Atoms, Molecules, and Ions Chapter 2 (Ch 2_Chandrakanthan_Fall 2024 PDF)

Atoms, Molecules, and Ions Chapter 2 Outline What constitutes an atom? 2.1 The Atomic Theory of Matter 2.2 The Discovery of Atomic Structure 2.3 The Modern View of Atomic Structure (Nuclear Chemistry & Isotopes) 2.4 Atomic Weig...

Atoms, Molecules, and Ions Chapter 2 Outline What constitutes an atom? 2.1 The Atomic Theory of Matter 2.2 The Discovery of Atomic Structure 2.3 The Modern View of Atomic Structure (Nuclear Chemistry & Isotopes) 2.4 Atomic Weights How are elements classified in the Periodic Table (PT)? 2.5 The Periodic Table 2.3 The Modern View of Atomic Structure (Nuclear Chemistry & Isotopes) How are binary ionic and molecular compounds named? 2.6 Molecules and Molecular Compounds (Ionic & Covalent Bonding) 2.7 Ions and Ionic Compounds (Oxidation Numbers) 2.8 Naming Inorganic Compounds 2.9 Some Simple Organic Compounds (Skip this section) What constitutes an atom? Sections 2.1 – 2.4 The Atomic Theory of Matter If we look at a piece of copper with a powerful microscope, we cannot see atoms. The copper appears to be continuous. What evidence is there to show that matter is composed of atoms that can not be seen? STM Images Scanning Tunneling Microscopy Silicon Graphite AFM Image of Pentacene Atomic Force Microscopy The Atomic Theory of Matter 5th Century BCE Democritus and Leucippus Aristotle All matter is made up of tiny, Rejected the idea of the atom indivisible particles Four basic elements atomos (meaning “indivisible” or “uncuttable”) The Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early nineteenth century, championed by John Dalton. In 1808 John Dalton published a paper on atomic theory, summarized by the following 4 postulates: The Atomic Theory of Matter Dalton’s Four Postulates The Atomic Theory of Matter Experiments in the eighteenth and nineteenth centuries led to an organized Atomic Theory by John Dalton in the early 1800s. Dalton’s theory explains several laws of chemical composition that were known at the time. These included: Law of Constant Composition (postulate 4) Law of Conservation of Mass (postulate 3) Law of Multiple Proportions (deduced from four postulates) The Atomic Theory of Matter Law of Constant Composition: In a given compound, the relative numbers and kinds of atoms are constant (Section 1.2). Elements combine in specific proportions: Example: water For ANY water molecule, it is ALWAYS made up of 2 hydrogen atoms and 1 oxygen atom in a 2:1 ratio The Atomic Theory of Matter Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction. The total mass of materials present after a chemical reaction is the same as the total mass present before the reaction. Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) The Atomic Theory of Matter Law of Multiple Proportions: If two elements (A and B) combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers. Ratio of oxygen to hydrogen is different for two compounds: 8:1 16 : 1  Concept Check Law of Conservation of Mass Hydrogen sulfide (H2S) is composed of two elements: hydrogen and sulfur. In an experiment, 6.500 g of hydrogen sulfide is fully decomposed into its elements as seen in the reaction below. 𝟐 𝑯𝟐𝑺 (𝒈) → 𝟐 𝑯𝟐 (𝒈) + 𝑺𝟐 (𝒈) If 0.384 g of hydrogen is obtained in this experiment, how many grams of sulfur must be obtained? Express your answer to THREE places past the decimal. The Discovery of Atomic Structure In Dalton’s view, the atom was the smallest particle possible. It took about 100 years to discover subatomic particles. Many discoveries led to the fact that the atom itself was made up of smaller particles. – Electrons and cathode rays – Radioactivity – Nucleus, protons, and neutrons The Discovery of Atomic Structure Atomic Models The Discovery of Atomic Structure Electrons Atomic Nucleus Discovered by