Week 3: Atoms and Elements CHEM 1010 PDF

CHEM 1010 General Chemistry I Week 3. Atoms and Elements  Week 3: Atoms and Elements Objectives: 1. Explain the structure of an atom 2. Give the symbol and name of the element mentioned in this course 3. Carry out calculations involving the determination of the number of...

CHEM 1010 General Chemistry I Week 3. Atoms and Elements  Week 3: Atoms and Elements Objectives: 1. Explain the structure of an atom 2. Give the symbol and name of the element mentioned in this course 3. Carry out calculations involving the determination of the number of protons, electrons and neutrons in a given isotope 4. Identify the groups and period to which an element belongs 5. Identify regions of the periodic table (metals, metalloids, nonmetals, alkali metals, alkaline earth metals, and transition metals) 6. Use isotopic masses and natural abundance to calculate the average atomic masses 2  If You Cut a Piece of Graphite… How far could you go? Could you divide it forever? Smaller pieces (far smaller than the eye could see) would eventually yield individual carbon atoms You cannot divide a carbon atom into smaller pieces and still have carbon Atoms are the key to connecting the macroscopic and microscopic worlds 3  Atoms An atom is the smallest identifiable unit of an element About 91 different naturally occurring elements + over 20 synthetic Salt Copper NOT an element An ELEMENT (made of sodium and chlorine) 4  2.3 Modern Atomic Theory and the Laws That Led to It The theory that all matter is composed of atoms grew out of observations and laws The three most important laws that led to the development and acceptance of the atomic theory were 1. the law of conservation of mass 2. the law of definite proportions 3. the law of multiple proportions 5  The Law of Conservation of Mass Antoine Lavoisier formulated the law of conservation of mass In a chemical reaction, matter is neither created nor destroyed The total mass of the substances involved in the reaction does not change The law is consistent with the idea that matter is composed of small, indestructible particles. 6 Concept Check 7  The Law of Definite Proportions Joseph Proust formulated the law of definite proportions in 1797 All samples of a given compound always have the same proportions of their constituent elements Also called the law of constant composition Illustration The decomposition of 18.0 g of water results in 16.0 g of oxygen and 2.0 g of hydrogen, or an oxygen-to- hydrogen mass ratio of: 16.0 g O Mass = ratio = 8.0 or 8:1 2.0 g H For ammonia (NH3), the mass ratio is 14.0 g:3.0 g i.e. 4.7:1 8  Concept Check We just saw that the mass ratio of nitrogen to hydrogen in ammonia is 4.7:1. If a sample of ammonia contains 10.0 grams of H, how many grams of N does it contain? a. 4.7 b. 9.4 c. 14 d. 47 9  The Law of Multiple Proportions In 1804, John Dalton published his law of multiple proportions When two elements A and B form two different compounds, the masses of B that combine with a fixed mass of A can be expressed in small whole number ratios An atom of A combines with either one, two, three, or more atoms of B (AB1, AB2, AB3, etc.). 10  The Law of Multiple Proportions Carbon monoxide and carbon dioxide are two compounds composed of the same two elements: carbon and oxygen. The mass ratio of oxygen to carbon in carbon dioxide is 2.67:1; therefore, 2.67 g of oxygen reacts with 1 g of carbon In carbon monoxide, however, the mass ratio of oxygen to carbon is 1.33:1, or 1.33 g of oxygen to every 1 g of carbon Mass oxygen to 1 g carbon in carbon dioxide 2.67 = = 2 Mass oxygen to 1 g carbon in carbon monoxide 1.33 Ratio of ratios! 11 Example 2.2 Law of Multiple Proportions Nitrogen forms several compounds with oxygen, including nitrogen dioxide and dinitrogen monoxide. Nitrogen dioxide contains 2.28 g oxygen to every 1.00 g nitrogen, while dinitrogen monoxide contains 0.570 g oxygen to every 1.00 g nitrogen. Show that these results are consistent with the law of multiple proportions. Solution Calculate the ratio of the mass of oxygen from one compound to the mass of oxygen in the other. Always divide the larger of the two masses by the smaller one. The ratio is a small whole number (4); these results are consistent with the law of multiple proportions. For Practice 2.2 Hydrogen and oxygen form both water and hydrogen peroxide. The decomposition of a sample of water forms 0.125 g hydrogen to every 1.00 g oxygen. The decomposition of a sample of hydrogen peroxide forms 0.0625 g hydrogen to every 1.00 g oxygen. Show that these results are consistent with the law of multiple proportions. 12  John Dalton and the Atomic Theory Dalton’s atomic theory explained the laws as follows: 1. Each element is composed of tiny, indestructible particles called atoms. 2. All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements. 3. Atoms combine in simple, whole-number ratios to form compounds. 4. Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms only change the way that they are bound together with other atoms. 