Lecture 3 The Periodic Table, Atomic Structure, Isotopes and Ions, Basic Nomenclature ( Ch 2)

1. Lecture 3The Periodic Table, Atomic Structure, Isotopes and Ions, Basic Nomenclature (Ch 2) Dr Harris Suggested HW: (Ch 2) 15, 19, 23, 28, 38, 49, 52, 58

2. Metallic vs Nonmetallic Elements • Elements can be metals, nonmetals, or semiconductors(we will discuss semiconductors later) • Metal characteristics • Malleable • Ductile • Conductive of electricity • Conductive of heat • Have luster and shine

3. Metallic vs Nonmetallic Character Br2 (L) Nonmetal characteristics • Most nonmetals are gases • Non conductive • Nonmetallic solids are not conductive, malleable, or ductile O2 (g) P(s) S(s)

4. Metals are all elements LEFT of the black line, except Hydrogen.

5. Elemental Groups and Diatomic Species • Elements listed down a column have very similar properties, and tend to behave the same way. These columns are called groups. We will revisit chemical groups later. • Certain elements are unstable, and hence, do not commonly exist as individual species, but as diatomic molecules • These include H, O, N, and all of the halogens (group 17) • H  H2 (Hydrogen gas) • O  O2 (Oxygen gas) • N  N2 (Nitrogen gas) • F, Cl, Br, I  F2, Cl2, Br2, I2 (Fluorine gas, chlorine gas, bromine gas and iodine gas)

6. Phases : Solids, Liquids, and Gases Solids • Atoms tightly bound • Fixed volume and shape (does not conform to container) • A chemical is denoted as solid by labeling it with (s) Au atoms Au (s)

7. Solids, Liquids, and Gases • Liquids • Atoms less tightly bound than solids • Has a definite volume, but not definite shape (assumes the shape of its container) • Denoted by (L) ex. H2O (L) Water molecule

8. Solids, Liquids, and Gases • Gases • Free atoms • No shape, no definite volume • Can be expanded or compressed (like engine piston) • Denoted by (g) ; ex. O2 (g) O2 molecules

9. Mixtures • Most substances in nature are mixtures, chemicals that exist together without actually bonding. Mixtures can be homogeneous (uniform throughout) or heterogeneous (non-uniform) Homogeneous Heterogeneous • water • fructose • CO2(g) • food coloring AIR • 78% N2 • 21% O2

10. Solutions • All solutionsare homogeneousfluids (free flowing), meaning that the composition is completely uniform throughout • Example: salt water. You add salt (NaCl) to water, it completely dissociates(dissolves; molecules split apart) and spreads evenly throughout the water. There is the same amount of salt everywhere in the glass • Salt is the solute. Water is the solvent. • When a substance dissolves in water, it forms an aqueous solution, labeled (aq): Salt water  NaCl(aq) • Air is another example of a solution

11. Law of Constant Composition

12. Example • We analyze 1.630 g of CaS and find that it’s 0.906 g Ca. Find the mass of S? Find the mass % of Ca and S? • First, we know that CaS is composed only of Ca and S. Therefore, all of its mass must come from Ca and S. We can find the mass of S. • mass Ca + mass S = total mass of CaS 0.906g + X = 1.630 g X = 0.724 g S • Now, we can find the mass % of each element mass % Ca + mass % S = 100%

13. Example contd. • Going back to the law of constant composition and our previous example of CaS, the results of our calculation show that all samples of CaS are 55.6% Ca and 44.4% S • This is true no matter how large the sample and regardless of its origin. This is the basis of the Law of Constant Composition.

14. Group Examples • A sample of KI has a mass of 2.00 g and is 23.5% K. What is the mass % I? What is the mass of I in the compound. • A sample of NaCl is 39.3% Na, and is found to contain 1.58 g Na. What is the total mass of NaCl? What is the mass of Cl?

