CHAPTER 2 Atoms, Molecules, Ions

Transcription

1 CHAPTER 2 Atoms, Molecules, Ions By 1800, two important laws were established. Law of conservation of mass (Lavoisier) CH O 2 CO H 2 O There is no change in mass when a chemical reaction occurs. Law of definite proportions (law of constant composition): (Proust) A compound always contains elements combined in the same proportion by mass. Basic postulates of Dalton's atomic theory: 1) Matter is composed of tiny particles called atoms. 2) All atoms of an element are identical in mass and other properties. 3) Atoms of different elements differ in mass and other properties. Lets look at examples: g water (H 2 O) g H g O by mass, H to O ratio is always 1: g methane (CH 4 ) g H g C by mass, H to C ratio is always 1:3 4) Compounds are composed of atoms of different elements combined in fixed proportions by mass. The numbers of atoms of elements in a compound is a small whole number ratio. 5) Atoms are indestructible. Atoms are neither created nor destroyed in chemical reactions; simply py rearranged to yield new substances. This theory explained the law of conservation of mass and the law of constant composition. ATOMIC THEORY OF MATTER John Dalton (180307) Developed the atomic theory of matter it came out of experimental observation. Basic idea: All matter, whether element, compound, or mixture, is composed of small, indivisible particles called atoms. Atoms are the basic building blocks of matter. Sir Humphrey Davy ( ): One of the most renowned scientists of the time. He did not accept Dalton s Atomic Theory. Consequently, it was not initially widely accepted. 1

2 DALTON & PREDICTIVE POWER Law of multiple proportions: When two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers. There are many examples of the law of multiple proportions. Dalton s atomic theory predicted the law of multiple proportions. This provided powerful support of the theory but it does not prove the theory. Sir Humphrey Davy is now convinced of Dalton s theory!! Consider compounds composed of only C and O: Fixed amount of C CO 1.00 g C 1.33 g O CO g C 2.66 g O Ratio of oxygen (O) between the two compounds: 2.66 = (small whole number ratio) Atom: The smallest particle of an element that retains the chemical properties of the element. After Dalton s atomic theory (prior to 1850) it was believed that atoms were simply indivisible balls of matter, indestructible, unchangeable. After 1850, evidence suggested that atoms are composed of even smaller particles called subatomic particles. Hence, atoms have a complex structure. Two compounds of sulfur (S) and oxygen (O): Fixed amount of S A 1.00 g S g O B 1.00 g S g O Ratio of oxygen (O) between the two compounds: Discovery of the First Subatomic Particle Phenomenon of an electrical discharge in an evacuated glass tube, gas discharge tube (Crooks Tube) = 1.5 or (small whole number ratio) Cathode ray tube (CRT) crude television tube. A SO 2 B SO 3 2

3 Cathode () Anode (+) Cathode Ray Tube fluorescent high volt age screen Cathode emits an invisible ray called a cathode ray. It causes gases in the tube to glow and yields a bright spot when it strikes a fluorescent screen (basis for TV screen). Figure 2.3 Cathode Ray Tube J.J. THOMSON S CRT TUBE CRT EXPERIMENTAL OBSERVATIONS Metal plates exposed to Cathode rays acquire a negative charge. J. J. Thomson (1897): Found that magnetic and electric fields bend Cathode rays. negatively charged particles same properties regardless of cathode material suggests that Cathode rays are a basic component of all matter Cathode rays are a beam of electrons. 3

