Topic 04 Bonding. Ch.08 Ions and their compounds Ch.09 Writing and naming formulas CH.10 Covalent Bonding CH.11 Molecular architecture

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1 Topic 04 Bonding Ch.08 Ions and their compounds Ch.09 Writing and naming formulas CH.10 Covalent Bonding CH.11 Molecular architecture Chemistry 10 T03D01

2 Molecular Interactions Intermolecular Forces Interaction between molecules that hold it together in a network. Hydrogen bonding, van der waals forces, etc Intramolecular Forces Forces that hold groups of atoms together and make them function as a unit Ionic and Covalent bonding

3 Intramolecular Forces: Bonding Forces that hold groups of atoms together and make them function as a unit. Ionic bonds transfer of electrons Covalent bonds sharing of electrons

4 IONIC CHARACTER T03 Calculating the Polarity of Bonds If two elements have the same electronegativity values, such as in N 2, O 2 or F 2 they will be completely NONPOLAR As soon as one element has a slightly smaller radius due to an increase in the nuclear charge (compared to electron shielding), that element then has a stronger attraction for electrons (electronegativity) and the molecule begins to become POLAR As one element becomes increasingly stronger, the IONIC CHARACTER increases as the bonds become more polar. Eventually, if the electronegativity difference increases enough, there is a transfer of the shared electron and the bond becomes IONIC As you can see, there are no discreet points. Even something slightly polar has ionic character NONPOLAR POLAR IONIC

Electronegativity The ability or affinity of an atom to attract

5 Bonding Ionic or Covalent How can we determine how two elements will form a bond? Electronegativity The ability or affinity of an atom to attract toward itself the electrons in a chemical bond. OR. Linus Pauling Placement on the Periodic Table

6 Table of Electronegativities

7 What type of bond? Rough range difference in electronegativity 0 is NonPolar Covalent Even sharing 0 > Diff < 1.8 is Polar Covalent Uneven sharing Diff > 1.8 is Ionic transfer of electrons H electron poor region F electron rich region e poor e rich H F +

8 Classify the following bonds as ionic, polar covalent, or nonpolar covalent: Cs Cl = 2.2 Ionic H S = 0.4 Polar Covalent N N = 0 NP Covalent

9 H Bond Determination: M/M = Metallic/Alloy M/NM = Ionic Nm/nm = Covalent Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Se Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds the stair Non metals He Metals Metaloids Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

10 Definition of an Ionic Bond Ionic bonding occurs when one or more electrons are transferred from the valence shell of one atom to another. The atom losing electrons forms a negative ion (Anion) The atom gaining electrons forma a positive ion (Cation) DEF: An ionic bond is the electrostatic interaction between oppositely charged ions + Separate atoms Covalent Molecule Ionic Compound (Ions)

11 Common Ions of Metals Metals lose electrons to form positively charged cations Low ionization energies due to large radius and electron shielding compared to nuclear charge The number of valence electrons determines the most common oxidation state Cations H + Li + Na + K + Mg 2+ Ca 2+ Ba 2+ Al 3+ Name Hydrogen Lithium Sodium Potassium Magnesium Calcium Barium Aluminum

12 Common Ions of NonMetals Nonmetals gain electrons to form negatively charged anions High ionization energies allow nonmetals to hold onto their electrons The electronegativity is high (compared to metals) as their atomic radius is small due to increased nuclear charge relative to electron shielding Easier to hold onto electrons that are closer Number of valence determined the charge Anions O 2 F Cl Br N 3 I S 2 P 3 Name Oxide Fluoride Chloride Bromide Nitride Iodide Sulfide Phosphide

13 Oxidation States of Transition Metals Sc Ti V Cr Mn Fe Co Ni Cu Zn Nomenclature of transition metals demonstrate the oxidation state: Copper (II) sulfate = Cu 2+ and SO 4 2 = CuSO 4 Copper (III) chloride: Cu 3+ and Cl = CuCl

14 Ions Ions form when atoms lose or gain electrons. Atoms with few valence electrons tend to lose them to form cations. Atoms with many valence electrons tend to gain electrons to form anions Na Mg N O F Ne Na + N 3 O 2 F Mg 2+ Cations Anions 14

15 15 Ionic Bonding Ionic bonds result from the attractions between positive and negative ions. Ionic bonding involves 3 aspects: 1. loss of an electron(s) by one element, 2. gain of electron(s) by a second element, 3. An electrostatic attraction between positive and negative ions

16 Positive ions attract negative ions forming ionic bonds. 16 Stable Octet Rule Atoms tend to either gain or lose electrons in their highest energy level to form ions Atoms prefer having 8 electrons in their highest energy level (unless H or He = 2) Examples Na atom 1 One electron extra Cl atom 7 One electron short of a stable octet Na + Ion 0 (8 below) Stable octet Cl Ion 8 Stable octet

How Ions are formed The formation of an ionic compound involves the reaction between a metal and a nonmetal For example in the formation of sodium chloride Sodium atom, Na 2.8.