13 The Structure of the Atom 14  Nucleus NUCLEUS: the hard center of an atom Contains protons and neutrons only Electrons stay outside the nucleus Even though it has 99% of the atom’s mass, it is very small 15  Electron Outside the nucleus Has a mass of almost 0 9.11x10-31 kg Has a negative charge The movement of electrons is responsible for almost all chemistry Electrons 16  Protons vs Neutrons Proton Neutron Inside the nucleus Inside the nucleus 1000x more mass than an 1000x more mass than an electron electron POSITIVE charge NEUTRAL (no) charge Neutrons Protons (n0) (p+) 17  2.6 Subatomic Particles The charge of the proton and the electron are equal in magnitude but opposite in sign The neutron has no charge 18  Elements: Defined by Their Numbers of Protons The most important number to the identity of an atom is the number of protons in its nucleus The number of protons defines the element The number of protons in an atom’s nucleus is its atomic number and is given the symbol Z 19 Periodic Table of the Elements 20 Periodic Table of the Elements  Periodic Table  Each element has its own unique name and CHEMICAL SYMBOL 1 or 2 letters o First letter is always capital o If there is a second letter, it is lower-case C O Co CO X  Periodic Table  CHEMICAL SYMBOLS may come from the element’s English or original Latin/Greek name C - Carbon Au - gold H - Hydrogen Ag - silver He - Helium Hg - mercury Au: gold (from Latin “Aurum”) Ag: silver (from Latin “Argentum”) Hg: mercury (from Greek “Hydrargyrum”) Note: You must remember the names and symbols of common elements. The list of these elements are available on your D2L account in the course resources module.  Concept Check Which element contains 50 protons in its nucleus? a. tin b. vanadium c. manganese d. fermium 24  Isotopes: When the Number of Neutrons Vary All atoms of a given element have the same number of protons They do not necessarily have the same number of neutrons For example: all neon atoms contain 10 protons, but they may contain 10, 11, or 12 neutrons All three types of neon atoms exist, and each has a slightly different mass 20 21 22 10 Ne10Ne10 Ne Atoms with the same number of protons but a different number of neutrons are called isotopes 25  Isotopes: When the Number of Neutrons Vary The relative amount of each different isotope in a naturally occurring sample of a given element is roughly constant These percentages are called the natural abundance of the isotopes Advances in mass spectrometry have allowed accurate isotopic mass measurements 20 21 22 10 Ne10Ne10 Ne Ne-20 Ne-21 Ne-22 neon-20 neon-21 neon-22 26  Isotopes: Varied Number of Neutrons Symbol Number of Number of A (Mass Natural Protons Neutrons Number) Abundance (%) 10 10 20 90.48 N e 20, or super 20 sub 10 N e Ne 10 11 21 0.27 N e 21, or super 21 sub 10 N e Ne 10 12 22 9.25 N e 22, or super 22 sub 10 N e Ne 27  Concept Check An argon isotope has a mass number of 40 (A = 40). How many neutrons does it contain? a. 40 b. 18 c. 22 d. 20 28  Concept Check Carbon has two naturally occurring isotopes: C-12 (natural abundance is 98.93%) and C-13 (natural abundance is 1.07%). If circles represent protons and squares represent neutrons, which image best represents the C-13 isotope? 29  Ions: Losing and Gaining Electrons The number of electrons in a neutral atom is equal to the number of protons in its nucleus (Z) Atoms can lose or gain electrons during chemical change and become charged particles called ions Positively charged ions, such as Na+, are called cations Negatively charged ions, such as F−, are called anions. 30  Concept Check How many electrons are present in the O2− anion? a. 2 b. 6 c. 8 d. 10 In light of the nuclear model of the atom, which statement is true? a. For a given element, the size of an isotope with more neutrons is larger than one with fewer neutrons. b. For a given element, the size of an atom is the same for all of the element’s isotopes. 31 2.7 Finding Patterns: The Periodic Law and the Periodic Table The Periodic Law In 1869, Mendeleev noticed that certain groups of elements had similar properties When elements are listed in order of increasing mass, these similar properties recurred in a periodic pattern To be periodic means to exhibit a repeating pattern Mendeleev summarized these observations in the periodic law: when the elements are arranged in order of increasing mass, certain sets of properties recur periodically 32  The Periodic Table Mendeleev organized the known elements in a table He arranged the rows so that elements with similar properties fall in the same vertical columns The table contained some gaps, which allowed him to predict the existence (and even the properties) of yet undiscovered elements Mendeleev predicted the existence of an element (eka-silicon) In 1886, eka-silicon was discovered by Clemens Winkler, who named it germanium 33  Modern Periodic Table Organization Elements are listed in order of increasing atomic number rather than increasing relative mass Contains more elements than Mendeleev’s original table 34  Classification of Elements Elements in the periodic table are classified as Metals Nonmetals Metalloids 35  Metals Metals, on the lower-left side and middle of the periodic table, share some common properties. Good conductors of heat and electricity Can be pounded into flat sheets (malleability) Can be drawn into wires (ductility) Often shiny Tend to lose electrons when they undergo chemical changes Chromium, copper, strontium, and lead are typical metals 36  Non-Metals Nonmetals lie on the upper-right side of the periodic table There are a total of 17 nonmetals Five are solids at room temperature (C, P, S, Se, and I) One is a