4 J. J. THOMSON measured the chargetomass ratio of electron charge = 1.76 X 10 8 C/g mass minus sign indicates negative charge of electron C Coulomb (unit of electric charge) awarded Nobel Prize in Physics for this work April, 1897 date of electron discovery RADIOACTIVITY Röntgen (1895): Observed that cathode rays caused some materials to also emit rays. emitted rays were able to pass through matter (thick boards of wood) not deflected by magnetic or electric fields Consequently, these rays were not composed of charged particles. These mysterious rays were coined Xrays. Becquerel (1896): Found that some compounds emitted rays spontaneously without cathode ray stimulation. MARIE CURIE Figure 2.5 Suggests term "radioactivity" describes spontaneous emission of particles and/or radiation from matter. Marie Curie was awarded two Nobel prizes (1903 in physics with Becquerel for radioactivity ); (1911 in chemistry for discovery of radium and polonium). Robert Millikan (1909): Determined the charge on an electron with the Oil Drop Experiment X C Now, electron mass can be determined: mass of electron charge 1 g 1.76 X 10 8 C ( charge ) = mass of l t electron ( 1.60 X 10 C) = 9.09 X 10 g The electron was well characterized by ERNEST RUTHERFORD Found that there are three types of rays produced by radioactivity. Alpha (α) rays: A beam of positively charged Alpha (α) rays: A beam of positively charged particles (α particles or Helium nuclei). Beta (β) rays: A beam of negatively charged particles (β particles or electrons). Gamma (γ) rays: High energy radiation; no charge; do not consist of particles, similar to Xrays. 4

5 Radioactive Rays Figure 2.6 Figure 2.7 STRUCTURE OF THE ATOM (Before 1908) We know that atoms contain negatively charged electrons are electrically neutral also contain positively charged matter Ernest Rutherford (1910) with Geiger and Marsden Studied scattering of a beam of α (alpha) particles by a very thin gold foil. Expected the dense α particles to undergo only very slight scattering. Observed small scattering of α particles, but some α particles had large scattering angles. Some α particles actually bounced back!!! J. J. Thomson (Early 1900 s before 1908) αparticle scattering experiment Model of the atom: composed of a uniform sphere of positively charged matter in which negatively charged electrons are embedded electrons only comprise a small fraction of atom s mass (ca. 1/2000 the mass of an atom) This is referred to as the plum pudding (English dessert) model for the structure of the atom. Figure 2.8 5

6 RUTHERFORD S MODEL OF THE ATOM (1911) Atom is composed of two distinct parts: 1) A dense central core (nucleus) that carries a positive charge. 2) Negatively charged electrons surround the nucleus and are far removed from the nucleus. The atom is mostly empty space!! This is called the nuclear model of the atom; the beginning of the modern view of atomic structure. αparticle scattering experiment Rutherford s αparticle scattering experiment Figure 2.8 Figure 2.8 Rutherford s αparticle scattering experiment: These results were inconsistent with Thomson s plumpudding model of atom. How? Should only see only small scattering angles for the αparticles. Based on this new information, Rutherford postulated a new model for the structure of the atom in

7 elect rons _ 1 oz hummingbird dense nucleus + Your body (150 lbs) (hydrogen nucleus) Modern View of the Structure t of the Atom Hummingbird is 2 miles from your body!!! The atom is composed of three subatomic particles: Protons: Positively charged particles that are part of the nucleus. Neutrons: Neutral particles that are part of the nucleus. Electrons: Negatively charged particles that are far removed from the nucleus. There are additional subatomic particles but these are not important to chemistry. Protons: Discovered by Rutherford in Neutrons: Discovered by James Chadwick in Neutrons and protons are part of the nucleus and have very similar masses. The neutron was difficult to discover and characterize because it is neutral (it is not affected by magnetic and electric fields as are charged particles). 7

8 Properties of Subatomic Particles (Table 2.1) Particle Charge Mass (g) Electron 1.60 X C X g Proton X C X g Neutron X g Note: Proton and electron have the same magnitude charge but opposite in sign. For convenience: Protons +1 Electrons 1 elect rons _ 1 oz hummingbird dense nucleus + Your body (150 lbs) (hydrogen nucleus) Hummingbird is 2 miles from your body!!! FACTS ABOUT ATOMS Atoms are electrically neutral, therefore: # protons = # electrons Protons and neutrons much more massive than electrons (ca. a factor of 1840). Essentially all of an atoms mass is found in the nucleus. What characterizes an element? The number of protons. Henry Moseley (1913) For example: 1 Proton H (hydrogen) 6 Protons C (carbon) 15 Protons P (phosphorus) Atomic number (Z): Number of protons, identifies the element chemical identity. FACTS ABOUT ATOMS Atoms are very small: Mass 1 X g Diameter (atom) 1 X m Diameter (nucleus) 1 X m (10,000 times smaller than the atomic radius) An atom is mostly empty space!!! Mass number (A): The total number of protons and neutrons for the atom. For example: Carbon: 6 protons and 6 neutrons Mass number of 12. Atoms of a particular element can have variable numbers of neutrons. 8

