CHEM 121 Introduction to Fundamental Chemistry. Summer Quarter 2008 SCCC. Lecture 2

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1 CHEM 121 Introduction to Fundamental Chemistry Summer Quarter 2008 SCCC Lecture 2 Could Stephanie, Liqingqing, Huong, Sophia and Yumiko see me after class for a few minutes. Thanks.

2 Matter, Measurements and Calculations Managing numbers Calculations Atoms and Molecules Symbols and Formulas Inside the Atom Isotopes Relative masses of atoms and molecules Isotopes and atomic weights Avogadro s number: the mole The mole and chemical formula s

3 People often break the rules for significant figures when writing whole numbers that end in trailing zeroes. e.g. The volume of a liquid may be written as 1000 ml. Using our rules this would mean that the volume was known to within 10mL

4 How can we write a measurement of 1000 ml taken with either of these two measuring cylinders that indicates which number is estimated? ANSWER: We can t (with what we have learnt so far).

5 To overcome this difficulty we use scientific notation. We write a number using scientific notation as follows: b a 10 The number a is called the pre-exponential term and is a number between 1 and 10 written so that it has 1 digit before the decimal point. The number b is called the exponent and may be any whole number.

6 How can we write a measurement of 1000 ml taken with either of these two measuring cylinders that indicates which number is estimated using scientific notation? ANSWER: 1.00 x 10 3 ml ANSWER: 1.0 x 10 3 ml

7 There is only a limited number of possible prefix modifiers and as you have probably already discovered they are not easy to remember. Scientific notation is a good method for writing very large and very small numbers. e.g x or 1.6 x 10-19

8 We often will combine two or more measurements using an equation to calculate a new quantity. e.g. Newton famously determined that the force acting on an object is equal to the product of the mass and acceleration. F = ma The bigger the apple the harder it hits!

9 In everday life we are used to taking ratios of measurements to give new quantities. e.g miles per hour, miles per gallon, dollars per gallon. A ratio commonly used in chemistry is the density. Density is the ratio of mass and volume. Density = mass volume or D d = m V What are the units for density?

10 As we have discussed previously mass is determined with a balance. How do we determine the volume of a substance? As we have seen previously for liquids we can use a measuring cylinder or similar device. How do we determine the volume of a solid?

11 If it is a regular shape we could use an algebraic relationship. e.g. the volume of a rectangular solid = length x width x height H L W What if it is an irregularly shaped solid?

12 We can determine the volume of irregularly shaped objects by displacement. How can we determine the volume of a gas? Gases fill whatever container they are placed in. So it s the volume of the container!

13 I measure the volume of a liquid to be 3.1 ml and its mass to be 15.5 g how should I express its density to the correct number of s.f s? There are two rules we need to remember: 1. When dividing or multiplying the answer has as many s.f s as the number with the fewest s.f s. e.g g 3.1 ml = 5.0 gml When adding or subtracting the answer has as many decimal places as the number with the fewest digits after the decimal place.

14 Sometimes when we perform calculations we will end up with more digits in our answer than the number of s.f s In this case we will need to round our answer. 1. If the first of the digits to be dropped is greater than 5 we increase the last s.f by 1. e.g. an answer of would be given to 2 s.f s as If the first of the digits to be dropped is less than or equal to 5 we leave the last s.f unchanged.

15 A ratio commonly used in chemistry is the density. Density is the ratio of mass and volume. Density = mass volume or D d = m V How do we express a calculated quantity such as density in a way that takes into account the uncertainty in the original measurements?

16 There are two rules we need to remember: 1. When dividing or multiplying the answer has as many s.f s as the number with the fewest s.f s. e.g g 3.1 ml = 5.0 gml When adding or subtracting the answer has as many decimal places as the number with the fewest digits after the decimal place. e.g m m = 27.3 m

17 Sometimes when we perform calculations we will end up with more digits in our answer than the number of s.f s In this case we will need to round our answer. 1. If the first of the digits to be dropped is greater than 5 we increase the last s.f by 1. e.g. an answer of would be given to 2 s.f s as If the first of the digits to be dropped is less than or equal to 5 we leave the last s.f unchanged.

18 Elements are pure substances made up of identical atoms. There are 115 known kinds of atoms. Each has its own symbol. The first character in an atoms symbol is an uppercase letter. Most atoms have a second character in their symbol which is always a lowercase letter. The symbol and names for the known atoms are given in Table 2.1.

19 The symbol and names for the known atoms are given in Table 2.1. DO NOT TRY AND MEMORIZE THIS!!

20 You can also find these symbols and names in the periodic table inside the front cover of your text.

21 What are the particles atoms made of? Atoms are composed of subatomic particles most of which exist outside the stable structure of atoms for very short periods of time. Three subatomic particles are of interest to chemists: 1. The proton (p) 2. The neutron (n) 3. The electron (e - )

22 1 amu = 1.67 x g Atoms have a dense heavy, positively charged nucleus containing the neutrons and protons. Outside the nucleus are the small, negatively charged electrons.

23 The electrons move rapidly around the nucleus. Most of an atom is empty space. If the nucleus was the size of a pea the closest electrons would be about 240 feet away. For a neutral atom the number of protons and electrons is equal.

24 What happens when we combine atoms? When two or more different types of atoms combine compounds are formed. In a compound formula we write the symbol for each atom type present in the compound. If more than one atom of a given type a subscript is used to indicate the number. e.g. NH 3 H 2 O SO 2 CH 4

25 Atoms have no overall charge so contain the same number of protons (+) as electrons (-). The number of protons in the nucleus of an atom is called the atomic number. It is given the symbol Z. All atoms of a given type or element have the same atomic number (Z). The periodic table arranges atoms in order of atomic number.

