Chem Chapter 2.notebook October 17, 2012

Transcription

1 Unit 1: Stoichiometry 1

2 Chapter 2: The Mole Atomic number the number of protons in an atom or ion Mass number the sum of the protons and neutrons in an atom 2

3 Isotope atoms which have the same number of protons and electrons but different numbers of neutrons Ex: isotope notation isotope name Carbon 12 (6p+6n) Carbon 14 (6p+8n) 3

4 You can convert an isotope name to isotope notation using the mass number and the atomic number for the element as shown in Example 1. Example 1: Write the isotope notation for oxygen 18 and list the number of protons, electrons, and neutrons in an atom of this isotope. Answer: Oxygen atomic number 8 mass number is 18 (from isotope name) Isotope notation: 18 8O Since the atomic number is 8, there are 8p+ and 8e. Subtracting the atomic number from the mass number gives the number of neutrons: 18 8 = 10. Thus, oxygen 18 has 10n. 4

5 You can convert isotope notation to an isotope name using the mass number and the symbol for the element as shown in Example 2. Example 2: Write the isotope name and list the number of protons, 238 electrons, and neutrons in an atom of this isotope: 92 U. Answer: The top number is the mass number. Therefore, the isotope name is uranium 238. Since the atomic number is 92, there are 92p+ and 92e Subtracting the atomic number from the mass number gives the number of neutrons = 146n. 5

6 Atomic Mass An atomic mass unit (amu) is defined as exactly 1/12 of the mass of a carbon 12 atom. Carbon 12 has 6p+ and 6n so each proton or neutron has a mass of 1 amu. The mass of an electron is insignificant in comparison about amu. 6

7 The atomic mass of chlorine 35 is 35 amu because its nucleus contains 17 p+ and 18 n. In other words, the atomic mass of chlorine 35 is based upon the mass unit defined using carbon 12. Atomic masses are relative units as opposed to measurable units. 7

8 Average Atomic Mass There are two naturally occurring isotopes of chlorine: chlorine 35 and chlorine 37. The atomic mass of this element is a combination of the two isotopes. The relative abundance of chlorine atoms in nature is 75% chlorine 35 and 25% chlorine 37. Average atomic mass is the weighted average of the atomic masses of the isotopes of an element. The average atomic mass is calculated by multiplying the percent abundance by the mass number for each isotope and then adding the products. 8

9 INCLUDEPICTURE \\d " Chem Chapter 2.notebook Exercise 1: Copy and complete the table Isotope Name sulfur 33 Isotope Notation Atomic #Mass # # protons # neutrons # electrons

10 Exercise 2: Write the isotope notation for each isotope of hydrogen. hydrogen 1 hydrogen 2 hydrogen 3 Why are the mass numbers for the three isotopes of hydrogen different? The average atomic mass of hydrogen is Which isotope of hydrogen do you think is most abundant in nature? Why? 10

11 Read: page Isotopes and atomic mass Questions: page 37: # 5 7 page 45: # 1 4 (#4 hint: Let one isotope be x, the other will be 1 x) page 46: #

12 The Mole (Avogadro's Number) Don't write this down Atoms are so small that it takes an incredibly huge number of them to make up a visible amount. An iron spike for example might contain more than 6.02 x atoms of iron. Dealing with such large numbers of atoms on a consistent basis can get to be tiresome after a while. That's where the mole comes in. 12

13 A mole (mol) is the number of carbon 12 atoms in exactly 12 grams of carbon mol of a substance contains 6.02 x particles of the substance. The constant 6.02 x mol 1 is called the Avogardo's constant. So instead of dealing with individual atoms and molecules, chemists deal with mole amounts. Unfortunately, there are no measuring devices that give readings in moles. Chemists measure things by mass or volume. 13

14 How big a number is it??? Don't write this down one mole of peas is enough to cover Earth and 250 more planets the same size as Earth one metre deep in green peas a stack of one mole of pennies is tall enough to reach Proxima Centauri (the second closest star to Earth) and back again 7448 times a mole of marbles spread over the Earth would cover it to a depth of 80 km if you own a mole of dollars and you spend a billion dollars a day, then you could spend that amount per day for over a trillion years before you run out of money. (Earth has only been around for 4.5 billion years i.e. that's 0.45% of a trillion years!) 14

