Chapter 2 Atoms and Elements

1. Chapter 2Atoms and Elements Principles of Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro Dr. Juana Mendenhall Morehouse College

2. LECTURE 3: Sept. 1, 2010 Objectives: State & interpret the law of conservation of mass State & describe the law of definite proportions State & explain the law of multiple proportions Explain Dalton’s Atomic Theory Define nucleus, protons, neutrons, and electrons Define Atomic mass unit (amu), Atomic number, and chemical symbol Define and categorize isotopes, mass number, and atomic number Define ions, anions, and cations

3. Law of Conservation of Mass In a chemical reaction, matter is neither created nor destroyed. Total mass of the materials you have before the reaction must equal the total mass of the materials you have at the end. total mass of reactants = total mass of products Antoine Lavoisier 1743–1794

4. Law of Definite Proportions All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements. Joseph Proust 1754–1826

5. Example 2.1 Show that two samples of carbon dioxide obey the Law of Definite Proportions g O1, g C1 g O2, g C2 O:C in each sample Given: Find: Sample 1: 25.6 g O and 9.60 g C Sample 2: 21.6 g O and 8.10 g C Proportion O:C Conceptual Plan: Relationships: All samples of a compound have the same proportion of elements by mass. Solution: Check: Since both samples have the same O:C ratio, the result is consistent with the Law of Definite Proportions.

6. Law of Multiple Proportions When two elements (call them A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small whole numbers. John Dalton 1766–1844

7. Dalton’s Atomic Theory Dalton proposed a theory of matter based on it having ultimate, indivisible particles to explain these laws. Each element is composed of tiny, indestructible particles called atoms. All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements. Atoms combine in simple, whole-number ratios to form molecules of compounds. In a chemical reaction, atoms of one element cannot change into atoms of another element. They simply rearrange the way they are attached.

8. Practice—Decide if each statement is correct according to Dalton’s Model of the Atom Copper atoms can combine with zinc atoms to make gold atoms. Water is composed of many identical molecules that have one oxygen atom and two hydrogen atoms.

9. Charges of Atoms two kinds of charge called + and – opposite charges attract + attracted to – like charges repel + repels + – repels – To be neutral, something must have no charge or equal amounts of opposite charges.

10. Some Problems How could beryllium have 4 protons stuck together in the nucleus? Shouldn’t they repel each other? If a beryllium atom has 4 protons, then it should weigh 4 amu; but it actually weighs 9.01 amu! Where is the extra mass coming from? Each proton weighs 1 amu. Remember, the electron’s mass is only about 0.00055 amu and Be has only 4 electrons—it can’t account for the extra 5 amu of mass.

11. Electrons Electrons are particles found in all atoms. Cathode rays are streams of electrons. The electron has a charge of −1.60 × 1019 C. The electron has a mass of 9.1 × 10−28 g.

12. Structure of the Atom Rutherford proposed that the nucleus had a particle that had the same amount of charge as an electron but opposite sign. based on measurements of the nuclear charge of the elements These particles are called protons. charge = +1.60 × 1019 C mass = 1.67262 × 10−24 g Since protons and electrons have the same amount of charge, for the atom to be neutral there, must be equal numbers of protons and electrons.

13. Relative Mass and Charge It is sometimes easier to compare things to each other rather than to an outside standard. When you do this, the scale of comparison is called a relative scale. We generally talk about the size of charge on atoms by comparing it to the amount of charge on an electron, which we call −1 charge units. A proton has a charge of +1 cu. Protons and electrons have equal amounts of charge, but opposite signs. We generally talk about the mass of atoms by comparing it to 1/12th the mass of a carbon atom with 6 protons and 6 neutrons, which we call 1 atomic mass unit. Protons have a mass of 1 amu. Electrons have a mass of 0.00055 amu, which is generally too small to be relevant.

