that – in its entirety – are unique for this specific substance. Among the most common things to report are colour, smell, taste state of aggregation at a specific set of temperature and pressure melting point, boiling point electrical conductivity pure water magnetic behaviour Fe vs. Al heat conductivity Cu vs. glass speed of sound H2 vs. Xe refractory index many spectroscopic properties Over time, chemists succeeded in isolating a startling number of substances in pure form and describing their properties in great detail. In doing so, they discovered that all pure substances could be further arranged in two large groups: 1. substances that decompose , i.e. upon application of harsh (but laboratory- Cu(OAc)2 ∆T↑ doable) physical conditions such as intense heat, electricity, radiation, etc. broke up into new compounds that – themselves in turn – could be isolated, purified and characterized in great detail 2. substances that did not break up into new ones While the former type of substance is commonly referred to as compound , the latter type is referred to as element. As a consequence, one can now provide two important definitions: 1. compounds are substances that contain different elements – they can be chopped up by physical and chemical methods 15  2 Getting to the Core – Atomic Theory Matter Homogeneous Separation Pure Separation Heterogeneous Mixture Substances Mixture Compounds Elements Figure 2.3: The family tree of substances. 2. elements are substances that cannot be turned into simpler substances by chemical (and most physical) methods Figure 2.3 summarizes the relationship in between the concepts discussed on the pre- ceding pages in a graphical manner. 2.2 Atomic Theory 2.2.1 An Old Concept – Historic Roots ow often can I divide a rock?". This was the question the ancient Greek "H philosopher Demokrit (cf. Figure 2.4 on page 17) was pondering over in the 5th century bc. By simply following through this experiment of thought, he concluded that all matter must – at one stage – be comprised of some very small units that could no longer be chopped up in smaller ones. Using the Greek term for indivisible – a-tomos – he named those basic building blocks atoms. Demokrit’s theory got shunned very soon after by another Greek philosopher, Aris- totle, and fell into oblivion again for several thousands of years. It was not before the 17th till the early 19th century ad that English scholars such as Robert Boyle (in 1661) and Isaac Newton (in 1687 and in 1704) got back to those early theories. The first real breakthrough, however, came with John Dalton between 1803 and 1808. Effectively, it was once again Lavoisier’s introduction of the rigorous and concise use of scales into chemistry that gave rise to the fundamental laws that, ultimately, proved the existence of smallest particles matter is made up from. 16  2 Getting to the Core – Atomic Theory Figure 2.4: From left to right: Demokrit, Isaac Newton and Aristotle. Source of photos: Wikipedia (users: Tomisti (Demokrit); Meidosensei (Isaac Newton); Jastrow (Aristotle)). 2.2.2 Masses, Masses, Masses lchemy – as explained earlier – was not a real natural science as it entailed A superstition, witchcraft and magic. Furthermore, it tried to create or alter matter from nothing such as in the case of the Philosopher’s Stone that was supposed to turn all metals into gold. Upon the introduction of keeping track of masses and weights by means of scales, Lavoisier was quickly able to refute several of the alchmists’ theories and allow other scientist to establish a set of important fundamental laws. 2.2.2.1 Conservation of Mass While alchemists thought that matter and substances can simply vanish, it was found Fe wool + O2 that masses, actually, did not change upon a chemical reaction if all reaction products C + O2 are taken into account. It should be pointed out that this observation has a tiny loop hole that is manifest in Einstein’s famous law that connects mass and energy: E = m × c2 (3) E energy m mass c speed of light However, these changes were below the detection limits of those times and, in general, one can state the Law of Conservation of Mass : In the wake of a chemical reaction, the sum of the masses of all products equals the sum of the masses of all starting materials. 