Chapter 2: Atoms, Molecules, and Ions PDF

AP* Chapter 2 Atoms, Molecules, and Ions Section 2.1 The Early History of Chemistry  (2.1) The early history of chemistry  (2.2) Fundamental chemical laws  (2.5) The modern view of atomic structure: An introduction  (2.6) Molecules and ion...

AP* Chapter 2 Atoms, Molecules, and Ions Section 2.1 The Early History of Chemistry  (2.1) The early history of chemistry  (2.2) Fundamental chemical laws  (2.5) The modern view of atomic structure: An introduction  (2.6) Molecules and ions  (2.7) An introduction to the periodic table  (2.8) Naming simple compounds  I Section 2.2 Fundamental Chemical Laws 1-Law of conservation of mass (Lavoisier) 2-Law of definite proportion (Proust) 3-Law of multiple proportion (Dalton) 3 Section 2.2 Fundamental Chemical Laws  Law of conservation of mass (Lavoisier):  Mass is neither created nor destroyed in a chemical reaction. 4  The law of mass conservation: mass remains constant during a chemical reaction.  Law of Mass Conservation The total mass of substances does not change during a chemical reaction. reactant 1 + reactant product 2 total mass = total mass calcium oxide + carbon calcium carbonate dioxide CaO + CaCO CO2 3 56.08 g + 100.08 44.00 g g Section 2.2 Fundamental Chemical Laws Important Laws  Law of definite proportion (Proust):  A given compound always contains exactly the same proportion of elements by mass. 7 aw of Definite (or Constant) Composition No matter the source, a particular compound is composed of the same elements in the same parts (fractions) by mass. Calcium carbonate (CaCO3) Analysis by Mass Mass Fraction Percent by Mass (grams/20.0 g) (parts/1.00 part) (parts/100 parts) 8.0 g calcium 0.40 calcium 40% calcium 2.4 g carbon 0.12 carbon 12% carbon 9.6 g oxygen 0.48 oxygen 48% oxygen 20.0 g 1.00 part by mass 100% by mass Section 2.2 Fundamental Chemical Laws f definite proportion (Proust): Example: The % of H in H2O is 11% and the % of O is 89%. These percentages have fixed values in all other samples of H2O regardless of their source. 9  Sample Problem Calculating the Mass of an Element in a Compound PROBLEM: Analysis of 84.2 kg of the uranium containing compound pitchblende shows it is composed of 71.4 kg of uranium, with oxygen as the only other element. How many grams of uranium can be obtained from 102. kg of pitchblende? PLAN: The mass ratio of uranium/pitchblende is the same no matter the source. We can use the ratio to find the answer. 8.65 x 104 g uranium  Sample Problem Calculating the Mass of an Element in a Compound 35 P74: A sample of chloroform is found to contain 12.0 g of carbon, 106.4 g of chlorine, and 1.01 g of hydrogen. If a second sample of chloroform is found to contain 30.0 g of carbon, what is the total mass of chloroform in the second sample? 299g Section 2.2 Fundamental Chemical Laws 3-Law of multiple proportion (Dalton) Section 2.5 The Modern View of Atomic Structure: An Introduction  The atom contains:  Electrons – found outside the nucleus; negatively charged.  Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge.  Neutrons – found in the nucleus; no charge; virtually same mass as a proton. 13 Section 2.5 The Modern View of Atomic Structure: An Introduction  The nucleus is:  Small compared with the overall size of the atom.  Extremely dense; accounts for almost all of the atom’s mass. 14 2.5 The Modern Reassessment of the Atomic Theory 1. All matter is composed of atoms. The atom is the smallest body that retains the unique identity of the element. 2. Atoms of one element cannot be converted into atoms of another element in a chemical reaction. Elements can only be converted into other elements in nuclear reactions. 