15 Molar Mass In chemistry, the conversion factor is molar mass (the mass of one mole). Molar mass is on the periodic table. 1 mole of iron atoms (6.02 x atoms) has a mass of g Molar mass values can be used to convert mole amounts to mass and vice versa. 15

16 16

17 A note before we do calculations... Sig Figs! Zeros: All zeros to the left of the first non zero digit are not significant Ex: has 3 All zeros in between on zero digits are significant Ex: 2002 has 4 Zeros after a non zero digit are not significant unless there is a decimal point Ex: 800 has has 4 Add/Subtract: Use the least # of decimal places Ex: = = 3 Mulitply/Divide: Use the least number of significant figures Ex: 300 x 2.50 = 750 =

18 Calculating Molar Mass To find molar mass: 1. Write the chemical formula 2. List the number of atoms of each element 3. Multiply this number by the molar mass of the element. Example: Calculate the molar mass of water. Answer: Chemical formula: H 2 O List the number of atoms of each element: 2 H,1 O Multiply each number by the molar mass of the element: Add the products together: 18

19 Example: Calculate the molar mass of barium hydroxide octahydrate. Answer: hydroxide octahydrate is Ba(OH) 2 8H 2 O 1 Ba 2 O 2 H 8 H 2 O The molar mass of barium hydroxide octahydrate is g/mol. 19

20 Calculate the molar mass of each compound. NaHCO 3 Sr(NO 3 ) 2 Al 2 (SO 4 ) 3 CuSO 4 FeSO 4 3H 2 O lead(ii) acetate magnesium chloride Answers: g/mol g/mol g/mol g/mol g/mol g/mol g/mol 20

21 Molar Volume The SI unit of volume is the litre. Amounts of gases can be compared in litres if the samples are at the same temperature and pressure because changes in pressure and/or temperature can effect the volume of a gas. Chemists compare gas volumes at 0 C and kpa. These conditions are known as Standard Temperature and Pressure or STP. Avogadro concluded: equal volumes of different gases at the same temperature and pressure, have the same number of particles (i.e. same number of moles) This very significant discovery led to the concept of molar volume: 1 mol of any gas at STP has a volume of 22.4 L. 21

22 Questions: What does STP mean? What is molar volume at STP? What is the volume of one mole of fluorine gas at STP? What is the volume of one mole of chlorine gas at STP? 22

23 Mass to Mole and Mole to Mass Conversions Formula: n is the number of moles m is the mass M is the molar mass Example: Determine the number of moles in a g piece of aluminum. 23

24 Example: Diamonds are pure carbon. Calculate the number of moles of carbon atoms in a 1.52 g diamond. 24

25 Example: Calculate the number of moles in a g sample of cerium nitrate, Ce(NO 3 ) 3? 25

26 Example: A sample of neon gas contains 2.00 mol of atoms. Calculate the mass of the sample. 26

27 Example: Calculate the mass of 2.40 mol of aluminum nitrate Al(NO 3 ) 3. 27

28 Mole to Volume Conversions Converting a gas volume to a number of moles requires use of this formula: n is moles Little v is volume Big V is molar volume (22.4 L) 28

29 Example: Calculate the number of moles of neon in a 6.81 L sample at STP conditions 29

30 Example: Calculate the volume of mol of sulfur trioxide gas at STP conditions. 30

31 Homework 1. Calculate the mass of 88.6 mol copper 2. Calculate the number of moles in 369 L of oxygen gas at STP conditions. 3. Calculate the volume of mol of carbon monoxide at STP conditions. #1 Answer: 5630 g Cu #2 Answer: 16.5 mol Answers #3 Answer: 6.61 L 31

32 Mole to Number of Particles Conversions Formula: # Particles = na n is the number moles A is Avogadro's number The particles could be atoms (for an element), molecules (for a molecular compound), or formula units (for ionic compounds) 32

33 Example: How many particles are there in each sample? a) 0.25 mol of helium (atoms) b) mol of carbon dioxide (molecules) c) 1.20 mol of sodium sulfide (formula units) 33

34 Example: Convert these numbers of particles to mole amounts. a) x argon atoms b) 1.35 x sodium chloride formula units 34

35 Example: a) What is the volume of 2.7 g of hydrogen gas (assume H 2 or diatomic hydrogen) at STP conditions? b) How many molecules are there in this sample? 35