17  2 Getting to the Core – Atomic Theory Table 2.2: The masses m of mecuryoxide, mercury and oxygen found upon thermal decompo- sition of the compound. All masses are provided in g. mmercuryoxide mmercury moxygen mmercury/moxygen 0.759 0.703 0.056 12.55 2.322 2.150 0.172 12.50 5.906 5.470 0.436 12.55 20.471 18.959 1.512 12.54 65.800 60.939 4.861 12.54 100.006 92.618 7.388 12.54 438.692 406.285 32.407 12.54 Table 2.3: The masses m of silveroxide, silver and oxygen found upon thermal decomposition of the compound. All masses are provided in g. msilveroxide msilver moxygen msilver/moxygen 1.098 1.022 0.076 13.45 2.552 2.376 0.176 13.50 7.813 7.274 0.539 13.50 26.111 24.308 1.803 13.48 59.999 55.856 4.143 13.48 114.067 106.191 7.876 13.48 509.404 474.233 35.171 13.48 2.2.2.2 Law of Constant Proportions When scientists started to characterize pure compounds taking into account the masses ∆T ↑ involved, an interesting observation was made. Upon heating mercuryoxide, scientists HgO −−−−→ found two products (a colourless gas and a silvery, metallic liquid) are obtained that – in turn – cannot be decomposed any further by chemical or physical means. Applying the definitions from Section 2.1.3 on page 15 it follows that mercuryoxide is a compound that decomposes to two elements – oxygen and mercury. Of course, the various scientists all over Europe started off with different amounts of mercuryoxide, thus yielding different amounts of the two elements. Comparing all the masses involved – and especially the ratio between the masses of the two products – as is done in Table 2.2 will show the observation as mentioned in the first sentence. ∆T ↑ A similar situation was encountered for the heating of silveroxide that decomposes AgO −−−−→ into two substances (a colourless gas and a darkish powder) that – themselves – can no longer be degraded into simpler materials. As discussed for mercuryoxide: silveroxide must be a compound that can be "chopped up" into two elements (oxygen and silver). The results of a similar series of experiments as for mercuryoxide is given in Table 2.3. Irrespective of the starting mass of mercuryoxide or silveroxide selected in the two separate experiments – the ratio of the masses of the two elements obtained upon de- 18  2 Getting to the Core – Atomic Theory Table 2.4: The masses m of carbonoxide № 1, carbon and oxygen found upon thermal decom- position of the compound (results Series A). All masses are provided in g. mcarbonoxide (№ 1) mcarbon moxygen mcarbon/moxygen 0.475 0.204 0.271 0.75 2.504 1.074 1.430 0.75 6.187 2.653 3.534 0.75 21.009 9.008 12.001 0.75 70.645 30.291 40.354 0.75 101.047 43.326 57.721 0.75 486.219 208.479 277.740 0.75 Table 2.5: The masses m of carbonoxide № 2, carbon and oxygen found upon thermal decom- position of the compound (results Series B). All masses are provided in g. mcarbonoxide (№ 2) mcarbon moxygen mcarbon/moxygen 0.632 0.172 0.460 0.37 1.989 0.543 1.446 0.38 5.207 1.421 3.786 0.38 22.287 6.082 16.205 0.38 68.303 18.639 49.664 0.38 104.000 28.381 75.619 0.38 462.775 126.288 336.487 0.38 composing the respective compound is found to be a constant value (the slight variations are due to errors of rounding). This principle could be confirmed for many other com- pounds and led to the formulation of the Law of Constant Proportions : In any specific compound, always the same elements are combined in always the same ratio of their masses. 2.2.2.3 Law of Multiple Proportions While working on analyzing compounds, some results seemed to contradict the Law of Constant Proportions. During the analysis of several carbonoxides – that fell apart into the elements carbon and oxygen – two series of mass relations were obtained as shown in Table 2.4 and Table 2.5. An even more confusing picture was obtained when compounds were analyzed that showed the presence of only the elements fluorine and sulfur – in this case, three different series of results as shown in Table 2.6 on page 20, Table 2.7 on page 20 and Table 2.8 on page 20 were obtained. Both sets of experiments – the one starting from carbonoxides and the one starting from sulfurfluorides – do not seem conclusive and in violation of the Law of Constant Proportions as there should only be one ratio for the masses of elements. The solution 19  2 Getting to the Core – Atomic Theory Table 2.6: The masses m of sulfurfluoride № 1, fluorine and sulfur found upon thermal decom- position of the compound (results Series A). All masses are provided in g. msulfurfluoride (№ 1) mfluorine msulfur mfluorine/msulfur 0.873 0.474 0.399 1.19 2.960 1.605 1.355 1.18 5.989 3.248 2.741 1.18 20.063 10.882 9.181 1.19 74.052 40.165 33.887 1.19 99.999 54.239 45.760 1.19 528.601 286.709 241.892 1.19 Table 2.7: The masses m of sulfurfluoride № 