liquid at room temperature (Br) Eleven are gases at room temperature (H, He, N, O, F, N e, C l, Ar, Kr, Xe, and Rn) Nonmetals as a whole tend to have these properties: Poor conductors of heat and electricity Not ductile and not malleable Gain electrons when they undergo chemical changes Oxygen, carbon, sulfur, bromine, and iodine are nonmetals 37  Metalloids Metalloids are sometimes called semimetals They are elements that lie along the zigzag diagonal line that divides metals and nonmetals They exhibit mixed properties Several metalloids are also classified as semiconductors 38  Periodic Table The periodic table can also be divided into: Main-group elements, whose properties tend to be largely predictable based on their position in the periodic table Transition elements or transition metals, whose properties tend to be less predictable based simply on their position in the periodic table 39  Periodic Table The periodic table is divided into vertical columns and horizontal rows Each vertical column is called a group (or family) Each horizontal row is called a period There are a total of 18 groups and 7 periods. The groups are numbered 1–18 (or the A and B grouping) Main-group elements are in columns labeled with a number and the letter A (1A–8A or groups 1, 2, and 13–18). Transition elements are in columns labelled with a number and the letter B (or groups 3–12) 40  Noble Gases The elements within a group usually have similar properties The group 8A elements, called the noble gases, are mostly unreactive The most familiar noble gas is probably helium, used to fill buoyant balloons Helium is chemically stable—it does not combine with other elements to form compounds—and is therefore safe to put into balloons Other noble gases are neon (often used in electronic signs), argon (a small component of our atmosphere), krypton, and xenon 41  Alkali Metals The group 1A elements, called the alkali metals, are all reactive metals A marble-sized piece of sodium explodes violently when dropped in water Lithium, potassium, and rubidium are also alkali metals 42  Alkaline Earth Metals The group 2A elements are called the alkaline earth metals They are fairly reactive but not quite as reactive as the alkali metals Calcium, for example, reacts fairly vigorously with water Other alkaline earth metals include magnesium strontium, and barium 43  Halogens The group 7A elements, the halogens, are very reactive nonmetals. They are always found in nature as a salt Fluorine, a pale-yellow gas Chlorine, a greenish-yellow gas with a pungent odor Bromine, a red-brown liquid that easily evaporates into a gas Iodine, a purple solid 44  Concept Check Which terms best apply to lead (Pb)? a. Transition element; metal b. Main-group element; nonmetal c. Halogen; nonmetal d. Main-group element; metal 45  Ions and the Periodic Table A main-group metal tends to lose electrons, forming a cation with the same number of electrons (isoelectronic) as the nearest noble gas A main-group nonmetal tends to gain electrons, forming an anion with the same number of electrons (isoelectronic) as the nearest noble gas In general, the alkali metals have a tendency to lose one electron and form +1 ions The alkaline earth metals (group 2A) tend to lose two electrons and form +2 ions The halogens (group 7A) tend to gain one electron and form −1 ions The oxygen family nonmetals (group 6A) tend to gain two electrons and form −2 ions 46  Ions and the Periodic Table For the main-group elements that form cations with a predictable charge, the charge is equal to the group number (e.g. Group 1A elements form +1 cations, 2A form +2 cations, etc.) For main-group elements that form anions with a predictable charge, the charge is equal to the group number minus 8 (e.g. Group 7A elements form -1 anions, 6A form -2 anions) Transition elements may form various different ions with different charges (e.g. Fe can form +2 and +3 cations, etc.) 47 Ions and the Periodic Table 48 2.8 Atomic Mass: The Average Mass of an Element’s Atoms Atomic mass is sometimes called atomic weight or standard atomic weight The atomic mass of each element is directly beneath the element’s symbol in the periodic table It represents the average mass of the isotopes that compose that element, weighted according to the natural abundance of each isotope 49  Example Naturally occurring chlorine consists of 75.77% chlorine-35 atoms (mass 34.97 amu) and 24.23% chlorine-37 atoms (mass 36.97 amu). Calculate its atomic mass Solution: Convert the percent abundance to decimal form and multiply each with its isotopic mass and add them. Cl-37 = 0.2423(36.97 amu) = 8.9578 amu Cl-35 = 0.7577(34.97 amu) = 26.4968 amu Atomic Mass Cl = 8.9578 + 26.4968 = 35.45 amu 50  Atomic Mass In general, we calculate the atomic mass with the following equation: 51  Concept Check Carbon has two naturally occurring isotopes: C-12 (natural abundance is 98.93%; mass is 12.0000 amu) and C-13 (natural abundance is 1.07%; mass is 13.0034 amu). Without doing any calculations, determine which mass is closest to the atomic mass of carbon. a. 12.00 amu b. 12.50 amu c. 13.00 amu d. 11.50 amu 52  Pearson Platform and eBook Readings/Exercises: Week 3: Atoms and Elements Readings: Chapter 2.3, 2.4, 2.5, 2.6, 2.7, 2.8 Pearson Platform Tutorial 2 (Practice Questions) Note: Tutorial questions will be available soon on the Pearson Platform 53