J.J. Thomson in in Center of the atom; made up of 1897; contributes negative positively-charged particles (protons) charge to the atom. and uncharged particles (neutrons). Neutrons Discovered by James Chadwick in 1932; contributes NO charge (neutral) to the atom. Protons Discovered by Ernest Rutherford in 1919; contributes positive charge. The Modern View of Atomic Structure While protons and neutrons make up most of the mass of an atom, the electrons make up the volume. For a Li atom: Angstrom (Å): 1 Å = 1 x 10-10 m e─ cloud 1” marble The Modern View of Atomic Structure Actual Relative Particle Symbol Charge Charge Mass (amu) 1 amu = 1.66054x10-24 g electron e— - 1.602 x 10-19 C -1 5.486 x 10-4 proton p+ + 1.602 x 10-19 C +1 1.0073 neutron n0 0 0 1.0087 Protons (+) and electrons (−) are the only atomic particles that have a charge. The charge of a proton is opposite in charge but equal in magnitude to that of an electron. Neutrons are neutral (hence, the name). Protons and neutrons have essentially the same mass. Electrons are much smaller. Every atom has an equal number of electrons and protons, so atoms have no net electrical charge. The number of neutrons can vary.  Concept Check Mass of an Electron If you had a bag of electrons weighing 1.0 lb, how many electrons would be in the bag? Hint: 1 lb = 453.59 g Hint: The mass of one e⎯ is 9.1093837015 x 10⎯31 kg. How are elements classified in the Periodic Table (PT)? Section 2.5 The Periodic Table Elements are arranged in order of atomic number, Z, and grouped according to similar properties. The Modern View of Atomic Structure Atomic Numbers, Mass Numbers and Isotopes Elements are represented by a one- or two-letter symbol (i.e. H or Mg). All atoms of the same element have the same number of protons, which is called the atomic number, Z. It is written as a subscript BEFORE the symbol. The mass number, A, is the total number of protons and neutrons in the nucleus of an atom. It is written as a superscript BEFORE the symbol. An atom can lose electrons (+ charge) or gain electrons (─ charge). The charge is written as a superscript AFTER the symbol. The Modern View of Atomic Structure Atomic Numbers, Mass Numbers and Isotopes Isotopes vary in mass (different number of neutrons) and abundance. The Modern View of Atomic Structure Atomic Numbers, Mass Numbers and Isotopes Often, only the mass number Examples: #p+ #no #e− (differing numbers of neutrons) and/or the charge (differing 24Na 11 13 11 number of electrons) is written for a specific isotope. boron-11 5 6 5 If you have access to a PT, this is sufficient information to determine 39Ca2+ 20 19 18 the number of electrons, protons, and neutrons. 35Cl− 17 18 18  Concept Check Isotopes Rank the following radioactive isotopes from the largest number of neutrons to the smallest. Isotope Mass # # of p+ # of n0 Rank barium-133 iodine-131 strontium-90 uranium-238 Atomic Weights The Atomic Mass Scale Atomic mass unit (amu) is exactly one-twelfth the mass of a carbon-12 atom. 1/12 mass of 12C atom = 1 amu The mass of one 12C atom = 12 amu (Exactly. This is the standard.) 1 amu = 1.66054 x 10-24 g 1 g = 6.02214 x 1023 amu What about the fact that different isotopes of an element have different numbers of neutrons, and hence, have different weights? Atomic Weights The Atomic Mass Scale Atomic mass on periodic table is the “weighted” average of abundance of an atom’s isotopes. 𝐴𝑡𝑜𝑚𝑖𝑐 𝑊𝑒𝑖𝑔ℎ𝑡 = ෍ [ 𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛𝑎𝑙 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 × 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑎𝑠𝑠 ] 𝑜𝑣𝑒𝑟 𝑎𝑙𝑙 𝑖𝑠𝑜𝑡𝑜𝑝𝑒𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 Example: For naturally occurring carbon, 12C = 12 amu; 98.89% 13C = 13.0034 amu; 1.11% Weighted Average: 0.9889 × 12 𝑎𝑚𝑢 + 0.0111 × 13.0034 𝑎𝑚𝑢 = 𝟏𝟐. 