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TYPES OF COMPOUNDS

TYPES OF COMPOUNDS. Chemical Family Resemblances. Binary salts. Binary salts are made of a metal and a nonmetal – only two different elements. Examples: NaCl, MnO 2 Binary salts are named with the name of the metal first, then the name of the nonmetal with the “-ide” ending. Example: K 2 O

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TYPES OF COMPOUNDS

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  1. TYPES OF COMPOUNDS Chemical Family Resemblances

  2. Binary salts • Binary salts are made of a metal and a nonmetal – only two different elements. Examples: NaCl, MnO2 • Binary salts are named with the name of the metal first, then the name of the nonmetal with the “-ide” ending. Example: K2O potassium oxide

  3. FORMULAS • The formula unit is the simplest ratio of ions in the salt. Ga2O3 2:3 ratio of gallium atoms to oxygen atoms 2 gallium atoms and 3 oxygen atoms make one formula unit

  4. formulas • Electrons and charge are conserved in a formula unit. • 2 gallium atoms have a total of 6 valence electrons and no charge • 3 oxygen atoms have a total of 18 valence electrons and no charge • so gallium oxide (Ga2O3) has 18+6=24 valence electrons and no charge

  5. conservation • Conservation of electrons and charge in gallium sulfide (Ga2S3)

  6. conservation

  7. oxidation numbers • Oxidation number of an ion is equal to the charge on an ion after it gains or loses electrons. • All atoms gain or lose electrons to try to attain a noble gas configuration (8 valence electrons) • Noble gases have no oxidation numbers

  8. oxidation numbers • Metals – lose all valence electrons, positive (+) oxidation numbers • Metals lose electrons so as to expose full valence shell in next lower level • Alkali metals and hydrogen are +1 • Alkaline earths are +2 • Aluminum and friends are +3

  9. oxidation states • Tin and lead are +2 or +4 • Transition metals vary • Nonmetals – gain electrons, negative (-) oxidation numbers • Enough electrons are gained to complete the valence shell • Oxygen is always –2, and sulfur is –2 unless with oxygen

  10. ternary salts • Halogens are –1 unless with oxygen • Nitrogen and phosphorus are –3 unless with oxygen or halogens • Ternary salts are composed of more than two elements • Ternary salts contain polyatomic ions • Polyatomic ions contain more than one atom example: CO3-2 carbonate

  11. polyatomic ions • Polyatomic anions have a (-) charge, and polyatomic cations a (+) charge • Polyatomic ions act as a unit – the subscripts of the formula may not be changed • Names and formulas • Most names end in “-ate” or “-ite”, which means the ion contains oxygen

  12. naming polyatomic ions • Examples: sulfate (SO4-2), sulfite (SO3-2) • The ending and prefix (if present) indicate the relative number of oxygen atoms in the formula. perchlorate ClO4– chlorate ClO3– chlorite ClO2– hypochlorite ClO–

  13. polyatomic cations • The “-ium” ending means a positive ion (hydronium, H3O+, and ammonium, NH4+) • Multiple ions are indicated by parentheses and a subscript • Example: magnesium hydroxide is Mg(OH)2 • Ammonium sulfide: (NH4)2S

  14. formulas with polyatomic ions • Formulas are made the same way as the binary salts, with the criss-cross method + -2 Na CO3 2 ( ) Ca+2 Ca OH- 2

  15. Naming ternary salts • Ternary salts are named with the metal name first, then the name of the polyatomic ion K3PO4 potassium phosphate

  16. Transition metal salts • Many transition and “other” metals have more than one oxidation number • These numbers are found on some periodic tables • Metals to know: Fe (+2, +3), Cu (+1, +2), Ag (+1), Zn (+2), Sn (+2, +4), Pb (+2, +4), Bi (+3, +5)

  17. transition metal salts • Oxidation number of transition metal is indicated by a Roman numeral in parentheses • FeCl3 is iron (III) chloride • Name these: CrO chromium (II) oxide Cr2O3 chromium (III) oxide CrO3 chromium (VI) oxide

  18. transition metal salts • The Roman numeral is not needed if there is only one oxidation state for the metal (i.e. Zn, Ag, Sc) • The Roman numeral is also used for “other” metal salts like tin (II) fluoride (SnF2, formerly used in toothpaste) • Transition metal salts are often brightly colored

  19. hydrates • Hydrates are salts that have water incorporated into the crystal structure • The water is usually associated with the cation • The number of water molecules in the crystal are specified in the formula MgCl2. 6H2O

  20. hydrates • The dot means they are not chemically bonded • Names of hydrates – “hydrate” plus a prefix is added to the salt name MgCl2. 6H2O is magnesium chloride hexahydrate • Prefix indicates the number of water molecules

  21. hydrate prefixes mono = 1 di = 2 tri = 3 tetra = 4 penta = 5 hexa = 6 hepta = 7 octa = 8 nona = 9 deca = 10

  22. Formation of hydrates • Hydrates can be formed when certain salts are crystallized from water. • Example – CuSO4. 5H2O {copper (II) sulfate pentahydrate} • Hygroscopic compounds become hydrates by taking water from the air.