3. All atoms of an element have the same number of protons and electrons, which determines the chemical behavior of the element. Isotopes of an element differ in the number of neutrons, and thus in mass number. A sample of the element is treated as though its atoms have an average mass. 4. Compounds are formed by the chemical combination of two or more elements in specific ratios. Section 2.5 The Modern View of Atomic Structure: An Introduction Nuclear Atom Viewed in Cross Section  General features of the atom The atom is antoday. electrically neutral, spherical entity composed of a positively charged central nucleus surrounded by one or more negatively charged The atomic nucleus consists of protons and neutrons. electrons.  Table 2.2 Properties of the Three Key Subatomic Particles Charg Mass Location e Name(Symb Relati Absolute(C Relative(amuAbsolute( in the ol) ve )* )† g) Atom Proton (p+) 1 +1.60218x10- 1.0072 1.67262x10- Nucleus + 19 7 24 Neutron (n0) 0 0 1.00866 1.67493x10- Nucleus 24 Outside Electron (e )- 1- -1.60218x10 - 0.0005485 9.10939x10 - Nucleus 19 8 28 * The coulomb (C) is the SI unit of charge. † The atomic mass unit (amu) equals 1.66054x10-24 g. 2.3 (P 44): Atomic Number, Mass Number, and Isotopes A Z X = The symbol of the atom or isotope X = Atomic symbol of the element A = mass number; A = Z + N Z = atomic number (the number of protons in the nucleus) N = number of neutrons in the nucleus Isotope = atoms of an element with the same number of protons, but a different number of neutrons Section 2.5 The Modern View of Atomic Structure: An Introduction Isotopes  Atoms with the same number of protons but different numbers of neutrons.  Show almost identical chemical properties; chemistry of atom is due to its electrons.  In nature most elements contain mixtures of isotopes. 20 Section 2.5 The Modern View of Atomic Structure: An Introduction Two Isotopes of Sodium Section 2.5 The Modern View of Atomic Structure: An Introduction EXERCISE! A certain isotope X contains 23 protons and 28 neutrons.  What is the mass number of this isotope?  Identify the element. Mass Number = 51 Vanadium 22 Section 2.6 Molecules and Ions Chemical Bonds  Covalent Bonds  Bonds form between atoms by sharing electrons.  Resulting collection of atoms is called a molecule. 23 Section 2.6 Molecules and Ions Covalent Bonding To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE 24 Section 2.6 Molecules and Ions Chemical Bonds  Ionic Bonds  Bonds form due to force of attraction between oppositely charged ions.  Ion – atom or group of atoms that has a net positive or negative charge.  Cation – positive ion; lost electron(s).  Anion – negative ion; gained electron(s). 25 Section 2.6 Molecules and Ions EXERCISE! A certain isotope X+ contains 54 electrons and 78 neutrons.  What is the mass number of this isotope? And what is the ion identity? 133,Cs(Caesium) 26 Section 2.6 Molecules and Ions EXERCISE!  Write the symbol for the atom that has an atomic no. of 9(fluorine ) and mass number 0f 19.  How many electrons and how many neutrons does this atom have 27 Section 2.7 An Introduction to the Periodic Table The Periodic Table  Metals vs. Nonmetals  Groups or Families – elements in the same vertical columns; have similar chemical properties  Periods – horizontal rows of elements 28  Chemist noted that the physical and chemical prosperities of certain groups of elements were similar to one another. These similarities with the need to arrange the large volume of available information about the structure and prosperities of elements led to the Development of the Periodic Table. Elements are arranged in the Periodic Table By Atomic Number in horizontal rows called PERIODS. and the Elements are arranged according to their Physical and Chemical Properties in columns called GROUPS. The modern periodic table.  The elements in the Periodic Table can be categorized as METALS, NON METALS, and METALLOIDS. -Metals: are good conductors of heat and electricity. Nonmetals: are poor conductors of heat and electricity. Metalloids: have properties that are intermediate between metals and nonmetals. Section 2.7 An Introduction to the Periodic Table Groups or Families  Table of common charges formed when creating ionic compounds. Group or Family Charge Alkali Metals (1A) 1+ Alkaline Earth Metals (2A) 2+ Halogens (7A) 1– Noble Gases (8A) 0 32 Section 2.7 An Introduction to the Periodic Table The periodic Table song https://www.youtube.com/watch ?v=VgVQKCcfwnU 33 Section 2.8 Naming Simple Compounds Naming Compounds  Binary Compounds  Composed of two elements  Ionic and covalent compounds included  Binary Ionic Compounds  Metal—nonmetal  Binary Covalent Compounds  Nonmetal—nonmetal 34 Section 2.8 Naming Simple Compounds Binary Ionic Compounds (Type I) 1. The cation is always named first and the anion second. 