The Parts: Atoms, Molecules

The Parts: Atoms, Molecules, Ions, Isotopes & Compounds The Atomic Theory of Matter I. Each element is composed of tiny, indestructible particles called atoms.* II. All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements.** III. Atoms combine in simple, whole-number ratios to form compounds. IV. Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms change only the way that they are bound together with other atoms. * At the time Dalton was putting together the theory there was no verifiable evidence that the atom was composed of subatomic particles (i.e., proton, neutron, electrons). ** The discovery of isotopes disproved that elemental atoms had the same mass. Atomic Theory of Matter: Dalton • Proposed by John Dalton (1776-1844), an English chemist. • It summarizes the experimental evidence of the times, which also supported the early philosophical ideas of Leucippus and Democritus about matter. • It explains the laws governing the behavior of matter at the microscopic and macroscopic level. • The Laws: • Law of conservation of mass • Matter is neither created or destroyed; it is constant in the universe. • Law of definite proportions (1799; Joseph Proust) • For a given compound, regardless of their source, have the same constituent element proportions • Law of multiple proportions (1804; John Dalton) • When one element combines with a fixed mass of another element, their combination is a fixed ratio of whole numbers. Atomic Models – Early Models • Greek (Philosophical Model) • The ultimate piece of matter is the atom. • Dalton (Solid Bullet/Billiard Ball Model) • Atomic Theory of Matter • Elements are composed of discreet particles of matter called the atom. • Atoms are tiny, indivisible, indestructible particles and that each one had a certain mass, size, and chemical behavior. • Thomson (Raisin/Plum Pudding Model) • The atom is composed of a positive cloud of matter in which electrons are embedded. • Explains the positive (+), negative (-) charged behavior of matter Dalton’s Model Reference credit: 2011modelsa.wikispaces.com Thomson’s Model Reference credit: www.chem1.com The Atom’s Subatomic Particles • All atoms are composed of the same subatomic particles: • Protons: positively charged (+) • Neutrons: uncharged (0) • Electrons: negatively charged (-) • The charge of the proton and the charge of the electron are equal in magnitude, but opposite in sign. The neutron has no charge. • Protons and neutrons have nearly identical masses. • The mass of the proton is 1.67262 × 10–27 kg. • The mass of the neutron is 1.67493 × 10–27 kg. • The mass of the electron is 9.1 × 10–31 kg. • The nucleus of an atom is dense & contains the mass of the atom. • Negatively charged electrons surround the nucleus. • The positively charged subatomic particle, PROTON is in the nucleus. Reference credit: imgarcade.com Atom Structure: Discoveries Electron: • J.J. Thomson (1890’s) experiments with cathode rays lead to the discovery of the electron as a negatively charged particle of matter. • Robert A. Millikan’s oil drop experiments (1909) determined the charge to mass ratio of electron. Nucleus and the Proton: • Ernest Rutherford’s Gold Foil Experiment (1909) led to the discovery of the atom’s nucleus & charge (positive). • The Experiment: • Aimed alpha particles (i.e., Helium nucleus which are positively charged) at an ultra thin piece of gold (Au) foil. • The Results: • Most of the particles did pass directly through the foil, but some particles were deflected: approximately 1 in 20,000 bounced back.   1kg  = 9.107 10 −31 kg Massof electron = 1.602 10 −19 C 11  1.759 10 C  The Atom as an Ion • Atom are neutrally charged pieces of matter when the number of protons EQUAL the number of electrons. • Example: Carbon has 6 protons (i.e., six positive charges) so it would have 6 electrons (i.e. six negative charges). • Atoms can be charged pieces of matter when they have a different number of electrons than protons. Charged atoms are called IONS. • If more protons than electrons – CATIONS. • Al3+ (Aluminum III ion) has thirteen protons but only ten electrons • If less protons than electrons – ANIONS. • O2- (oxide ion) has eight protons and ten electrons. Problem: 1. Indicate whether the following ions are cations or anions. 