2, fluorine and sulfur found upon thermal decom- position of the compound (results Series B). All masses are provided in g. msulfurfluoride (№ 2) mfluorine msulfur mfluorine/msulfur 0.950 0.668 0.282 2.37 3.126 2.199 0.927 2.37 5.700 4.009 1.691 2.37 22.734 15.989 6.745 2.37 75.528 53.120 22.408 2.37 110.653 77.824 32.829 2.37 512.673 360.570 152.103 2.37 Table 2.8: The masses m of sulfurfluoride № 3, fluorine and sulfur found upon thermal decom- position of the compound (results Series C). All masses are provided in g. msulfurfluoride (№ 3) mfluorine msulfur mfluorine/msulfur 0.734 0.573 0.161 3.56 2.852 2.226 0.626 3.61 6.241 4.871 1.370 3.56 20.743 16.190 4.553 3.56 78.001 60.880 17.121 3.56 104.908 81.881 23.027 3.56 502.955 392.557 110.398 3.56 20  2 Getting to the Core – Atomic Theory Table 2.9: The relation between the two ratios obtained for the mass relations for the two series of carbonoxide decompositions. mcarbon/moxygen (Series A) mcarbon/moxygen (Series B) ratioSeries A/ratioSeries B 0.75 0.37 2.03 0.75 0.38 1.97 0.75 0.38 1.97 0.75 0.38 1.97 0.75 0.38 1.97 0.75 0.38 1.97 0.75 0.38 1.97 Table 2.10: The relation between the three ratios obtained for the mass relations for the three series A, B and C of sulfurfluoride decompositions. mfluorine/msulfur (A) mfluorine/msulfur (B) mfluorine/msulfur (C) ratioC/ratioA ratioB/ratioA ratioC/ratioB 1.19 2.37 3.56 2.99 1.99 1.50 1.18 2.37 3.61 3.06 2.01 1.52 1.18 2.37 3.56 3.02 2.01 1.50 1.19 2.37 3.56 2.99 1.99 1.50 1.19 2.37 3.56 2.99 1.99 1.50 1.19 2.37 3.56 2.99 1.99 1.50 1.19 2.37 3.56 2.99 1.99 1.50 to this seemingly bad problem is easy once one focuses on the ratios in between the two series of results obtained for the carbonoxides on the one hand and the three series of results obtained for the sulfurfluorides on the other hand as shown in Table 2.9 and Table 2.10. As can be seen from the last column in Table 2.9 and the fourth and fifth column in Table 2.10, the ratio in between the ratios in the various series is always an integer – the slight deviations stem from errors of rounding. As a consequence, the Law of Multiple Proportions could be stated: If two elements A and B combine to two (or more) different compounds, the ratios of the masses of these two elements – mA/mB – are in relation to each other by means of integer numbers. The reason for the values in the last column in Table 2.10 – that are right in the middle between two integer numbers – will be discussed at a later stage when talking about empirical formulae of sum. The Law of Multiple Proportions has been confirmed on a large number of compounds as has been the Law of Constant Proportions. Both principles turned out very important for the development of modern atomistic theories. 21  2 Getting to the Core – Atomic Theory Figure 2.5: From left to right: John Dalton, George Stoney, Joseph Thomson and Ernest Rutherford. Source of photos: Wikipedia (users: Materialscientist (Dalton, Rutherford); QWerk (Thomson); Seanwal111111 (Stoney)). 2.2.3 Dalton’s and Rutherford’s Atomic Model ased on the Law of Constant Proportions and – especially – on the Law of Mul- B tiple Proportions early chemists re-considered the construction plan of matter. The first concept to be be spelt out in greater detail stems from John Dalton (cf. Figure 2.5). Dalton set up a number of postulates to account for the findings: 1. Elements consist of extremely small particles – the atoms. All atoms of one element are identical. The atoms of different elements are different. 2. Atoms cannot be destroyed, created from scratch or change into the atom of an- other element. This accounted for the Law of Conservation of Mass. 3. During a chemical reaction atoms are either connected with one another or sepa- rated from one another or re-arranged amongst each other. 4. A chemical compound is the combination of various atoms. A specific compound contains always the same types of atoms, combined in a fixed ratio. This ac- counted for both, the Law of Constant Proportions as well as the Law of Multiple Proportions.1 Although Dalton’s theory of atoms called for the same indivisible smallest particles as Demokrit has done, experimental data gathered from around the second half of the 19th century onwards left no doubt that atoms were not the smallest particles but – themselves – had to be composed of even smaller building blocks. Apart from the discovery of natural radioactivity that gave rise to the notion that atoms are not necessarily stable but can disintegrate several planned experiments proved to provide crucial insight. Without going into too much detail: experiments with electricity, specifically beams of particles in evacuated tubes, George Stoney and Joseph Thomson (cf. Figure 2.5) could 1 To be precise: the Law of Multiple Proportions was stated by Dalton based on this last postulate, i.e. it was the result of his theory, not a pre-requisite! 