Unit 3 Language of Chemistry Part 1. Zumdahl: Chapter 4 Holt: Chapter 3. ATOMS: The Building Blocks of Matter. Objectives Law of conservation of mass Law of definite proportions Law of multiple proportions Dalton’s Atomic Theory How Dalton’s Atomic Theory relates to 1, 2, & 3.
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ATOMS: The Building Blocks of Matter Objectives Law of conservation of mass Law of definite proportions Law of multiple proportions Dalton’s Atomic Theory How Dalton’s Atomic Theory relates to 1, 2, & 3
Atomic Theory Foundations Law of Conservation of Mass – mass is neither created or destroyed during a chemical or physical change Law of Definite Proportions – a compound contains the same proportions by mass regardless of the size of the sample Example: NaCl – always 39.34% Na & 60.66% Cl
Atomic Theory Foundations Law of Multiple Proportions – if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers Example CO2 and CO : ratio of oxygen is always 2:1
JJ Thomson Showed atoms could emit negative particles “plum pudding” model Electrons were embedded in a positively charged spherical cloud
Rutherford • Shoots alpha particles (Helium atoms) at gold foil • Expected to pass right through • Particles are deflected • Leads to idea of a dense positively charged center with e- orbiting around it
Dalton’s Atomic theory • All matter is composed of atoms • Atoms of an element have the same size, mass and properties; atoms of a different element have different sizes, masses and properties • Atoms cannot be divided, created or destroyed • Atoms of different elements combine in simple whole number ratios • Chemical reactions combine, separate, or rearrange atoms
Modern Atomic Theory • Atoms can be divided • Atoms of the same element can have different masses • All else remains the same
Structure of the Atom Objectives Discovery of the Electron Rutherford’s Experiments Protons, Neutrons, Electrons
Atomic Structure • Electron = no mass; negative charge • Proton mass = hydrogen atom; positive • Neutron mass = hydrogen atom; no charge • Dalton’s Model • JJ Thompson’s Plum Pudding Model
The Electron • Mass of 9.109 x 10-31 kg • Negative charge
The Proton • Mass = 1.673 x 10-27 kg • Positive charge
Isotopes II.A.2(c) – compare the characteristics of isotopes of the same element IV.B.2(d) – calculate the weighted average atomic mass of an element from isotopic abundance, given the mass of each contributor
Isotopes Def: atoms of the same element that have different masses Example: hydrogen protium – 1 proton in nucleus deuterium – 1 proton; 1 neutron tritium – 1 proton; 2 neutrons *Nuclide – general term for any isotope
Average Mass Number Def: the weighted average of the atomic masses of the naturally occurring isotopes of an element Like calculating a “weighted” grade (decimal % of each isotope x mass of that isotope)
The Periodic Table IV.B.2(c) – use the periodic table to determine the atomic number, atomic mass, mass number, and number of protons, electrons and neutrons in isotopes of elements
Atomic Number • Atomic Number • # of protons in nucleus • Element Symbol • Element Name • Atomic Weight • Electron Configuration 3 Li Lithium 6.941 [He]2s1
Calculating Protons, Electrons & Neutrons IV.B.2(c) – use the periodic table to determine the atomic number, atomic mass, mass number, and number of protons, electrons and neutrons in isotopes of elements
Mass Number Def: the number of protons and neutrons in the nucleus of an isotope mass # - atomic # = # of neutrons Example – oxygen Mass # (16) – atomic # (8) = # of neutrons (8)
Formulas II.A.2(a) – use the IUPAC symbols of the most commonly referenced elements III.A.1(a) – distinguish between chemical symbols, empirical formulas, molecular formulas and structural formulas III.A.1(b) – interpret the information conveyed by chemical formulas for numbers of atoms of each element represented III.A.(d) – provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids
Zumdahl – Ch 5 Holt – Ch 7 III.A.1(d) - provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids NOMENCLATURE
Oxidation Numbers • Pure element have an oxidation number of zero • Fluorine has an oxidation of -1 in all compounds • Oxygen has an oxidation number of -2 except when in a peroxide when its oxidation number is -1 • Hydrogen has an oxidation number of +1 except when bonded to a metal • The algebraic sum of the oxidation numbers of all atoms in a neutral compound is zero • The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion • These rules apply to covalently bonded atoms
Oxidation Numbers Cont’d • Group 1 = +1 • Group 2 = +2 • Transition Metals = varies • Group 15 = varies (-3 usually) • Group 16 = varies (-2 usually) • Group 17 = varies (-1 usually)
General Rules for Naming Chemical Compounds • Cation (pronounced cat-ion) • is named first • has a positive charge/oxidation number • name does NOT change • Anion (pronounced an-ion) • is named second • has a negative charge/oxidation number • ending is “-ide”
Binary Ionic CompoundsType I • Metal is from Group 1 or 2 Examples: NaCl –sodiumchloride KF –potassiumfluoride CaBr2–calciumbromide Li2O –lithiumoxide
Binary Ionic Compounds Type II • Metal is a transition metal • Many metals form more than one type of positive ion • Iron: Fe2+ or Fe3+ • Copper: Cu1+ or Cu2+ • Lead: Pb2+ or Pb4+ • Tin: Sn2+ or Sn4+ • Roman numeral indicates charge on the metal ion (NOT the number of atoms) • Look at 2nd element to determine charge on transition metal
Examples of Type II EX: FeCl2 –iron(II) chloride FeCl3–iron(III) chloride CuO –copper(II) oxide Cu2O –copper(I) oxide SnF4–tin(IV) fluoride SnO2–tin(IV) oxide SnO –tin(II) oxide **NOTE: sum of oxidation numbers = 0
Exceptions to the Rule Aluminum – always Al3+ Cadmium – always Cd2+ Zinc – always Zn2+ Silver – always Ag+ No Roman numerals needed
Binary Covalent Compounds orBinary Molecular Compounds • Formed between two non-metals • first element name doesn’t change • Second element ends with “-ide” • Don’t use “mono-” on first name • Use prefixes to indicate number of atoms of element