15. Basic Laws of Chemical Reactions • Law of conservation of mass: Atoms are not created or destroyed in a chemical reaction, only rearranged • Law of constant composition: Specific molecules have fixed ratios of atoms • One N atom combines with 3 H atoms to make ammonia gas, NH3 (g). It does not matter how much N or H you have. • So, if you had 10 N atoms and 1000 H atoms, you would make 10 NH3molecules, with 970 unused H atoms left over. • Some elements combine in multiple ratios, like C and O • * CO  carbon monoxide • * CO2  carbon dioxide

16. Dalton’s Atomic Theory • Matter is composed of atoms • The atoms of a given element are exactly the same in every way • The atoms of different elements differ in mass and physical properties • Compounds are composed of two or more atoms of different elements bonded together. • In a reaction, atoms are rearranged, separated or recombined to form new substances. No atoms are created or destroyed, and the atoms themselves DO NOT CHANGE

17. Atomic Mass Units • Each element in the periodic table is assigned a relativemass, called its atomic mass, given in atomic mass units (1 amu = 1.66 x 10-24 g) • Remember the CaS example. The total mass of CaS is 55.6% Ca and 44.4% S. If all the mass is from Ca and S, then the ratio of the mass % must also be the ratio of the atomic masses • This suggests that a Ca atoms weighs 1.25 times as much as an S atom • To develop the periodic table, atom masses are determined by using the carbon-12 isotope as a reference

18. Atomic Mass Units

19. Calculating Molecular Mass • What is molecular mass of H2O to 3 sig figs? • H = 1.0079 amu • O = 15.9994 amu • What is the molecular mass of NaHCO3 to 3 sig figs? • Na = 22.990 amu • C = 12.01 amu

20. % Mass Calculations Using Atomic Mass • What is the % mass of Na in Na2S? • How many grams of sulfur atoms are there in a 4.50 g sample of Na2S?

21. Atomic Structure • Atoms were initially thought to be indivisible, and the smallest particles in existence • Toward the end of the 19th century, it was discovered that atoms consist of even smaller sub-atomic particles • G. JohnstoneStoney hypothesized that electricity exists in discrete units, or small packets of equally distributed charge. He called these electric units electrons • J.J. Thompson was one of the first scientists to proved the existence of such particles. He is most famous for his cathode-ray experiment

22. Cathode Rays • Electrical discharge (current) in evacuated tubes had been previously observed when high voltages were applied between two electrodes with slits in them. • These rays of current are called cathode rays, because they flowed from the cathode (negative end) to the anode (positive end) • Scientist initially thought that these rays were caused by atoms of the electrodes.

23. Discovery of The Electron • J.J. Thomson modified the typical cathode ray tube experiment in several ways: • He used various materials to compose the cathode. It was found that the nature of the beams did notdepend at all on what the cathode was composed of, and the mass of the cathode never changed. Hence, these rays were NOT caused by atoms or heavy particles. • He applied electrical fields to the cathode ray. He found that the cathode rays were actually beams of negatively charge particles that are MUCH lighter than atoms, and are the same in every element. This marked the discovery of the electron.

24. Radioactivity • Shortly after the discovery of electrons, Antoine Henri Becquerel discovered radioactivity, the process by which atoms spontaneously break apart. • By monitoring emissions from a radioactive material under an electric field, Ernest Rutherford found that radiation consists of three types of radiation • alpha (α) particles • beta (β) particles masses below are amu • gamma (ϒ) particles charges are relative

25. The “Plum Pudding” Model • Since atoms are electrically neutral, and electrons are negatively charged, Thomson knew that a positively charged sub-atomic particle must also exist • Thomson envisioned the atom as electrons evenly distributed in a “sea” of positive charge, the so-called “plum pudding” model (shown to the right) • Ernest Rutherford then formed an experiment to see how a high energy beam of positively charged α particles would interact with atoms positive charge electrons

26. Discovery of the Nucleus • A very thin sheet of gold film was placed in front of a beam of α-particles • Most of the α-particles passed through the gold film • However, some α-particles were scattered at large angles. How could this happen? • Rutherford realized that most of the mass of an atom is concentrated in a small positively-charged region. He called this region the nucleus.

27. The Atom Nucleus Electron cloud α particle • Most of the volume of an atom is empty space, which electrons spread sparsely throughout. The α-particles easily pass through this space. • However, the nucleus is very dense, and very positively charged. The α-particles that approach this region are strongly repelled.

28. Protons and Neutrons Comprise the Nucleus • Later experiments showed that protonsand neutronsreside in the nucleus and comprise the bulk of the mass of an atom • Protons and electron have equal but opposite charge Charges shown in table are relative to the charge of a proton. A proton has an actual charge of 1.602 x 10-19 C, an electron has a charge of -1.602 x 10-19 C. Opposite charges attract! Like charges repel!!