9 For example: C: 6 protons + 7 neutrons = 13 C: 6 protons + 8 neutrons = 14 Neutrons only affect the mass of the atom, not the chemical properties. C 12; C 13; C 14 These are called isotopes. Isotopes: Atoms of the same element that have different mass numbers different numbers of neutrons. The # of neutrons only affect the mass of the atom, not the chemical properties. Atoms of the same element that have different mass numbers have different # s of neutrons. The nucleus of a specific isotope is called a nuclide. Important Radioactive Isotopes and Their Uses Cobalt60 ( 60 Co) Cancer radiation therapy Cesium137 ( 137 Cs + ion) Food Sterilization (as cesium chloride) Iodine129 ( 129 I ion) Thyroid radiation therapy (as NaI) Technetium99 ( 99 Tc) Medical imaging of internal organs Chemical Symbols: Isotopes Example: Carbon Isotopes mass number atomic number Naturally occurring 12 6 radioactive C 6 C 6 C These carbon isotopes have identical chemical properties. ISOTOPES J. J. Thomson Discovered isotopes ( ) Ne20, Ne22 Discovery of isotopes is in conflict with Dalton's atomic theory (Postulate 2). Modify theory to include isotopes: "all atoms of a given element have the same number of protons and chemical properties". 9

10 For any atom: Mass number = # of protons + # of neutrons and # of protons = # of electrons If we know two of the former, the third can be determined. For example, how many neutrons are in the O18 nuclide? All oxygen atoms have 8 protons. Number of neutrons = 18 8 = th century Scientists recognized that some elements had very similar chemical properties. Mendeleev (Russian) and Meyer (German) 1869 Elements were arranged by increasing atomic number (mass originally) in a horizontal row. Elements with similar chemical properties arranged in columns called Groups. Proposed an extensive tabulation of the elements based on the regular, periodic recurrence of properties; known as the periodic table How many protons, neutrons, and electrons are in an atom of Tungsten ( 184 W)? Atomic number, Z = 74, 74 protons For an atom, # protons = # electrons; 74 electrons Mass number, A = 184, = 110 neutrons Figure 2.10 TimeLine for Discovery of Elements Three categories for the elements: METALS NONMETALS METALLOIDS 10

11 METALS Metals, Nonmetals, Metalloids located left and center of Table good conductors of heat and electricity luster (shiny) ductile (drawn into wire) malleable (pounded into sheets) All metals are solids at room temperature except mercury (Hg) which is liquid. NONMETALS located to the right of the Table typically have properties opposite to those for metals poor conductors of heat and electricity solids are brittle (not malleable or ductile) dull surface (not shinny) Periodic Table Group A elements representative or main group elements Group B elements transition (metal) elements Two rows below the table inner transition (metal) elements Lanthanide elements (5871) Actinide elements (90103) METALLOIDS located to rightcenter of Table divide nonmetals and metals properties intermediate between metals and nonmetals. Most elements are metals Of the first 92 elements 67 metals 17 nonmetals 8 metalloids Periodic Table Group numbers 1A, 2A, 3A...8A 1B, 2B, 3B...8B Some groups have special names Group 1A alkali l metals (not hydrogen) Group 2A alkaline earth metals Group 6A chalcogens Group 7A halogens Group 8A noble (rare) gases 11

