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Chapter 2: Atoms and Molecules

Chapter 2: Atoms and Molecules. Chapter 2 Objectives. LEARNING OBJECTIVES/ASSESSMENT When you have completed your study of this chapter, you should be able to: 1. Use symbols for chemical elements to write formulas for chemical compounds. (Section 2.1; Exercise 2.4)

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Chapter 2: Atoms and Molecules

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  1. Chapter 2:Atoms and Molecules

  2. Chapter 2 Objectives • LEARNING OBJECTIVES/ASSESSMENT • When you have completed your study of this chapter, you should be able to: • 1. Use symbols for chemical elements to write formulas for chemical compounds. (Section 2.1; Exercise 2.4) • 2. Identify the characteristics of protons, neutrons, and electrons. (Section 2.2; Exercises 2.10 and 2.12) • 3. Use the concepts of atomic number and mass number to determine the number of subatomic particles in isotopes and to write correct symbols for isotopes. (Section 2.3; Exercises 2.16 and 2.22) • 4. Use atomic weights of the elements to calculate molecular weights of compounds. (Section 2.4; Exercise 2.32) • 5. Use isotope percent abundances and masses to calculate atomic weights of the elements. (Section 2.5; Exercise 2.38) [just the qualitative part – no need to do calc] • 6. Use the mole concept to obtain relationships between number of moles, number of grams, and number of atoms for elements, and use those relationships to obtain factors for use in factor‐unit calculations. (Section 2.6; Exercises 2.44 a & b and 2.46 a & b) • 7. Use the mole concept and molecular formulas to obtain relationships between number of moles, number of grams, and number of atoms or molecules for compounds, and use those relationships to obtain factors for use in factor‐unit calculations. (Section 2.7; Exercise 2.50 b and 2.52 b)

  3. Questions for Chapter Two Example exercises include: 2.5 – 2.6 2.9 – 2.22 2.24 – 2.25 2.29 – 2.31 2.41 – 2.48 2.51 – 2.56

  4. SYMBOLS & FORMULAS • A unique symbol is used to represent each element. • Formulas are used to represent compounds. • ELEMENTAL SYMBOLS • A symbol is assigned to each element. The symbol is based on the name of the element and consists of one capital letter or a capital letter followed by a lower case letter. • Some symbols are based on the Latin or German name of the element.

  5. CHEMICAL ELEMENTS & THEIR SYMBOLS

  6. Elements and their symbols • Should know first 20 elements • Also know • Fe – iron • Cu – copper • Zn – zinc • Br – bromine • Ag – silver • Sn – tin • I – iodine • Au – gold • Hg – mercury • Pb - lead

  7. COMPOUND FORMULAS • A compound formula consists of the symbols of the elements found in the compound. Each elemental symbol represents one atom of the element. If more than one atom is represented, a subscript following the elemental symbol is used.

  8. COMPOUND FORMULAS EXAMPLES • Carbon monoxide, CO • one atom of C • one atom of O • Water, H2O • two atoms of H • one atom of O • Ammonia, NH3 • one atom of N • 3 atoms of H

  9. ATOMIC STRUCTURE • Atoms are made up of three subatomic particles, protons, neutrons, and electrons. • The protons and neutrons are tightly bound together to form the central portion of an atom called the nucleus. • The electrons are located outside of the nucleus and thought to move very rapidly throughout a relatively large volume of space surrounding the small but very heavy nucleus.

  10. SUBATOMIC PARTICLES • Protons are located in the nucleus of an atom. They carry a +1 electrical charge and have a mass of 1 atomic mass unit (u). • Neutrons are located in the nucleus of an atom. They carry no electrical charge and have a mass of 1 atomic mass unit (u). • Electrons are located outside the nucleus of an atom. They carry a -1 electrical charge and have a mass of 1/1836 atomic mass unit (u). They move rapidly around the heavy nucleus.

  11. Powers of 10 http://www.powersof10.com/film

  12. SUBATOMIC PARTICLE CHARACTERISTICS

  13. ATOMIC STRUCTURE REVIEW • Which subatomic particles are represented by the pink spheres? • Which subatomic particles are represented by the yellow and blue spheres? • What structure do the yellow and blue spheres form?

  14. ATOMIC & MASS NUMBERS • ATOMIC NUMBER OF AN ATOM • The atomic number of an atom is equal to the number of protons in the nucleus of the atom. • Atomic numbers are represented by the symbol Z. • MASS NUMBER OF AN ATOM • The mass number of an atom is equal to the sum of the number of protons & neutrons in the nucleus of the atom. • Mass numbers are represented by the symbol A.

  15. ATOMIC & MASS NUMBERS APPLICATION • Based on the information given above, what is the atomic number and mass number of fluorine? • Note: The periodic table does not show the mass number for an individual atom. It lists an average actual mass (atomic weight) for a collection of atoms!

