CHEMISTRY - BROWN 13E CH.2 - ATOMS, MOLECULES & IONS.

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2 CONCEPT: GROUP NAMES AND CLASSIFICATIONS Ever wonder where did this periodic table ever come from? At the end of the 18 th century, Lavoisier compiled a list of the 23 elements known at the time. In 1869, Dmitri Mendeleev coined the term Periodic Table. Today the total is 114 and still counting! Now, to understand chemistry fully it will be imperative that you memorize and learn the different portions of the Periodic Table. Phase Differences At room temperature (between 20 o C to 25 o C), all elements are except: Mercury and bromine are. Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine and the Noble Gases are. Page 2

3 CONCEPT: CHARGE DISTRIBUTIONS OF THE PERIODIC TABLE A majority of the elements on the periodic table are reactive because they all want to be like the. They have the perfect number of electrons in their outer atomic shells. 1. Metals tend to electrons to become positively charged ions called. Metals that have ONLY one charge are referred to as metals. Metals that have MORE THAN one charge are referred to as metals. 2. Nonmetals tend to electrons to become negatively charged ions called. Page 3

4 CONCEPT: ELEMENT SYMBOLS Some of the names and symbols for the elements are easy to recognize like Aluminum is Al, but some others aren t. EXAMPLE 1: Identify the elements by their given symbols. a. Au b. Hg c. Pb d. Fe e. Ag Some elements exist in nature connected to their exact double. We call these chemical Siamese twins. To recall them just remember this funny phrase: Have No Fear Of Ice Cold Beer Some elements exist in nature as monoatomic elements such as &. Some elements exist in nature as polyatomic molecules such as &. Page 4

5 CONCEPT: MASS CONVERSIONS The is the chemical unit for the amount of a substance. One mole (1 mol) contains x entities, which is known as. Entities means, or. We use when dealing with a single, individual element. We use or when dealing with more than one element or a compound x atoms of Fe is equal to 1 mole of Fe and has a mass of amu Atoms Moles Grams EXAMPLE: Determine the mass (in grams) found in 7.28 x nitrogen atoms x molecules of H2O is equal to 1 mole of H2O and has a mass of amu Molecules Moles Grams EXAMPLE: Determine how many molecules of carbon dioxide, CO2, are found in 75.0 g CO2. Page 5

6 CONCEPT: MASS CONVERSIONS (PRACTICE) PRACTICE 1: If the density of water is 1.00 g/ml at 25 o C calculate the number of water molecules found in 1.50 x 10 3 µl of water. PRACTICE 2: Calculate the number of oxygen atoms found in g CuSO4 5 H2O. PRACTICE 3: The density of the sun is 1.41 g/cm 3 and its volume is 1.41 x m 3. How many hydrogen molecules are in the sun if we assume all the mass is hydrogen gas? PRACTICE 4 (CHALLENGE): A cylindrical copper wire is used for the fences of a house. The copper wire has a diameter of in. How many copper atoms are found in cm piece? The density of copper is 8.96 g/cm 3. ( V = π r 2 h ). Page 6

7 CONCEPT: ATOMIC MASS Whether you call it atomic mass or weight both terms tell us the combined mass of the protons and neutrons in an element. The atomic masses listed for the elements on the periodic table are the of their isotopes. Isotopes are elements with the number of protons, but number of neutrons. Atomic Mass = [(Mass of Isotope 1) x (Fractional Abundance 1)] + [(Mass of Isotope 2) x (Fractional Abundance 2)] EXAMPLE 1: Antimony has two common isotopes. If one of the isotopes 121 Sb has an isotopic mass of amu and a natural abundance of 57.25%, what is the isotopic mass (to 4 significant figures) of the other isotope? The atomic mass of antimony is g/mol. EXAMPLE 2: The atomic mass of an imaginary element A is amu. If element A consists of two isotopes that have atomic masses of 250 and 253 respectively, what is the natural abundance of each isotope? Page 7

8 CONCEPT: MASS SPECTROMETRY Mass spectrometry involves the,, and of gaseous ions according to their mass to charge ratios. Page 8