𝟎𝟏𝟏 𝒂𝒎𝒖  Concept Check Atomic Weight Naturally-occurring chlorine is 75.78% 35Cl (atomic mass = 34.969 amu) and 24.22% 37Cl (atomic mass = 36.966 amu). What is the atomic weight (in amu) of chlorine? Enter your answer to TWO places past the decimal. The Periodic Table There are 38 radioactive elements, most occurring for elements with Z ≥ 83. They either have no stable, naturally-occurring isotope, OR are entirely synthetic as all synthetic elements have no stable isotopes. Elements with Z < 83 have isotopes (stable nucleus) and most have at least one radioisotope (unstable nucleus).  The Periodic Table Three Main Categories of Elements: 1. Representative (Main Group) Elements—Groups 1A, 2A, 3A – 8A. Most abundant, show the strongest trends every 8 elements. 2. Transition Metals (d-block elements)—Groups 3B – 2B (middle of PT). 3. Lanthanide and Actinide Series Rare earth metals— (Inner Transition Metals, f-block La series, Sc, Y elements)—elements 57 – 71, 89 – 103 (follows La, Ac). The Periodic Table The rows on the periodic chart are periods. Numbered 1 – 7 Columns are groups (or families). Labeled with a number (1 – 8) and a letter (A or B) Elements in the same group have similar physical & Some PTs use Roman numerals while others chemical properties. number groups from 1 – 18 with A and B labels omitted (IUPAC convention). The Periodic Table Common Group Names The Periodic Table The elements can also be classified as metals, nonmetals, and metalloids. See detailed Periodic Table Handout on Canvas! The Periodic Table Metals: Summary of Properties Shiny 'metallic' appearance (high luster) Most are solids at room temperature (except mercury) High melting points High densities Large atomic radii Lose electrons easily ✓Low ionization energies ✓Low electronegativities Malleable Ductile Good thermal and electrical conductors The Periodic Table Nonmetals: Summary of Properties Most are gases at room temperature ✓Hard and brittle solids ✓Low or no metallic luster Low melting points Gain electrons easily ✓High ionization energies ✓High electronegativities Good insulators ✓Poor thermal conductors ✓Poor electrical conductors The Periodic Table Metalloids: Summary of Properties Elements on the step-like line are metalloids (except Al, Po, and At). Their properties are sometimes like metals and sometimes like nonmetals. Often make good semiconductors.  Concept Check Periodic Table Which statement does NOT describe the relationship of O to S? A. Both are metals. B. They are in the same group. C. Both are likely to form ions with a (-2) charge. D. Both are main group elements. E. They are chalcogens. How are binary ionic and molecular compounds named? Sections 2.6 – 2.8 Molecules and Molecular Compounds Molecules and Chemical Formulas Molecular compounds are composed of atoms chemically bonded together. The elements in molecular compounds are almost always nonmetals. The bonds between atoms are called covalent bonds – bonding through sharing of electrons (much more on this in Chap. 8) Molecules and Molecular Compounds Molecules and Chemical Formulas Molecular formulas give the exact number of atoms of each element in a compound. Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound. Examples: Molecular Empirical C6H12O6 CH2O H2O2 HO C6H6 CH Ionic compounds are written as empirical formulas. Molecules and Molecular Compounds Molecules and Chemical Formulas Molecular (or monatomic) elements are composed of only one atom. Diatomic elements are composed of molecules that contain two identical atoms. H2 N2 O2 F2 Know for Exam! Have No Fear Of Ice Cold Beer Cl2 Br2 I2 Molecules and Molecular Compounds Molecules and Chemical Formulas The subscript to the right of each element symbol indicates the number of atoms of that element in one molecule of the compound. O2 (2 O atoms) CH4 (1 C atom, 4 H atoms) N2O (2 N atoms, 1 O atom) The order in which a molecular formula is written does not necessarily indicate the order of the bonded atoms. Naming Inorganic Compounds Names and Formulas of Binary Molecular Compounds To Name a Molecular Compound: CO2 = carbon dioxide Greek prefixes are added the first and second element to indicate the number of each atom. EXCEPTION: The prefix mono- is never used with the first element. 