Chapter 3 PDF

Preview

Summary

Summary of Chapter 3, including lecture and book notes...

Description

Chemical compound - 2 or more different elements chemically joined by a bond Chemical bonds as a transformation of matter (new substances created) Electron-level interactions alter electron ownership, forming chemical bonds

Ionic bonds - result from electron transfer from one atom to another, mutual attraction between newly formed cation (+) and anion (-), found in ionic compounds, form between metals and nonmetals

Covalent bonds - result from electrons being shared between 2 or more atoms, mutual attraction between the nucleus (+) and the electrons (-), form between nonmetals and nonmetals Found in molecular compounds (ie. ionic bonds in Group 1A)

Ionic bonds as formed due to the electrostatic interactions (forces) between cations and anions Electrostatic forces - attraction and repulsion between ions ie. NaCl (solid) yields Na+ (positive, aqueous, metal) and Cl (negative, aqueous, nonmetal) Crystalline array of alternating cations and anions (found in ionic substances)

Ionic compounds - hard/brittle solids, high melting/boiling points, high density, tend to dissolve in water, aqueous solutions that conduct electricity Conductivity as a result of dissolution in water, which causes ions to separate

Electrolyte - a substance that dissociates (a physical change) and produces ions when dissolved in water

Resulting solution conducts electricity (ie. NaCl, KCl, acids [despite not being ionic compounds]) Extent of disassociation and ion production in water varies Extent of Dissociation

Strength of Electrolyte

Conductivity

Complete dissociation

Strong electrolyte

Strong

Incomplete dissociation

Weak electrolyte

Poor

No dissociation

Non-electrolyte

None

Strong electrolytes produce soluble salts Weak electrolytes produce acids (ie. vinegar) Non-electrolytes produce table sugar Strength is dependent on charge and size (as charge increases, strength increases; as size increases, strength decreases)

Ionic substances are always compounds as they are composed of 2 or more different elements Covalent bonds can form both elements and compounds Electrostatic intentions in ions defined by Coulomb’s Law (energy as a proportional quantity) Energy = ɑ | charge 1 x charge 2| / internuclear separation^2 (| = absolute value) As charge increases, interaction decreases; as size increases, interaction decreases

As a result of Coulomb's Law, charge plays a greater role than the internuclear separation in determining the strength of electrostatic forces in ionic compounds (ie. KBr vs. K2O - KBr = 1, K2O = 2; K20 has a stronger bond)

Covalent bonds as the sharing of electrons between 2 or more nonmetals (synonymous with a pair of electrons shared between 2 nuclei) Covalent substances exist as molecules (ie. O2 as a molecule, NH3 as both a molecule and a compound) Ionic substances exist as compounds (ie. NaCl as a compound) Molecular substances as being gases, liquids, or solids, having a low melting/boiling point, low density, tend to not dissolve in water, aqueous solutions may or may not conduct electricity Differences between ionic and molecular (covalent) compounds: Ionic Compounds -

Molecular Compounds

Electrons transferred Ionic bonds form Comprised of a metal and a nonmetal or a metal and a polyatomic atom Formula as a smallest whole number ratio

-

Electrons shared Covalent bonds form Comprised of a nonmetal and a nonmetal Formula as the actual number of atoms present

Ionic compounds as being comprised of cations (metals, charged +1 through +7) and anions (nonmetals, charged -1 through -3) Both ions can be polyatomic (ie. LiF, MnO2)

Position on the Periodic Table reveals ion charge; some elements are variable D block/transition metals form more than one cation Each cation has a different charge (ie. Cu+ or Cu2+) Some only form 1 (ie. Ag+, Cd2+, Zn2+) Some non-transition metals form cations of variable charge (ie. Sn2+ or Sn4+)

Polyatomic ions - ions containing 2 or more atoms that carry an overall charge (ie. Carbonate Co3^-2, Sulfite - SO3^-2, Sulfate - SO4^-2) Comprised of 2 or more nonmetals covalently bonded as molecular ions If an atom carries a charge, it creates an ion (which has a fixed charge) (ie. CaCO3 s breaks down into Ca2+ aq and CO3^-2 aq) Polyatomic Cations -

H3O+ (Hydronium) NH4+ (Ammonium)

Polyatomic Anions -

OH- (Hydroxide) NO3- (Nitrate) CN- (Cyanide) ClO3- (Chlorate) SO4 -2 (Sulfate) CO3 -2 (Carbonate) PO4 -3 (Phosphate)

Many polyatomic ions are oxoanions; anions that are made up of oxygen and a nonmetal Some oxoanions come in pairs (with the number of O2 differing by 1) (ie. NO2- and NO3-, SO4 -2 and SO3 -2, ClO3 - and ClO2 -) Base name ends in -ate; 1 less O2 results in the base name ending in -ite (ie. ClO4- is perchlorate, ClO3- is chlorate, ClO2- is chlorite, ClO- is hypochlorite)