9 CONCEPT: STRUCTURE OF THE ATOM We learned that the basic functional unit in chemistry is the. Now it s time to go into an atom to figure out its components: subatomic particles. In the center of an atom there is the, It contains the subatomic particles: and. Spinning around it we find the third subatomic particle: the. PROTONS are charged subatomic particles. ELECTRONS are charged subatomic particles.! NEUTRONS are charged subatomic particles. ATOMIC NUMBER equals the number of and determines of an element. ATOMIC MASS equals the number of in an element. EXAMPLE: Identify the unknown element. a. Element X (8 protons, 8 electrons, 8 neutrons) b. Element Y (35 protons, 36 electrons, 46 neutrons) c. Element Z (12 protons, 10 electrons, 13 neutrons) Page 9

10 CONCEPT: MODERN ATOMIC THEORY According to the Law of in a reaction matter is neither created nor destroyed. Originated in 1789 by Antoine Lavoisier. CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g) According to the Law of all samples of a compound, no matter on their origin or preparation has the same ratio in terms of their elements. Originated in 1797 by Joseph Proust. CO 2 Mass Ratio = (12.0gC) (32.0gO) = According to the Law of when two elements (A & B) form different compounds, the masses of element B that combine with 1 g of A are a ratio of whole numbers. Originated in 1804 by John Dalton. NO Mass Ratio = (16.0gO) (14.0g N) =1.143 NO 2 Mass Ratio = (32.0gO) (14.0g N) = The ratio of the two mass ratios obtained then gives us a whole number: = 2.0 Page 10

11 CONCEPT: MODERN ATOMIC THEORY (PRACTICE) EXAMPLE 1: A g sample of iodine reacts with g of chlorine to form iodine pentachloride, ICl5. If iodine pentachloride is the only product formed calculate its mass. EXAMPLE 2: Two samples sodium fluoride decompose into their constituent elements. The first sample produces 15.8 kg of sodium and 20.1 kg of fluorine. If the second sample produces g of sodium, how many grams of fluorine were also produced? PRACTICE: Which of the following is an example of the law of multiple proportions? a. A sample of bromine (Br) contains equal amounts of its two isotopes. b. Two different samples of H2O have the same mass ratio. c. The atomic mass of sodium (Na) is amu. d. Two different compounds composed of sulfur (S) and oxygen (O) have different mass ratios: 2.48 g O: 1 g S and 1.24 g O: to 1 g S. Page 11

12 CONCEPT: THOMSON CATHODE RAY TUBE EXPERIMENT J.J. Thomson s cathode ray tube experiments led to the discovery of the. Apply an Electric Field When an electric field is applied across the cathode ray tube, the cathode ray is attracted to the plate with a charge. Applying a Magnetic Field A moving charged body behaves like a tiny magnet, and it can interact with an external magnetic field. The electrons are by the magnetic field. Determining the Charge-To-Mass Ratio In 1897, JJ Thomson, an English Physicist, determined the charge-to-mass ratio of an electron by adjusting the electric field so that the deflection (θe) was the same as the deflection (θb), and was able to calculate the charge-to-mass ratio of an electron using the following equation: e / m ratio = Eθ E B 2 l Thomson determined the charge-to-mass ratio of an electron to be x 10 8 coulombs per gram, meaning it was approximately 2000 times lighter than hydrogen, the lightest known atom. e / m ratio = Eθ E B 2 l = coulombs per gram Page 12

13 CONCEPT: RUTHERFORD GOLD FOIL EXPERIMENT The experiment also called the Rutherford Gold Foil experiment helped to discover that any given atom had a positively charged center called the. It is there where most of the atom s mass was concentrated. Subatomic Particle Charge Mass Relative Absolute Relative (in amu) Absolute (in kg) Proton (p + ) x C x Neutron (n o ) x Electron (e ) x C 5.49 x x Page 13

14 CONCEPT: MILLIKAN OIL DROP EXPERIMENT In 1913 Robert Millikan and Harvey Fletcher discovered the charge of an electron as being. The charge of an electron When an oil droplet is suspended, mass x acceleration (m x g) due to gravity is exactly counterbalanced by the electric force applied. The electric force applied equals the applied electric field E times the charge on the drop (q). Making them equal to one another: The mass of an electron By using his discovered charge and then the charge-to-mass ratio determined by Thomson s cathode ray tube experiment we are able to calculate the mass of electron. Page 14

