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Chemical Nomenclature

Chemical Nomenclature. Ions and Ionic Compounds. Ions and Ionic Compounds. Remember an ion is an atom that has lost or gained electrons Cations – positive – lost electrons Anions – negative – gained electrons

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Chemical Nomenclature

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  1. Chemical Nomenclature Ions and Ionic Compounds

  2. Ions and Ionic Compounds • Remember an ion is an atom that has lost or gained electrons • Cations – positive – lost electrons • Anions – negative – gained electrons • Ionic Compounds form when 2 or more atoms are joined by the loss and gain of electrons • ALWAYS cation + anion • Cation is always first and anion is always second.

  3. Ionic Charge • Remember… • for the representative elements you can use the location on the periodic table to determine how many electrons will be lost or gained. • Ex. Hydrogen is +1 • Ex. Fluorine is -1 • For the transition elements you need more information.

  4. Transition Metal Ions • Transition Elements show their charge by the roman numeral that follows their name • Ex. Copper (I) is Cu+1 • Transition elements always lose electrons so their charge is ALWAYS positive.

  5. Naming Ions • The name of cations (positive ions) is the SAME as the name of the element. • Ex. K+1 is potassium • Transition metals do need the charge as a Roman numeral following the name. • So, if the element is in the middle – it needs a roman numeral! • Ex. Cu+1 is copper (I) • The name of anions (negative ions) is different. Anions end in –ide. • Ex. F-1 is fluoride (fluorine)

  6. Write the formula for each ion. • Copper(II) + Sulfur = • Potassium + Nitrogen = • Barium + Sulfur =

  7. Write the formula for each ion. • Lithium + Oxygen = • Calcium + Nitrogen = • Copper(II) + Iodine =

  8. Write the formula for each ion. • Potassium + Sulfate = • Magnesium + Hydroxide = • Ammonium + Sulfite =

  9. Write the formula for each ion. • Calcium + Phosphate = • Aluminum + Nitrate = • Potassium + Chromate =

  10. Write the formula for each ion. • Rubidium + Perchlorate = • Potassium + Permanganate = • Lead(IV) + Hydroxide =

  11. Combining ions in formulas • Compound must be neutral • Charge must equal ZERO. • The Sum of the cation and anion charges must be zero. • Ex. (+1) + (-1) = 0 • Ex. (+2) + (-2) = 0 • We needed one cation and one anion to make the sum ZERO.

  12. Sum must = ZERO • If the charges do notadd up to be zero, then you must add additional cations or anions so that the sum does equal zero. • Ex. (+1) + (-2) ≠ 0 • so… you must add another (+1) ion. • (+1) + (+1) + (-2) = 0 • We needed TWO (+1) cations and ONE (-2) anion to make the sum ZERO.

  13. Formulas The number of cations and anions you need to make the sum ZERO is the ratio of cations to anions in an ionic compound. This ratio is called the FORMULA UNIT for the ionic compound. The ratio is represented in a formula as subscripts for the cation and anion. Since the charges add up to equal zero, NO CHARGE should be written in the formula.

  14. Write the formula for each ion. • Copper(II) + Sulfur = • Cu+2 + S-2 = CuS • Potassium + Nitrogen = • K+1 + N-3 = K3N • Barium + Sulfur = • Ba+2 + S-2 = BaS

  15. Write the formula for each ion. • Lithium + Oxygen = • Calcium + Nitrogen = • Copper(II) + Iodine =

  16. Write the formula for each ion. • Potassium + Sulfate = • Magnesium + Hydroxide = • Ammonium + Sulfite =

  17. Write the formula for each ion. • Calcium + Phosphate = • Aluminum + Nitrate = • Potassium + Chromate =

  18. Write the formula for each ion. • Rubidium + Perchlorate = • Potassium + Permanganate = • Lead(IV) + Hydroxide =

  19. Naming Ionic Compounds • Name cation first • Remember Transition Metals should have a roman numeral representing the charge in the name. • Name the anion second. • Elements that are anions will always end –ide. • Do not change the ending of a polyatomic ion. • Ex. Fluorine  Fluoride (ending change) • Ex. Sulfate  Sulfate (no change)

  20. Ionic Bonds • Cation + Anion • Ions are joined by the transfer of electrons • Creates Electrostatic forces (attraction of opposite charges) that hold the ions together • Ionic Compounds are composed of a continuous arrangement of oppositely charged ions. • NOT a single separate unit

  21. Ionic Bonds • Bonds between atoms will be ionic when there is a LARGE difference in electronegativity between the atoms. • > 1.7 ∆EN • Metal + Nonmetal • Opposite sides of the periodic table – large differences in electronegativity

  22. Properties of Ionic Compounds • Metal + Nonmetal (usually) • Crystalline SOLIDS at room temp. (most) • Crystal – repeating geometric pattern • Brittle • HIGH melting points • HIGH boiling points • Some are so high it takes extreme conditions to get them to change to gas • Conduct electrical currents when melted or dissolved in water

