Introduction to Chemistry

1. 1 Introduction to Chemistry

2. CHAPTER OBJECTIVES To recognize the breadth, depth, and scope of chemistry To understand what is meant by the scientific method To be able to classify matter To understand the development of the atomic model To know the meaning of isotopes and atomic masses To become familiar with the periodic table and to be able to use it as a predictive tool

3. Chemistry: Principles, Patterns, and Applications, 1e 1.1 Chemistry in the Modern World

4. 1.1 Chemistry in the Modern World Chemistry is the study of matter and the changes that material substances undergo Of all the scientific disciplines, it is the most extensively connected to other fields of study Disciplines that focus on living organisms and their reactions with the physical world rely heavily on chemistry Practical applications of chemistry are everywhere Chemistry affects our daily lives

5. Chemistry: Principles, Patterns, and Applications, 1e 1.2 The Scientific Method

6. 1.2 The Scientific Method A procedure that searches for answers to questions and solutions to problems Consists of making observations,formulatinghypotheses, and designing experiments,which leads to additional observations, hypotheses, and experiments in repeated cycles

7. Observations Can be qualitative or quantitative Qualitative observations Describe properties or occurrences in ways that do not rely on numbers Quantitative observations Measurements that consist of both a number and a unit

8. Hypotheses A tentative explanation for the observations May not be correct, but it puts the scientist’s understanding of the system being studied into a form that can be tested

9. Experiments Tests the validity of the hypothesis Are systematic observations or measurements made under controlled conditions, in which the variable of interest is clearly distinguished from any others If experimental results are reproducible, they are summarized in a law.

10. Experiments Law – A verbal or mathematical description of a phenomenon that allows for general predictions – Describes what happens and not why – Unlikely to change greatly over time unless a major experimental error is discovered Theory – Attempts to explain why nature behaves as it does –Is incomplete and imperfect, evolving with time to explain new facts as they are discovered

11. Chemistry: Principles, Patterns,and Applications, 1e 1.3 A Description of Matter

12. 1.3 A Description of Matter Matter — anything that occupies space and possesses mass Mass — quantity of matter an object contains — does not depend on location of the object • Weight — a force caused by the gravitational attraction that operates on the object — depends on the location of an object

13. 1.3 A Description of Matter Three distinct states of matter: 1. Solids — relatively rigid and have fixed shapes and volumes —volumes of solids independent of temperature and pressure 2. Liquids—have fixed volumes but flow to assume the shape of their containers —Volumes of liquids independent of temperature and pressure 3. Gases —have neither fixed shapes nor fixed volumes and expand to fill their containers completely — Depend strongly on temperature and pressure

14. Pure Substances and Mixtures Pure Chemical Substance — any matter that has a fixed chemical composition and characteristic properties Mixture — combinations of two or more pure substances in variable proportions in which the individual substances retain their identity 1. Homogeneous mixtures a. All portions of a material are in the same state, have no visible boundaries, and are uniform throughout b. Also called solutions 2. Heterogeneous mixtures a. Composition of a material is not completely uniform

15. Pure Substances and Mixtures  Homogeneous mixtures —can be separated into their component substances by physical processes that rely on differences in some physical property 1. Distillation — uses differences in volatility, a measure of how easily a substance is converted to a gas at a given temperature 2. Crystallization — separates mixtures based on differences in solubility, a measure of how much of a solid substance remains dissolved in a given amount of a specified liquid Heterogeneous mixtures —components can be separated by simple means such as filtration

16. Pure Substances and Mixtures Most mixtures — can be separated into pure substances that may be either elements or compounds  Element —a substance that cannot be broken down into simpler ones by chemical changes Compound — contains two or more elements and has chemical and physical properties that are usually different from those of the elements of which it is composed. Can be broken down into their elements by chemical processes

17. Pure Substances and Mixtures •Overall organization of matter and the methods used to separate mixtures are summarized here

18. Properties of Matter Properties used to describe material substances can be classified as either physical or chemical. Physical properties: Characteristics that scientists can measure without changing the composition of the sample under study (mass, color, volume, amount of space occupied by the sample). Chemical properties: Describe the characteristic ability of a substance to react to form new substances (flammability and corrosion).

