Chapter 2 Atoms, Molecules, and Ions

1. CHEMISTRYThe Central Science 9th Edition Chapter 2Atoms, Molecules, and Ions David P. White Chapter 2

2. The Atomic Theory of Matter • John Dalton: • Each element is composed of atoms • All atoms of an element are identical. • In chemical reactions, the atoms are not changed. • Compounds are formed when atoms of more than one element combine. • Dalton’s law of multiple proportions: When two elements form different compounds, the mass ratio of the elements in one compound is related to the mass ratio in the other by a small whole number. Chapter 2

3. The Atomic Theory of Matter Chapter 2

4. The Discovery of Atomic Structure • The ancient Greeks were the first to postulate that matter consists of indivisible constituents. • Later scientists realized that the atom consisted of charged entities. • Cathode Rays and Electrons • A cathode ray tube (CRT) is a hollow vessel with an electrode at either end. • A high voltage is applied across the electrodes. Chapter 2

5. The Discovery of Atomic Structure • Cathode Rays and Electrons • The voltage causes negative particles to move from the negative electrode to the positive electrode. • The path of the electrons can be altered by the presence of a magnetic field. • Consider cathode rays leaving the positive electrode through a small hole. • If they interact with a magnetic field perpendicular to an applied electric field, the cathode rays can be deflected by different amounts. Chapter 2

6. The Discovery of Atomic Structure • Cathode Rays and Electrons • The amount of deflection of the cathode rays depends on the applied magnetic and electric fields. • In turn, the amount of deflection also depends on the charge to mass ratio of the electron. • In 1897, Thomson determined the charge to mass ratio of an electron to be 1.76  108 C/g. • Goal: find the charge on the electron to determine its mass. Chapter 2

7. The Discovery of Atomic Structure Cathode Rays and Electrons Chapter 2

8. The Discovery of Atomic Structure • Cathode Rays and Electrons • Consider the following experiment: • Oil drops are sprayed above a positively charged plate containing a small hole. • As the oil drops fall through the hole, they are given a negative charge. • Gravity forces the drops downward. The applied electric field forces the drops upward. • When a drop is perfectly balanced, the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate. Chapter 2

9. The Discovery of Atomic Structure Cathode Rays and Electrons

10. The Discovery of Atomic Structure • Cathode Rays and Electrons • Using this experiment, Millikan determined the charge on the electron to be 1.60  10-19 C. • Knowing the charge to mass ratio, 1.76  108 C/g, Millikan calculated the mass of the electron: 9.10  10-28 g. • With more accurate numbers, we get the mass of the electron to be 9.10939  10-28 g. Chapter 2

11. The Discovery of Atomic Structure • Radioactivity • Consider the following experiment: • A radioactive substance is placed in a shield containing a small hole so that a beam of radiation is emitted from the hole. • The radiation is passed between two electrically charged plates and detected. • Three spots are noted on the detector: • a spot in the direction of the positive plate, • a spot which is not affected by the electric field, • a spot in the direction of the negative plate. Chapter 2

12. The Discovery of Atomic Structure Radioactivity Chapter 2

13. The Discovery of Atomic Structure • Radioactivity • A high deflection towards the positive plate corresponds to radiation which is negatively charged and of low mass. This is called b-radiation (consists of electrons). • No deflection corresponds to neutral radiation. This is called g-radiation. • Small deflection towards the negatively charged plate corresponds to high mass, positively charged radiation. This is called a-radiation. Chapter 2

14. The Discovery of Atomic Structure • The Nuclear Atom • From the separation of radiation we conclude that the atom consists of neutral, positively, and negatively charged entities. • Thomson assumed all these charged species were found in a sphere.