2. Determine the number of protons and electrons are in each of the following ions? a. F– b. Mg2+ c. N3– d. W6+ e. V3+ Atomic Structure Discoveries • Elements are distinguished by their atomic number (Z) (Henry Moseley, 1913). • Elemental atoms are identified by the number of protons; their atomic number (Z). Isotopes of Hydrogen • ISOTOPES: • Elemental atoms can have different masses due to the number of neutrons in the atom's nucleus (Frederick Soddy, 1921). • NEUTRON: • Un charged (i.e. neutral) subatomic particle having relatively the same mass a proton and located in the nucleus (James Chadwick, 1932). Reference credit:www.ck12.org Isotopes & Isotopic Abundances • 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. Symbol Atomic Number Number of Protons Number of Neutrons Atomic Mass Unit (amu) % Natural Abundance H-1 1 1 0 1.0078 99.988 1 1 1 2.0141 0.0110 1 1 2 3.01605 trace protium H-2 deuterium H-3 tritium Mass Spectroscopy: Experimentally Determining the % Natural Abundance of Isotopes • Mass Spectrometry is an instrumental technique that separates particles according to their masses. • The percent abundance and masses of element isotopes are determined using a mass spectrophotometer. Reference credit: OpenStax2017 Isotopes: Elemental Atoms that differ in Mass • The identity of an element is by it atomic number (Z) or the number of protons. They however do no necessarily have the same number of neutrons. • The sum of the number of neutrons and protons in an atom is its mass number and is represented by the symbol A. A = # protons (p+) + # neutrons (n) • Isotopic symbol Problem: How many protons, neutrons, and electrons are each of the following isotopes? a. S-32 b. 19F c. 76 333As d. 238 2+ 92U Problem: Assuming the element symbols are correct, which of the following isotopic symbols are correct? Explain why. a. 6 14C b. O-17 c. 2H d. 147N e. S-6 Atomic Mass (amu) of Elements • The atomic mass of an element is the averaged mass of the element’s isotopes. • It is a weighted value based on the element’s natural abundance of each isotope. 𝒂𝒕𝒐𝒎𝒊𝒄 𝒎𝒂𝒔𝒔 𝒐𝒇 𝒆𝒍𝒆𝒎𝒆𝒏𝒕 = ෍[(𝒇𝒓𝒂𝒄𝒕𝒊𝒐𝒏𝒂𝒍 𝒂𝒃𝒖𝒏𝒅𝒂𝒏𝒄𝒆 𝒐𝒇 𝒊𝒔𝒐𝒕𝒐𝒑𝒆 𝒏) 𝒙 (𝒎𝒂𝒔𝒔 𝒐𝒇 𝒊𝒔𝒐𝒕𝒐𝒑𝒆 𝒏)] 𝒏=𝒊 • Examples: • Lithium (Li) has two isotopes: 6Li = 7.50% abundant and 7Li = 92.5% The atomic mass for lithium is 6.93 amu [(6.00 amu x 0.0750) + (7.00 amu x 0.925)] = 6.93 amu • Silicon (Si) has three isotopes: 28Si = 92.23%, 29Si = 4.67% and 30Si = 3.10% The atomic mass for silicon is 28.10 amu [(28.00 amu x 0.9223) + (29.00 amu x 0.0467) + (30.00 amu x 0.0310)] = 28.10 amu Problem: The element silver (Ag) has an atomic mass of 107.868 amu. It has two isotopes - Ag-109 (108.905 amu) with an abundance of 48.1600% and Ag-107. Determine the amu for Ag-107. The Modern Atom Models: Bohr & the Quantum • Bohr Model (1913) • Proposed an atomic model that tried to resolve the atomic paradox of the Rutherford model (solar system). • Why classical physics could not explain the behavior of the electron orbiting the nucleus. • Schrodinger Quantum Atom (1927) • Views the electron as a particle of light, specifically focusing on the electron’s wave behaviors to explain the chemical and physical properties of matter. • Einstein: photoelectric effect & wave duality • Planck: E= hu or E = hc/l • De Broglie: l = h /(mass x velocity) Reference credit: socratic.org Duality Relationship of Matter • At the microscopic level (i.e., subatomic particles such as the electron), matter and energy are a singular entity that possess characteristics & behaviors of each