19. Properties of Matter Physical properties can be extensive or intensive 1. Extensive properties a. Vary with the amount of the substance, b. Include mass, weight, and volume. 2. Intensive properties a. Do not depend on the amount of the substance, b. Include color, melting and boiling point, electrical conductivity, and physical state at a given temperature, c. Determine a substance’s identity, d. Have an important intensive property called density (d), a ratio of two extensive properties, mass and volume density = mass d = m volume V

20. Chemistry: Principles, Patterns, and Applications, 1e 1.4 A Brief History of Chemistry

21. In fourth centuryB.C., ancient Greeks proposed that matter consisted of fundamental particles called atoms. Over the next two millennia, major advances in chemistry were achieved by alchemists. Their major goal was to convert certain elements into others by a process called transmutation. 1.4 A Brief History of Chemistry

22. • Beginnings of modern chemistry were seen in the sixteenth and seventeenth centuries, where great advances were made in metallurgy, the extraction of metals from ores. • In the seventeenth century, Boyle described the relationship between the pressure and volume of air and defined an element as a substance that cannot be broken down into two or more simpler substances by chemical means. Modern Chemistry

23. During the eighteenth century, Priestley discovered oxygen gas and the process of combustion where carbon-containing materials burn vigorously in an oxygen atmosphere. In the late eighteenth century, Lavoisier discovered respiration and wrote the first modern chemistry text. His most important contribution was the law of conservation of mass, which states that in any chemical reaction, the mass of the substances that react equals the mass of the products that are formed. He is known as the father of modern chemistry. Modern Chemistry

24. In 1803, Dalton proposed that elements consist of individual particles called atoms.His atomic theory ofmatter contains four hypotheses: 1. All matter is composed of tiny particles called atoms. 2. All atoms of an element are identical in mass and fundamental chemical properties. 3. A chemical compound is a substance that always contains the same atoms in the same ratio. 4. In chemical reactions, atoms from one or more compounds or elements redistribute or rearrange in relation to other atoms to form one or more new compounds. Atoms themselves do not undergo a change of identity in chemical reactions. The Atomic Theory of Matter

25. Dalton’s atomic theory is essentially correct, with four minor modifications: 1. Not all atoms of an element must have precisely the same mass. 2. Atoms of one element can be transformed into another through nuclear reactions. 3. The composition of many solid compounds are somewhat variable. 4. Under certain circumstances, some atoms can be divided (split into smaller particles). The Atomic Theory of Matter

26. The Law of Multiple Proportions Dalton could not use his theory to determine the elemental compositions of chemical compounds because he had no reliable scale of atomic masses. Dalton’s data led to a general statement known as the law of multiple proportions. Law states that when two elements form a series of compounds, the ratios of the masses of the second element that are present per gram of the first element can almost always be expressed as the ratios of integers.

27. Gay-Lussac attempted to establish the formulas of chemical compounds by measuring, under constant temperature and pressure conditions, the volumes of gases that reacted to make a given chemical compound, together with the volumes of the products if they were gases. Gay-Lussac’s results were explained by Avogadro’s hypothesis, which proposed that equal volumes of different gases contain equal numbers of gas particles when measured at the same temperature and pressure. Avogadro’s Hypothesis

28. Chemistry: Principles, Patterns, and Applications, 1e 1.5 The Atom

29. 1.5 The Atom Each element is chemically unique. To understand why they are unique, you need to know the structure of the atom (the smallest particle of an element) and the characteristics of its components.

30. 1.5 The Atom Atoms consist of electrons, protons, and neutrons. 1. Electrons and protons have electrical charges that are identical in magnitude but opposite in sign. Relative charges of 1 and +1 are assigned to the electron and proton, respectively. 2. Neutrons have approximately the same mass as protons but no charge—they are electrically neutral. 3. The mass of a proton or a neutron is about 1836 times greater than the mass of an electron. Protons and neutrons constitute the bulk of the mass of the atom.

31. 1.5 The Atom

32. The Electron In 1897, Thomson demonstrated that cathode rays could be deflected, or bent, by magnetic or electric fields, which indicated that the cathode rays consisted of charged particles. Measuring the extent of the deflection of the cathode rays, Thomson calculated the mass-to-charge ratio of the particles. Since like charges repel each other and opposite charges attract, Thomson concluded that the particles had a net negative charge. These particles are called electrons. Millikan calculated the charge on a single electron and determined the mass of an electron: mass X charge = mass charge

33. Radioactivity 1896, Becquerel discovered that certain minerals emitted a new form of energy. Becquerel’s work was extended by Pierre and Marie Curie, who used the word radioactivity to describe the emission of energy rays by matter. Rutherford, building on the Curies’ work, showed that compounds of elements emitted at least two distinct types of radiation. One was readily absorbed by matter and consisted of particles that had a positive charge and were massive compared to electrons. These particles were called α particles. Particles in the second type of radiation were called β particles and had the same charge and mass-to-charge ratio as electrons. A third type of radiation, γ rays, was discovered later and found to be similar to a lower energy form of radiation called X -rays.