  23. Formation of Hydrates • Example – sodium carbonate becomes sodium carbonate decahydrate (Na2CO3. 10H2O) • Deliquescent compounds take enough water from the air to form concentrated solutions – examples: calcium chloride (CaCl2), sodium hydroxide (NaOH)

  24. More about polyatomic ions • Bonding • Polyatomic ions form ionic bonds with metals • They are held together with covalent bonds • Formal oxidation states can be assigned to each atom in the ion using the oxidation state rules

  25. Oxidation state rules • Oxygen is always -2 • Hydrogen is always +1 • Sulfur is -2 unless with oxygen • Nitrogen and phosphorus are -3 unless with oxygen • Halogens are -1 unless with oxygen • When with oxygen oxidation states of other atoms vary

  26. Oxidation states of atoms in polyatomic ions • The sum of all the oxidation states must add to the charge of the ion • Carbonate CO3-2 • oxygen – always (–2) charge • total negative charge = (-6) • overall charge is (-2), so carbon must be (+4)

  27. Oxidation states of atoms in polyatomic ions • Try these: • Arsenate AsO4-3 O is -2, As is +5 • Cyanate NCO-1 O is -2, N (not next to O) is -3, so C must be +4

  28. Dot structures of polyatomic ions • All valence electrons must be counted, with extras added for a negative charge. Carbonate CO3-2

  29. Dot structures of polyatomic ions • Make single bonds between all atoms

  30. Dot structures of polyatomic ions • Pair all electrons and make double bonds where necessary to fulfill the octet rule

  31. Dot structures of polyatomic ions • Add brackets and the charge

  32. Dot structures of polyatomic ions • For positive ions, leave out valence electrons for positive charge • Example: Ammonium (NH4+) has 5 + 4 -1 = 8 valence electrons

  33. Molecular substances • Made of molecules, which are loosely held together • Tend to be liquids, gases or low melting solids • Liquids can be purified by distillation • Solids can be purified by recrystallization • Most are insulators

  34. Molecular substances • Molecular elements • Most nonmetals are molecular • Diatomic gases – H2, N2, O2, F2, Cl2, Br2, I2 (BrINClHOF) • Bromine also exists as a liquid, and iodine exists as a solid

  35. allotropes • Many elements exist in more than one molecular form • oxygen (O2) and ozone (O3) • carbon: • charcoal, soot (random arrangement) • graphite (flat sheets) • diamond (three dimensional crystal lattice)

  36. allotropes • fullerenes (hollow balls) • linear acetylenic carbon • -(-CC-CC-CC-)x- • phosphorus (P4): • black (three dimensional, semiconductor) • red (concatenated P4 molecules, used in matches) • white (individual P4 molecules, unstable in air) • Different allotropes have different properties

  37. Formulas and names of small molecules • Many have common names (i.e. water, ammonia) • Systematic names use prefixes for each element – same set of prefixes as for hydrates • P2O5 – diphosphorus pentoxide • N2O – dinitrogen monoxide

  38. Formulas and names of small molecules • “mono” is not used for the first element in a compound • CO2 – carbon dioxide • CO – carbon monoxide • SO3 – sulfur trioxide • CCl4 – carbon tetrachloride

  39. Organic compounds • Covalent carbon containing compounds – usually also contain H; may also contain O, N, S, halogens, P • Many names are derived from alkane names • Alkanes are hydrocarbons (containing only C and H) with all single bonds

  40. alkanes • Alkanes are named for the number of carbons in a chain • CH4 – methane C2H6 – ethane • C3H8 – propane C4H10 – butane • C5H12 – pentane C6H14 – hexane • C7H16 – heptane C8H18 – octane • C9H20 – nonane C10H22– decane • General formula for an alkane: CnH2n+2

  41. alkanes • Carbon always makes 4 bonds in organic compounds, and hydrogen makes only 1 bond • Oxygen makes 2 bonds (two lone pairs), and nitrogen makes three bonds (one lone pair)

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