17 How Ions are formed The formation of an ionic compound involves the reaction between a metal and a nonmetal For example in the formation of sodium chloride Sodium atom, Na Chlorine atom, Cl The ionic bond occurs when the valence electron from the 3 rd shell of Na are transferred to the 3 rd shell of Cl Sodium ion, Na Chloride ion, Cl Each ion now has a stable Noble Gas electron arrangement

18 Ionic Bonding Examples: Transferring electons Li + F Li + F Ionization of Lithium: [He] Li Li + + e Electron Affinity of Fluorine: [Ne] e + F F F Ca Ca 2+ F F F

19 Lewis Dot I II III IV V VI VII VIII Diagrams Transition metals Lewis Dot Diagrams Metal Metalloids are Nonmetals used in both ionic and covalent bonding

20 Lattice Structures Ionic bonds result from the attractions between positive and negative ions. Ionic bonding involves 3 aspects: 1. loss of an electron(s) by one element, 2. gain of electron(s) by a second element, 3. attraction between positive and negative

21 Ionic Bonding The shape and form of the crystal lattice depend on several factors: The size of the ions The charges of the ions The relative numbers of positive and negative ions

22 Ionic Structures + Type of Bonding Examples Composition Nature of Bonding Physical State (STP) Hardness Melting Point Electrical Conductivity (Molten) Solubility NaCl, MgO, CaF 2, Na 2 CO 3, etc Ions Ionic Strong electrostatic attraction between oppositely charged ions Solids Hard and brittle; undergo cleavage High Conductor Usually soluble in polar solvents; insoluble in nonpolar solvents

23 Physical Properties of Ionic Compounds The brittle nature of ionic compounds, and the ability to cleave along planes (think grinding salt) results from the mutual repulsion present when ions are displaced by an applied force Applied Force Repulsion of like ions

24 Ionic Lattice and Conductivity The attractive forces of opposite ions are much stronger than the repulsions of like ions, therefore the an ionic lattice requires large amounts of energy to break up. Results in very high melting and boiling points Conductivity of Ionic Compounds: When in a solid lattice, the electrostatic forces do not allow for the conduction of electricity When molten (liquid) or dissolved in aqueous solution, the lattice is broken up and ions are free to move and transmit charge

25 Writing Ionic Compound Formulas Example: Magnesium sulfide 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. 4. Write the formula without charges Mg 2+ S 2 balanced! MgS

26 Writing Ionic Compound Formulas Example: Potassium Oxide 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. 4. Write the formula without charges K + O Not balanced! K 2 O

27 Writing Ionic Compound Formulas Example: Beryllium Nitride 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. 4. Write the formula without charges Be 2+ N Not balanced! Be 3 N 2

28 first element Metal = Ionic NO prefixes The simplest whole number ratio is generally the ionic formula. (empirical formula) Type I: Binary NaCl Type II: Polyatomic Na 2 CO 3 Type III: Transition FeCl 3 MgBr 2 Li 2 S Mg(NO 3 ) 2 Li 3 PO 4 Ni 3 (PO 3 ) 2 ZrSO 4

29 There are many Polyatomic Ions (1+) CATION (1) ANION (2) ANION (3) ANION Ammonium NH + 4 Chlorate ClO 3 Carbonate CO 2 3 Phosphate PO 3 4 Hydronium H 3 O + Chlorite ClO 2 Sulfate SO 2 4 Phosphite PO 3 3 Hydroxide OH Sulfite SO 2 3 Citrate C 6 H 5 O 3 7 Nitrate NO 3 Hydrogen phosphate HPO 2 4 Ferricyanide Fe(CN) 3 6 Nitrite NO 2 Oxalate C 2 O 2 4 Arsenate AsO 3 4 Acetate CH 3 COO or C 2 H 3 O 2 Chromate CrO 2 4 Borate BO 3 3 Cyanide CN Dichromate Cr 2 O 2 7 Hydrogen bicarbonate HCO 3 Thiosulfate S 2 O 2 3 Hydrogen bioxalate HC 2 O 4 Molybdate MoO 2 4 Hypochlorite ClO Peroxydisulfate S 2 O 2 8 Hydrogen bisulfide HS Silicate SiO 2 3 Hydrogen bisulfate HSO 4 Tartrate C 4 H 4 O 2 6 Hydrogen bisulfite HSO 3 Peroxidate O 2 2 Dihydrogen biphosphate H 2 PO 4 Permanganate MnO 4 Iodate IO 3 Bromate BrO 3 Thiocyanate SCN Periodate IO 4 Perchlorate ClO 4 Perbromate BrO 4

30 Polyatomic Ions Covalent Compounds that contain a charge and will therefore ionically bond though an electrostatic attraction with opposite charges H H N + H H O + H Ammonium NH 4 + Ex: NH 4 Cl H H Hydronium H 3 O + Ex: H 3 OCl (HCl) O O N + O Nitrate NO 3 Ex: AgNO 3 O O P O O Phosphate PO 4 3 Ex: K 3 PO 4 O Sulfate SO 2 4 MgSO 4 O S O O O H O O Hydrogen Carbonate HCO 3 Ex: KHCO 3 H O Hydroxide OH Ex: NaOH O O O Carbonate CO 3 2 Ex: Na 2 CO 3

Made the FIRST periodic table

Made the FIRST periodic table 1869 Mendeleev organized the periodic table based on the similar properties and relativities of certain elements Later, Henri Moseley organized the elements by increasing