15 CONCEPT: CHADWICK NEUTRON EXPERIMENT In 1920, Ernest Rutherford stated that the nucleus must contain neutral, massive particles. In the early 1930s with experiments designed by Walter Bothe as well as Mr. and Mrs. Joliot it was determined that bombarding with alpha particles would produce high-energy radiation. In 1932, James Chadwick modified the earlier experiments and determined that the unknown neutral particles in the nucleus were the. By examining the motion of these neutral and unknown particles, Chadwick was able to determine the velocity of the protons. Through he determined that the mass of the neutral particles were nearly identical to the mass of a proton. Relative (in amu) Absolute (in kg) Proton (p + ) x Neutron (n o ) x His equation to prove the existence of this neutral particle can be written as: Page 15

16 CONCEPT: ATOMIC PROPERTIES AND CHEMICAL BONDS Before we examine the types of chemical bonding, we should ask why atoms bond at all. Generally, the reason is that ionic bonding the potential energy between positive and negative ions. Generally, the reason covalent bonds form is to follow the rule, in which the element is then surrounded by 8 valence electrons. There are three models of chemical bonding: In bonding, metals connect to non-metals. transfers an electron to the, creating ions with opposite charges that are attracted to each other. Li F Li F Li F In bonding, non-metals connect to non-metals. In it the nonmetals electron pairs between their nuclei. Cl Cl In bonding, metal atoms pool their valence electrons to form an electron sea that holds the metal-ion together Page 16

17 CONCEPT: CHEMICAL BONDS (PRACTICE) EXAMPLE: Describe each of the following as either a(n): atomic element, molecular element, molecular compound or ionic compound. atomic element molecular element molecular compound ionic compound a. Iodine b. NH3 c. Graphite d. Na3P e. Ag2(SO4)2 Page 17

18 CONCEPT: THE IONIC-BONDING MODEL The central idea of ionic bonding is that the metal transfers an electron(s) to a nonmetal. The metal then becomes a(n) (positive ion). and the nonmetal becomes a(n) (negative ion). Their opposite charges cause them to combine into a crystalline solid. PRACTICE: Determine the molecular formula of the compound formed from each of the following ions. a. K+ & P3- b. Sn4+ & O2- c. Al3+ & CO32- Page 18

19 CONCEPT: COMMON POLYATOMIC IONS Polyatomic ions are compounds made up of different elements, usually only, and possess a. Singly Charged Cation (Positive Ion) NH4 + Ammonium Doubly Charged Anions (Negative Ions) CO3 2 Carbonate Doubly Charged Cation (Positive Ion) CrO4 2 Chromate Hg2 2+ Mercury (I) Cr2O7 2 Dichromate Singly Charged Anions (Negative Ions) O2 2 Peroxide CH3CO2 or C2H3O2 Acetate SO4 2 Sulfate CN Cyanide Sulfite OH Hydroxide MnO4 Permanganate Triply Charged Anions (Negative Ions) NO3 Nitrate PO4 3 Phosphate Nitrite Phosphite Doubly & Singly Charged Anions (Negative Ions) HPO4 2 H2PO4 HCO3 HSO4 Hydrogen Phosphate Dihydrogen Phosphate Hydrogen Carbonate or Bicarbonate Hydrogen Sulfate or Bisulfate Page 19

20 CONCEPT: POLYATOMIC IONS w/ HALOGENS Polyatomic ions containing halogens are sometimes referred to as halogens or halogen. These compounds share 4 common characteristics: These compounds use the same system for naming: PRACTICE: Name each of the following compounds. a. BrO4 b. FO2 c. ClO d. IO3 Page 20

21 CONCEPT: NAMING MOLECULAR COMPOUNDS Features: & Because molecular compounds combine in different proportions to form different compounds, we must use numerical prefixes. Rules for Naming: a. The first nonmetal is named normally and uses all numerical prefixes except. b. The second nonmetal keeps its base name but has its ending changed to. EXAMPLE: Write the formula for each of the following compounds. a. Disulfur monobromide b. Iodine Tetrachloride PRACTICE: Give the systematic name for each of the following compounds: a. CO b. N2S4 c. IO5 Page 21

