KTT 111/3 – Inorganic Chemistry I

1. Chapter 1: Atoms and Elements KTT 111/3 – Inorganic Chemistry I Dr. Farook Adam August 2005

2. Chapter 1: Atoms and Elements • Chemistry is a science that studies the composition and properties of matter • Matter isanything that takes up space and has mass • Mass is a measure of the amount of matter in a sample • Chemistry holds a unique place among the sciences because all things are composed of chemicals • A knowledge of chemistry will be valuable whatever branch of science you study

3. Chemistry is constantly changing as new discoveries are made by researchers • Researchers use a commonsense approach to the study of natural phenomena called the scientific method • A scientific study normally: • Begins with a question about nature • Involves a search of the work of others • Requires observing the results of experiments • Often results in a conclusion, or a statement based on what is thought about a series of observations

4. Experiments provide empirical facts • Facts are called data • A broad generalization based on the results of many experiments is called a (scientific) law • Laws are often expressed as mathematical equations • Laws summarize the results of experiments

5. Theoretical models attempt to explain why substances behave as they do • A hypothesis is a tentative explanation • A theory is an experimentally tested explanation of the behavior of nature The scientific method is dynamic: observations lead to laws, which suggest new experiments, which may lead to or change a hypothesis, which may produce a theory.

6. Chemical substances are comprised of atoms • Atoms combine to form molecules which can be represented in a number of ways, including: • Using chemical symbols and lines for “connections” • A 3-D ball-and-stick model • A 3-D space-filling model

7. Characteristics or properties of materials distinguish one type of substance from another • Properties can be classified as physical or chemical • Physical properties can be observed without changing the chemical makeup of the substance • Chemical properties involve a chemical change and result in different substances • Chemical changes are described by chemical reactions

8. Properties can also be described as intensive or extensive • Intensive properties are independent of sample size • Examples: sample color and melting point • Extensive properties depend on sample size • Examples: sample volume and mass • In general, intensive properties are more useful in identifying a substance • Matter is often classified by properties

9. The three common physical states of matter have different properties: • Solids have a fixed shape and volume • Particles are close together and have restricted motion • Liquids have indefinite shape but fixed volume • Particles are close together but are able to flow • Gases have indefinite shape and volume • Particles are separated by lots of empty space

10. Elements are substances that cannot be decomposed by chemical means into simpler substances • Each element is assigned a unique chemical symbol • Most are one or two letters • First letter is always capitalized • All remaining letters are lowercase • Names and chemical symbols of the elements are listed on the inside front cover of the book

11. Compounds are substances formed from two or more different elements combined in a fixed proportion by mass • The physical and chemical properties of a compound are, in general, different than the physical and chemical properties of the elements of which it is comprised • Elements and compounds are examples of pure substances whose composition is the same, regardless of source

12. A mixture consists of varying amounts of two or more elements or compounds • Homogeneous mixtures or solutions have the same properties throughout the sample • Heterogeneous mixtures consist of two or more phases • Matter can be classified:

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14. We take for granted the existence of atoms and molecules • The concept of the atom had limited scientific usefulness until the discovery of two important laws: the Law of conservation of mass and the Law of Definite Proportions • These laws summarized the results of the experimental observations of many scientists

15. Law of Conservation of Mass: • No detectable gain or loss of mass occurs in chemical reactions. Mass is conserved. • Law of Definite Proportions: • In a given chemical compound, the elements are always combined in the same proportions by mass. • In the sciences mass is measured in units of grams (symbol, g) • One pound equals 453.6 g

16. The laws of conservation of mass and definite proportions provided the experimental foundation for the atomic theory • Dalton’s Atomic Theory (~1803 AD): • Matter consists of tiny particles called atoms. • Atoms are indestructible. In chemical reactions, the atoms rearrange but they do not themselves break apart. • In any sample of a pure element, all the atoms are identical in mass and other properties. • The atoms of different elements differ in mass and other properties. • In a given compound the constituent atoms are always present in the same fixed numerical ratio.

