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UNIT 2

UNIT 2. Atoms, Molecules , and Ions Electronic Structure. Molecules. This is a molecule of water. Chemistry happens among the electrons of atoms. Bonds occur between atoms as a result of interactions among the electrons.

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UNIT 2

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  1. UNIT 2 Atoms,Molecules, and Ions Electronic Structure

  2. Molecules This is a molecule of water. • Chemistry happens among the electrons of atoms. • Bonds occur between atoms as a result of interactions among the electrons. • When the interaction results in shared electrons, the bond is said to be covalent and the entity formed is a molecule. • Formed from nonmetals. H O H H – O – H Structural formula H2O chemical (molecular) formula

  3. Special case: Diatomic Molecules • When two atoms of the same element form a covalent bond, the result is a diatomic molecule. • Seven nonmetal elements exist as diatomic molecules in nature: H2, N2, O2, F2, Cl2, Br2, I2 H + H H – H (H2)

  4. Elements that exist as diatomic molecules

  5. A molecular formula shows the number and type of elements in a molecule, e.g., the molecular formula for glucose is C6H12O6. An empirical formula gives only the relative number and type of elements, e.g., the empirical formula for glucose is CH2O. Note: When only one atom of an element is present, no subscript is written. Molecular Formulas The molecular formula for water is H2O. This is also the empirical formula for water.

  6. Types of Formulas • Molecular formula: CH4 • Empirical formula: ? • Structural formulas show the order in which atoms are bonded. • Perspective drawings also show the three-dimensional array of atoms in a compound.

  7. Practice • Give the empirical formula for the substance called diborane, whose molecular formula is B2H6. • Answer: BH3

  8. Practice • Give the empirical formula for the following molecular formulas: • Al2Br6 • C8H10 • C6H4Cl2 AlBr3 C4H5 C3H2Cl

  9. Ions • Atoms of an element with different numbers of neutrons are called isotopes. • Atoms of an element with different numbers of electrons are called ions: • ions with more electrons than the atom are negatively charged and called anions. • ions with fewer electrons than the atom are positively charged and called cations. atomion symbol 23Na name sodium-23 atomic number 11 mass number 23 protons 11 neutrons 12 electrons 11 23Na+ sodium-23 ion 11 23 11 12 10

  10. Ions • When atoms lose or gain electrons, they become ions. • Cations are positive and are formed by elements on the left side of the periodic chart: metals (plus H) • Anions are negative and are formed by elements on the right side of the periodic chart: nonmetals

  11. Periodic Properties – Atom vs. Ion Size • Trends to know: • Cations (+) are smaller than their parent atoms. • Electrons are removed from the outer shell. • Anions (-) are larger than their parent atoms. • Electron-electron repulsion causes the electrons to spread out more in space.

  12. Sizes of Ions - Trends • Ions increase in size as you go down a column.

  13. Ionization Energy = amount of energy required to remove an electron of a gaseous atom or ion to form a cation or more positively charged cation. • The first ionization energy is that energy required to remove first electron. • The second ionization energy is that energy required to remove second electron, etc.

  14. Ionization of Gaseous Sodium: Na (g)  Na+ (g) + e- • The higher the ionization energy increases, the harder it is to remove an electron.

  15. Ionization Energy • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap.

  16. Trends in First Ionization Energies • As one goes down a column, less energy is required to remove the first electron. • valence electrons are farther from the nucleus. Within each row, the ionization energy increases from left to right

  17. Ionization Energy Which element has the higher ionization energy, Br or Ca? Which one will lose an electron easier? • Br has the higher ionization energy • further to the right • Ca will lose an electron easier because its ionization energy is lower.