12 Modern Periodic Table Elemental Forms of Important Nonmetallic Molecules: H 2, O 2, N 2, F 2, Cl 2 RT (room temp.) Br 2 RT I 2 RT All are diatomic molecules (made up of two atoms) Other nonmetals have more complex forms For example: P 4, S 8 solids Group 8A elements (noble or rare gases) exist in nature as free, single atoms. Most other elements exist as molecules or ions in nature. Atoms combine in two ways: 1) By sharing electrons to yield molecules. 2) By transferring electrons to yield ions resulting in the formation of ionic compounds. Elemental oxygen: Two forms O 2 and O 3 O 2 and O 3 are allotropes of the element oxygen. Allotropes: Different forms of an element Carbon: Three allotropes (graphite, diamond, fullerenes (1985)). Sulfur: S 2, S 4, S 6, S 8 molecules Molecules Formed by the combination of nonmetals with nonmetals or metalloids. Have a definite shape (geometry, Chapter 10) Atoms are combined in definite proportions by mass (constant composition). All are neutral species. Molecular Compounds Molecules composed of two or more different nonmetals or metalloids. HCl, CO diatomic molecular compounds CO 2, NH 3, H 2O polyatomic molecular compounds SiO 2, SiH 4 polyatomic molecular compounds containing a metalloid 12

13 CHEMICAL FORMULAS Molecular, Empirical, and Structural Denote elements in a compound by chemical symbols and the ratios in which the atoms are combined by subscripts. Molecular formula: A chemical formula that denotes the exact number of atoms of each element in a molecule. Examples: H 2 O, NH 3, CH 4, C 2 H 6, C 6 H 6 Subscripts denote the exact number of atoms of that element in one molecule. Ions and Ionic Compounds Ions: Species containing a net charge by either gaining electrons (negative charge) or by losing electrons (positive charge) Cation: Positively charged ion, attracted to the negative cathode () in an electrolytic cell. Anion: Negatively charged ion, attracted to the positive anode (+) in an electrolytic cell. Metals: Tend to lose electrons cations. Nonmetals: Tend to gain electrons anions. Empirical formula (simplest): Gives the smallest wholenumber ratio of atoms of each element in the compound. Compare molecular and empirical formulas: Molecular l Empirical i C 6 H 6 CH C 2 H 4 CH 2 CH 4 CH 4 N 2 O 5 N 2 O 5 Monatomic Ions Ions with only one atom; metals and nonmetals Monatomic metal ions (Cations) Examples: Na +, Ca 2+, K +, Mg 2+, Al 3+ How are they formed? By the loss of one or more e s M M n+ + ne (n = # of electrons) Na Na + + e Ca Ca e Structural Formulas: Show which atoms are connected (chemically bonded) together. Examples: H 2 O H O H H 2 O 2 H O O H CO 2 O C O H H N 2 H 4 N N H H Monatomic nonmetal ions (Anions) F, Cl, O 2, N 3 How are they formed? By gaining one or more e s Examples: X+ne X n F + e F N+3e N 3 13

14 N 7 protons and 7 electrons Common Stable Monatomic Ions & Charges N 3 7 protons and 10 electrons (7 + 3) Mg 12 protons and 12 electrons Mg protons and 10 electrons (12 2) S 16 protons and 16 electrons S 2 16 protons and 18 electrons (16 + 2) The periodic table is very useful for determining the charges for stable monatomic ions. Predict Charges for Stable Monatomic Ions The number of electrons Group 1A or 2A metal atoms loose is equal to the Group number. Group 1A: Lose 1e to form M + ions. Polyatomic Ions Ions that consist of atoms chemically bonded together in a molecular sense and carry a net charge (positive or negative). Examples: NO 3, SO 2 4, PO 3 4, CO 2 3, NH + 4 Group 2A: Lose 2e to form M 2+ ions. The number of electrons that a nonmetal atom gains is equal to 8 minus the Group number. Group 7A: Gain 1e to form X ions (8 7 = 1) Group 6A: Gain 2e to form X 2 ions (8 6 = 2) Group 5A: Gain 3e to form X 3 ions (8 5 = 3) Figure 2.11 (Page 54) Common stable monatomic ions and their chargesknow!!! Ionic Compounds Composed of cations and anions interacting in appropriate proportions to yield an overall neutral compound. e.g., common table salt ionic compound composed of Na cations and Cl anions A crystal lattice structure is formed. 14