22  2 Getting to the Core – Atomic Theory Table 2.11: Important properties of atomic building blocks. Particle Symbol Location Charge (in Coulomb) Mass (in kg) Electron e− hull −1.602 × 10−19 9.109 383 56(11) × 10−31 Proton p core +1.602 × 10−19 1.672 621 777(74) × 10−27 Neutron n core 0 1.674 927 471(21) × 10−27 deduce that atoms must contain negatively and positively charged sub-particles. While the negative particles were named electrons , the positive particles were named protons. A later experiment by Ernest Rutherford (cf. Figure 2.5 on page 22) showed that the positive charge within an atom had to be concentrated in a very small region with ample "space" in between. The spot of positive charge became known as the atomic core – also known as nucleus – while the "empty space" – the atomic hull – is taken up by the electrons. 2.2.4 The Basic Construction Plan of an Atom part from the two particles found in atoms mentioned in the previous section A – the electron and the proton – a third particle was required to be present in atoms as to account for the unexpectedly high masses of some atoms (this point will be discussed in more detail in the following section). This particle – the neutron, originally already postulated by Ernest Rutherford in 1920 – was eventually detected by James Chadwick (cf. Figure 2.6) in 1932. Some important properties of these three fundamental particles are summarized in Table 2.11. With this knowledge at hand, a first fitting de- scription of atoms was possible with the small nucleus containing positively-charged particles – the protons – and non-charged particles – the neutrons – and all negatively-charged particles – the electrons – present in the much bigger atomic hull. One can, therefore, state that the biggest part of all matter we know is actually "nothing"! Dalton’s consideration that the atoms of the various elements must be different compared to each other can now be taken from a postulate to a proven fact by basing the differences among these atoms on variations in the numbers of those sub- atomic particles present in exactly those atoms, i.e. the atoms of different elements have differ- Figure 2.6: James Chadwick. ent amounts of protons, neutrons and electrons in Source of photo: Wikipedia (user: their cores and hulls. As protons and neutrons Materialscientist). reside in the nucleus, they are often referred to collectively as nucleons. It is now possible to ad- 23  2 Getting to the Core – Atomic Theory dress the various chemical elements known not only by their name or their symbol in a qualitative manner but also in a quantitative manner that provides information about the number of protons, neutrons and electrons present in a specific atom. To do this, the symbol of the element (cf. Section ?? on page ??) is endowed with a superscript and a subscript number (both preceding the elemental symbol itself). While the super- script number A contains the total number of nucleons present in the core, the subscript number Z only identifies the number of protons: A ZX (4) A number of nucleons (protons + neutrons) Z number of protons X chemical symbol of element Example: Determine the number of protons and neutrons in 79 35Br! The number of protons is 35 while the number of neutrons is 79 − 35 = 44. One can even quantify the number of electrons in an atom based on its symbol: as the charge of an electron is equal to the charge of a proton – just the signage is different – the number of electrons in a neutral atom must be equal to the number of protons in the core. Example: Determine the number of protons, neutrons and electrons in 238 92U! The number of protons is 92 while the number of neutrons is 238 − 92 = 146. As the given atom is neutral, the number of electrons must equal the number of protons, i.e. there are 92 electrons in the given atom. In case of charged atoms – so-called ions – the number of electrons will be greater or smaller than the number of protons. Example: Determine the number of protons, neutrons and electrons in 31 3– 15P ! The number of protons is 15 while the number of neutrons is 31 − 15 = 16. As the given atom has a charge of −3, the number of