  18. Electron Affinity • The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity. Cl (g) + e- Cl- (g) • The electron affinity becomes increasingly negative as the attraction between an atom and an electron increases • more negative electron affinity = more likely to gain an electron and form an anion

  19. Electron Affinity • Trends: • Halogens have the most negative electron affinities. • Electron affinities become increasing negative moving from the left toward the halogens. • Electron affinities do not change significantly within a group. • Noble gases will not accept another electron. • To do so would require adding an electron to a new electron shell (significantly higher in energy)

  20. So: Why do ions form to begin with?

  21. Ionic Compounds Bonds occur between atoms as a result of interactions among the electrons. When the interaction is to strip electrons, the resulting bond is said to be ionic and the entity formed is an ionic compound. “Atoms” – now ions – are held together by electrostatic interactions, ionic bonds. This is a crystal of NaCl. Na+ Cl- NaCl formula unit AND empirical formula

  22. Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

  23. Ionic Compounds • Ionic compounds are made of cations and anions, held together by electrostatic attraction: • opposite electrical charges attract each other • like electrical charges repel each other. • Ionic compounds do not exist as discrete molecules, but as structured aggregates (crystals). • In NaCl, an ionic compound, Na exists as Na+ and Cl exists as Cl-. • Important: The overall charge of the ionic compound is ZERO! Equal negative and positive charges.

  24. Properties of Molecular Compounds Properties of Ionic Compounds • held together by covalent bonds • form discrete molecules • soft • low melting point • generally nonconductive • includes all organic compounds • held together by ionic bonds • do NOT form discrete molecules • hard, rigid, brittle • high melting points • conductive when melted or when dissolved in water

  25. Identify Which of the following compounds would you expect to be ionic: N2O, Na2O, CaCl2, SF4? Which of the following compounds are molecular: CBr4, FeS, P4O6, PbF2?

  26. Ion Charges • Metal ions typically have a positive charge. • Group 1A metals always have a +1 charge: • Li+, Na+, K+, etc. • Group 2A metals always have a +2 charge: • Mg2+, Ca2+, Ba2+, etc. • Some metal ions can form differently charged ions (Fe2+ and Fe3+)

  27. Ion Charges • Nonmetal ions typically have a negative charge. • Group 7A nonmetals typically have a -1 charge: • F-, Cl-, Br-, etc. • Group 6A nonmetals typically have a -2 charge: • O2-, S2-, Se2-, etc. !!! Knowing what the Groups mean and knowing where the metal/nonmetal boundary is on the periodic table is a BIG help when dealing with ions and ionic compounds !!!

  28. Practice • Give the chemical symbol, including mass number, for the ion with 22 protons, 26 neutrons, and 19 electrons: Metal or nonmetal? Anion or cation?

  29. Practice • Give the chemical symbol, including mass number, for the ion of sulfur that has 16 neutrons and 18 electrons.

  30. Predict the Charges Predict the charge expected for the most stable ion of barium and for the most stable ion of oxygen. • assume that these elements form ions that have the same number of electrons as the nearest noble-gasatom. 56Ba -- nearest noble gas: 54Xe 8O -- nearest noble gas: 10Ne

  31. Polyatomic Ions • A group of atoms that is covalently bonded yet still has an overall charge is a polyatomic ion. • NO3- • SO42- • PO43- • ClO2- 3- O O P O O phosphate ion !!! You are responsible for knowing the names, symbols, and correct charges for the ions listed in Unit 2 of the syllabus !!!

  32. Nomenclature (Naming) • Inorganic compounds • molecular compounds • ionic compounds • hydrates • acids • oxoacids • acid salts

  33. Naming Binary Molecular Compounds • Name the element farther to the left on the periodic table. (Exception: O is written last unless it’s bonded to F.) • If both elements are in the same group, name the lower one first. • Give the anion name of the second element (ie replace the end of element name with “ide” ending, eg oxygen -> oxide, chlorine->chloride). • Greek prefixes give the number of each atom in the formula. mono- is not used on the first element listed

  34. Nomenclature of Binary Compounds • If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. N2O5: dinitrogen pentoxide pentaoxide

  35. Name them! • Name the following compounds: • (a) SO2 • (b) PCl5 • (c) N2O3

  36. Naming Binary Ionic Compounds • Name the cation (metal). • Name the anion (nonmetal). • replace the end of the nonmetal with –ide • oxygen becomes oxide • fluorine becomes fluoride • sulfur becomes sulfide • more: NaCl BaI2 Ba3P2 K2S sodium chloride barium iodide barium phosphide potassium sulfide