15. The Discovery of Atomic Structure • The Nuclear Atom • Rutherford carried out the following experiment: • A source of a-particles was placed at the mouth of a circular detector. • The a -particles were shot through a piece of gold foil. • Most of the a-particles went straight through the foil without deflection. • Some a-particles were deflected at high angles. • If the Thomson model of the atom was correct, then Rutherford’s result was impossible. Chapter 2

17. The Discovery of Atomic Structure • The Nuclear Atom • In order to get the majority of -particles through a piece of foil to be undeflected, the majority of the atom must consist of a low mass, diffuse negative charge - the electron. • To account for the small number of high deflections of the -particles, the center or nucleus of the atom must consist of a dense positive charge. Chapter 2

18. The Discovery of Atomic Structure • The Nuclear Atom • Rutherford modified Thomson’s model as follows: • assume the atom is spherical but the positive charge must be located at the center, with a diffuse negative charge surrounding it.

19. The Modern View of Atomic Structure • The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons). • Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus. • There can be a variable number of neutrons for the same number of protons. Isotopes have the same number of protons but different numbers of neutrons. • Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons. Chapter 2

20. The Modern View of Atomic Structure Chapter 2

21. The Modern View of Atomic Structure • Isotopes, Atomic Numbers, and Mass Numbers • Atomic number (Z) = number of protons in the nucleus. • Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). • By convention, for element X, we write ZAX. • Isotopes have the same Z but different A. • We find Z on the periodic table. Chapter 2

22. Atomic Weights • The Atomic Mass Scale • 1H weighs 1.6735 x 10-24 g and 16O 2.6560 x 10-23 g. • We define: mass of 12C = exactly 12 amu. • Using atomic mass units: • 1 amu = 1.66054 x 10-24 g • 1 g = 6.02214 x 1023 amu Chapter 2

23. Atomic Weights • Average Atomic Masses • Relative atomic mass: average masses of isotopes: • Naturally occurring C: 98.892 %12C + 1.108 %13C. • Average mass of C: • (0.98892)(12 amu) + (0.0108)(13.00335) = 12.011 amu. • Atomic weight (AW) is also known as average atomic mass (atomic weight). • Atomic weights are listed on the periodic table. Chapter 2

24. The Periodic Table • The Periodic Table is used to organize the 114 elements in a meaningful way. • As a consequence of this organization, there are periodic properties associated with the periodic table. Chapter 2

25. The Periodic Table

26. The Periodic Table • Columns in the periodic table are called groups (numbered from 1A to 8A or 1 to 18). • Rows in the periodic table are called periods. • Metals are located on the left hand side of the periodic table (most of the elements are metals). • Non-metals are located in the top right hand side of the periodic table. • Elements with properties similar to both metals and non-metals are called metalloids and are located at the interface between the metals and non-metals. Chapter 2

27. The Periodic Table • Some of the groups in the periodic table are given special names. • These names indicate the similarities between group members: • Group 1A: Alkali metals. • Group 2A: Alkaline earth metals. • Group 6A: Chalcogens. • Group 7A: Halogens. • Group 8A: Noble gases. Chapter 2

28. Molecules and Molecular Compounds • Molecules and Chemical Formulas • Molecules are assemblies of two or more atoms bonded together. • Each molecule has a chemical formula. • The chemical formula indicates • which atoms are found in the molecule, and • in what proportion they are found. • Compounds formed from molecules are molecular compounds. • Molecules that contain twoatoms bonded together are called diatomic molecules. Chapter 2

29. Molecules and Molecular Compounds Molecules and Chemical Formulas Chapter 2

30. Molecules and Molecular Compounds • Molecular and Empirical Formulas • Molecular formulas • give the actual numbers and types of atoms in a molecule. • Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4. Chapter 2

31. Molecules and Molecular Compounds • Molecular and Empirical Formulas • Molecular formulas Chapter 2

32. Molecules and Molecular Compounds • Molecular and Empirical Formulas • Empirical formulas • give the relative numbers and types of atoms in a molecule. • That is, they give the lowest whole number ratio of atoms in a molecule. • Examples: H2O, CO2, CO, CH4, HO, CH2. Chapter 2

33. Molecules and Molecular Compounds • Picturing Molecules • Molecules occupy three dimensional space. • However, we often represent them in two dimensions. • The structural formula gives the connectivity between individual atoms in the molecule. • The structural formula may or may not be used to show the three dimensional shape of the molecule. • If the structural formula does show the shape of the molecule, then either a perspective drawing, ball-and-stick model, or space-filling model is used. Chapter 2