2. A monatomic cation takes its name from the name of the parent element. 3. A monatomic anion is named by taking the root of the element name and adding –ide. 35 Section 2.8 Naming Simple Compounds Binary Ionic Compounds (Type I)  Examples: KCl Potassium chloride MgBr2 Magnesium bromide CaO Calcium oxide 36 Section 2.8 Naming Simple Compounds 1. Give the chemical formula for the following: Lithium phosphide Aluminium Sulfide 2. Give the chemical name for the following: MgI2 CaO2 Section 2.8 Naming Simple Compounds Binary Ionic Compounds (Type II)  Metals in these compounds form more than one type of positive ion (TRANSITION METALS).  Charge on the metal ion must be specified.  Roman numeral indicates the charge of the metal cation.(I,II,III,IV,V,VI,VII,VIII,IX,X).  Transition metal cations usually require a Roman numeral.  Elements that form only one cation do not need to be identified by a roman numeral (Main group elements except Pb and Sn). 38 Some common monatomic ions of the elements. Can you see any patterns? -The Charge of Cations formed by atoms Group 1A, 2A, and 3A equals the Number of their group. -The Charge of Anions formed by atoms in Group 4A through 7A equals their group number minus 8. -Transition elements (B groups) could Have more than one charge of their cations Section 2.8 Naming Simple Compounds Binary Ionic Compounds (Type II)  Examples: CuBr FeS PbO2 40 Section 2.8 Naming Simple Compounds Section 2.8 Naming Simple Compounds Binary Ionic Compounds (Type II)  Examples: CuBr Copper(I) bromide FeS Iron(II) sulfide PbO2 Lead(IV) oxide 42 Common Monoatomic Ions Common ions are in red. Cations Anions Charge Formula Name Charge Formula Name H+ hydrogen H- hydride Li+ lithium F- fluoride +1 Na+ sodium -1 Cl- chloride K+ potassium Br- bromide Cs+ cesium I- iodide Ag+ silver Mg2 magnesium + O2- oxide Ca 2+ calcium S2- sulfide +2 Sr2 strontium -2 + Ba2+ barium Zn2+ zinc Cd2+ cadmium +3 Al3+ aluminum -3 N3- nitride Some Metals That Form More Than One Monatomic Ion Element Ion Formula Systematic Name Common Name Co+2 cobalt(II) Cobalt Co+3 cobalt (III) Copper Cu+1 copper(I) cuprous Cu+2 copper(II) cupric Fe+2 iron(II) ferrous Iron Fe+3 iron(III) ferric Pb+2 lead(II) Lead Pb+4 lead(IV) Sn+2 tin(II) stannous Tin Sn+4 tin(IV) stannic (partial table) Some Common Polyatomic Ions Formula Name Formula Name Cations NH4+ ammonium H 3 O+ hydronium Common Anions CH3COO- acetate CO3-2 carbonate CN- cyanide CrO4-2 chromate OH- hydroxide Cr2O7-2 dichromate ClO4- perchlorate ClO3- chlorate PO4-3 phosphate ClO2- chlorite MnO4- permanganate ClO- hypochlorite SO4-2 sulfate NO2- nitrite SO3-2 sulfite NO3- nitrate Section 2.8 Naming Simple Compounds Polyatomic Ions  Examples of compounds containing polyatomic ions: NaOH Sodium hydroxide Mg(NO3)2 Magnesium nitrate (NH4)2SO4 Ammonium sulfate 46 -Monatomic ions: are ions that consist of just one atom of positive or negative charge. -How can we Indicate the Charge of monatomic ions????? -Because Nobel Gases are stable during to their electronic distribution as we will discuss in the coming chapters, metals try to lose electrons to have the same number of electrons of the corresponding NOBEL GAS. While Nonmetals try to gain electrons to get the same number of electrons of their corresponding NOBEL GAS. -The relation between atoms and their corresponding Nobel Gas is shown in the next figure. The relationship between ions formed and the nearest noble gas. Formulas of Ionic Compounds: In order to write the chemical formula of Ionic Compound you must make