15 Cl Na + Cl Na + Na + Cl Na + Cl Cl Na + Cl Na + Na + Cl Na + Cl Every Na + ion is surrounded by Cl ions and every Cl ion is surrounded by Na + ions. Opposite charges attract!!! CHEMICAL FORMULAS OF IONIC COMPOUNDS Formulas represent the smallest wholenumber ratio of the ions in the compound (Empirical formula). All ionic compounds are represented by empirical formulas (formula unit). Always electrically neutral, therefore, the amount of positive charge must equal amount of negative charge. + = Ionic Compounds (NaCl) An extended 3d array is formed where the number of Na cations equals the number of Cl anions. Consequently, discrete molecular units are not formed (i.e., we do not have individual NaCl molecules). Write Formulas For Ionic Compounds 1) Ionic compounds are always neutral. 2) Cation is always written first then anion. 3) Subscripts denote number of ions of each type required to yield a neutral compound (smallest wholenumber ratio); referred to as a formula unit. MAKING AN IONIC COMPOUND Al 3+ +O 2 IONIC COMPOUND Na + Cl NaCl FORMATION OF AN IONIC COMPOUND Al 3+ + O 2 Al 2 O 3 For charge balance, the subscript of the cation is equal to the charge of the anion; the subscriptp of the anion is equal to the charge of the cation. 2 (3+ charges) + 3 (2 charges) = zero or a neutral compound Na + + CO 3 2 Na 2 CO 3 15

16 Mg 2+ + O 2 Mg 2 O 2 NAMING INORGANIC COMPOUNDS INCORRECT MgO is the smallest wholenumber ratio of ions simplify!!! We need systematic rules for naming compounds. Four categories of compounds: 1) Ionic compounds 2) Acids 3) Molecular compounds 4) Hydrates Write the Chemical Formulas for the ionic compounds formed from the following cations and anions: A.) K + and Br B.) Ca 2+ and NO 3 C.) Ag + and SO 4 2 D.) Fe 3+ and SO 4 2 Ionic Compounds: Named from the IONS Cation first Anion second MONATOMIC METAL CATIONS Named from the elements Examples: Na + sodium ion Ca 2+ calcium ion Mg 2+ magnesium ion Naming Inorganic Compounds (Chemical Nomenclature) Formula Name Over 13 million compounds are known. Some have common names. TRANSITION METAL IONS (Group B elements) Can form stable ions with variable charges. For example, Iron forms ions with variable charge: Examples: H 2 O water NH 3 ammonia AsH 3 arsine SiH 4 silane Fe 2+ and Fe 3+ Use the Stock system to name these ions. 16

17 Stock System: element name(charge) ion Roman numerals Fe 2+ iron(ii) ion Fe 3+ iron(iii) ion Polyatomic Anions: Many contain oxygen; classified as oxoanions; have suffixes ate and ite. Formula OH CN O 2 2 CO 2 3 HCO 3 MnO 4 CrO 4 2 Cr 2 O 7 2 PO 4 3 Ion Name hydroxide ion cyanide ion peroxide ion carbonate ion hydrogen carbonate ion (bicarbonate ion) permanganate ion chromate ion dichromate ion phosphate ion POLYATOMIC CATIONS Two common ions. NH 4+ ammonium ion H 3 O + hydronium ion Name ends in ium plus ion. Polyatomic Anions (cont.) Formula Ion Name SCN thiocyanate ion C 2 H 3 O 2 acetate ion NO 2 nitrite ion NO 3 nitrate ion SO 2 3 sulfite ion SO 2 4 sulfate ion ClO 2 chlorite ion ClO 3 chlorate ion Table 2.3 Polyatomic ions and names Know!!! ANION NAMES Monatomic nonmetal anions Know Table 2.2 Name derived from stem (root) of element name followed by the suffix ide plus ion. Element Ion name Ion Name F fluorine fluoride ion Cl chlorine chloride ion O 2 oxygen oxide ion S 2 sulfur sulfide ion Oxoanions: For a homologous series (e.g., NO 3, NO 2 ) same central atom, variable number of oxygen atoms: ite suffix anion with less oxygen ate suffix anion with more oxygen If more than two oxyanions in series, use prefixes: hypo (even less oxygen than ite) per (even more oxygen than ate) 17

