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Unit 2: Atoms, Molecules, and Ions PART 1 An overview of the evolution of the Atomic Model from John Dalton to the Modern Theory. But first, a little ancient history: http://www.meta-synthesis.com/webbook/35_pt/pt_database.php 1. John Dalton’s Atomic Theory – 1803 Dalton stated a group of assumptions to explain the nature and behavior of chemical systems. These became known as Dalton’s Atomic Theory and he proposed this the year 1803. The primary difference between Dalton’s theory and previous ones was that Dalton’s was based on reproducible laboratory evidence. Dalton’s Atomic Theory Four Assumptions: 1. All substances are composed of small, dense particles called ATOMS. 2. Atoms of a given substance are identical in mass, size and shape. 3. An atom is the smallest part of an element that enters into a chemical reaction. 4. Molecules are produced by a combination of atoms. Dalton based his theory on the work of a number of scientists who developed the following laws: 1. The Law of Conservation of Mass • Or the Law of Conservation of Matter • The law states that in ordinary chemical reactions, the mass of the system remains constant. Zn + S → ZnS 65.4g + 32.1 g = 97.5 g 2. The law of conservation of energy • The heat lost by the system (reaction) is equal to the heat gained by the surroundings (or, in ordinary chemical reactions, the energy of the system remains constant). • ENDOTHERMIC: heat energy is absorbed during a chemical reaction (it gets colder) • EXOTHERMIC: heat energy is released during a chemical reaction (it gets hotter!) • 95% of all reactions are exothermic!! • Heat + nitrogen + oxygen → nitric oxide • q + N2 + O2 → 2NO (q is one of the symbols used for heat or energy) If q is written on the left side, this indicates the reaction is endothermic (since E is being absorbed). If q is written on the right side of the equation, the reaction is exothermic (since heat is being released to the surroundings) 3. The law of definite composition • When elements combine and form specific compounds, they do so in definite proportions by mass. Zn + S → ZnS 65.4g + 32.1 g = 97.5 g Zinc and sulfur will always combine in a definite fixed ratio of 65.4 parts to 32.1 parts by mass. If Zn or S were present in any other ratio, the one in excess would remain unchanged or unused. The excess would remain unreacted!! 4. The law of multiple proportions • When two elements combine and form more than one compound, the masses of one element that combine with a fixed mass of the other are in the ratio of small, whole numbers. • Look at the formulas of C to O in your notes: notice that in none of those formulas do you see C1.2O2.67 or C.98O3.11 • Formulas are always WHOLE NUMBERS! C O C/O Ratio CO 12g 16g 16/12 = 1.3 1.3/1.3 =1 CO2 12g 32g 32/12 =2.6 2.6/1.3 =2 CO3 12g 48g 48/12 =4.0 4.0/1.3 =3 5. Guy-Lussac’s Law of combining volumes • Gases react chemically with a volume of small, whole numbers!! • 1 volume H2 + 1 volume Cl2 → 2 volumes HCl • S • 1L + + • 2H2 + • 2L + O2 1L → = SO2 1L O2 1L → = 2H2O 2 Liters • Dalton’s Atomic Theory is often called the Billiard Ball Model. His theory, however, fails to explain many different types of behavior in chemical reactions. “Billiard Ball Model” Discovery of the electron - 1897 • JJ Thomson studied electrical discharges in partially, evacuated tubes called cathode –ray tubes Diagrams: Initial: After: • J.J. Thomson- In a series of experiments with a cathode ray tube (CRT) in 1897, discovered that negatively charged particles of matter could be removed from atoms. • This indicated that atoms were not indivisible but were composed of even smaller particles. • The discovered particle was the electron. • Since the ray was attracted to the positive electrode, Thomson called the stream of particles electrons. • Thomson also determined the charge-to-mass ratio of the electron to be: e/m = -1.76 x 108 Coulombs/gram (where e = charge of the e- in Coulombs and m = mass in grams) One of the key findings was that the charge of the electron was negative!! Thomson’s Concept of Atoms • Thomson proposed a “new and improved” model using Dalton’s theory as a foundation (remember – The idea that matter was composed of atoms was not