29. Atomic Number • The number of protons in an atom is called the atomic number. An element is defined by its atomic number. (ex. only carbon has 6 protons) • For a given element, the number of protons DOES NOT CHANGEduring a reaction. • In a neutralatom, the number of protons is equal to the number of electrons.

30. Mass Number • The mass number of an element is the sum of the protons and neutrons. • The mass numbers listed on the periodic table are average values. • The reason for these averages is that elements exist in nature as multiple isotopes.

31. Isotopes • Isotopes are variations of elements with the same number of protons but different numbers of neutrons. • For example, the most common isotope of hydrogen contains one proton, one electron, and no neutrons (99.985% of all hydrogen atoms). • An isotope of hydrogen is deuterium, which has one neutron (.0115 %) • A third isotope, tritium, has two neutrons (~ 0%). mass number hydrogen atomic number deuterium tritium

32. Average atomic mass is obtained using the % abundance and the isotope mass.

33. Example • Using these values, calculate the average atomic mass of Hydrogen. Does it match the value given in the periodic table?

34. Group Problem • The 3 main isotopes of Carbon are and Fill in the table below with the correct number of protons, neutrons, and electrons. • Then, calculate the average atomic mass. Does your value match that of the periodic table?

35. How To Determine % abundance • The average atomic mass of Boron, as given on the periodic table, is 10.811 amu • It is known that two isotopes of Boron exist, • (isotopic mass: 10.013 amu) • (isotopic mass: 11.009 amu) • What are the % abundances of each isotope?

36. Continued. • We know that: • We also know that the total % abundance must be 100% • We allow • Therefore: • Solving for x, we find that

37. Group Example • Oxygen has three isotopes. Given that the abundance of 18O is 0.205%, calculate the abundances of the other two isotopes. • (isotopic mass: 15.994 914 amu) • (isotopic mass: 16.999 131 amu) • (isotopic mass: 17.999 160 amu)

38. Ions • Thus far, we’ve learned than an element is essentially defined by it’s atomic number • Each element has an exact number of protons. • For example, Hydrogen has only one proton. If you force a second proton onto the atom, you no longer have hydrogen… you now have Helium. • We have also learned that atoms of a particular element can have variations in the number of neutrons. Atoms of the same element with varying numbers of neutrons are called isotopes. • Next, we will discuss ions.

39. Ions • Ions are electrically charged atoms, resulting from the gain or loss of electrons. • Positively charged ions are called cations. You form cations when electrons are lost • Negatively charged ions are called anions. You form anions when electrons are gained • A cation is named by adding the word “ion” to the end of the element name • Anions are named by adding the suffix –ide to the end of an element

40. Ions Lithium ion Sodium ion Magnesium ion Aluminum ion Chloride Sulfide Oxide Phosphide

41. Ions Gains 2 electrons Sulfur-32 Sulfide-32 Loses 3 electrons Aluminum-27 Aluminum ion-27

42. Nomenclature • Example: KF • K is a metal. F is a non metal. • We write the name as: Potassium Fluoride • There are special rules to naming molecules. In this chapter, you will see two types of molecules: IONICand COVALENT • Ionic bonds are formed between metals and nonmetals • To name an ionic compound, you do the following • write the name of the metal • Follow it with the name of the nonmetal, but change the ending of the name of the nonmetal to –ide Na2S = Sodium Sulfide; MgO = Magnesium Oxide

43. Nomenclature SF4 = sulfur tetrafluoride N2O = dinitrogenmonoxide P2O3 = diphosphorustrioxide SiC = silicon carbide • Covalent bonds are formed between nonmetals and nonmetals • To name a covalent compound, you do the following • Include a prefix corresponding to the first nonmetal. Do not use mono on the first nonmental. • Use a prefix to name the second nonmetal. We only use –mono for oxygen. Drop the ending of the second nonmetal and replace it with –ide

44. Rules of Hydrogen • Hydrogen is strange. It’s a nonmetal, but can sometimes react as metals do, and is listed on the left (metallic) side of the periodic table. 1. If hydrogen is listed first, and the nonmetal is of group 16, treat it as a metal and name the compound accordingly (except water). Ex. H2S = hydrogen sulfide Compounds of group 17 are acidsand are named as such. We drop –gen and end the second nonmetal with the suffix “–icacid” HCl = hydrochloric acid ; HF = hydrofluoric acid 2. If hydrogen is listed last, treat the molecule as covalent. Change the ending to –ide and use the appropriate prefix. MgH2= magnesium hydride; AsH3= arsenic trihydride