electrons must surpass the number of protons by three, i.e. there are 15 + 3 = 18 electrons in the given atom. Example: Determine the number of protons, neutrons and electrons in 88 2+ 38Sr ! The number of protons is 38 while the number of neutrons is 88 − 38 = 50. As the given atom has a charge of +2, the number of electrons must undercut the number of protons by two, i.e. there are 38 − 2 = 36 electrons in the given atom. It should be noted at the end of this section that Dalton’s first statement about the equality of all atoms of one specific element is not quite correct. For quite a large number of elements, variations are known that have the same number of protons (and in case of neutral atoms, therefore, the same number of electrons) but different numbers of neutrons. Such variations of atoms within one element are referred to as isotopes. Examples would be the isotopes of hydrogen 11H, 21H and 31H or the chlorine isotopes 24  2 Getting to the Core – Atomic Theory Table 2.12: Relative atomic masses based on hydrogen for chlorine, nitrogen and carbon based on hydrogenchloride HCl, ammonia NH3 and methane CH4. Compound Content hydrogen Content other element relative atomic mass of other element HCl 2.76 % 97.24 % 97.24/2.76 × 1 = 35.232 NH3 17.76 % 82.24 % 82.24/17.76 × 3 = 13.892 CH4 25.12 % 74.88 % 74.88/25.12 × 4 = 11.924 35 and 37 17Cl 17Cl. Sets of atoms that differ in the number of protons but have an identical number of neutrons are known as isotones such as 146C and 168O. 2.2.5 Atomic Masses onducting Demokrit’s experiment of thought about chopping up a rock into C ever smaller pieces in mind will already provide an idea about how small the world of atoms and – even worse – the world of protons, neutrons and electrons is. Irrespective of the fact that these particles are so small, they, nevertheless, have a mass. The laws about constant and multiple proportions that tied the composition of compounds to the ratios of masses of elements they are composed of is proof of that. But because of their small size it is not possible to weigh individual atoms. However, based on the mass ratios for compounds when analyzing compounds, one can assign relative masses. As an example, water shall be used as a model compound. Shortly we will see that water is a compound in which two hydrogen atoms are bonded to one oxygen atom – or as chemists would write it: H2 O. The analysis of water (done in the way as shown for the laws of mass proportions) shows that its mass consists of 11.19 % hydrogen and 88.81 % oxygen, i.e. the mass ratio of oxygen to hydrogen is 88.81/11.19 = 7.937. If the atomic mass of hydrogen is now arbitrarily set to be ’1’ it would mean that the relative atomic mass – i.e. relative to hydrogen as the basis – of oxygen would have to be 2 × 7.937 = 15.874. Taking a similar approach, the mass ratios and formulae of sum of hydrogenchloride, ammonia and methane were able to provide the relative atomic masses for chlorine, nitrogen and carbon as listed in Table 2.12. For practical reasons (not all elements form well-defined compounds with hydrogen), the reference point was later pegged to oxygen (with O = 16) and eventually – oxygen is a gas that is difficult to handle – shifted to carbon (with 126C = 12) by stating that the relative atomic mass unit u corresponds to 1/12 of the mass m of the carbon isotope 12 6C: 1 u = 1/12 × m 126C (5) u relative atomic mass unit m mass 12 6C the carbon isotope 126C 25  2 Getting to the Core – Atomic Theory Table 2.13: Selected relative atomic masses based on different reference points. 12 Hydrogen = 1 Oxygen = 16 6C = 12 Hydrogen 1.000 1.008 1.008 Chlorine 35.232 35.457 35.453 Oxygen 15.874 16.000 15.999 Nitrogen 13.892 14.008 14.007 Carbon 11.924 12.011 12.000 Table 2.14: The masses of the atomic building blocks proton, neutron and electron (in kg and u). Particle Mass (in kg) Mass (in u) Electron 9.109 383 56(11) × 10−31 5.485 799 090 70(16) × 10−4 Proton 1.672 621 777(74) × 10−27 1.007 276 466 812(90) Neutron 1.674 927 471(21) × 10−27 1.008 664 915 88(49) As a consequence, one could provide the relative atomic mass of oxygen as 15.874 u, carbon as 11.924 u, nitrogen as 13.892 u and chlorine as 35.232 u. The absolute value for u is very small – one u corresponds to just 1.660 538 921 × 10−27 kg. Although it does not seem to be of significance, Table 2.13 shows that the change of reference for the relative atomic mass unit has had a small impact on the mass values reported. Upon some experimentation whose details lie beyond