universally accepted (Mendeleev) • Thomson proposed that atoms consist of a solid bulk of positive charge with electrons dispersed throughout. His model is known as the Plum Pudding Model. Diagram of Thomson’s Atom “Plum Pudding Model” A brief aside to discuss Radioactivity • The French scientist Henri Becquerel found that a piece of Uranium produced its image on a photographic plate. So, as a third arm was growing out of his ribs, he figured out that this U had some weird energy coming from it (kind of like a chicken patty sandwich) • RADIOACTIVITY: the spontaneous emission of radiation (particles with lots of energy!!) • Some materials are radioactive and may produce different types of radiation. • These “particles” suggested that there may be other parts to an atom. Some types of radiation Type of Radiation Symbol Definition Alpha particle α Like a helium atom (2 protons fused together); +2 charge; mass is 7300 times an e- Beta particle β High speed/High energy electrons Gamma ray γ “high energy” light; very damaging to our DNA; most damaging of the 3 Discovery of the Nucleus • ERNEST RUTHERFORD – most famous for his discovery of the nucleus. • Rutherford used a radioactive source that emitted alpha (α)particles (these were his “bullets”), a piece of VERY THIN gold foil, a fluorescent screen, and a lead block. • His experiment would be similar to if you took your genuine Red Ryder BB gun and shot BB’s at a Kleenex. What would you expect to happen? Java – Rutherford scattering http://www.mhhe.com/physsci/chemistry/essen tialchemistry/flash/ruther14.swf Gold foil experiment Rutherford’s 3 assumptions 1. The slightly deflected α particle had a close encounter with the “positive center” of the atom. 2. Most of the atom is empty space (because most of the α particles passed through) 3. The alpha particles that were completely deflected hit head on with the nucleus (because likes repel – both were + charged) Ernest Rutherford - 1911 • Rutherford discovered that the positive charge and the mass was concentrated in the center of the atom (called the NUCLEUS). • He postulated that the electrons were moving at fast speeds around the nucleus but were contained by a certain boundary. Diagram of Rutherford’s Atom “Empty Space Model” Discovery of protons Henry Moseley – 1913 – Discovered that the number of positive charges in an atom is equal to the element number. This indicated that there was a particle in the nucleus that was the source of + charge. In 1920, Rutherford named the particle “proton”. Discovery of neutrons • James Chadwick – 1932 – Explained the difference between the observed mass of atomic nuclei and the number of + charges (also considering spin) by proposing the presence of particles with masses similar to those of protons but with no charge - the neutron. The Bohr model of the atom Niels Bohr – 1913 – Considered that the Rutherford model of the atom was unstable and the spectra of atoms (discrete bands of light absorbed and emitted by atoms) to propose a new model with the electrons confined to specific energy levels (sometimes called shells or orbits). Neil Bohr’s Concept of the Atom - 1913 • Bohr proposed that electrons are arranged in definite energy levels(shells) and follow a definite orbit. Lowest energy level was nearest the nucleus • Bohr’s concept or model is known as the “satellite” or “solar system” model of the atom . Diagram of Bohr’s Atom • Solar system model Modern Concept of the Atom – 1920’s to present • The modern theory states that the electrons have wave-like properties as they travel around the central nucleus. The + charged nucleus is surrounded by electrons with definite energy levels (called orbitals). • The paths of the e- are described in terms of the probability of being found in certain regions. The e- do not follow a prescribed path. • The modern concept of the atom is known as the “wave-mechanical” model of the atom. (Much more on this later!!!!) Diagram of the Modern Atom “wave mechanical” model Evolution of the atomic model We now have a workable model of the atom: It is composed of three particles, proton, neutron and electron. Name Symbol Mass Charge Location Proton p+ 1.67x10-27 kg + nucleus Neutron n0 1.67x10-27 kg None Nucleus Electron e- 9.1x10-31 kg - Energy levels • The mass of a neutron is actually very slightly more than that of a proton however, in chemistry we generally consider them to be the same. • We use a convenient unit to express this mass, the amu (atomic mass unit). • An amu is defined as 1/12 the mass of a carbon-12 atom. • We generally consider the masses of both p+ and n0 to be 1 amu. • The mass of an e- is so much less than the other particles that we considered it to be zero in calculating the mass of an atom. So the mass of an atom, in amu’s, is simply the number of protons plus the number of neutrons. This is sometimes called “mass number” atomic mass = #p+ + #n0 The atomic number for an atom is simply its number of protons. atomic number = #p+ The total charge on an atom is determined by the number of p+ and e-. Since these particles have charges of equal magnitude and opposite sign, their charges cancel. When an atom has the same number of p+ and e- , the total charge must be zero – a neutral atom. charge = #p+ - #e- • An ION is an atom with different numbers of p+ and e-. • More p+ - positive charge – CATION • More e- - negative charge – ANION Note: The only way to form an ion is by gaining or losing electrons.!!!!!! Java – build an atom A,Z,X Method (atomic number, atomic weight, etc.) MASS NUMBER (or atomic weight) – the Total number of protons and neutrons In the atom. SYMBOL ATOMIC NUMBER – the number of protons (and also = to the number of electrons if the atom is neutral!) What if it is an ION ? • An ion is an atom with a net positive or a net negative charge. +1 A positive ion means that the element has lost electrons (in this case 1 electron). Na now has one more proton than electrons. A negative ion means that element has gained electrons (it now has more electrons than protons). Isotopes • isotopes are atoms with the same number of protons but different numbers of neutrons. • Examples of the isotopes of Hydrogen: • protium deuterium tritium • • p+ = no = p+ = no = p+ = no = Write the symbol of the element which has 20 protons, 18 electrons and a mass number of 40, and then write the formula of a different isotope of this element. IONS • ION: an atom with a positive or negative charge. • CATION: a positive ion Na+1 •ANION: Cl-1 Cu+2 Al+3 a negative ion C2H3O2-1 BO3-3 SO4-2 Introduction to the fabulous PERIODIC TABLE!!! • Identify the following sections on a blank periodic table: – Metals nonmetals – Hydrogen alkali metals – Halogens metalloids – Al family Carbon family – Oxygen family – Rare Earth elements – Lanthanide Series noble gases alkali earth metal transition met. Nitrogen family Groups IA-VIIIA Actinide series Four Classifications of Elements 1. METALS a. shiny luster b. good conductors; poor insulators c. malleable – can be hammered into thin sheets d. ductile – can be drawn into a thin wire e. All are solids except Ga and Hg Properties of Metal (con’d) f. many metals have 1 – 3 electrons in the outer shell g. all metals lose electrons during chemical change. Cations! 2. Nonmetals a. very brittle but pretty colors b. poor conductors; good insulators c. nonmetals are solids, liquids and gases d. all nonmetals have 5 – 7 e- in the outer shell e. nonmetals gain e- in chemical reactions. Anions 3. Metalloids • Also called semimetals or semiconductors • Metalloids have properties of both metals and nonmetals. 4. Noble or Inert Gases a. all are gases (no way!) b. the noble gases generally do not form compounds c. they are unreactive – they have 8 electrons in the outer shell (except Helium which only has and only needs 2 electrons) SPECIFIC TYPES OF METALS 1. ALKALI METALS – very active metals that form ions with a +1 charge. Group IA.(1) 2. ALKALI EARTH METALS – reactive (but not as much as Group IA) metals that form +2 ions. These are the Group IIA (2). These are often called the Fireworks Metals!! 