other. • Matter has the properties and behaviors of energy in the form of light (i.e., wavelength & frequency). • Energy has particulate characteristic of matter (i.e., has mass & occupies space). • The energy of a photon (particle of light) was directly proportional to its frequency, but inversely proportional to its wavelength (photoelectric effect). • Einstein proposed that the light energy was delivered to the atoms in packets, called quanta (photons). • Planck’s experiments with black body radiation discovered that energy is absorbed and released in discrete packets of energy (hu); quanta. • De Broglie experiments related wavelength (energy) to matter. • The theory that describes the duality concept, specifically the electron, is Quantum Mechanics Theory. • Concept of Wave Duality as noted by Heisenberg: “ Light and matter are both single entities, and the apparent duality arises in the limitations of our language.” Images of Duality Reference credit: sciencekinds.co.nz Reference credit: molillusions.com Reference credit:express.co.uk Reference credit: amsi.org.au A Look at E = hu • Max Planck work with blackbody radiation of objects lead to the discovery that absorb and emitted radiation energy (light) is proportional to its frequency (n) in distinct units called QUANTA (hn). • Mathematically this relationship can be described as: E = hn • Where h is Planck’s constant having a value of 6.6262 x 10-34 J·s Reference credit: chemwiki.ucdavis.edu Reference credit: becuo.com PROBLEM: Lithium atoms when radiated with energy of 2.961 x 10-19 J emit photons of light. 1. Determine the frequency & wavelength (nm) of a photon of light. 2. Determine the energy associated with 1 mole of photons. 3. What is the color of light (visible range) observed? c = 3.00 x 108 m/s h = 6.63 x 10-34 J s Schrodinger: Quantum Atom • Views the electron as a particle of light, specifically focusing on the electron’s wave behaviors to explain the chemical and physical properties of matter. – The electron no longer is viewed as a small piece of matter orbiting around the nucleus, but as a cloud of probability that is spreads-out over the orbit area of the atom. • Using Mathematics (Hy = EY )he viewed electrons as a circular threedimensional stationary wave described as the wave function: • The wave function Y: • Describes an area that the electron can be found in relationship to the nucleus. • It does not denote the EXACT location of the electron in an orbital but gives the probability of knowing the location of the electron. • When Y is solved, four solution values results referred to as Quantum numbers: n, l, ml & ms Reference credit: OpenStax 2017 Reference credit: universe-review.ca Heisenberg Uncertainty Principle Quantum Atom Connection • The Uncertainty Principle explains why the wave-duality nature of the electron makes it “fundamentally impossible to simultaneously determine with exactness both the momentum (Dmv) and placement (position) of an electron“ as it moves around the nucleus of the atom. • Makes it difficult to know electron’s exact position & velocity simultaneously as it travels/orbits around the nucleus. • Heisenberg: Dx (position) Dmv (momentum) > h/4p Reference credit: suppiya.worldpress.com Quantum Atom & Quantum Numbers • • Picture the quantum atom as having its nucleus as the center apex on a three-dimensional Cartesian (x, y , z axis) probability plot. Cartesian Coordinate Plot The four Quantum numbers (n, l, ml & ms ) describing the electron’s position relative to the atom’s nucleus on this plot • Each value describes an allowed energy state for an electron. • According to the Pauli Exclusion Principle no two electrons may have the same set of quantum numbers. • Each orbital type (l & ml) can only have TWO electrons with opposite spins. Reference credit: flickr.com Quantum Numbers: The electron’s address Principle quantum number (n) is the energy level (shell). • Its numerical values are n = 1, 2, 3, etc. Angular momentum (l) describes the orbital shape (subshell). • The numerical values of l are determined by (n-1). Orientation/magnetic quantum number (ml) indicates the axis location of the orbital. • Its numerical values are determined by +/- value of l » ml explains why there is only 1 type of s orbital, but 3 types of p orbitals, 5 types of d orbitals & 7 types of f orbitals. Spin quantum number (ms) describes the spin of the electron in an orbital. • Electrons can spin with a magnetic field or against a magnetic field. » The numerical values for spin quantum number are either + ½ or – ½ Reference: chem.libretexts Problem: 1. Determine the number and types of subshells when n = 3. 2. Calculate the total number of electrons that occupy when n = 2. 3. Identify these orbitals: a. b. Quick Overview of the Periodic Table • Periodic Table: • Several attempts to organized elements based on their physical and chemical properties. • Dimitri Mendeleev in Russia (1869) and Lothar Meyer in Germany (1870) independently recognized that there was a periodic relationship among the properties of the elements known at that time. • They argued the elements periodic behavior was a function of their mass. • Modern Table: • An organizational chart of the elements as arranged by their atomic number. Mendeleev’s elemental table Reference credit: msnucleus.org Modern Table & Periodic Law • Periodic Law states that the properties of the elements are periodic functions of their atomic numbers. • Today’s periodic table the elements are arranged in increasing order of their atomic numbers with groups atoms with similar properties in the same vertical column. • Periods or series are the horizontal rows. • Groups and families are the vertical columns. Reference credit: becuo.com Periodic Table: Groupings Reference credit: pinstake.com Chemical Symbol of Elements • Elements are arranged according to atomic number in the periodic table. • Each element in the table is represented by a chemical symbol, • Symbols have one or two letters for named elements • Three-letter symbols are used to describe elements predicted or recently discovered elements whose name has not yet be certified by IUPAC. • See atomic numbers greater than 112. • The chemical symbol: • is an abbreviation of the elements name – English or Latin version. • For example: • The chemical symbol for carbon is C. • The chemical symbol for iron is Fe (ferrum) • Only the first letter of a chemical symbol is capitalized. Problem: Use the periodic table to identify each of the following. a. the fifth element in the fourth period of transition metals. b. the element of the fourth period that is also in group 6. c. the alkaline earth in period six. d. the first element of group/column 8. e. the third period halogen. f. the first alkali metal. g. the second inner transition metal in period seven. What is a Chemical Formula? • A chemical formula represents the composition of the substance. • Types of chemical formulas • Condensed: indicates the type and number of each element in the molecule. • Structural: gives the same information as a molecular formula but indicates how the atoms are connected. • Empirical: is the simplest ratio of atom-to-atom present in the compound. • In a chemical formula, the chemical symbols indicate the type of atoms & the subscripts after the symbol indicate the number of each atom type. • Subscripts are only used if there is more than ONE atom of a given type present in the molecule. • Parenthesis are used to indicate a group of atoms – polyatomic. • Examples: • C is the element formula/symbol for carbon. • H2O is the molecular formula for water. • NaCl is the chemical formula for sodium chloride. • Fe(NO3)3 is the chemical formula for iron (III) nitrate. Structural formula for glucose Reference credit: chemicalformula.org Empirical Formula: Atom to