36 1. Calculate the mass of each sample. a) mol silver b) mol tin(iv) silicate c) 4.23 mol Ba(BrO 3 ) 2 d) mol diphosphorus pentaoxide 2. Calculate the number of moles in each sample of gas at STP conditions. a) 65 L of helium b) 67 L of carbon dioxide 3. Calculate the volume of each amount of gas at STP conditions. a) 2.7 mol of krypton b) mol of dinitrogen monoxide #1 Answers: #2Answers: 3.73 g Ag 2.9 mol 1.0 x 10 1 g ( molar Answers mass was g/mol) 3.0 mol 1660 g (molar mass was g/mol) 9328 g (molar mass was g/mol) #3 Answers: 60. L L 36

37 4. Which sample in each pair contains the greater number of particles (i.e. greater number of moles)? a) 25.0 g of aluminum or 25.0 g of copper b) 10.0 L of oxygen gas or 10.0 L of nitrogen gas (both at STP). c) 15.0 L of hydrogen gas (at STP) or g of calcium chloride #4 Answer: Answers a) mol Al > mol Cu b) mol O 2 = mol N 2 c) mol H 2 < mol CaCl 2 37

38 Qualitative Analysis versus Quantitative Analysis Qualitative analysis is carried out to identify the composition of a substance or a mixture. The focus is on determining which chemical species are present in a sample. Quantitative analysis involves identifying the quantity of each species is present in a sample. 38

39 Percentage Composition This is a type of quantitative analysis used to find the percent by mass of each element in a substance and is a key step in the determination of the compound's chemical formula. By the way, percent composition analysis is not restricted to chemistry. Fish buyers determine the percentage composition of a catch of capelin before deciding how much to pay for a catch. They offer higher prices for catches that have higher percentages of female capelin. 39

40 Calculating Percentage Composition Step 1: Interpret the subscripts in terms of moles Step 2: Convert moles to mass Step 3: Calculate the molar mass of the compound Step 4: Calculate the percentage composition of each element Example: AlBr 3 One mole of AlBr 3 contains 1 mol of Al ions and 3 mol of Br ions. Converting moles to mass, one mole of AlBr 3 contains g of Al ions and g of Br ions g g = g. 40

41 Example: Calculate the percentage composition of ethanol, C 2 H 5 OH. 41

42 Once percentage composition is known, then a chemical formula for the compound can be determined. This is a really important point. Whether the substance is one that is known to exist or one that as never been discovered, a key step in identifying it (i.e. coming up with a chemical formula) is to determine its percent composition. Two kinds of chemical formulas: 1. empirical formula is a lowest whole number ratio of the elements in a compound. 2. molecular formula shows the number of atoms of element in the compound; it is not necessarily a lowest whole number ratio. 42

43 Using percent composition to determine an empirical formula: Step 1: Assume the mass of the sample is g and multiply it by the percentage composition for each element Step 2: Use the mole formula to convert each mass to moles Step 3: Write a ratio of the mole values, and divide the larger mole value by the smaller mole value Step 4: From the mole value, determine the empirical formula. 43

44 Example: Answer: A prospector finds a rock that may contain rutile an ore of titanium. A crystal from the sample is analyzed and found to be 59.94% titanium and 40.06% oxygen by mass. Determine the empirical formula of the crystalline substance. Step 1: Step 2: Step 3: Step 4: The ratio of titanium to oxygen (1 Ti : 2 O) is a whole number ratio, so the empirical formula for the compound is TiO 2. 44

45 Remember... Percent to mass Mass to mole Divide by small Multiply 'til whole 45

46 Using percent composition to determine a molecular formula: Step 1: Step 2: Step 3: *Note: Assume a sample mass equal to one mole of the compound and convert the percentage composition data to mass amounts. Convert the mass amount for each element to moles. Round off the mole amounts and use them as numerical subscripts in the chemical formula. In order to find the molecular formula of an unknown, the molar mass of the compound must be known. Nowadays, molar mass is usually determined by mass spectroscopy. 46

47 Example: An unknown was subjected to a combustion analysis and found to be 85.60% carbon and 14.40% hydrogen. Mass spectrometry indicated that its molar mass was g/mol. Determine the molecular formula of the compound. Answer: Step 1 Step 2 Step 3 Mole ratio is C : 4.00 H The molecular formula is C 2 H 4. 47

48 48