  16. ATOMIC & MASS NUMBERS APPLICATION • Consider Chlorine – 35 • Write the symbol for this atom • What is the mass number and atomic number • How many p, n and e- are in this atom • An atom of carbon contains 8 neutrons. • How many p, n and e- are in this atom • What is the name of this atom? • Write its formula • What is the mass number and atomic number • An atom has 19 protons and 20 neutrons • How many e- are in this atom • What is the name of this atom? • Write its formula • What is the mass number and atomic number

  17. ISOTOPES • Atoms of most elements can have different numbers of neutrons. These are called Isotopes. They have the same atomic number but different mass numbers. • Because they have the same number of protons in the nucleus, all isotopes of the same element have the same number of electrons outside the nucleus.

  18. ISOTOPE SYMBOLS • Isotopes are represented by the symbol , where Z is the atomic number, A is the mass number, and E is the elemental symbol. • An example of an isotope symbol is . This symbol represents an isotope of nickel that contains 28 protons and 32 neutrons in the nucleus. • Isotopes are also represented by the notation: Name-A, where Name is the name of the element and A is the mass number of the isotope. • An example of this isotope notation is magnesium-26. This represents an isotope of magnesium that has a mass number of 26. 60 28 Ni

  19. RELATIVE MASSES • The extremely small size of atoms and molecules makes it inconvenient to use their actual masses for measurements or calculations. Relative masses are used instead. • Relative masses are comparisons of actual masses to each other. For example, if an object had twice the mass of another object, their relative masses would be 2 to 1.

  20. ATOMIC MASS UNIT (u) • An atomic mass unit is a unit used to express the relative masses of atoms. One atomic mass unit is equal to 1/12 the mass of a carbon-12 atom. • A carbon-12 atom has a relative mass of 12 u. • An atom with a mass equal to 1/12 the mass of a carbon-12 atom would have a relative mass of 1 u. • An atom with a mass equal to twice the mass of a carbon-12 atom would have a relative mass of 24 u.

  21. ATOMIC WEIGHT • The atomic weight of an element is the relative mass of an average atom of the element expressed in atomic mass units. • Atomic weights are the numbers given at the bottom of the box containing the symbol of each element in the periodic table. • According to the periodic table, the atomic weight of nitrogen atoms (N) is 14.0 u, and that of silicon atoms (Si) is 28.1 u. This means that silicon atoms are very close to twice as massive as nitrogen atoms. Put another way, it means that two nitrogen atoms have a total mass very close to the mass of a single silicon atom.

  22. MOLECULAR WEIGHT • The relative mass of a molecule in atomic mass units is called the molecular weight of the molecule. • Because molecules are made up of atoms, the molecular weight of a molecule is obtained by adding together the atomic weights of all the atoms in the molecule. • The formula for a molecule of water is H2O. This means one molecule of water contains two atoms of hydrogen, H, and one atom of oxygen, O. The molecular weight of water is then the sum of two atomic weights of H and one atomic weight of O: • MW = 2(at. wt. H) + 1(at. wt. O) • MW = 2(1.01 u) + 1(16.00 u) = 18.02 u

  23. Determine the molecular weight • CO2 • Al2Cl3

  24. ISOTOPES & ATOMIC WEIGHTS • Many elements occur naturally as a mixture of several isotopes. • The atomic weight of elements that occur as mixtures of isotopes is the average mass of the atoms in the isotope mixture. • The average mass of a group of atoms is obtained by dividing the total mass of the group by the number of atoms in the group. • The text book shows how to calculate the average relative mass of any element but we’ll settle for just the general concept. Consider chlorine and carbon.

  25. THE MOLE CONCEPT Using quantities based on AMU is not practical but amounts at the gram level are practical. The mole concept was developed in order to use the same relative weights of the elements used in terms of AMU but instead based on grams. • One mole of any element is a sample of the element with a mass in grams that is numerically equal to the atomic weight (in AMUs) of the element.

  26. THE MOLE CONCEPT • THE MOLE CONCEPT APPLIED TO ELEMENTS • The number of atoms in one mole of any element is called Avogadro's number and is equal to 6.022x1023 . • A one-mole sample of any element will contain the same number of atoms as a one-mole sample of any other element. • EXAMPLES OF THE MOLE CONCEPT • 1 mole Na = 22.99 g Na = 6.022x1023 Na atoms • 1 mole Ca = 40.08 g Ca = 6.022x1023 Ca atoms • 1 mole S = 32.07 g S = 6.022x1023 S atoms • 1 mole of O2 = 16.00 g O = 1.2044x1024 O atoms