18 Example: ClO ClO 2 ClO 3 ClO 4 hypochlorite ion chlorite ion chlorate ion perchlorate ion Naming ACIDS Acids: Compounds that yield hydronium ions, H 3 O +, when dissolved in water. HCl + H 2 O H 3 O + + Cl HNO 3 + H 2 O H 3 O + + NO 3 Shorthand: H 2 O HNO 3 H + + NO 3 Naming Ionic Compounds Cation named first, followed by anion: NaCl sodium chloride CaCl2 calcium chloride Ca(NO3)2 calcium nitrate BaSO4 barium sulfate Al 2 (SO 4 ) 3 aluminum sulfate Note: Ion is not included in the name. Naming Acids: Acids based on anions whose name ends in ide: Add hydro prefix Change ide suffix to ic acid Cl (chloride ion) HCl (hydrochloric acid) CN (cyanide ion) HCN (hydrocyanic acid) S 2 (sulfide ion) H 2 S (hydrosulfuric acid) Transition metal ionic compounds specify charge by using the Stock system. FeO; Fe 2+ and O 2 ; iron(ii) oxide Acids from oxoanions: Based on anion root name; simply change the suffix: For ate suffix; change to ic acid For ite suffix; change to ous acid Fe 2 O 3 ; Fe 3+ and O 2 ; FeCl 3 ; Fe 3+ and Cl ; FeSO 4 ; Fe 2+ Fe 2 (SO 4 ) 3 ; Fe 3+ iron(iii) oxide iron(iii) chloride iron(ii) sulfate iron(iii) sulfate Examples: SO 2 4 (sulfate ion) H 2SO 4 (sulfuric acid) SO 2 3 (sulfite ion) H 2SO 3 (sulfurous acid) NO 3 (nitrate ion) HNO 3 (nitric acid) NO 2 (nitrite ion) HNO 2 (nitrous acid) ClO 4 (perchlorate ion) HClO 4 (perchloric acid) 18

19 Binary Molecular Compounds Binary compounds: Contain only two elements Complication: Acids whose anion ends in ide. If these compounds are not in water (i.e., pure compounds), then we name them as binary molecular compounds. Rules similar to those for binary ionic compounds, e.g., NaCl (sodium chloride). Given Formula: Name first as the element, the second as an anion (root with ide suffix). Cl (chloride ion) CN (cyanide ion) S 2 (sulfide ion) HCl (hydrogen chloride) in water (hydrochloric acid) HCN (hydrogen cyanide) in water (hydrocyanic acid) H 2 S (hydrogen sulfide) in water (hydrosulfuric acid) Specify number of atoms of each element by using Greek prefixes (Table 2.4): Greek Prefixes Prefix Meaning Mono 1 Di 2 Tri 3 Tetra 4 Penta 5 Hexa 6 Prefix mono never used for the first element. Hydrates Hydrate: A compound with a specific number of water molecules that are strongly adhered. CuSO 4 copper(ii) sulfate (anhydrous) CuSO. 4 5H 2 O copper(ii) sulfate pentahydrate t BaCl 2 barium chloride (anhydrous) BaCl 2. 2H 2 O barium chloride dihydrate Anhydrous: Dry, without water When the prefix ends in a or o and the anion name begins with a vowel, the a or o of the prefix is often omitted. What are the chemical names or formulas of the following compounds? Formula CO CO 2 PCl 3 NO 2 N 2 O 5 SO 3 SO 2 Name carbon monoxide carbon dioxided phosphorus trichloride nitrogen dioxide dinitrogen pentoxide sulfur trioxide sulfur dioxide (NH 4 ) 2 SO 4 Hydroiodic acid (hydrogen iodide) KMnO 4 Mercury(II) nitrite Fe(NO 3 ) 3 Potassium dichromate 19

20 Note: An ion and the corresponding neutral species have completely different chemical properties. NO 2 NO 2 SO 2 3 SO 3 Na Na + Fe Fe 3+ H H + 20