22 CONCEPT: IONIC COMPOUNDS In the early days of chemistry, newly discovered compounds were given fancy names such as morphine, quicklime and muriatic acid. Since then thousands of new compounds have been discovered and named under a system called. Metals tend to electrons to become positively charged ions called. Nonmetals tend to electrons to become negatively charged ions called. Page 22

23 CONCEPT: NAMING BINARY IONIC COMPOUNDS Features: & Rules for Naming: a. The metal is named and written first. If the metal is a transition metal we must use a to describe its positive charge. b. The nonmetal keeps its base name but has its ending changed to. EXAMPLE: Provide the molecular formula or name for each of the following compounds. a. Calcium phosphide b. CoO PRACTICE: Provide the molecular formula or name for each of the following compounds. a. AlBr3 b. Lead (IV) sulfide c. SnO2 Page 23

24 CONCEPT: NAMING IONIC COMPOUNDS w/ POLYATOMICS Features: & Rules for Naming: a) The metal keeps its name and is named and written first. If the metal is a transition metal we must use a to describe its positive charge. b) Name the polyatomic as you would normally. EXAMPLE: Write the formula for each of the following compounds: a. Iron (III) Acetate b. Copper (I) phosphate c. Strontium Carbonate d. Ammonium Nitrite EXAMPLE: Give the systematic name for each of the following compounds: a. Pb(CrO4)2 b. Ga(ClO4)3 c. Mn(HSO4)2 d. Ba(CN)2 Page 24

25 CONCEPT: NAMING IONIC HYDRATES Features: & CuSO 4 5 H 2 O Rules for Naming the Ionic Compound portion: a. The metal is named normally and written first. If the metal is a transition metal we must use a to describe its positive charge. b. The nonmetal keeps the first part of its name but has its ending changed to. c. Name the polyatomic as you would normally. Rules for Naming the H2O portion: a. The H2O portion will be called. b. To describe the number of H2O molecules use these prefixes. EXAMPLE: Write the formula for each of the following compounds. a. Calcium carbonate hexahydrate b. Lead (IV) Sulfate pentahydrate PRACTICE: Give the systematic name for each of the following compounds: a) K2Cr2O7 3 H2O b) Sn(SO3)2 4 H2O Page 25

26 CONCEPT: NAMING ACIDS 1. BINARY ACIDS Features: + Rules for Naming: a. The prefix will be. b. Use the base name of the nonmetal. c. The suffix will be. EXAMPLE: Write the formula for each of the following compounds: a. Hydroiodic acid b. Hydroselenic acid c. Hydrofluoric acid PRACTICE: Give the systematic name for each of the following compounds: a. HBr b. H2S c. HCN 2. OXOACIDS or OXYACIDS Features: + Rules for Naming: a. If the polyatomic ion ends with ate then change the ending to. b. If the polyatomic ion ends with ite then change the ending to. EXAMPLE: Give the systematic name or formula for each of the following compounds: a. H2CO3 b. Nitric acid c. H2SO4 PRACTICE: Give the systematic name or formula for each of the following compounds: a. Hypobromous acid b. HClO3 c. Acetic acid Page 26

27 3. How many molecules of hexane are contained in 55.0 ml of hexane? The density of hexane is g/ml and the molar mass is g/mol. Page 27

28 4. How many SO3 ions are contained in mg of Na2SO3? The molar mass of Na2SO3 is g/mol. Page 28

29 5. What mass of phosphorus pentafluoride, PF5, has the same number of fluorine atoms as 50.0 g of oxygen difluoride, OF2? Page 29

30 6. How many bromide ions are there in 4.50 moles of gallium bromide? Page 30

31 7. How many moles of oxygen atoms are required to combine with 3.05 moles of Pb to create lead (IV) phosphate? Page 31

32 8. How many cations are there in g of lithium nitride? Page 32

33 10. Which of the following amounts would contain the least atoms? a) 10.0 g Sr b) 10.0 g Br c) 10.0 g Mg d) 10.0 g Li Page 33

34 11. Which of the following amounts have the most molecules? a) 15.0 g N2 b) 15.0 g Br2 c) 15.0 g O2 d) 15.0 g I2 Page 34