Atoms, Molecules, and Ions Chapter 2

1. Atoms, Molecules, and IonsChapter 2

2. 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.

3. 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.

4. The Discovery of Atomic Structure 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. 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.

5. The Discovery of Atomic Structure Cathode Rays and Electrons

6. The Discovery of Atomic Structure • Cathode Rays and Electrons • 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. • 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.

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

8. The Discovery of Atomic Structure Millikan Oil Drop Experiment

9. 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.

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.

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.

12. The Discovery of Atomic Structure Radioactivity

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 (He atom).

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’s a-particle experiment:

16. 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.

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. Atoms are mostly empty space!!

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.

20. The Modern View of Atomic Structure

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 Isotopes have the same Z but different A.

22. The Modern View of Atomic Structure Isotopes, Atomic Numbers, and Mass Numbers

23. Periodic Table

24. 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.

25. The Periodic Table

26. 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.

27. Molecules and Molecular Compounds • Molecules and Chemical Formulas • Molecules are assemblies of two or more atoms bonded together. • The chemical formula indicates • which atoms are found in the molecule, and in what proportion they are found. • Compounds formed from molecules are called molecular compounds.

28. 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. • 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.

29. Molecules and Molecular Compounds Picturing Molecules 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.

30. Molecular Naming Two non-metals Same side of Table

31. Naming 2 nonmetals Binary Molecular Compounds The most metallic element is usually written first (i.e., the one to the farthest left on the periodic table). Exception: NH3. If both elements are in the same group, the lower one is written first. Greek prefixes are used to indicate the number of atoms.

32. Naming Inorganic Compounds Names and Formulas of Binary Molecular Compounds

33. Ionic Compounds Opposite sides of Periodic Table MetalsNonmetals

34. 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.

35. Ions and 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.

36. 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.

37. 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 / molecule can lose more than 1 electron.

38. Naming Inorganic Compounds • Names and Formulas of Ionic Compounds • Name the cation then anion for the ionic compound. • Example: BaBr2 = barium bromide.

39. Ions and Ionic Compounds Predicting Ionic Charge The number of electrons an atom loses is related to its position on the periodic table. Metals tend to form cations whereas non-metals tend to form anions.

40. Naming Inorganic Compounds • Names and Formulas of Ionic Compounds • 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 (usually Transition Metals). • Examples: Cu+ = copper(I); Cu2+ = copper(II).

41. Naming Inorganic Compounds • Names and Formulas of Ionic 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.

42. Naming Inorganic Compounds Names and Formulas of Ionic Compounds Polyatomic anions containing oxygen with more than two members in the series are named as follows (in order of decreasing oxygen): per-…..-ate ClO4-perchlorate -ate ClO3- chlorate -ite ClO2- chlorite hypo-….-ite ClO -hypochlorite

43. Naming Inorganic Compounds • Names and Formulas of Ionic Compounds • Polyatomic anions containing oxygen with additional hydrogensare 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 dihydrogenphosphate anion.

44. Naming Inorganic Compounds • Names and Formulas of Acids • The names of acids are related to the names of anions: • -idebecomes hydro-….-icacid; • -ate becomes -ic acid; • -ite becomes -ous acid.

45. PERIODIC TRENDS

46. Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity

47. Effective Nuclear Charge, Z* (kernel Charge) • Z* is the nuclear charge experienced by the outermost electrons. • Explains why Energy(2s) < Energy(2p) • Z* increases across periodic table

48. Orbitals in Many Electron Atoms Effective Nuclear Charge • Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nuclear charge. • The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons. • As the average number of screening electrons (S) increases, the effective nuclear charge (zeff) decreases. • zeff = z - S

49. Effective Nuclear Charge, Z* • Atom Z* Experienced by 2s Electrons in Valence Orbitals • Li +1.28 • B +2.58 • C +3.22 • N +3.85 • O +4.49 • F +5.13 Increase in Z* across a period [Values calculated using Slater’s Rules]

50. Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity