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General Chemistry (Chem110) Dr. Rima Alharthy Assistant professor in Medicinal and organic synthesis University of Nottingham, UK [email protected] Building 4 room 204 Office hours (Mondays and Wednesdays 8:00-9:30 a.m.) Chapter 2 (p1-p31) Chapter 2: Atoms, molecules and ions 2.1-The Atomic theory 2.3-Atomic number, mass number and isotopes 2.4-The periodic table 2.5-Molecules and ions 2.6-Chemical formulas 2.7-Naming compounds Chapter 2: Dalton’s Atomic Theory (1808) • Elements are composed of extremely small particles called atoms. Examples: Na, Cl, Ca and H2 • All atoms of a given element are identical, having the same size, mass and chemical properties. • The atoms of one element are different from the atoms of all other elements. Chapter 2: Dalton’s Atomic Theory (1808) • Compounds are composed of atoms of more than one element. Examples: H2O, NaCl and CaCO3 • Chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. Chapter 2: The structure of the atom Neutron ± Proton + Electron • Protons and electrons are the only particles that have a charge. • Protons and neutrons have essentially the same mass. • The mass of an electron is so small we ignore it. Chapter 2: Atomic number and mass number Atomic number (Z) Number of protons in nucleus Mass number (A) = Number of protons + number of neutrons = atomic number (Z) + number of neutrons Atomic number Mass number 11 Na 22.98 Sodium Element symbol Element name Example 2.1 Give the number of protons, neutrons and electrons in each of the following species: Element Mass Number Atomic Number Number of electrons Number of protons Number of neutrons 20 11 Na 22 11 Na 17 8 O 14 6 C Chapter 2: isotopes Isotopes Atoms that have the same atomic number (i.e protons and electrons) but different mass numbers (i.e. neutrons) Examples Active, used in nuclear weapons 235 92 U 238 92 U Not active Chapter 2: isotopes Some applications of isotopes • Nuclear reactors and weapons. • Radiation oncology. • Spectroscopy. • Chemical labelling technique. Properties of isotopes • The chemical properties of an element are determined primarily by the protons and electrons. • Neutrons do not take a part in chemical changes under normal conditions. Chapter 2: Periodic table Alkaline Metal Group 1A Alkaline Earth Metal Group 2A Main group elements 1A-7A Group Halides 7A Noble Gases 8A Transition metals 1-8B Period Lanthanides Actanides Chapter 2: molecules Molecule An aggregate of two or more atoms in a definite arrangement held together by chemical forces (bonds). Diatomic molecule polyatomic molecule contains only two atoms contains more than two atoms H2, N2, O2, Br2, HCl, CO O3, H2O, NH3, CH4 Chapter 2: ions Ion is an atom, or group of atoms, that has a net positive or negative charge. Cation Na - 1 electron 11 electrons Anion + 1 electron Na+ Cl 10 electrons 17 electrons If a neutral atom loses one or more electrons it becomes a cation has a positive charge Cl- 18 electrons If a neutral atom gains one or more electrons it becomes an anion. Anion has a negative charge Chapter 2: ions Cation Monoatomic cation Polyatomic cation Na+, K+, Mg2+ NH4+ Anion Monoatomic anion Cl-, Br-,S2- Polyatomic anion OH- Chapter 2: Chemical formulas Molecular formula Shows the exact number of atoms of each element in the smallest unit of a substance Empirical formula Shows the simplest whole-number ratio of the atoms in a substance Molecular H 2O C6H12O6 empirical H 2O CH2O O3 O N 2H 4 NH2 Chapter 2 Formula of ionic compounds Ionic compounds Consist of a combination of cations (metal) and anions (non metal) The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl Chapter 2: Ionic formula Example 2.4 Write the formula of a) magnesium nitride containing Mg2+ and N3- ions. b) Chromium sulfate containing Cr3+ and SO42- ions. c) Titanium oxide containing Ti4+ and O2- ions. Chapter 2 Naming compounds Chemical nomenclature of inorganic compounds 1- Ionic compounds 2- molecular compounds 3- acids and bases 4- hydrates Chapter 2: Naming ionic compounds Ionic Compounds Formed of metal + nonmetal Rules of naming ionic compounds • Starts with the cation (metal) original name. • End with anion (nonmetal), add -ide to element name Examples Mg(OH)2 magnesium hydroxide KNO3 potassium nitrate BaCl2 barium chloride K 2O potassium oxide Chapter 2: Naming ionic compounds Cation Anion Aluminium Al3+ Ammonium NH4+ Barium Ba2+ Cadmium Cd 2+ Calcium Ca2+ Cesium Cs+ Chromium(III)Cr3+ Cobalt (II) Co2+ Copper (I) Cu+ Copper (II) Cu2+ Hydrogen H+ (Proton) Iron (II) Fe2+ Iron (III) Fe3+ Lead(II) Pb2+ Lithium Li+ Magnesium Mg2+ Manganese (II) Mn2+ *Mercury (I) Hg22+ Mercury (II) Hg2+ Potassium K+ Rubidium Rb+ Silver Ag+ Sodium Na+ Strontium Sr2+ Tin (II) Sn2+ Tin (IV) Sn4+ Zinc Zn2+ Bromide Br Carbonate (CO3)2Chlorate (ClO3)Chloride ClChromate (CrO4)2Cyanide CNDichromate (Cr2O7)2Dihydrogen phosphate (H2PO4)Fluoride FHydride HHydrogen carbonate or bicarbonate (HCO3)Hydrogen phosphate (HPO4)2Hydrogen sulfate or bisulfate (HSO4)Hydroxide OHIodide INitrate (NO3)Nitride NNitrite (NO2)Oxide O2Permenganate (MnO4)Peroxide O22Phosphate (PO4)3Sulfate (SO4)2Sulfide S2Sulfite (SO3)2Thiocyanate (SCN)- Chapter 2: Naming ionic compounds • Transition metal ionic compounds can form more than one cation. Therefore, old Roman nomenclature system (Stock system) applies. Examples FeCl2 Iron (II) chloride FeCl3 Iron (III) chloride Cr2S3 Chromium (III) sulfide Chapter 2: Example 2.5 Name the following compounds: a) Cu(NO3)2 b) KH2PO4 c) NH4ClO3 d) PbO e) Li2SO3 Chapter 2: Example 2.6 Write chemical formulas for the following compounds a) Mercury (I) nitrite b) Cesium sulfide c) Calcium phosphate d) Rubidium sulfate e) Barium hydride Chapter 2: naming molecular compounds Molecular binary compounds Formed of nonmetal + nonmetal or nonmetals + metalloids Examples H2O, NH3 and CH4 Rules of naming molecular compounds • Starts with element nonmetal • End with –ide for nonmetal • Omit prefix (eg. mono for first element. • For oxides omit the ending ‘a’ and ‘o’ in the prefix. Chapter 2: naming molecular compounds Examples HI hydrogen iodide SO2 sulfur dioxide N2Cl4 dinitrogen tetrachloride NO2 nitrogen dioxide N 2O dinitrogen monoxide Chapter 2: Example 2.7 Name the following molecular compounds: a) SiCl4 b) P4O10 c) NF3 d) Cl2O7 Chapter 2: Example 2.8 Write chemical formulas for the following compounds a) Carbon disulfide b) Disilicon hexabromide c) Sulfur tetraflouride d) Dinitrogen pentoxide Chapter 2: naming acids and bases Acid a substance that yields hydrogen ions (H+) when dissolved in water. Example HCl •Pure substance •Dissolved in water hydrogen chloride hydrochloric acid Chapter 2: naming acids and bases Rules of naming acids • Starts with Hydro • Ends with element + ic + acid Chapter 2: Naming bases Base a substance that yields hydroxide ions (OH-) when dissolved in water. Rules of naming basse • Starts with element • Ends with hydroxide Examples NaOH Sodium hydroxide KOH Potassium hydroxide Ba(OH)2 Barium hydroxide