the scope of this module not only the masses of full atoms but also the masses of the sub-atomic particles could be determined and are given – in kilograms – in Table 2.11 on page 23. For comparison, the masses could also be provided in units of u as done in Table 2.14. One can easily see that the mass of an electron is by a factor of roughly 2000 smaller than that of the two nucleons. It is obvious that the mass of each atom is simply the sum of the masses of its sub- atomic building blocks. However, two observations must be mentioned in this aspect as they have profound implications: 1. The mass of any atom consisting of more than just one nucleon – i.e. everything heavier than the lightest element isotope of hydrogen, 11H – is less than the sum of the masses of the individual building blocks. This – apparent! – discrepancy is known as the mass defect. The resting mass of the 42He nuclide (mind: that excludes the electrons!) is 4.00 151 u. As the resting mass of a proton is 1.007 276 u and that of a neutron is 1.008 665 u one would expect a total mass of 4.03 188 u. Therefore, the mass defect of a 42He nuclide is 0.03 037 u. The "lost" mass is, of course, not really lost but has been transformed into energy as briefly mentioned in Section 2.2.2.1 on page 17. 26  2 Getting to the Core – Atomic Theory Question: Where did the energy go to? 2. Even if the mass defect is taken into account one cannot simply resort to the expected mass of an atom for practical applications. As was mentioned earlier, many elements are present as mixtures of isotopes. Whenever one has to deal with many atoms – i.e. everytime "one does chemistry" – one has to take the different masses of these isotopes as well as their natural abundance into account by working with the average atomic mass. Natural bromine occurs as the isotopes 79 81 35Br and 35Br. Their respective masses are 78.918 336 1 u and 80.916 289 u, their natural abundances are 50.69 % and 49.31 %. Therefore, it follows that the average atomic mass of bromine is 78.918 336 1 u × 50.69/100 + 80.916 289 u × 49.31/100 = 79.903 525 u. It should be pointed out that the average atomic mass of an element is, as a consequence, not the mass any of the atoms of this specific element would have in reality! You should have learned the following things from this section: 1. What states of aggregation exist and how they interconvert 2. How phase diagrams can be used to determine the state of aggregation of a substance 3. Which types of mixtures exist in chemistry 4. How mixtures are separated into pure compounds 5. What properties can be recorded for pure compounds 6. Which observations were made for the relation of masses of elements entailed in specific compounds 7. Which atomic models were set up by Demokrit, Dalton and Rutherford 8. Which particles are found in atoms 9. What information can be gained from the full elemental symbol 10. Why average atomic masses deviate from the pure sum of the atomic constituent particles 11. How the average atomic masses of elements are calculated 27  2 Getting to the Core – Atomic Theory To practice you could answer the following questions: 1. Name the states of aggregation and the processes that convert them among each other! 2. Do you know compounds that lack certain states of aggregation? 3. It is an urban legend that ice skating is possible thanks to the presence of a slippery film of liquid water on the ice that is generated by the high pressure the focusing of a person’s body weight on the thin blades of the skating shoes excerts. Can you rationalize this (false) urban legend on grounds of the phase diagram of water as shown in Figure 2.1 on page 12? 4. Solid carbondioxide is also referred to as "dry ice" because upon warming up it does not create a puddle of liquid carbondioxide. What does that tell you qualitatively about the pressure carbondioxide’s triple point is located at with regards to ambient conditions? 5. Name the various heterogeneous mixtures one can encounter for substances in different states of aggregation! 6. Why is there no specific term for a heterogeneous mixture of two gases? 7. Name several real-life examples each for all homogeneous and heterogeneous mix- tures! 8. How would you separate the following mixtures (are these homogeneous or hete- rogeneous?): a) mud (stones and silt in water) b) milk (water in oil) c) ocean water (salt in water) d) air (oxygen in nitrogen) 9. The distillation of red wine can yield water (Tb = 100 ◦ C) and alcohol (Tb ≈ 78 ◦ C). Let us neglect all the solid residues that remain in the end. How does distillation achieve this? 10. Name five ways how one could characterize water as a pure compound and what values/properties you would expect for your chosen ways! 11. State the Law of Conservation of Mass! Is it really correct in each and every aspect? 