34. Molecules and Molecular Compounds Picturing Molecules Chapter 2

35. Ions and Ionic Compounds • When an atom or molecule loses electrons, it becomes positively charged. • For example, when Na loses an electron it becomes Na+. • Positively charged ions are called cations. Chapter 2

36. Ions and Ionic Compounds • When an atom or molecule gains electrons, it becomes negatively charged. • For example when Cl gains an electron it becomes Cl-. • Negatively charged ions are called anions. • An atom or molecule can lose more than one electron. Chapter 2

37. Ions and Ionic Compounds • In general: metal atoms tend to lose electrons to become cations; nonmetal ions tend to gain electrons to form anions. • Predicting Ionic Charge • The number of electrons an atom loses is related to its position on the periodic table. Chapter 2

38. Ions and Ionic Compounds Predicting Ionic Charge Chapter 2

39. Ions and Ionic Compounds • Ionic Compounds • The majority of chemistry involves the transfer of electrons between species. • Example: • To form NaCl, the neutral sodium atom, Na, must lose an electron to become a cation: Na+. • The electron cannot be lost entirely, so it is transferred to a chlorine atom, Cl, which then becomes an anion: Cl-. • The Na+ and Cl- ions are attracted to form an ionic NaCl lattice which crystallizes. Chapter 2

40. Ions and Ionic Compounds Ionic Compounds Chapter 2

41. Ions and Ionic Compounds • Ionic Compounds • Important: note that there are no easily identified NaCl molecules in the ionic lattice. Therefore, we cannot use molecular formulas to describe ionic substances. • Consider the formation of Mg3N2: • Mg loses two electrons to become Mg2+; • Nitrogen gains three electrons to become N3-. • For a neutral species, the number of electrons lost and gained must be equal. Chapter 2

42. Ions and Ionic Compounds • Ionic Compounds • However, Mg can only lose electrons in twos and N can only accept electrons in threes. • Therefore, Mg needs to lose 6 electrons (2  3) and N gain those 6 electrons (3  2). • I.e., 3Mg atoms need to form 3Mg2+ ions (total 3  2+ charges) and 2 N atoms need to form 2N3- ions (total 2  3- charges). • Therefore, the formula is Mg3N2. Chapter 2

43. Naming Inorganic Compounds • Naming of compounds, nomenclature, is divided into organic compounds (those containing C) and inorganic compounds (the rest of the periodic table). • Cations formed from a metal have the same name as the metal. • Example: Na+ = sodium ion. • If the metal can form more than one cation, then the charge is indicated in parentheses in the name. • Examples: Cu+ = copper(I); Cu2+ = copper(II). Chapter 2

44. Naming Inorganic Compounds • Cations formed from non-metals end in -ium. • Example: NH4+ ammonium ion. Chapter 2

45. Naming Inorganic Compounds Chapter 2

46. Naming Inorganic Compounds • Monatomic anions (with only one atom) are called -ide. • Example: Cl- is chloride ion. • Exceptions: hydroxide (OH-), cyanide (CN-), peroxide (O22-). • Polyatomic anions (with many atoms) containing oxygen end in -ate or -ite. (The one with more oxygen is called -ate.) • Examples: NO3- is nitrate, NO2- is nitrite. Chapter 2

47. Naming Inorganic Compounds • Polyatomic anions containing oxygen with more than two members in the series are named as follows (in order of decreasing oxygen): • per-….-ate • -ate • -ite • hypo-….-ite Chapter 2

48. Naming Inorganic Compounds Chapter 2

49. Naming Inorganic Compounds • Polyatomic anions containing oxygen with additional hydrogens are named by adding hydrogen or bi- (one H), dihydrogen (two H), etc., to the name as follows: • CO32- is the carbonate anion • HCO3- is the hydrogen carbonate (or bicarbonate) anion. • H2PO4- is the dihydrogen phosphate anion. Chapter 2

50. Naming Inorganic Compounds • Name the anion then cation for the ionic compound. • Example: BaBr2 = barium bromide. Chapter 2