the total charge of the compounds equals zero (Electrically neutral) To do so……the following equation must be satisfied: (Number of cationic atoms x their charge) + (Number of anionic atoms x their charge) =0 Example: The chemical formula of the ionic compound formed by the combination between Mg+2 and Cl-1 is MgCl2 (1x2)+(2x-1)=0  Determining Formulas of Binary Ionic Compounds PROBLEM: Write empirical formulas for the compounds named in Sample Problem 2.7: (a) magnesium (b) cadmium iodide nitride (c) strontium (d) cesium sulfide fluoride PLAN: Compounds are neutral. We find the smallest number of each ion which will produce a neutral formula. Use subscripts to the right of the element symbol. SOLUTION : (a) Mg2+ and N3-; three Mg2+(6+) and two N3-(6-); Mg3N2 (b) Cd2+ and I-; one Cd2+(2+) and two I-(2-); CdI2 (c) Sr2+ and F-; one Sr2+(2+) and two F-(2-); SrF2 (d) Cs+ and S2-; two Cs+(2+) and one S2- (2-); Cs2S  Determining Formulas of Binary Ionic Compounds PROBLEM: Write the formulas and the names of the compounds that will result from the following combinations? (a) Sodium and (b) Magnesium and sulfate Sulfur (c) Mercury(I) and Oxygen PLAN: Compounds are neutral. We find the smallest number of each ion which will produce a neutral formula. Use subscripts to the right of the element symbol. SOLUTION : (a) Sodium will form Na+ and Sulfate is SO4-2; Two Na+(2+) and one SO4-2 (2-); Na2SO4 (b) Magnesium will form Mg2+ and Sulfur forms S2- ;One Mg2+(2+); and one Sulfide (-2); MgS (c) Mercury(I) is Hg22+ and Oxygen will form O2- ; one Hg22+ (2+) and one O2- (2-); Hg2O Section 2.8 Naming Simple Compounds Binary Covalent Compounds (Type III)  Formed between two nonmetals. 1. The first element in the formula is named first, using the full element name. 2. The second element is named as if it were an anion (changing the end of its name by ide). 3. Prefixes (mono, di, tri,…..etc.) are used to denote the numbers of atoms present. 4. The prefix mono- is never used for naming the first element. 52 Section 2.8 Naming Simple Compounds Prefixes Used to Indicate Number in Chemical Names 53 Section 2.8 Naming Simple Compounds Binary Covalent Compounds (Type III)  Examples: CO2 Carbon dioxide SF6 Sulfur hexafluoride N2O4 Dinitrogen tetroxide 54 -The Perfix Mono is Generally Omitted for the first element. EXAMPLE: Write the name of the following compounds: -CO2 : -CO: --SO2: --SO3: --NO2: -NO: --N2O5: -NF3: -Cl2O5: -The Perfix Mono is Generally Omitted for the first element. EXAMPLE: Write the name of the following compounds: -CO2 : Carbon Dioxide Not Monocarbon Dioxide -CO: Carbon Monoxide --SO2: Sulfur Dioxide --SO3: Sulfur Trioxide --NO2: Nitrogen Dioxide -NO: Nitrogen Monoxide --N2O5: Dinitrogen Pentoxide -NF3: Nitrogen Trifluoride -Cl2O5: Dichlorine Pentoxide EXAMPLE: Write the Chemical Formula for the following compounds: -Boron trichloride: BCl3 -Sulfur tetrafluoride: SF4 -Tetraphosphorus DecaSulfide: P4S10 -Carbon tetrachloride: CCl4 Section 2.8 Naming Simple Compounds Flowchart for Naming Binary Compounds 58 Section 2.8 Naming Simple Compounds Overall Strategy for Naming Chemical Compounds 59 Section 2.8 Naming Simple Compounds EXERCISE! Which of the following compounds is named incorrectly? a) KNO3 potassium nitrate b) TiO2 titanium(II) oxide c) Sn(OH)4 tin(IV) hydroxide d) PBr5 phosphorus pentabromide e) CaCrO4 calcium chromate 60 Exceptions: Some molecular compounds containing hydrogen do not usually conform to the nomenclature gaudiness. B2H6: Diborane SiH4: Silane NH3: Ammonia PH3: Phosphine H2O: Water H2S: Hydrogen sulfide. (hydrosulfuric acid) Naming Hydrates : Hydrate: an ionic compound contains loosely bonded water. The name of a hydrate follows a set pattern: the name of the ionic compound followed by a numerical prefix and the suffix “-hydrate.” Example:. CaCl2 2H2O Calcium chloride dihydrate The chemical name of the following compounds are: ZnSO4.7H2O CuSO4.5H2O The chemical formula of the following are: Cobalt(II)chloride hexahydrate Sodium carbonate decahydrate