12. Two different compounds containing the elements phosphorus and bromine only were analyzed by seven chemists each. The results are given in the following two tables (all masses in gram). In each case, one of the researchers made a mistake 28  2 Getting to the Core – Atomic Theory when writing down the results. Justify which chemist made the mistake and how you can tell that! Chemist mphosphorusbromide mphosphorus mbromine A 0.766 0.088 0.678 B 2.307 0.264 2.043 C 8.102 0.927 7.175 D 22.744 2.602 20.142 E 61.080 8.988 52.092 F 124.983 14.299 110.684 G 511.654 58.539 453.115 Chemist mphosphorusbromide mphosphorus mbromine A 1.295 0.235 1.060 B 2.766 0.199 2.567 C 7.753 0.558 7.195 D 28.024 2.016 26.008 E 60.000 4.316 55.684 F 121.381 8.732 112.649 G 499.732 35.951 463.781 13. State the Law of Multiple Proportions on grounds of the preceding question! 14. Which particles can be found in an atom? 15. Why does the occurrence of isotopes contradict Dalton’s atomic model? 16. Complete the following table! You may consult the Periodic Table of the Elements given in Figure ?? on page ?? or at the end of the lecture notes! Symbol Z A Protons Neutrons Electrons Te 52 125 Cs 55 133 P3 – 31 Bi 209 56 138 Sn 70 50 Kr 84 48 Sc3+ 24 8 8 10 7 7 10 1 3 2 29  2 Getting to the Core – Atomic Theory 17. Natural gallium consists of the isotopes 69 71 31Ga and 31Ga. If their natural abun- dance is 60.108 % (for the lighter isotope) and 39.892 % (for the heavier isotope), respectively, and their masses are 68.9255736 u and 70.9247013 u what is the av- erage atomic mass of gallium? 18. Natural silver is found in form of two isotopes – 107 109 47Ag and 47Ag. While the atomic mass of the lighter isotope is 106.906 u, the heavier one has an atomic mass of 108.905 u. What is the percentage distribution of these two isotopes if the average atomic mass of silver is determined to be 107.868 u? 19. Bromine occurs in nature as two isotopes, 79 81 35Br and 35Br. Hydrogen can be reacted with bromine to form hydrogenbromide, HBr. Given the two different masses of the two bromine isotopes – should there not be two different masses of bromine one would have to react with hydrogen to yield this product? How can the Law of Constant Proportions maintain its validity after all? 20. Why does it matter which atom/element/isotope is used as the basis for the atomic mass unit u? References: 1. State of Aggregation, Phase Diagrams — A. Burrows, J. Holman, A. Parsons, G. Pilling, G. Price: Chemistry3 — introducing inorganic, organic and physical chemistry, Oxford University Press, 2009, 810–821. J. C. Kotz, P. M. Treichel, J. R. Townsend: Chemistry & Chemical Reac- tivity (International Edition), 8th Edition, Brooks/Cole, Cengage Learning, 2012, 7; 606–607. http://en.wikipedia.org/wiki/Phase_diagram (last accessed: Jan 12th , 2025). 2. Mixtures, Separation Techniques, Physical Properties of Pure Substances — J. C. Kotz, P. M. Treichel, J. R. Townsend: Chemistry & Chemical Reac- tivity (International Edition), 8th Edition, Brooks/Cole, Cengage Learning, 2012, 8–14. http://en.wikipedia.org/wiki/Mixture (last accessed: Jan 12th , 2025). http://en.wikipedia.org/wiki/Separation_process (last accessed: Jan 12th , 2025). 3. Atomic Theories, Laws of Mass and Proportions — A. Burrows, J. Holman, A. Parsons, G. Pilling, G. Price: Chemistry3 — introducing inorganic, organic and physical chemistry, Oxford University Press, 2009, 13–14; 72–77. 30  2 Getting to the Core – Atomic Theory J. C. Kotz, P. M. Treichel, J. R. Townsend: Chemistry & Chemical Reac- tivity (International Edition), 8th Edition, Brooks/Cole, Cengage Learning, 2012, 51; 335–343. http://en.wikipedia.org/wiki/Atomic_theory (last accessed: Jan 12th , 2025). 4. Atomic Weight, Isotopes — A. Burrows, J. Holman, A. Parsons, G. Pilling, G. Price: Chemistry3 — introducing inorganic, organic and physical chemistry, Oxford University Press, 2009, 14–18. J. C. Kotz, P. M. Treichel, J. R. Townsend: Chemistry & Chemical Reac- tivity (International Edition), 8th Edition, Brooks/Cole, Cengage Learning, 2012, 52–57. http://en.wikipedia.org/wiki/Relative_atomic_mass (last accessed: Jan 12th , 2025). http://en.wikipedia.org/wiki/Isotope (last accessed: Jan 12th , 2025). 31