Atoms, Molecules, and Ions Chapter 2 Atoms, Molecules, and Ions

Laws of Chemical Combination Established from scientific observations Used to establish other scientific theories Include the following 3 laws: Law of Conservation of Mass Law of Definite Proportions Law of Multiple Proportions

Law of Conservation of Mass Lavoisier-mid 1700s Total mass remains constant during a chemical reaction Entire mass of reactants must be accounted for in the products Also called a mass balance 11.1g H2 + 88.9g O2 = 100.0g H20

Law of Definite Proportions Joseph Proust - Late 1700s All samples of a compound have the same composition Same mass proportions of elements present (Same atomic ratios of elements present) Water 11.1% H 88.9% O (2 atoms H, 1 atom O) Also called law of constant composition

Law of Definite Proportions Copper Carbonate, CH2Cu2O5 All samples of a compound have the same composition

Dalton’s Law of Multiple Proportions John Dalton- Early 1800s When two or more different compounds of the same two elements are compared, the masses of one element that combine with a fixed mass of the second element are in the ratio of small whole numbers. Example: CO: 12g Carbon reacts with 16g Oxygen CO2: 12g Carbon reacts with 32g Oxygen CO2 has 2X as much oxygen as CO Ratio = 32:16= 2:1 2:1 is ratio of small whole numbers in Dalton’s Law

Atomic Theory of Matter All matter is composed of atoms All atoms of a given element are alike in mass, but atoms of different elements differ in mass Compounds are formed when atoms of different elements unite in fixed proportions. A chemical reaction involves atomic rearrangement No atoms are created or destroyed

The Atom and Sub-Atomic Particles Proton large, positively charged particle in nucleus Electron small, negatively charged particle in orbit around the nucleus Neutron large, neutral particle in nucleus Elements are not electrically charged Must have equal numbers of protons and electrons electron proton neutron nucleus

Atomic Symbols Atomic Number (Z) Protons in a nucleus Determines element identity Located lower left on symbol Mass Number (A) # protons + # neutrons Determines isotope identity Located upper left on symbol

Isotopes Elements with the same number of protons and electrons, but differing number of neutrons Used in chemistry for structure identification or to follow a particular molecule through a reaction Example: hydrogen and deuterium Water, H2O H has 1 proton and 0 neutrons in nucleus Most abundant Heavy water, D2O D (deuterium) has 1 proton &1 neutron in nucleus Occurs 1/6700 molecules 1 H 2

Atomic Mass Units (AMU) Weighted average of the masses of the naturally occurring isotopes of an element Mass of protons & neutrons in nucleus Mass proton or neutron: ~ 1.66 x 10-24 g /particle Particle mass independent of element Electrons mass ignored Atoms are difficult to weigh due to extremely small size Use AMU- atomic mass units (u) Internationally based on pure isotope C-12 C-12 atom contains 6 protons + 6 neutrons = 12 particles 1 u = weight of 1 particle = 1.66 x 10-24 g C-12 then has a mass of 12.000000 u

Atomic Mass Calculations Percent abundance isotope mass ÷ total mass Fractional abundance % abundance ÷ 100 Isotope contribution fractional abundance x atomic mass Atomic mass sum of all mass contributions from isotopes

Calculating the Atomic Mass for C Look up mass of isotopes C-12 12.00000u C-13 13.00335u Look up % Abundances C-12 98.89% C-13 1.11% Calculate contribution from Isotopes C-12 0.9889 x 12.00000u = 11.87 u C-13 0.0111 x 13.00335u = 0.144 u Add up contributions to determine atomic mass Contribution of C-12 + Contribution of C-13 11.87 u + 0.144 u = 12.01 u

Mendeleev’s Periodic Table 1869 Arranged the known elements in order of increasing atomic weight from left to right and from top to bottom in groups. Elements with similar properties are placed in the same column. Used table to predict properties of some undiscovered Mendeleev elements

Germanium: Prediction vs. Observation

The Modern Periodic Table

Metals Characteristics Metallic, Conductive, Malleable, Location: Left side of main table Characteristics Metallic, Conductive, Malleable, Ductile, Luster, Solids (not Hg) , Positive ionic charge

Non-metals Location: Right side of table Characteristics: Nonmetallic Brittle solids Gas, liquids or solids Negative or uncharged ions

Noble Gases Location: Last column on the right side of table Characteristics Nonreactive Gases Extremely stable Difficult to ionize

Halogens Location: Last column before noble gases Characteristics Reactive 1- charge most common Diatomic Stable ionized in water

Metalloids Location: Border between metals and nonmetals Characteristics Metal/nonmetal characteristics Solids Variable ionic charges Semi-conductors

Lanthanides and Actinides Table Location: Bottom 2 rows below table Characteristics Very reactive, Often radioactive, Unstable Solids, Difficult to measure, and Positive ions

Molecules and Chemical Formulas A group of 2 or more atoms held together with covalent bonds Chemical Formula Symbolic representation of molecular composition Three different types of chemical formula Empirical: ratio of atoms Molecular: types and numbers of atoms Structural: relationship of atoms in molecule

Chemical Formulas Acetic Acid: 2C 2O 4H Empirical Formula Shows ratio of atoms CH2O Molecular Formula Shows number and type of atoms C2H4O2 Structural Formula Shows relationship between atoms CH3COOH

Naming Binary Compounds Name has 2 words, one for each element: N2O4 1st word is for 1st element: N: Nitrogen 2nd word is stem of element name O: Oxygen change ending to “–ide” O: Oxide Use prefixes to designate # of atoms 2N: Dinitrogen 4O: Tetroxide Put it together N2O4: Dinitrogen tetroxide

Prefixes for Molecular Compounds Memorize!