1. Prefix-first atom (**The prefix mono- is NOT included on the first atom.) 2. Prefix-second atom root-ide Drop “o” or “a” from the prefix if the element begins with a vowel.  Concept Check Naming Molecular Compounds Name the following molecular compounds. Molecular Compound Name N2O5 NO SF6 P2O5 SiO2 Cl2O7 Naming Inorganic Compounds Names and Formulas of Binary Molecular Compounds Common names retained by IUPAC system: water H2O dihydrogen monoxide ammonia NH3 nitrogen trihydride hydrogen peroxide H2O2 dihydrogen dioxide bleach NaClO sodium hypochlorite vinegar CH3COOH acetic acid  Concept Check Naming Molecular Compounds Which of the following chemical formulas is correctly paired with its name? A. N2H4 binitrogen tetrahydride B. CCl4 monocarbon tetrachloride C. P4O10 tetraphosphorus decoxide D. Cl2O7 chlorine oxide E. H2O hydrogen dioxide F. NH3 vinegar Ions and Ionic Compounds Ions are atoms or groups of atoms that have an electrical charge, called cations and anions. Ions and Ionic Compounds Typically, ionic compounds are composed of a metal cation (like Na+) and a non-metal anion (like Cl─). Formula is always written with the cation first. Example: sodium chloride, NaCl Ionic bonds form from the electrostatic attraction between the positive cations and negative anions.  Concept Check Cations and Anions Classify the following descriptions as belonging to either a cation or an anion. Possible Answer Choices: Cation Anion gained electrons lost electrons metal nonmetal negatively-charged ion positively-charged ion written first… written second… Ions and Ionic Compounds Predicting Ionic Charges Certain elements have a predictable charge based on their position in the periodic table. Ions and Ionic Compounds Predicting Ionic Charges Ionic compounds are neutral, so the ion charges add to zero. 𝑵𝒂+ + 𝑪𝒍− → 𝑵𝒂𝑪𝒍 sodium chloride 𝑴𝒈𝟐+ + 𝑪𝒍− → 𝑴𝒈𝑪𝒍𝟐 magnesium chloride Example: What is the formula for calcium nitride? 3 2+ 2 3─ ___Ca + ___N __ → Ca __ __ 3 N __ 2 Always check that these subscripts are in the LOWEST whole-number ratio! Ions and Ionic Compounds Predicting Ionic Charges Oxidation Number is a number assigned to each atom in a compound to represent the “charge” each would have if the electrons were divided among the atoms. Rules for assigning oxidation numbers are on Nomenclature Handout on Canvas. Rules for Assigning Oxidation Numbers (ONs): 1. Pure elements (including those that exist as diatomics) are assigned a zero for their ON. N2 (g) N 0 Hg (l) Ca (s) N 0 ON = 0 ON = 0 Charge 0 Ions and Ionic Compounds Predicting Ionic Charges Rules for Assigning Oxidation Numbers (ONs): 2. Monatomic ions have an ON equal to their common charge. Metals always have a positive charge. Group 1A metals have an ON = +1 Group 2A metals have an ON = +2 Other metals often have a charge equal to their group number on the Periodic Table (NOTE: This is not true for elements like the transition metals that have more than one oxidation number.). Na+ (aq) Zn2+ (aq) Cr6+ (aq) ON = +6 ON = +1 ON = +2 Cr3+ (aq) ON = +3 Ions and Ionic Compounds Predicting Ionic Charges Some metals (i.e. most transition metals) don’t have a common charge. Identify the charge of the anion to which it is bound. Use a Roman numeral to indicate the charge. Examples: Cr2O3 CrO3 Cr +3 O -2 Cr +6 O -2 Cr +3 O -2 O -2 O -2 O -2 Charge +6 + -6 =0 Charge +6 + -6 =0 Ions and Ionic Compounds Predicting Ionic