Formulas of ionic compounds: Charge neutrality is key (cation charges must equal anion charges; proportional charge will create a neutral ion; use the smallest whole number ratio of ions in the formula) (ie. NaCl = Na+ and Cl-; Na in Group 1A (metal, +1 cation), Cl in Group 7A (nonmetal, -1 anion) Polyatomic ions are inseparable units; use parentheses to indicate their quantity and charge when necessary Diatomic elements are ions in compounds (ie. Cl2 is represented as Cl and Cl)

Name the cation (metal), then the anion (nonmetal) Name of cation and the name of anion; change the anion’s ending to -ide, -ite, or -ate (ie. NaCl = Sodium Chloride, Na2O = Sodium Oxide, NH4Cl = Ammonium Chloride, LiNO3 = Lithium Nitrate) Never use di- or tri- suffixes in naming of ionic compounds Indicate charge of a cation of any transition metal as a Roman numeral (variable charges of these elements) Fe2+ = Fe (II) Mn(OH)2 = Manganese (II) Hydroxide

Copper (I) phosphate has 4 ions; a central polyatomic phosphate (PO4^-3) surrounded by 3 Cu+ ions Ide ending - element as anion (ie. sodium chloride) Ite or Ate ending - polyatomic ion

Nomenclature of Molecular (covalent) Compounds: 2 non-metals with a shared pair of electrons between them (ie. N2O = dinitrogen monoxide, P2O5 = diphosphorus pentoxide, PCL3 = phosphorus trichloride, N2O3 = dinitrogen trioxide all non-metals) When naming binary molecular compounds, use the name of the first element and an -ide ending on the second element Use prefixes before each element to indicate the quantity of each atom present Prefix + 1st element + Prefix + 2nd element with -ide ending

Numerical Value

Prefix

1

Mono

2

Di

3

Tri

4

Tetra

5

Penta

6

Hexa

7

Hepta

8

Octa

9

Nona

10

Deca

Use the name to write out the formula (ie. carbon disulfide = CS2, iodine pentafluoride = IF5, barium bromide = BaBr2, chlorine dioxide = ClO2, copper (II) perchlorate = Cu(ClO4)2, dinitrogen trioxide = N2O3) Naming Acids: Acid - a molecular compound that releases H+ ions when placed in water Properties of acids: corrosive, sour taste, react with metals to form H2 (elemental, diatomic hydrogen gas), react with carbonates to form gaseous CO2, H20, and a salt, turns blue litmus paper red Ionization - molecular compounds (such as acids) release H+ ions in water, producing ions Ion production means that acids in water are electrolytes (ie. HCl (s) yields H+ (aq) and Cl- (aq) in water) Acid ionization of HCl leaves H as a single proton; attaches to H2O to form hydronium ion (H3O+; HCl (g) broken down in H20 yields HCl (aq); HCl (aq) then yields H+ (aq) and Cl- (aq)

in water, thus leading to the dissociation reaction of H+ (aq) and H2O (l) which yields H3O+ (aq)) If ionization (aka dissociation) is complete, a strong acid or electrolyte is formed The six strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO4 If ionization is incomplete, a weak acid or electrolyte is formed As such, all other acids are weak (ie. citric acid and acetic acid) Binary acids are composed of hydrogen and a nonmetal (ie. HF, HCl, HBr, HI) Hydro + base name of non-metal (with an -ic suffix) + acid = acid name Oxoacids contain hydrogen and an oxoanion (a polyatomic anion consisting of oxygen and a nonmetal) The addition of hydrogen offsets the charge of the oxoanion to form an acid (see table below)

Oxoanion Name

Oxoanion Formula

Acid Formula

Acid Name

Sulfate

SO4^-2

H2SO4

Sulfuric acid

Sulfite

SO3^-3

H2SO3

Sulfurous acid

Nitrate

NO3^-

HNO3

Nitric acid

Nitrite

NO2^-

HNO2

Nitrous acid

Carbonate

CO3^-2

H2CO3

Carbonic acid

Hypochlorite

ClO^-

HClO

Hypochlorous acid

Oxoanions ending in -ate take on the base name of the oxoanion with an -ic suffix and acid Oxoanions ending in -ite take on the base name of the oxoanion with an -ous suffix and acid Think sulfate forming sulfuric acid and sulfite forming sulfurous acid

Bases produce hydroxide ions (OH-, polyatomic) when placed in water (either because they contain a hydroxide ion or they react with H2O to form one) (ie. NaOH (s) yields Na+ (aq) and OH- (aq) in H2O; NH3 (s) yields NH4+ (aq) and OH- (aq) However, not all compounds with a hydroxide ion are bases (must be bound to a metal and in an aqueous state) Bases with a hydroxide are classified as strong bases (ie. Group 1A metals - LiOH, NaOH, KOH, RbOH, CSOH; Group 2A metals starting with calcium - Ca(OH)2, Sr(OH)2, Ba(OH)2) Non-hydroxide bases are weak (ie. ammonia and its derivatives - codeine/morphine etc.) To name any compound, analyze its contents (think flowchart) Molecules can be represented by molecular formulas (H2O2), empirical formulas (HO), the Lewis structure, structural formulas, the ball and stick model, and the space filling model...