Chemical Formulas for Binary Compounds Choose the first element symbol Boron Trifluoride Use the one farthest to left (metal) B Lowest in group if in same group Then write the other atom such as a nonmetal F Remember the subscripts Count # of atoms of each element 1B, 3F BF3

Ionic Compounds (salts) Atoms in a reaction may gain or lose electrons Both atoms become charged and are called ions Anion: atom gaining the electron is negatively charged Cation: atom losing the electron is positively charged The net charge is 0 Cations & anions balance out over the entire compound No distinct molecular units Positive charge of one ion attracts all nearby negative charges

Predicting Ionic Charge from the Periodic Table Metals, Metalloids and Nonmetals Goal: Get to column 8A by going to the right or left Right: Count each box as -1 until reaching 8A Left: Count each box as +1 until reaching 8A in previous row The correct charge is usually the smallest number Left Side (metals): Li1+ or Li-7 Be2+ or Be6- Right Side (nonmetals): O2-, F1-, He0, Ne0 Center (metalloids): B3+ (B5-), C4+ C4- , N5+, N3-

Predicting Ionic Charge from the Periodic Table Group B (Transition Metals) Charge is positive, but unpredictable Charge on a transition metal is designated with a Roman numeral when naming the compound Iron (III) oxide contains Fe3+

Naming Ions Add the word ION after element names Sodium Chloride: NaCl Table Salt Na Cl loses 1 e- gains 1 e- Na+ Cl- Sodium ion Chloride ion Net Charge = 0 = (+1) + (-1)

Monotomic and Polyatomic Ions Atoms can lose or gain electrons singly or as a group Monoatomic ions Lose or gain electrons singly The total charge is on a single atom ex: Na+ Polyatomic ions Lose or gain electrons as a group The total charge is spread over 2 or more atoms ex: SO4 2-

Polyatomic Ions Memorize the following polyatomic ions Ammonium NH4 + Hydronium H3O+ Phosphate PO4 3- Acetate CH3COOHydroxide OH- Nitrate NO3 - Cyanide CN- Sulfate SO4 2- Permanganate MnO4 - Chlorate ClO3 Carbonate CO3 2- Perchlorate ClO4

Acids, Bases and Salts Acid Characteristics taste sour sting the skin turn litmus paper from blue to red react with metals to produce ionic salts and hydrogen gas Base Characteristics taste bitter feel slippery on the skin turn litmus red to blue react with acids to become neutral, often form ionic salts

Arrhenius Acids and Bases Compound that ionizes in water to form a solution of H+ ions and anions Base Compound that ionizes in water to form a solution of OH- ions and cations Neutralization Reaction between Arrhenius acid & base H+ + OH- =H2O and cation + anion = salt Uses Identify acids and bases Write formulas

Naming Arrhenius Acids Binary Acids (HX) Contain hydrogen and 1 other type of atom Dissolved in water, so change hydrogen to hydro Change “–ide” ending to “–ic acid” example HCl: hydrochloric acid Ternary and Oxoacids- Binary acid that also contains oxygen Oxoacids have oxygen as one of the atoms Change “–ate” ending to “–ic acid” example Nitric Acid: HNO3

Common acids to memorize Hydrochloric Acid: HCl Sulfuric Acid: H2SO4 Nitric Acid: HNO3 Carbonic Acid: H2CO3 Phosphoric Acid: H3PO4 Perchloric Acid: HClO4

Naming Arrhenius Bases Named the same way as ionic compounds All Arrhenius bases contain OH (hydroxide) in the Name the metal first, then use word “hydroxide” Memorize: Sodium hydroxide: NaOH Potassium hydroxide: KOH Ammonium hydroxide: NH4OH

Polyatomic Ions with Oxygen Suffixes -ate Most common form Chlorate ClO3- -ite 1 less oxygen than “ate” Chlorite ClO2- Prefixes Hypo- 1 less oxygen than ite ion Hypochlorite ClO per 1 more oxygen than -ate ion. Perchlorate ClO4-

Hydrogen attached to ion Use numerical prefixes for # of hydrogen atoms mono, di, tri, etc. Then add the word “hydrogen” before ion name H2PO4- Dihydrogen phosphate ion

Hydrates A compound that is associated with a fixed number of water molecules is called a hydrate Naming Hydrates Use numerical prefixes for # of water molecules Then add the word “hydrate” A dot shows H2O association Water molecules are part of the mass Copper(II)sulfate pentahydrate, CuSO4 ∙ 5H2O