34. Radioactivity Three kinds of radiation – α particles, β particles and γ rays 1. Distinguished by the way they are deflected by an electric field and by the degree to which they penetrate matter 2. α particles and β particles are deflected in opposite directions; α particles are deflected to a much lesser extent because of their higher mass-to-charge ratio. 3. γ rays have no charge and are not deflected by electric or magnetic fields. 4. α particles have the least penetrating power, and γ rays are able to penetrate matter readily.

35. The Atomic Model Rutherford’s results strongly suggested that both the mass and positive charge are concentrated in a tiny fraction of the volume of the atom, called the nucleus. Rutherford established that the nucleus of the hydrogen atom was a positively charged particle, which he called a proton. Also suggested that the nuclei of elements other than hydrogen must contain electrically neutral particles with the same mass as the proton. The neutron was discovered in 1932 by Rutherford’s student Chadwick. Because of Rutherford’s work, it became clear that an α particle contains two protons and neutrons—the nucleus of a helium atom.

36. Chemistry: Principles, Patterns, and Applications, 1e 1.6 Isotopes and Atomic Masses

37. 1.6 Isotopes and Atomic Masses Atoms of different elements exhibit different chemical behavior. Identity of an element is defined by its atomic number. (Z) isthe number of protons in the nucleus of an atom of the element. The atomic number is therefore different for each element. Known elements are arranged in order of increasing Z in the periodic table.

38. 1.6 Isotopes and Atomic Masses

39. 1.6 Isotopes and Atomic Masses The chemistry of each element is determined by its number of protons and electrons. In a neutral atom, the number of electrons equals the number of protons. Symbols for elements are derived directly from the element’s name. Nuclei of atoms contain neutrons as well as protons. The number of neutrons is not fixed for most elements, unlike protons. Atoms that have the same number of protons, and hence the same atomic number, but different numbers of neutrons are called isotopes.

40. 1.6 Isotopes and Atomic Masses Isotopes 1. All isotopes of an element have the same number of protons and electrons, which means they exhibit the same chemistry. 2. Isotopes of an element differ only in their atomic mass, which is given by the mass number (A),the sum of the numbers of protons and neutrons.

41. 1.6 Isotopes and Atomic Masses • Atomic mass 1. The mass of any given atom is not simply the sum of the masses of its electrons, protons, and neutrons. 2. Atoms are too small to measure individually and do not have a charge. 3. The arbitrary standard that has been established for describing atomic mass is the atomic mass unit (amu), defined as one-twelfth of the mass of one atom of 12C. .

42. 1.6 Isotopes and Atomic Masses 4. Most elements exist as mixtures of several stable isotopes. The weighted average is of the masses of the isotopes is called the atomic mass. 5. Electrons added or removed from an atom produce a charged particle called an ion,whose charge is indicated by a superscript after the symbol for the element. .

43. 1.7 Introduction to the Periodic Table Chemistry: Principles, Patterns,and Applications, 1e

44. 1.7 Introduction to the Periodic Table The single most important learning aid in chemistry Summarizes huge amounts of information about the elements so that you can predict many of their properties and chemical reactions

45. 1.7 Introduction to the Periodic Table

46. 1.7 Introduction to the Periodic Table Elements are arranged in seven horizontal rows, in order of increasing atomic number from left to right and from top to bottom. Rows are calledperiods and are numbered from 1 to 7. Elements with similar chemical properties form vertical columns, calledgroups, which are numbered from 1 to 18. Groups 1, 2, and 13 through 18 are the main group elements. Groups 3 through 12 are in the middle of the periodic table and are the transition elements. The two rows of 14 elements at the bottom of the periodic are the lanthanides and actinides.

47. The heavy orange zigzag line running diagonally from the upper left to the lower right through Groups 13–16 divides the elements into metals ( in blue, below and to the left of the line) and nonmetals (in bronze, above and to the right). Elements colored in gold that lie along the diagonal line are semimetals and exhibit properties intermediate between metals and nonmetals. Metals, Nonmetals, and Semimetals

48. Metals - Good conductors of electricity and heat - Ductile - Malleable - Lustrous - In chemical reactions, metals lose electrons to form positively charged ions - Vast majority of known elements are metals - All are solids except for mercury, which is a liquid at room temperature and pressure Metals, Nonmetals, and Semimetals

49. Nonmetals - Poor conductors of heat and electricity - Not lustrous - Can be gases, liquids, or solids - Solid nonmetals are brittle - Tend to gain electrons in reactions with metals to form negatively charged ions - Share electrons in reactions with other nonmetals Semimetals - Exhibit properties intermediate between metals and nonmetals Metals, Nonmetals, and Semimetals

50. Chemistry of the Groups Elements with similar chemical behavior are in the same group. Elements of Group 1 are alkali metals. Elements of Group 2 are the alkaline earths. Elements of Group 17 are the halogens. Elements of Group 18 are the noble gases.