Atom Ratio • An empirical formula is the ratio of atom to atom in a molecule. Structural formula for glucose • Example: Glucose • Molecular formula for glucose is C6H12O6. • A molecular formula represents the actual number of atoms in the molecule. • The empirical formula for glucose is (CH2O)n where n = 6 • The subscript n represents the number of (CH2O) units in a glucose molecule. • An empirical formula represents the simplest whole-number ratio of the number of atoms (or ions) in the compound. Reference credit: chemicalformula.org • Using percent by weight of each element present in the compound an empirical formula can be derived. Ionic, Molecular & Metallic Compounds • Ionic compounds are molecules composed of a metal & a nonmetal. • Bonding type is ionic (i.e., cation + anion or transfer of electrons from one element – the metal to another element - nonmetal). • Examples: • CaCl2 (calcium chloride) & CuBr (copper (I) bromide) • Molecular compounds are molecules composed of two or more nonmetals. • Bonding type is covalent (i.e., electrons are shared between atoms to make a bond.) • Examples: • CO2 (carbon dioxide) & NH3 (ammonia) • Metallic compounds (i.e., alloys) are molecules composed of metal elements. • Bonding type is metallic (i.e., sharing of electrons by metals throughout the compound). • Examples: • Brass, Steel & Gold (Au) Overview of Molecular Compounds • Molecular compounds are formed between two or more nonmetal elements. • Exist as discrete, neutral molecules. • Bonding is: • the result of the sharing of electrons between the two nonmetals. • Covalent in nature. • The sharing of the electrons may be relatively equal or may be unequal (polar covalent). • Examples: CO P2O5 Cl2 carbon monoxide diphosphorous pentaoxide chlorine • General Properties: • Often exist as gases. • Have low-boiling liquids and low-melting solids. Overview on Ionic Compounds • Ionic compounds are formed when a metal & nonmetal combine. • Bonding is: • the result of the transfer of electrons from metal to nonmetal. • ionic (cation + anion) attraction. • Examples: Mg2+ + 2 NO3- → Mg(NO3)2 magnesium nitrate 3 Fe2+ + 2 PO43- → Fe3(PO4)2 iron(II) phosphate • General Properties: • typically, solids have high melting and boiling points. • Non-conductive in solid form, but if soluble in water conductive • Conductive in molten form • The formula of an ionic compound must have a ratio of ions such that the numbers of positive and negative charges are equal. Predicting Formulas for Ionic Compounds • Ionic Compounds are composed: • metal + nonmetal elements • metal + nonmetal polyatomic ions • cations + anions • Molecular formulas of ionic compounds are “neutral” or net charge of zero. • The formula of an ionic compound must have a ratio of ions such that the numbers of positive and negative charges are equal. • Predict the molecular formulas • Example #1: What is the molecular formula for Mn+3 & O-2 (oxide ion) • Mn+3 has an oxidation state of +3 & oxide ion has an oxidation state of –2 • Criss-cross the absolute value of the respective oxidations states to derive the molecular formula • Mn2O3 Mn: +3 x 2 = +6 O: -2x3=-6 0 Problem #1: Predict the molecular formula for V+4 & the sulfite (SO3-2) ion. Problem #2: Predict the molecular formula for Sr+2 & nitride ion (N- 3). Naming Compounds Naming compounds nonmetal + nonmetal metal + nonmetal (ionic) (covalent) metal metal forms only one cation: Use the name of the element nonmetal Metal has more than one cation form: Use the name of the element followed by () with charge given as a Roman Numeral Single anion: Element's stem name + suffix "ide" added Polyatomic ion: Use the name of the polyatimc ion 1st nonmetal 2nd nonmetal Element name with prefix indicatin the number of that element. Element's stem name +suffix "ide" added *Mono is not used; start with di Nomenclature: Ionic Compounds Metal with Nonmetal elements form