14. Summary of Subatomic Particles

15. Elements Each element has a unique number of protons in its nucleus. The number of protons in the nucleus of an atom is called the atomic number. The elements are arranged on the Periodic Table in order of their atomic numbers. Each element has a unique name and symbol. symbol either one or two letters one capital letter or one capital letter and one lowercase

16. Structure of the Nucleus Soddy discovered that the same element could have atoms with different masses, which he called isotopes. There are two isotopes of chlorine found in nature, one that has a mass of about 35 amu and another that weighs about 37 amu. The observed mass is a weighted average of the weights of all the naturally occurring atoms. The percentage of an element that is one isotope is called the isotope’s natural abundance. The atomic mass of chlorine is 35.45 amu.

17. Atomic Mass We previously learned that not all atoms of an element have the same mass. isotopes We generally use the average mass of all an element’s atoms found in a sample in calculations. However, the average must take into account the abundance of each isotope in the sample. We call the average mass the atomic mass.

18. Isotopes All isotopes of an element are chemically identical. undergo the exact same chemical reactions All isotopes of an element have the same number of protons. Isotopes of an element have different masses. Isotopes of an element have different numbers of neutrons. Isotopes are identified by their mass numbers. protons + neutrons

19. Isotopes • Atomic Number • Number of protons • Z • Mass Number • Protons + Neutrons • Whole number • A • Abundance = relative amount found in a sample

20. Which of these are isotopes of Nitrogen:

21. LECTURE 4: September 3, 2010 Objectives: Use periodic table to identify properties of elements Calculate atomic mass, mole, and number of particles (atoms, things)

22. Neon Symbol Number of Protons Number of Neutrons A, Mass Number Percent Natural Abundance Ne-20 or 10 10 20 90.48% Ne-21 or 10 11 21 0.27% Ne-22 or 10 12 22 9.25%

23. Example 2.3b How many protons, electrons, and neutrons are in an atom of ? atomic number symbol # p+ # e- atomic & mass numbers symbol # n0 Given: Find: therefore, A = 52, Z = 24 # p+, # e−, # n0 Conceptual Plan: Relationships: in neutral atom, # p+ = # e− mass number = # p+ + # n0 Solution: Z = 24 = # p+ # e− = # p+ = 24 A = Z + # n0 52 = 24 + # n0 28 = # n0 Check: for most stable isotopes, n0≥ p+

24. Practice—Complete the Table

25. Practice—Complete the Table

26. Reacting Atoms When elements undergo chemical reactions, the reacting elements do not turn into other elements. Statement 4 of Dalton’s Atomic Theory This requires that all the atoms present when you start the reaction will still be there after the reaction. Since the number of protons determines the kind of element, the number of protons in the atom does not change in a chemical reaction. However, many reactions involve transferring electrons from one atom to another.

27. Charged Atoms When atoms gain or lose electrons, they acquire a charge. Charged particles are called ions. When atoms gain electrons, they become negatively charged ions, called anions. When atoms lose electrons, they become positively charged ions, called cations.

28. Ions and Compounds • Ions behave much differently than the neutral atom. • e.g., The metal sodium, made of neutral Na atoms, is highly reactive and quite unstable. However, the sodium cations, Na+, found in table salt are very nonreactive and stable. • Since materials like table salt are neutral, there must be equal amounts of charge from cations and anions in them.

29. A Useful Tool in Describing Chemical Compounds Non-metals tend to gain electrons. Cl + e- Cl- Metals tend to lose electrons. Na  Na+ + e- Ion: form with different # electrons & protons Anion = extra electron(s) Cation = fewer electron (s)

30. Atomic Structures of Ions Nonmetals form anions. For each negative charge, the ion has 1 more electron than the neutral atom. F = 9 p+ and 9 e−, F─ = 9 p+ and 10 e− P = 15 p+ and 15 e−, P3─ = 15 p+ and 18e− Anions are named by changing the ending of the name to -ide. fluorine F + 1e− F─fluoride ion oxygen O + 2e− O2─ oxide ion