Charges Rules for Assigning Oxidation Numbers (ONs): OF2 O +2 3. The ON of a nonmetal depends on what other F -1 elements it is bonded to, but can be assigned in the following order: F -1 Fluorine (F) has an ON of -1, ALWAYS! Charge 0 Hydrogen (H) has an ON of +1 (Note: H has an ON of -1 when bound to a metal). SO2 S +4 Oxygen has an ON of -2 (Note: O has an ON of +2 when bound to F, O has an ON of -1 in peroxide and -1/2 in superoxides). O -2 Halogens (Group 7A) have an ON of -1. O -2 The most electronegative atom is assigned a charge first Charge 0 (equal to that of its common charge) if all the atoms in a compound are non-metals. Ions and Ionic Compounds Predicting Ionic Charges Rules for Assigning Oxidation Numbers (ONs): 4. The sum of all the ONs in a compound/ion must add up to the overall charge of the compound/ion. For neutral compounds, this charge is zero. P +5 LiH H2O2 H +1 PO43- O -2 Li +1 H +1 O -2 H -1 Charge 0 O -1 O -2 O -1 O -2 Charge 0 Charge -3 When assigning oxidation numbers, always go through each of these steps IN ORDER!  Concept Check Oxidation Number What is the oxidation number for the iron atom in the compound, Fe3(PO4)2? Enter your answer as a whole number. Include the sign. Fe3(PO4)2 Fe PO4 Fe PO4 Fe Charge + =0 Ions and Ionic Compounds Polyatomic Ions The other common type of ionic compound is composed of a metal cation and a polyatomic anion. See Polyatomic Ion Handout on Canvas Polyatomic ions are molecular compounds with a net charge. Metal cation-polyatomic anion compounds are also held together by electrostatic attraction between the (+) and (−) charges. Example: sodium sulfide, sodium sulfite, sodium sulfate Na2S Na2SO3 Na2SO4  Concept Check Writing an Ionic Formula What is the correct formula for an ionic compound composed of ammonium and sulfur? Swap and Drop Method: 1. Write each ion’s symbol A. NH4S2 along with their charges. B. NH4S 2. The charge (without the C. (NH4)4S2 sign) of one ion becomes the subscript D. (NH4)2S of the other ion. E. NH8S 3. Reduce subscripts to the smallest whole number that keep the ratio of ions. Naming Inorganic Compounds Names and Formulas of Ionic Compounds To Name an Ionic Compound: 1. Write the name of the cation metal. If the cation is a polyatomic ion, write the name of the polyatomic ion (i.e. NH4+, ammonium). 2. Name the anion: a) If the anion is an element, use the “root” name and add –ide (i.e. chlorine becomes chloride). b) If the anion is a polyatomic ion, write the name of the polyatomic ion (i.e. carbonate, CO32-). 3. If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses (i.e. Fe3+, iron(III)).  Concept Check Naming Ionic Compounds Name the following ionic compounds. Ionic Compound Name Na2CO3 Al(NO3)3 NH4Cl AgCl Fe2O3 KI  See handout for a list of strong Naming Inorganic Compounds and weak bases to MEMORIZE! Names and Formulas of Acids Strong Acids (Inorganic) Weak Acids oxyacids binary acids HCl hydrochloric acid Acids that are not one of these 7 HBr hydrobromic acid are considered weak. A few examples include: HI hydroiodic acid CH3COOH = acetic acid H2SO4 sulfuric acid H3PO4 = phosphoric acid H2S = hydrosulfuric acid HNO3 nitric acid HCN = hydrocyanic acid HClO4 perchloric acid HCOOH = formic acid HF = hydrofluoric acid HClO3 chloric acid HNO2 = nitrous acid  See handout for a list of strong Naming Inorganic Compounds and weak bases to MEMORIZE! Names and Formulas of Bases Strong Bases Weak Bases The alkali and heavy alkaline earth Some examples of a weak base metal hydroxides are strong include: bases. NH3 = ammonia LiOH C5H5N = pyridine NaOH CH3NH2 = methylamine KOH Ca(OH)2 RbOH Sr(OH)2 CsOH Ba(OH)2 Nomenclature Flow Chart (Handout)  Concept Check Nomenclature Review Categorize and name the following compounds. Compound Category Name CO HCl KOH P4S10 H2 H2O END OF CHAPTER 2 SLIDES