  27. THE MOLE CONCEPT (continued) • THE MOLE CONCEPT APPLIED TO COMPOUNDS • The number of molecules in one mole of any compound is called Avogadro's number and is numerically equal to 6.022x1023. • A one-mole sample of any compound will contain the same number of molecules as a one-mole sample of any other compound. • One mole of any compound is a sample of the compound with a mass in grams equal to the molecular weight of the compound. • EXAMPLES OF THE MOLE CONCEPT • 1 mole H2O = 18.02 g H2O = 6.022x1023 H2O molecules • 1 mole CO2 = 44.01 g CO2 = 6.022x1023 CO2 molecules • 1 mole NH3 = 17.03 g NH3 = 6.022x1023 NH3 molecules

  28. THE MOLE CONCEPT (continued) • THE MOLE AND CHEMICAL CALCULATIONS • The moleconcept can be used to obtain factors that are useful in chemical calculations involving both elements and compounds. One mole quantities of six metals; top row (left to right): Cu beads (63.5 g), Al foil (27.0 g), and Pb shot (207.2 g); bottom row (left to right): S powder (32.1 g), Cr chunks (52.0 g), and Mg shavings (24.4 g). One mole quantities of four compounds: H2O (18.0 g); small beaker NaCl (58.4 g); large beaker aspirin, C9H8O4, (180.2 g); green (NiCl2· 6H2O) (237.7 g).

  29. MOLE CALCULATIONS • The mole-based relationships given earlier as examples for elements provide factors for solving problems. • The relationships given earlier for calcium are: 1 mole Ca= 40.08 g Ca = 6.022x1023 Ca atoms • Any two of these quantities can be used to provide factors for use in solving numerical problems. • Examples of two of the six possible factors are: and

  30. MOLE CALCULATION EXAMPLE • Calculate the number of moles of Ca contained in a 15.84 g sample of Ca. • The solution to the problem is: • We see in the solution that the g Ca units in the denominator of the factor cancel the g Ca units in the given quantity, leaving the correct units of mole Ca for the answer.

  31. MOLE CALCULATIONS (continued) • The mole concept applied earlier to molecules can be applied to the individual atoms that are contained in the molecules. • An example of this for the compound CO2 is: 1 mole CO2 molecules = 1 mole C atoms + 2 moles O atoms 44.01 g CO2 = 12.01 g C + 32.00 g O 6.022x1023 CO2 molecules = 6.022x1023 C atoms + (2) 6.022x1023 O atoms • Any two of these nine quantities can be used to provide factors for use in solving numerical problems.

  32. MOLE CALCULATION EXAMPLES • Example 1: How many moles of O atoms are contained in 11.57 g of CO2? • Note that the factor used was obtained from two of the nine quantities given on the previous slide.

  33. MORE MOLE CALCULATION EXAMPLES • Example 2: How many CO2 molecules are needed to contain 50.00 g of C? • Note that the factor used was obtained from two of the nine quantities given on a previous slide.

  34. MORE MOLE CALCULATION EXAMPLES • Example 3: What is the mass percentage of C in CO2? • The mass percentage is calculated using the following equation: • If a sample consisting of 1 mole of CO2 is used, the mole-based relationships given earlier show that: 1 mole CO2 = 44.01 g CO2 = 12.01 g C + 32.00 g O

  35. MORE MOLE CALCULATION EXAMPLES (continued) • Thus, the mass of C in a specific mass of CO2 is known, and the problem is solved as follows:

  36. MORE MOLE CALCULATION EXAMPLES • Example 4: What is the mass percentage of oxygen in CO2? • The mass percentage is calculated using the following equation: • Once again, a sample consisting of 1 mole of CO2 is used to take advantage of the mole-based relationships given earlier where: 1 mole CO2 = 44.01g CO2 = 12.01 g C + 32.00g O

  37. MORE MOLE CALCULATION EXAMPLES (continued) • Thus, the mass of O in a specific mass of CO2 is known, and the problem is solved as follows: • We see that the % C + % O = 100% , which should be the case because C and O are the only elements present in CO2.

  38. MOLE CALCULATIONS HELP

  39. Number of Units of Parts Number of Units of A Number of Units of Parts Av. # Av. # Av. # Formula ratio Formula ratio Moles A Moles Parts of A Moles Parts of A Molar Mass Parts Molar Mass Parts Molar Mass A Grams Parts Grams A Grams Parts

  40. Mole calculations • How many moles of O atoms are contained in 11.57 g of CO2? [same question as earlier] • How many moles of Al are in 2.5 mol of Al2Cl3? • In the compound Al2Cl3, how many moles of Al would be combined with 0.06 mol of Cl? • How many mol of Al2Cl3 would be found in 33.5 g of Al2Cl3? • How many g of Al would be found in 33.5 g of Al2Cl3?

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