31. Atomic Structures of Ions Metals form cations. For each positive charge, the ion has 1 less electron than the neutral atom. Na atom = 11 p+ and 11 e−, Na+ ion = 11 p+ and 10 e− Ca atom = 20 p+ and 20 e−, Ca2+ ion = 20 p+ and 18e− Cations are named the same as the metal. sodium Na  Na+ + 1e− sodium ion calcium Ca  Ca2+ + 2e− calcium ion

32. Practice—Complete the Table

33. Practice—Complete the Table

34. = Alkali Metals = Alkali Earth Metals = Noble Gases = Transition Metals = Halogens = Lanthanides = Actinides

35. Noble Gases + 1 +3 ±4 -3 -2 -1 +2 -1 Transition Metals Main Group Lanthanides and Actinides The Periodic table

36. Example 2.5 If copper is 69.17% Cu–63 with a mass of 62.9396 amu and the rest Cu–65 with a mass of 64.9278 amu, find copper’s atomic mass isotope masses, isotope fractions avg. atomic mass Given: Find: Cu–63 = 69.17%, 62.9396 amu Cu–65 = 100 − 69.17%, 64.9278 amu atomic mass, amu Conceptual Plan: Relationships: Solution: Check: The average is between the two masses, closer to the major isotope.

37. Counting Atoms by Moles If we can find the mass of a particular number of atoms, we can use this information to convert the mass of an element sample into the number of atoms in the sample. The number of atoms we will use is 6.022 x 1023 and we call this a mole. 1 mole = 6.022 × 1023things (atoms)

38. Relationship between Moles and Mass 6.02214 x 1023 atoms = 1 mol 6.02214 x 1023 atoms = Molar Mass of element Molar mass = g of element over 1 mol

39. Example 2.6 Calculate the number of atoms in 2.45 mol of copper mol Cu atoms Cu Given: Find: 2.45 mol Cu atoms Cu Conceptual Plan: Relationships: 1 mol = 6.022 × 1023 atoms Solution: Since atoms are small, the large number of atoms makes sense. Check:

40. Chemical Packages—Moles mole = number of particles equal to the number of atoms in 12 g of C-12 1 atom of C-12 weighs exactly 12 amu. 1 mole of C-12 weighs exactly 12 g. The number of particles in 1 mole is called Avogadro’s Number =6.0221421 × 1023. 1 mole of C atoms weighs 12.01 g and has 6.022 × 1023 atoms. The average mass of a C atom is 12.01 amu.

41. Mole and Mass Relationships 1 mole sulfur 32.06 g 1 mole carbon 12.01 g

42. Example 2.7 Calculate the moles of carbon in 0.0265 g of pencil lead g C mol C Given: Find: 0.0265 g C mol C Conceptual Plan: Relationships: 1 mol C = 12.01 g Solution: Check: Since the given amount is much less than 1 mol C, the number makes sense.

43. Practice—Calculate the moles of sulfur in 57.8 g of sulfur

44. g S mol S Practice—Calculate the moles of sulfur in 57.8 g of sulfur Given: Find: 57.8 g S mol S Conceptual Plan: Relationships: 1 mol S = 32.07 g Solution: Check: Since the given amount is much less than 1 mol S, the number makes sense.

45. Example 2.8 How many copper atoms are in a penny weighing 3.10 g g Cu mol Cu atoms Cu Given: Find: 3.10 g Cu atoms Cu Conceptual Plan: Relationships: 1 mol Cu = 63.55 g, 1 mol = 6.022 × 1023 Solution: Since the given amount is much less than 1 mol Cu, the number makes sense. Check:

46. Practice—How many aluminum atoms are in a can weighing 16.2 g

47. g Al mol Al atoms Al Practice—How many aluminum atoms are in a can weighing 16.2 g Given: Find: 16.2 g Al atoms Al Conceptual Plan: Relationships: 1 mol Al = 26.98 g, 1 mol = 6.022 × 1023 Solution: Check: Since the given amount is much less than 1 mol Al, the number makes sense.