17. Support for Dalton’s Atomic Theory: The Law of Multiple Proportions Whenever two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers. Each molecule has one sulfur atom, and therefore the same mass of sulfur. The oxygen ratio is 3 to 2 by both mass and atoms: Sample experimental data: Mass Mass Compound Size S O Sulfur dioxide 2.00 g 1.00 g 1.00 g Sulfur trioxide 2.50 g 1.00 g 1.50 g

18. It follows from Dalton’s Atomic Theory that atoms of an element have a constant, characteristic atomic mass or atomic weight • For example, for any sample of hydrogen fluoride: • F-to-H atom ratio: 1 to 1 • F-to-H mass ratio: 19.0 to 1.00 • This is only possible if each fluorine atom is 19.0 times heavier than each hydrogen atom

19. It turns out that most elements in nature are uniform mixtures of two or more kinds of atoms with slightly different masses • Atoms of the same element with different masses are called isotopes • For example: there are 3 isotopes of hydrogen and 4 isotopes of iron • Chemically, isotopes have virtually identical chemical properties • The relative proportions of the different isotopes are essentially constant

20. A uniform mass scale for atoms requires a standard • For atomic mass units (amu, given the symbol u) the standard is based on carbon: • 1 atom of carbon-12 = 12 u (exactly) • 1 u = 1/12 the mass of 1 atom of carbon-12 (exactly) • This definition results in the assignment of approximately 1 u for the mass of hydrogen (the lightest atom)

21. Example: Naturally occurring chlorine is a mixture of two isotopes. In every sample of this element, 75.77% of the atoms are chlorine-35 and 24.23% are chlorine-37. The measured mass of chlorine-35 is 34.9689 u and that of chlorine-37 is 36.9659 u. Calculate the average atomic mass of chlorine. Abundance Mass Isotope (%) (u) Contribution Chlorine-35 75.77 34.9689 0.7577 * 34.9689 = 26.50 u Chlorine-37 24.23 36.9659 0.2423 * 36.9659 = 8.957 u (Rounded) Total = 35.46 u The average mass of 1 atom of chlorine in nature is 35.46 u.

22. Experiments have been performed that show atoms are comprised of subatomic particles • There are three principal kinds of subatomic particles: • Proton – carries a positive charge, found in the nucleus • Electron – carries a negative charge, found outside the nucleus, about 1/1800 the mass of a proton • Neutron – carries no charge, found in the nucleus, a bit heavier than a proton, about 1800 times heavier than an electron

23. An element can be defined as a substance whose atoms all contain the identical number of protons, called the atomic number (Z) • Isotopes are distinguished by mass number (A): • Atomic number, Z = number of protons • Mass number, A = (number of protons) + (number of neutrons) • For charge neutrality, the number of electrons and protons must be equal

24. Mass number, A (protons + neutrons)  Chemical Symbol  Atomic number, Z (number of protons)  235 U 92 • This information can be summarized • Example: For uranium-235 • Number of protons = 92 ( = number of electrons) • Number of neutrons = 143 • Atomic number (Z) = 92 • Mass number (A) = 92 + 143 = 235 • Chemical symbol = U • Summary for uranium-235:

25. The Periodic Table summarizes chemical and physical properties of the elements • The first Periodic Tables were arranged by increasing atomic mass Dmitri Mendeleev (Rusia) 1869 Julius Lothar Meyer (German) Both these researchers drew out the first periodic table independently of each other.

26. Mendeleev have been honored as the first person to arrange the elements in the form of a table because he reported his findings to the Russian Chemical Society a few months earlier than Meyer!!!!

27. The Modern Periodic table is arranged by increasing atomic number: • Elements are arranged in numbered rows called periods • The vertical columns are called groups or families (group labels vary)

28. Modern Periodic Table with group labels and chemical families identified Lanthanides Actinides Note: Placement of elements 58 – 71 and 90 – 103 saves space

29. Some important classifications: • A groups = representative elements or main group elements • I A = alkali metals • II A = alkaline earth metals • VII A = halogens • VIII = noble gases • B groups = transition elements • Inner transition elements = elements 58 – 71 and 90 – 103 • 58 – 71 = lanthanide elements • 90 – 103 = actinide elements

30. Classification as metals, nonmetals, and metalloids

31. Metals • Tend to shine (have metallic luster) • Can be hammered or rolled into thin sheets (malleable) andcan be drawn into wire (ductile) • Are solids at room temperature and conduct electricity • Nonmetals • Lack the properties of metals • React with metals to form (ionic) compounds • Metalloids • Have properties between metals and nonmetals

32. The End Do the exercises at the end of chapter ONE. For your own practice only!! You do not have to pass up this assignment. Any questions?