neutral charged compounds. • Alkali & Alkaline Earth Metals • Metal name first followed by nonmetal • Non metal ending is exchanged with “IDE.” • Example: • MgCl2 magnesium chloride • KNO3 potassium nitrate • Alkali (+1) & Alkaline Earth (+2) metals oxidation states are known. • That is why their charge is NOT indicated in the formula name! Nomenclature: Ionic Compounds Metal with Nonmetal elements form neutral compounds. Metalloid with Nonmetal elements form neutral compounds. • Transitions Metals, Inner Transition Metals & p-block (main group) Metals • Metal name first followed by nonmetal. • Metal oxidation state is indicated by (Roman Numeral) • Nonmetal ending is exchanged with “IDE.” • Example: • MnBr4 manganese (IV) bromide • Fe2O3 iron (III) oxide • SnF2 tin (II) fluoride • SiO2 silicon (IV) oxide Monatomic, Polyatomic Ion & Oxyanions: What are they? • Monatomic ions are ions formed from only one atom. • Examples: Cl- is chloride ion & Si4+ is the silicon (IV) ion • Polyatomic ions are electrically charged molecules. • They are a group of bonded atoms having an overall charge. • Examples: SO42- is the sulfate ion & NH4+ is the ammonium ion • Oxyanions are polyatomic ions that contain a nonmetal along with one or more oxygen atoms. • Examples:CO32- is the carbonate ion & NO2- is the nitrite ion Problem: Correct these names. a. NaCl sodium(I) chloride b. MnO manganese(I) oxide c. Ca(ClO3)2 calcium dichlorate d. NH4Cl nitrogen trihydride chloride e. ZnBr2 zinc(II) bromite Problem: Name these compounds Write formulas for these compounds. a. NaNO3 a. Iron (II) nitrate b. Cu2S b. Cobalt(III) hypochlorite c. (NH4)2SO4 c. Magnesium phosphate Nomenclature: Molecular Compounds Nonmetal with Nonmetal or Metalloid with Nonmetal form neutral compounds • Name the most “metallic” of the nonmetals first. • Change the ending of the nonmetal to “IDE.” • Indicate the number of atoms of the nonmetal by Latin prefixes • Mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca • Examples: • S2Cl4 disulfur tetrachloride • SiO2 silicon Compounds with Common Names Water H2O Ammonia NH3 Problem: Name the compound whose molecules contain the following elements. a. one atom of silicon and four atoms of chlorine b. one atom of sulfur and three atoms of oxygen c. one atom of sulfur and two atoms of bromine d. two atoms of phosphorus and three atoms of sulfur 47 Nomenclature: Naming Binary acids • Several binary molecular compounds when dissolved in water form an acid. • Examples: • Name of Gas Name of Acid • HF(g), hydrogen fluoride HF(aq), hydrofluoric acid • HCl(g), hydrogen chloride HCl(aq), hydrochloric acid • HBr(g), hydrogen bromide HBr(aq), hydrobromic acid • HI(g), hydrogen iodide HI(aq), hydroiodic acid • H2S(g), hydrogen sulfide H2S(aq), hydrosulfuric acid • When naming these compounds, the: • word hydrogen is changed to the prefix hydro. • other nonmetallic element name is modified by adding the suffix –ic • word “acid” is added as a second word. Summary of Acid Nomenclature Anion Suffix -ide -ate -ite Acid Name Hydro___ic acid ___ic acid ___ous acid Examples – Cl = chloride HCl = hydrochloric acid – ClO4 = perchlorate HClO4 = perchloric acid – ClO2 = chlorite HClO2 = chlorous acid – CN = cyanide HCN = hydrocyanic acid – ClO3 = chlorate HClO3 = chloric acid – ClO = hypochlorite HClO = hypochlorous acid Nomenclature: Hydrates • Hydrates are ionic compounds containing a specific number of waters for each formula unit. • water of hydration often “driven off” by heating • in formula, attached waters follow ∙ • CoCl2∙6H2O • in name, attached waters indicated by prefix+hydrate after name of ionic compound • CoCl2∙6H2O = cobalt(II) chloride hexahydrate • CaSO4∙½H2O = calcium sulfate hemihydrate Prefix No. of Waters hemi ½ mono 1 di 2 tri 3 tetra 4 penta 5 hexa 6 hepta 7 octa 8 Problem: Write formulas for these compounds. a. anhydrous zinc sulfate b. zinc sulfate heptahydrate