  23. Chemical Nomenclature Molecules and Covalent Bond

  24. Molecular Compounds • Compounds that form when atoms SHARE electrons • Forms a covalent bond • NEVER contain ions • NEVER have charges • Contain only nonmetals

  25. Properties of Covalent Compounds • Nonmetal + Nonmetal • Can be solids, liquids, or gases at room temperature • State is determined by the bond strength (compounds with stronger bonds tend to be solids) • LOW melting points • LOW boiling points • POOR conductors of electricity under any conditions • Can be Polar or Nonpolar • Depends on how the electrons are shared • Polar compounds are better conductors

  26. Covalent Compounds Formulas and Names • Since molecules do not have charges, their names and formulas are determined differently. • Two elements may form covalent bonds in more than one way creating DIFFERENT chemical compounds • Ex. H2O and H2O2 • Other examples: • CH4 and C2H6 • C6H12O6 and C12H22O11

  27. Law of Multiple Proportions • Whenever two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers. • Ex. H2O and H2O2 • contain the same mass of Hydrogen • The mass of Oxygen is in a ratio of 1:2

  28. Names and Formulas • Since two elements can combine in more than one way, we must use information besides their identities to determine their formulas and names. • Scientists perform quantitative analysis to determine the mass ratios of elements to determine their formulas • We can use electron dot diagrams (Lewis structures) to help us determine possible arrangements

  29. Names and Formulas Prefixes Mono- Di- Tri- Tetra- Penta- Hexa- Hepta- Octa- Nona- Deca- • Names • Elements are named in the order they appear in the formula. • The final element’s ending is changed to –ide. • Prefixes are added to each element to identify the number of each atom in the formula • EXCEPTION: Mono- is not used for the first element. If there is only 1 of the first element, no prefix is added.

  30. Names and Formulas • The prefixes in the name can be used to write the formula for a given compound. • Symbols for each element are written in the order they appear. • Subscripts are added based on the prefix of each element.

  31. Names and Formulas: Examples • CO2 • Carbon dioxide • CCl4 • Carbon tetrachloride • N2O5 • Dinitrogenpentaoxide

  32. Names and Formulas: Examples • Carbon Monoxide • CO • Sulfur dioxide • SO2 • Diphosphorustrisulfide • P2S3

  33. Acids • Acids are molecular compounds that behave more like ionic compounds • ALWAYS contain Hydrogen • Hydrogen is always the first element in the formula • Can contain polyatomic ions • Since they form compounds like ionics but consist of only nonmetals, there are special rules for naming and writing formulas for acids.

  34. Acid Formulas • Formulas for acids are determined the same as ionic compounds. • Sum of the charges must equal ZERO • Hydrogen is always +1 • Element or polyatomic ion it combines with must be NEGATIVE • Example: Hydrogen + Chlorine • H+1 + Cl-1 = HCl • Example: Hydrogen + Sulfate • H+1 + SO4-2 = H2SO4

  35. Acid Names • Two different ways to name acids • Hydrogen + Element • Ex. HCl • Hydrogen + Polyatomic Ion • H2SO4

  36. Naming - Hydrogen + ElementHCl • Name begins with Hydro- • Ex. Hydro- • Root of the element • Ex. Hydrochlor- • Ends with –ic acid • Ex. Hydrochloricacid • Practice • HF • H2S

  37. Naming – Hydrogen + Polyatomic Ion • Name begins with polyatomic ion Root • Sulf- • Name ends with –ic acid or –ous acid • -ic acid is used when the polyatomic ion name ends with –ide or –ate • Sulfuricacid (H2SO4) • -ous acid is used when the polyatomic ion name ends with –ite • Sulfurousacid (H2SO3) • Practice • HClO3 • HCN • HClO2 • H3PO3

  38. Chemical Nomenclature Metallic Bonds

  39. Metallic Compounds • Metal + Metal • Consists of positive metal ions with a “sea of electrons” • Electrons are free floating – are not attached to any one atom or ion • HIGH melting points • HIGH boiling points • HIGH conductivity • Malleable, Ductile, High luster (shine) • All properties are a result of the “free” electrons.

  40. Additional Practice • sodium hydroxide • sodium bromide • barium hydroxide • calcium oxide • lithium sulfide • carbon monoxide • sulfur dioxide • iron (II) sulfate • silver (I) chloride • copper (II) hydroxide • ammonium sulfide • nickel (II) fluoride • mercury (I) sulfate • iron (III) oxide • magnesium phosphate • nickel (II) carbonate • diphosphorouspentoxide • aluminum phosphate • nitrogen dioxide • phosphorus trichloride • dinitrogenpentoxide • germanium tetrachloride • scandium bromide • bromine monoiodide • antimony pentasulfide

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