  37. Naming Metal Ions When More Than One Ion is Possible • Two methods • Stock system (Roman numeral is the charge of the cation) • Fe2+ is iron(II) • Fe3+ is iron(III) • Sn2+ is tin(II) • Sn4+ is tin(IV) • Classic (-ic, -ous) system • -ic is for the ion with the higher charge • -ous is for the ion with the lower charge • Fe2+ is ferrous • Fe3+ is ferric • Sn2+ is stannous • Sn4+ is stannic

  38. Naming Binary Ionic Compounds • MgCl2 • CuS • Cu2S • Fe2O3 • Na2O • magnesium chloride • copper(II) sulfide or cupric sulfide • copper(I) sulfide or cuprous sulfide • iron(III) oxide or ferric oxide • sodium oxide

  39. Writing Formulas for Binary Ionic Compounds • The overall ionic compound MUST BE electrically neutral (have a net charge of 0). • If you do not know the charges of the ions in the compound, you will not be able to write the correct formula for the compound! Write the formula for potassium fluoride. 1. Write the two elements K F 2. Write their charges K+ F- 3. If the charges are equal and opposite, then just put the two elements together: KF Note: there are NO charges in the formula!

  40. Writing Formulas for Binary Ionic Compounds Write the formula for silver oxide. 1. Write the two elements Ag O 2. Write their charges Ag+ O2- 3. When the charges are different, perform a swap: Ag+ O2- Ag2O • Naming inorganic compounds: • molecular compounds • ionic compounds • hydrates • acids • oxoacids • acid salts

  41. Writing Formulas for Binary Ionic Compounds calcium iodide titanium(II) nitride lead(IV) chloride iron(III) oxide CaI2 Ti3N2 PbCl4 Fe2O3

  42. Naming Polyatomic Ions and Polyatomic Oxyanions • Polyatomic ions – memorize list in your syllabus. The names and formulas for other polyatomic ions will be provided to you. • Polyatomic oxyanions • sulfate: SO42- (more O’s, -ate) • sulfite: SO32- (less O’s, -ite) • perchlorate: ClO4- (one more O, per-) • chlorate: ClO3- • chlorite: ClO2- • hypochlorite: ClO- (one less O, hypo-)

  43. Writing Formulas for Ionic Compounds with Polyatomic Ions Write the formula for magnesium sulfate. 1. Write the two ions with their charges. Mg2+ SO42- 2. If the charges are equal and opposite, put the two ions together, DO NOT include the charges in the formula. MgSO4

  44. Writing Formulas for Ionic Compounds with Polyatomic Ions Write the formula for ammonium sulfate. 1. Write the two ions with their charges. NH4+ SO42- 2. If the charges are not equal and opposite, do the “swap.” NH4+ SO42- (NH4)2SO4 Note the parentheses!

  45. Writing Formulas for Ionic Compounds with Polyatomic Ions sodium hydroxide magnesium hydroxide aluminum hydroxide NaOH Mg(OH)2 Al(OH)3 aluminum phosphate sodium phosphate ammonium phosphate calcium phosphate AlPO4 Na3PO4 (NH4)3PO4 Ca3(PO4)2

  46. Practice • Name the following compounds: • (a) K2SO4 • (b) Ba(OH)2 • (c) FeCl3

  47. Naming Hydrates • Hydrates are ionic compounds with a specific number of water molecules attached. • Name the ionic compound, use the Greek prefix to indicate the number of waters of hydration, and end with the word “hydrate.” • CuSO4•5H2O copper(II) sulfate pentahydrate • BaCl2•2H2O barium chloride dihydrate

  48. Acids • Are compounds that produce hydrogen ions (H+) in water. They are molecular compounds that ionize in water. “aq” is short for aqueous which means water. • HCl(g)  H+(aq) + Cl-(aq) • or, more correctly, the hydronium ion (H3O+) • HCl(g) + H2O H3O+(aq) + Cl-(aq) In general, acids are hydrogen containing compounds with a formula of HX, H2X, H3X, etc. where X can be an anion from Group 6A or 7A or X can be an oxyanion.

  49. Acid Nomenclature • If the anion in the acid ends in -ide, change the ending to -icacid and add the prefix hydro- . • HCl: hydrochloric acid • HBr: hydrobromic acid • HI: hydroiodic acid

  50. Acid Nomenclature • If the anion in the acid ends in -ate, change the ending to -icacid. • HClO3: chloric acid • HClO4: perchloric acid

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