26 Neutrons have no electrical charge. For a given atom type the number of neutrons may vary. The sum of the number of neutrons and protons in an atom is called the mass number. It is given the symbol A. Atoms with the same number of protons but different numbers of neutrons in their nucleus are called isotopes. Does anyone know of any medical application of isotopes?

27 We can distinguish isotopes with the following notation: Where: A ZE E is the atom s symbol A is the mass number Z is the atomic number e.g C, 6 13 C, 6 14 C, 11 H, 12 H and 13 H As all atoms with a given symbol have the same atomic number sometimes we skip writing Z. e.g C (carbon-12), 6 13 C (carbon-13) and 6 14 C (carbon-14)

28 If we assign a mass of 12 atomic mass units (u or sometimes amu) to a carbon-12 atom then we can compute the relative atomic weight for any other atom. Atomic weights or atomic masses are given in your periodic table.

29 Some atoms naturally occur as a mixture of isotopes. e.g. 11 H, 12 H and 13 H The atomic weights given in the periodic table take this into account. They are the average atomic weight taking into account the amounts of each isotope present.

30 To calculate the relative mass of a molecule we sum together the atomic weights of the atoms that make up the molecule. This is called the molecular weight (MW, M). e.g. determine the molecular weight of H 2 O MW H 2 O = 2x(1.008) + 1x(15.999) = u

31 we defined the relative atomic weight of each kind of atom as follows: assign a mass of 12 atomic mass units to a carbon-12 atom then compute the relative atomic weight for all other atom types. Atomic mass units are very small. Not very practical!!! 1 amu = 1.67 x g

32 Consider the following ozone has the formula O 3. Two molecules of ozone contains 6 oxygen atoms. Oxygen has the formula O 2. What is the mass of three molecules of O 2? Three molecules of O 2 has the same mass as two molecules of O 3 as they each contain a total of six oxygen atoms. In chemistry we are often more interested in the number of particles but we can only measure the mass easily.

33 So we have two problems: 1. It is difficult to count particles as small as molecules and atoms. 2. Atomic mass units are not very practical. We need a system that allows us to switch between mass (in units which we can measure) and number of atoms or molecules.

34 We can define a constant called Avogadro s constant (N A ): N A is the number of atoms needed to give a mass in grams of an atom equal to its atomic weight in atomic mass units. e.g. An Avogadro s number of carbon-12 atoms would weigh 12 g. An Avogadro s number of hydrogen atoms would weigh 1.008g.

35 One Avogadro s number of something is called a mole (mol). N A = x A mole is often compared to a dozen. Just as a dozen of something means there are 12 of them, a mole of something means there are x of them.

36 We can define the molar mass or molecular weight of a substance can be defined as: the mass of one mole of that substance. We give molar mass the symbol M and it has units gmol -1. For an atom the molar mass is equal to the atomic weight we find on the period table.

37 For a molecule the molecular weight is equal to the sum of the atomic weights of the atoms that make it up. The mathematical relationship between mass and moles is: m = nm This can be summarized by the following diagram: moles

38 The key equation to remember for mole calculations is: m = nm Where: M = molecular weight (gmol -1 ) n = number of moles (mol) m = mass of sample (g)

39 For a mole of a molecule the number of moles of each atom is determined by how many of that atom are in each molecule. e.g. One mole of H 2 O contains: One mole of oxygen atoms Two moles of hydrogen atoms In 5 moles of H 2 SO 4 how many moles of oxygen atoms is there? 20 moles of O atoms.

40 In 5 moles of H 2 SO 4 how many grams of oxygen atoms is there? Number of moles of O atoms = 5 x 4 = 20 moles m = nm m = 20 mol x gmol -1 = g

41 In 50g of H 2 SO 4 how many moles of oxygen atoms is there? How many moles of H 2 SO 4 do we have in 50g m = nm n = m/m (we need to find M for H 2 SO 4 ) M = 2 x x M = gmol -1 of H 2 SO 4 (now we can find n for H 2 SO 4 ) n = 50 g gmol -1 n = 0.51 mol of H 2 SO 4 (now we can find n for O) number of moles of O = 4 x number of moles of H 2 SO 4 number of moles of O = 4 x 0.51 mol = 2.0 mol of O atoms

42 In 50g of H 2 SO 4 how many grams of oxygen atoms is there? There are 2.0 mol of O atoms in 50g of H 2 SO 4 m = nm m = 2.0 mol x gmol -1 m = 32 g of oxygen atoms

43 Calculate the number of moles of oxygen atoms in 5 grams of H 2 SO 4. n = m/m M = 1.008x x = gmol -1 n = m/m = 5/ = 0.05 moles of H 2 SO moles of oxygen

44 In 5 grams of H 2 SO 4 what percentage of the total mass will be oxygen atoms? 0.2 moles of oxygen (from previous question ) m = nm m = 0.2 x = g of O % O = (3.1998/5)x100% = 64% Oxygen atoms by mass

45 HOME WORK 1.) REVISE CHAPTER 2. 2.) STUDY EXAMPLES IN CHAPTER 2 AND MAKE SURE YOU UNDERSTAND THEM. 3.) PRACTICE 2-3 PROBLEMS FOR EACH SECTION. AND CHECK YOUR ANSWERS. 4.) PREPARE FOR LAB ON TUESDAY. 5.) READ CHAPTER 3.

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