1. HALOGENS – Group VIIA (17). The word Halogen means “salt forming”; all halogens are very reactive and react with metals to form salts. Gain 1 e- to form -1 ions. 1. NOBLE OR INERT GASES – Group VIIIA (18); nonreactive gases/elements Chemical reactions are represented by both words and symbols •WORD: Zinc and sulfur yields zinc sulfide • EQUATION Chemical materials and their reactions are usually designated by formulas and equations. This is the language of chemistry Formulas …learning the language • CHEMICAL FORMULA: The symbol for the elements are used to indicate the types of atoms present and the subscripts are used to show the relative numbers of atoms. • Example of a chemical formula: CO2 STRUCTURAL FORMULAS • A formula showing the individual bonds (using lines to show the bonds). • H2O CH4 BINARY IONIC COMPOUNDS (FORMULAS) a. Binary compounds are composed of 2 elements b.The components of a binary ionic compound are a monoatomic cation and a monoatomic anion. (what does monoatomic mean?) c. Binary compounds end in the suffix –ide d.Ionic compounds must be electrically neutral. Writing Formulas e. In writing a formula, we must exactly balance the positive charge of the cation with the negative charge of the anion (the net charge should be 0). f. Methods for determining subscripts: visualization LCM Criss-cross-reduce Writing Binary Ionic Formulas Ex: Potassium chloride • Calcium bromide • Iron III oxide • Calcium sulfide Rules for writing ionic formulas 1.Write the symbol for the + ion first. Do not write the charges in the formula. 2.Use subscripts to show the number of each ion required to give a net charge of zero. 3.Subscripts give the smallest ratio between the ions. 4.Use parentheses only with polyatomic ions to show more than one of them in the formula. 5.Do not write “1” as a subscript. No subscript means 1. a. Binary ionic compounds are named by writing the name of the cation followed by the anion (ending in –ide). b. When the cation has more than one possible ionic charge, it is important to use the Roman numeral.. c. DON’T overkill the use of the Roman numerals. The representative (Group A) metals only have one charge so a Roman numeral is NOT needed. Go by your ion sheet. Naming Binary Ionic Compounds • AlI3 • FeO • Cu2S • CaSe 3. Ternary Ionic Compounds a. Ternary ionic compounds contains atoms of three or more different elements. b. Ternary ionic compounds usually contain one or more polyatomic ions. c. First, write down the symbol of the ions d. Second, use the charges to determine subscripts e. Third, use parenthesis whenever a polyatomic ion needs to be taken 2, 3 or 4 times. Examples of Ternary Ionic Compounds Calcium nitrate Potassium sulfate Magnesium hydroxide Ammonium sulfide Examples…ternary ionic • Aluminum bicarbonate • Chromium III benzoate • LiCN • Sr(H2PO4)2 More examples…ternary • NH4C2H3 O2 • Fe(ClO3)3 4. Binary Molecular Compounds a. Composed of two or more nonmetallic elements b. Most of the elements that form binary molecular compounds are not charged atoms (not ions). c. When two nonmetallic elements combine, they often do so in more than one way. CO CO2 CO3 CO4 Binary Molecular Compounds d. Binary molecular compounds end in –ide. e. Greek prefixes are used to show how many atoms of each element are present in each molecule. f. Note that the vowel at the end of the prefix “mono” is dropped when the name of the element begins with a vowel; monoxide not monooxide! g. The prefix mono is omitted if there is just a single atom of the first element in the name. Greek Prefixes GREEK PREFIX NUMBER 1 2 3 4 5 6 7 8 9 10 A simple method for naming binary covalents! Number Name Number Name-ide Never start with mono. Always end with –ide. Examples of Binary Molecular… • N2O • PCl3 • SF6 • N4O7 Binary Molecular Compounds • Dinitrogen tetrahydride • Diphosphorus trioxide • Carbon tetrachloride • Nonapotassium monophosphide Types of Combinations of Atoms 1. FREE ELEMENT – one type of atom; no charge example: 2. MONOATOMIC ION – one type of atom; charged example: 3. COMPOUNDS/MOLECULES: more than one type of atoms; no charge example: 4. POLYATOMIC IONS: more than one type of atom but with a charge. example: Chem-Nerds of the world unite and identify these: • H2 Mg • SO4 -2 Mg+2 • CaCl2 • Cu+1 • Acid names and formulas…good idea to know these!! Names of Acids That Do Not Contain Oxygen Acid Name HF hydrofluoric acid HCl hydrochloric acid HBr hydrobromic acid HI hydroiodic acid HCN hydrocyanic acid H2S hydrosulfuric acid Names of Some Oxygen-Containing Acids Acid Name HNO3 nitric acid HNO2 nitrous acid H2SO4 sulfuric acid H2SO3 sulfurous acid H3PO4 phosphoric acid HC2H3O2 acetic acid